Category: CBSE Notes

By going through these CBSE Class 11 Chemistry Notes Chapter 14 Environmental Chemistry, students can recall all the concepts quickly.

Environmental Chemistry Notes Class 11 Chemistry Chapter 14

→ Environmental Pollution: Atmospheric, Tropospheric & stratospheric.

→ Water Pollution: Causes of water pollution & International standards of drinking water.

→ Soil Pollution: Pesticides-Intecticides & Herbecides & Fungicides

→ Industrial wastes: Recycling of wastes

→ Strategies to Control Environmental Pollution: Waste management.

→ Green Chemistry: Green Chemistry in day to day life.

→ Environmental Pollution: Environmental pollution is the effect of undesirable changes in the surrounding that have harmful effects on animals & plants & human beings.

→ Atmospheric pollution: Tropospheric & stratospheric.

→ Global warming: About 75 % of solar energy reaching the earth is absorbed by the earth’s surface, which increases its temperature. The rest of the heat radiates back to the atmosphere. Some of the heat is trapped by gases such as CO₂, CH₄ & O₃ & C.F.Cs. & water vapours in the atmosphere. Thus they add to the heating of the atmosphere. This causes Global Warming.

→ Green House Effect: Atmosphere traps the sun’s heat near the earth’s surface & keeps it warm. This is called the Greenhouse effect because this makes the earth perfect for life.

→ Acid Rain: When the pH of rainwater drops below 5.6, it becomes acidic. Acid rain is caused by the presence of oxides of N & O in the atmosphere.

Acid rain is harmful to agriculture, trees & plants & washes away nutrients needed for their growth,

→ Classical Smog: It is a mixture of smoke, fog & SO₂ chemically it is a reducing mixture & so it is called reducing smog.

→ Photochemical smog: The main components of the photochemical smog result from the action of sunlight on unsaturated hydrocarbons & nitrogen oxides produced by automobiles & factories. Photochemical smog has a high concentration of oxidizing agents & is, therefore, called oxidizing smog.

→ Ozone hole: In October 1979, A zone of lowered ozone concentration was detected over Antarctica which occurs mainly during September-October & gets replenished during November-December. This is called the ozone hole.

→ B.O.D.: Biochemical oxygen demand.

→ C.O.D.: Chemical oxygen demand.

→ Green Chemistry: It involves processes & products that reduce or eliminate the use or the generation of hazardous substances.

Chapter In Brief:
Environment Pollution:
It is defined as the effect of undesirable changes in our surroundings due to natural sources or human activity that have harmful effects on plants, animals and human beings. A.substance, which causes pollution to one or more components of the ecosystem are called pollutants. Pollutants can be solid, liquid, or gas.

Pollutants are of two types.
1. Non-Biodegradable Pollutants: The materials which do not undergo degradation or degrade very slowly in the environment like DDT, plastic materials, heavy metals, many chemicals, nuclear wastes etc.

2. Biodegradable Pollutants: The materials which are easily decomposed by the micro-organisms either by nature or by suitable treatment are called biodegradable pollutants. Examples; Domestic waste like discarded vegetables, cow dung etc.

→ Atmospheric Pollution: The lowest region of the atmosphere in which human beings along with other organisms live is called the troposphere. It extends up to the height of- 10 km from sea level. Above the troposphere between 10-15 km above sea level lies the stratosphere. Atmospheric pollution is generally studied as tropospheric and stratospheric pollution.

→ Tropospheric Pollution: It occurs due to the presence of undesirable solid or gaseous particles in the air.
1. Gaseous air pollutants: These are oxides of sulphur, nitrogen and carbon, hydrogen sulphide, hydrocarbons, ozone etc.

2. Particulate pollutants: These are dust, mist, fumes, smoke and smog”.
1. Gaseous air pollutants
(a) Oxides of sulphur are produced when sulphur-containing fossil fuel is burnt. SO₂ is poisonous to both animals and plants. Particulate matter present in polluted air catalyses the oxidation of SO₂ to SO₃.
2SO₂(g) + O₂(g) → 2 SO₃(g)

The reaction can also be promoted by ozone and hydrogen peroxide.
SO₂(g) + O₃(g) → SO₃(g) + O₂(g)
SO₂(g) + H₂O₂ → H₂SO₄

(b) Oxides of Nitrogen: At high altitudes when lightning strike’s N₂ and O₂ present in the air combine to form nitric oxide.

Rate of formation of NO₂ Then NO reacts with O₃ is faster.
NO(g) + O₃(g) → NO(g) + O₂(g)
The irritant red haze in the traffic and congested places is due to oxides of nitrogen. NO₂ is a lung irritant.

(c) Hydrocarbons are formed by incomplete combustion of fuel used in automobiles. They are carcinogenic.

(d) Oxides of carbon
1. Carbon monoxide (CO) is one of the most serious air
pollutants. It is highly toxic. It blocks the delivery of oxygen to the organs and tissues. It is produced as a result of incomplete combustion of carbon
C(s) + \(\frac{1}{2}\)O₂(g) → CO(g)
Insufficient quantity

It binds to haemoglobin for which it has 200-300 times more affinity than oxygen and forms carboxy-haemoglobin thus the oxygen-carrying capacity of blood is greatly reduced.

2. Carbon dioxide (CO₂): It is released into the atmosphere by respiration and burning of fossil fuels. It is also emitted during volcanic eruptions and limestone burning to produce lime for cement plants. Normally a percentage of CO₂ in the atmosphere is only 0.03 by volume. The increased amount of CO₂ in the air is mainly responsible for global warming.

Global Warming and Greenhouse Effect:
Some of the heat from solar energy is trapped by gases such as carbon dioxide, methane (CH₄), ozone, chlorofluorocarbons compounds (CFCs) and water vapours in the atmosphere. They add to the heating of the atmosphere. This causes global warming.

The natural greenhouse where there is an ecological balance of CO₂, water vapours which are continuously being converted to carbohydrates and a fresh supply of oxygen during photosynthesis by a lot of green plants, flowers, vegetables makes the earth perfect for life. Sun’s warmth is retained on the earth giving sufficient heat to the soil and plants which emit infrared radiations.

If the amount of CO₂ crosses the delicate 0.03% the natural greenhouse balance is disturbed. It is a major contributor to global warming. CFCs are damaging the ozone layer. Other contributors to global warming are CH₄, N₂O, and ozone. Cutting down forests and trees, burning fossil fuels, production of fertilizers, man-made industrial chemicals etc. add to global warming.

Acid Rain:
Due to the presence of CO₂ in the atmosphere, the pH of rainwater is 5.6.
H₂O(l) + CO2(g) ⇌ H₂CO₃(aq)
H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃ (aq)

When the pH of rainwater falls below 5.6, it is called Acid Rain Various human activities release oxides of Nitrogen and Sulphur in the atmosphere. SO₂ and NO₂ are thus major contributors to acid rain.
2SO₂(g) + O₂(g) + 2H₂O(l) → 2H₂SO₄ (aq)
4NO₂(g) + O₂(g) + 2H₂O(g) → 4HNO₃(aq)

Acid rain is harmful:

1. For plants, trees and agriculture as it washes away nutrients needed for their growth.
2. For human beings and animals as it causes respiratory ailments.
3. An aquatic ecosystem is disturbed when this add rain reaches water objects like rivers, lakes etc.
4. It corrodes water pipes resulting in the leaching of heavy metals such as iron, lead and copper into drinking water.
5. Acid rain damages buildings and other structures made of stone or metal. The Taj Mahal is being corroded by acid rain.
   CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂

As a result, the beautiful white marble of which the Taj Mahal was built centuries ago is rendered lusterless and is being slowly eaten away.

Particulate Pollutants:
They are of two types.
1. The Viable Particulates: They are bacteria, fungi, moulds, algae etc. which are minute living organisms that are dispersed in the atmosphere. They cause allergy to human beings and bring disease to plants.

2. Non-Viable Particulates: They are classified according to their nature and size as follows:
(a) Smoke particulates are solid or a mixture of solid and liquid particles formed during combustion of organic matter,
(b) Dust particles. Over 1 mm in diameter produced during the crushing, grinding and attribution of solid materials, sawdust, sand and cement particles, fly ash, dust storm etc. are the typical examples of this type of particulate emission.
(c) Mists are produced by particles of spray liquids and by condensation of vapours in air, Examples are H2S04 mist, herbicides and insecticides.
(d) Fumes coming out from chemical plants, sublimation, distillation, generally organic solvents, metals and metal oxides form fume particles.

The effect of particulate pollutants are largely dependent on particle size, They are all dangerous to human health.

Smogs:
The word smog is derived from smoke and fog. This is the major cause of air pollution.

Smogs are of two types:
(a) Classical Smog
(b) Photochemical Smog

(a) Classical Smog: It occurs in a cool humid climate. It is a mixture of smoke, fog, and SO2. Because of its reducing nature, it is also called Reducing Smog.

(b) Photochemical Smog: It occurs in a warm, dry and sunny climate. The main components result from the action of sunlight on unsaturated hydrocarbons and oxides of nitrogen. Since it has a high concentration of oxidizing agents, it is called Oxidizing Smog.

Formation of Photo Chemical Smog:
When both unsaturated hydrocarbons (unburnt fuels) and nitric oxide (NO) build-up to sufficiently high levels, a chain reaction occurs in the presence of sunlight.

Ozone is a toxic gas. Both NO₂ and O₃ are oxidising agents.

They further produce chemicals like acrolein (CH₂ = CH CHO) and peroxyacetyl nitrate (CH₃ CO₃ NO₂) or PAN

→ Effects of Photochemical Smog: Photochemical smog causes serious health problems. Both O₃ and PAN present in photochemical smog cause extreme irritation to the eyes. O₃ and NO irritate nose and throat and their high concentration causes a headache, chest pain, dryness of throat, cough and difficulty in breathing. It causes extensive damage to plant life, causes corrosion of metals, building materials, cracking of rubber etc.

→ Control: Photochemical smog can be controlled by controlling—the production of oxides, of nitrogen, controlling the burning of fossil fuels, using catalytic converters in automobiles etc. Certain plants also metabolise NO and therefore their plantation can help.

Stratospheric Pollution
Formation and Breakdown of Ozone: A large amount of ozone (O₂) is present in the stratosphere. O₃ protects us from the harmful UV rays coming from the sun.

→ Reactions occurring in the stratosphere: The main reactions occurring in the stratosphere. In this region ozone is formed in two steps: In the first step ultraviolet radiation coming from the sun have sufficient: energy to split dioxygen into two oxygen atoms. In the second step, the oxygen atoms react with more dioxygen to form ozone.


This ozone absorbs ultraviolet radiation and breaks down into dioxygen and an oxygen atom. Heat is given off, which warms up the stratosphere. Thus there is a dynamic equilibrium between the production and decomposition of ozone molecules.

This ozone layer is acting like a protective layer for the life on the earth. The following reactions take place in this context:

Freons are introduced in the atmosphere from aerosol sprays in which they function as propellants and from refrigerant equipment in which they act as coolants:

This is reactive chlorine destroys the ozone through the following reactions:

The Ozone Hole:
The depletion of the ozone layer by two major compounds [NO₂ and CH₄] is called the ozone hole.

In the summer season, NO₂ and CH₄ react with chlorine monoxide (CIO₂) and chlorine atoms (Cl) forming chlorine sinks, preventing much ozone depletion.

Whereas in winter, a special type of clouds, called polar stratospheric clouds are formed over Antarctica. On the surface of these clouds, chlorine nitrate (CIONO₂) formed gets Hydrolysed to form hypochlorous acid. It reacts also with HCl (above) to give Cl₂
CIONO₂(g) + H₂O(g) → HOCl(g) + HNO₃(g)
CIONO₂(g) + HCl(g) → Cl₂(g) + HNO₃(g)

When sunlight returns to Antarctica in the spring, the sum’s warmth breaks up the clouds and HOCl and Cl₂ are photolysed to gaseous chlorine atoms.

The chlorine-free radicals thus formed, initiate the chain reaction for ozone depletion. One CFC molecule destroys one lac molecules of O₃.

Effects of Depletion of Ozone Layer:

1. The most serious effect is the development of an ozone hole which will allow UV radiations from the sun to pass through the stratosphere and reach the earth causing among other ailments skin cancer.
2. UV rays cause; damage to the cornea and lens of the eye and may cause cataract or even blindness. ,
3. It affects plants chlorophyll, proteins, and causes harmful mutation.
4. O₃ depletion will affect the climate badly. It will upset the heat balance of the earth.
5. It will lead to ecological imbalance which will adversely affect both man and animal.

Water Pollution: Major Water Pollutants

Water Pollutants	Sources
1.. Plant nutrients.	1. Chemical fertilizer
2. Sediments.	2. Erosion of soil agriculture and strip mining.
3. Microorganisms.	3. Domestic sewage.
4. Toxic heavy metals.	4. Industries and chemical tãctories.
5. Organic wastes.	5. Domestic sewage, animal waste, decaying animals and plants and discharge from food processing factories.
6. Pesticides.	6. Chemicals used for killing insects, fungi, and weeds.
7. Radioactive substances.	7. Mining of Uranium containing minerals.
8. Heat	8. Water used for cooling in industries.

Causes of Water Pollution:
1. Pathogens: They are disease-causing agents. They include bacteria and other organisms that enter the water from domestic sewage and animal excreta. Human excreta contain bacteria that cause gastrointestinal diseases.

2. Organic Wastes: The other major water pollutant is organic matter such as leaves, grass, trash etc. Excessive phytoplankton growth within the water is also a cause of water pollution. These are biodegradable. Decomposition of organic matter by bacteria in water causes depletion of dissolved oxygen in the water which is very essential for aquatic life.

3. Chemical Pollutants: Water-soluble inorganic chemicals like cadmium, nickel and nickel salts are an important class of water pollutants. They are dangerous to human life because our body cannot excrete them. Slowly and slowly they damage kidneys, central nervous system, liver etc. Petroleum products like oil spills in oceans, polychlorinated biphenyls (PCBs), though biodegradable are another source of organic chemicals that pollute water.

International Standards For Drinking Water:

1. Fluoride: A deficiency of fluoride in water for drinking is harmful to man and causes tooth decay. Soluble fluoride is added to drinking water to bring its concentration to 1 ppm. Its concentration in water is more than 2 ppm and has harmful effects on teeth, bones.
2. Lead: The upper prescribed limit for the presence of lead in water is 50 ppb. Lead damages the kidney, liver and reproductive system.
3. Sulphate: Excessive sulphate (> 500 ppm) in drinking water causes a laxative effect. At moderate levels it is harmless.
4. Nitrate: The maximum limit of nitrate in drinking water is 50 ppm. Excessive presence can cause disease such as methemoglobinemia (blue baby syndrome)
5. Other metals: The maximum concentration of some common metals recommended in drinking water are given below:

Maximum Prescribed Concentration of Some Metals in Drinking Water:

Metal	Maximum Concentration (ppm or mg dm-3)
Fe	0.2
Mn	0.05
Al	0.2
Cu	3.0
Zn	5.0
Cd	0.005

Soil Pollution
Insecticides, pesticides and herbicides used for the protection of crop cause soil pollution.

Pesticides
Commonly used pesticides which are toxic substances are DDT, Aldrin, and Dieldrin. When their concentration increases beyond a limit, they cause serious metabolic and physiological disorders in higher animals. They are water-insoluble and non—biodegradable. With the passage of time, insects & pests have become immune to these pesticides and insecticides. More potent and biodegradable products like organophosphates and carbonates have beef put to use.

These chemical are extremely harmful to humans and agriculture felid work who spray them.

Nowadays, the pesticide industry is using Herbicides such as sodium chlorate (NaCIO₃), sodium arsenite (Na₃AsO₃) and many others. But all of them are not environmental friendly. Most herbicides are toxic to mammals,

Industrial Waste:
Industrial solid wastes are sorted out as biodegradable and non- Biodegradable wastes. The former is generated by cotton mills, food processing units, paper mills and textile factories, whereas the latter is generated by thermal power plants, iron and steel plants, industries producing Al, Zn and Cu.

The disposal of non-degradable industrial solid waste should be done suitably, otherwise, it may cause a serious threat to the environment. There should be proper management of both domestic and industrial wastes.

Green Chemistry:
Green chemistry as an alternative tool for reducing pollution: One way to protect our environment from chemical effluents and waste is to use green chemistry. By green chemistry, we mean producing the chemicals of our daily needs using such reactions and chemical process which neither use toxic chemicals nor emit such chemicals into the atmosphere. Although it is a very challenging task.

Green chemistry does not employ toxic reagents and severe conditions but uses mild and environmentally friendly reagents, such as sunlight, microwaves, sound waves and enzymes. The use of sunlight and ultraviolet radiation is very useful to produce some useful products which cannot be obtained by simple chemical reactions microwaves and sound waves are used for chemical reactions giving extraordinary results which are not possible by simple chemical reactions.

Enzymes are environmentally friendly reagents. These work in aqueous solutions and at ambient temperatures. These methods are used for preparing medicines and certain antibiotics. For examples, semi-synthetic penicillins such as ampicillin, and Amoxycillin, have been prepared by using this technique.

By using these methods, green chemistry will help us to keep our environment pollution-free.

Green Chemistry in Day-to-Day Life:
1. Dry cleaning of clothes: In places of tetrachloroethene (C₁₂C = CC₁₂) which was earlier used for dry-cleaning nowadays liquefied carbon dioxide along with a suitable detergent is used. It will cause less harm to groundwater. Hydrogen peroxide (H₂O₂) is also nowadays used for bleaching clothes which gives better results.

2. Bleaching of paper: Cl₂ gas which was earlier used for bleaching paper-has been replaced by hydrogen peroxide (H₂O₂) with a suitable catalyst.

3. Synthesising chemicals: Acetaldehyde (CH₃CHO) is now prepared by one-step oxidation of ethylene (H₂C = CH₂) in the presence of an anionic catalyst in an aqueous medium with a yield of 90% on a commercial scale.

In a nutshell, Green chemistry is a cost-effective approach, which involves

1. Reduction in cost
2. Reduction in energy consumption
3. Reduction in waste production.

By going through these CBSE Class 11 Chemistry Notes Chapter 13 Hydrocarbons, students can recall all the concepts quickly.

Hydrocarbons Notes Class 11 Chemistry Chapter 13

→ Classification-classification of hydrocarbons.

→ Alkanes-Nomenclature. isomerism, preparation, properties of alkanes, conformations.

→ Alkenes-structure of double bonds, Nomenclature, Isomerism, preparation and properties.

→ Alkynes-Nomenclature of isomerism, the structure of the triple bond, preparation & properties.

→ Aromatic hydrocarbons-Nomenclature & isomerism structure of benzene, Aromaticity, preparation & properties.

→ Directive influence of a functional group in mono-substituted benzene.

→ Carcinogenicity & Toxicity Benzene of polynuclear hydrocarbons.

→ Hydrocarbons: Hydrocarbons are the compounds of carbon & hydrogen only. Hydrocarbons are mainly obtained from coal & petroleum.

→ Petrochemical: Petrochemicals are the prominent starting material used for the manufacture of a large number of commercially important products.

→ L.P.G.: Liquified petroleum gas

→ C.N.G.: Compressed natural gas.

→ Classification of hydrocarbons: Saturated, unsaturated, cyclic (alicyclic) & Aromatic

→ Important reactions of Alkanes: Free radical substitution, combustion, oxidation & aromatization.

→ Alkenes & Alkynes: Undergo mainly addition reactions, (electrophilic additions).

→ Aromatic hydrocarbons: Despite having unsaturation undergo mainly electrophilic substitution reactions

→ Conformation Isomerism: Alkanes show conformational isomerism due to free rotation along with the C – C sigma bonds. Out of staggered of the eclipsed conformations of ethane, staggered conformations are more stable as hydrogen atoms are farthest apart.

→ Geometrical Isomerism: Alkanes exhibits geometrical isomerism (cis-trans) due to restricted rotation around the carbon-carbon double bond

→ Huckel Rule: Benzene of benzenoid compounds show aromatic character. Aromaticity, the property of being aromatic is possessed by compounds having specific electronic structure characterized by Huckel Rule (4n + 2) π electron rule.

→ Carcinogenic property: Some of the polynuclear hydrocarbons having fused benzene ring system have carcinogenic property.

→ Activation & deactivation of benzene ring: The nature of groups or substituents attached to the benzene ring is responsible for activation or deactivation of the benzene ring towards further electrophilic substitution and also for orientation of the incoming group.

Friedel-crafts Reaction:
1. Friedel craft alkylation reaction

2. Friedel craft acylation reaction


Structure of Benzene

Kekule’s Structure

→ Markownikov Rule: The rule states that the negative part of the addendum gets attached to that carbon atom which possesses a lesser number of hydrogen atoms as:
CH₃ – CH = CH₂ + HBr →

2-Bromopropane (Main product)
(ii) CH₃ CH₂ CH₂Br
1-Bromopropane (Minor product)

→ Lindlar’s Catalyst: Partially deactivated palletised charcoal is known as Lindlar’s Catalyst.

Chapter In Brief:
Hydrocarbons are the compounds of carbon and hydrogen only, Alkanes, Alkenes, alkynes, and aromatic compounds constitute hydrocarbons. Alkanes are saturated hydrocarbons containing carbon-carbon single bonds. Alkenes are unsaturated hydrocarbons containing at least one

double bonds, whereas alkynes are unsaturated hydrocarbons containing at least one — C ≡ C — triple bond.

Alkanes: Earlier known as paraffin, the general formula of their homologous series is C_nH_2n+2.

Methane, the first member is having a tetrahedral shape according to VSEPR Theory. It is multiplanar in which a carbon atom lies at the centre and four hydrogen atoms lie at the four corners of a regular tetrahedron. H-C-H bond angle is 109.5°. In alkenes, C-C and C-H bond lengths are 154 pm and 112 pm respectively. C-C and C – H bonds are formed by head-on the overlapping of sp³ hybrid orbitals of carbon and Is atomic orbitals of hydrogen atoms.

Nomenclature & Isomerism in Alkanes:
The first three members of the alkane family namely methane, ethane and propane have only one structure but higher alkanes can have more than one structure.
e.g. C₄H₁₀ have the following two structures


They are called Chain Isomers
Similarly, C₅H₁₂ have the following three structures
1. CH₃-CH₂-CH₂-CH₂-CH₃
: Pentane (n-pentane) b.p. 309 K
2.
: 2-Methyl butane (isopentane) b.p 301K
3.
: Dimethylpropane (neopentane) b.p. 282 K.

1, 2, 3 are the chain isomers of pentane. They differ in their boiling points and other properties, though they have the same molecular formula. This difference in properties is due to the difference in their structures, they are termed Structural Isomers.

Preparation of Alkanes:
Petroleum and natural gas are the main sources of alkanes. However, alkanes can be prepared by the following methods.
1. From unsaturated hydrocarbon by hydrogenation.

2. From alkyl halides:
1. Alkyl halides (except fluorides) on reduction with zinc and dilute hydrochloric acid give alkanes,

2. By Wurtz reaction: Alkyl halides on treatment with sodium in dry ether give higher alkanes. This method is used to prepare higher alkanes containing an even number of carbon atoms.

3. From carboxylic acids
1. Sodium salts of fatty acids on heating with soda-lime [a mixture of NaOH + CaO] give alkanes. The process is called decarboxylation [Removal of a molecule of CO₂]

2. Kolbe’s electrolytic method: An aqueous solution of sodium or potassium salt of a carboxylic acid on electrolysis gives alkanes containing an even number of carbon atoms.

Properties Of Alkanes
(A) Physical Properties:

1. Alkanes are almost non-polar due to the covalent nature of C-C and C—H bonds and due to very little difference of electronegativity between C and H atoms. Therefore, they are insoluble in water but soluble in organic solvents.
2. Due to weak van der Waals forces, the first four members (from C₁ to C₄) are gases. The next thirteen (C₅ to C₁₇) are liquids and those containing 18 carbon atoms or more solids at 298 K.
3. They are colourless and odourless.
4. Their boiling points increase with the increase in molecular mass as shown in the table below.

Table: Variation of melting point and boiling point in alkanes

It is due to fact that intermolecular van der Waals forces increase with the increase in molecular size or surface area of the molecules. For example, among the isomeric pentanes.

B. Pt of n-pentane is highest (309.1 K), whereas that of 2, 2- dimethyl propane is the lowest (282.5 K). With the increase in the number of branched chains, the molecule attains the shape of a sphere. This results in decreased surface area and hence weaker intermolecular van der Waals forces thus lowering the boiling points.

Chemical Properties Of Alkanes
1. Substitution Reaction: Halogenation

Order of reactivity of halogens is F₂ > > Cl₂ > Br₂ > I₂
Rate of replacement of hydrogens of alkanes is: 3° > 2° > 1°
Fluorination is too violent to be controlled.

Bromination is similar. Iodination is very slow and a reversible reaction. It can be carried out in the presence of some oxidising agents like HNO₃ or HIO₃.
CH₄ + I₂ ⇌ CH₃I + HI
HIO₃ + 5HI ⇌ 3I₂ + 3H₂O
Substitution of halogens in alkanes proceeds via a free-radical mechanism.

2. Combustion: Alkanes on heating in the presence of air or oxygen are completely oxidised to carbon dioxide and water with the evolution of a large amount of heat.
CH₄(g) + 2O₂ → CO₂(g) + 2H₂O (l); Δ_cH°=- 890 kJ mol^-1
C₄H₁₀(g) + 6\(\frac{1}{2}\)O₂(g) → 4 CO₂(g) + 5H₂O (1); Δ_cH° = -2876 kJ mol^-1

Due to the evolution of large amount of heat during combustion, alkanes are used as fuels.
During incomplete combustion in insufficient supply of air or Oxygen, carbon black is formed.

3. Controlled Oxidation: In a regulated supply of air or oxygen at high pressure and in the presence of suitable catalysts, alkanes give a variety of oxidation products.

(iv) Ordinarily alkanes resist oxidation but alkanes having tertiary H atoms can be oxidised to corresponding alcohols by KMnO₄.

4. Isomerisation: n-Alkanes on heating in the presence of anhydrous aluminium chloride and hydrogen chloride gas isomerises to branched-chain alkanes.

5. Aromatisation: n-alkanes having six or more C atoms on heating to 773 K at 10-20 atmospheric pressure in the presence of oxides of V, Mo or Cr supported over alumina gel dehydrogenated and cyclised to benzene and its homologues. This reaction is termed Aromatisation or reforming.

6. Reaction with steam

7. Pyrolysis: Higher alkanes on thermal decomposition give lower alkanes, & a mixture of alkanes. The process is also called Cracking.

Conformations:
Alkanes contain C-C sigma (a) bonds. Free-rotation around C – C bond is possible. Such different Spatial, arrangement of atoms obtained by rotation around the C — C bond is called Conformations or Conformers or Rotamers.

Ethane (C₂H₆) has two major conformational isomers amongst several spatial arrangements differing from each other by a small energy barrier.

One is called eclipsed form which is less stable as it is associated with more energy [due to repulsion of electrons] and the other is called staggered form which is more stable as it is associated with lower energy. Any other intermediate confrontation is called a skew form.

Eclipsed and staggered forms of ethane (C₂H₅) can be represented by Sawhorse and Newman Projections as shown below of all the conformations of ethane.
1. Sawhorse Projections

Sawhorse projections of change

2. Newman’s Projections

Newman projections of ethane

The staggered form has the least torsional strain and the eclipsed form the maximum torsional strain. The energy difference between the two extreme forms is of the order of 12.5 kJ mo^-1 which is very small. These forms have not been separated.

Alkenes. Alkenes are unsaturated hydrocarbons containing at least one

(double bond) or —C = C— (Triple bond)
They are also called Olefins. The general formula of alkenes is C_nH_2n.

Structure of Double Bond
Cabon atoms constituting a double bond undergo sp² hybridisation. The double bond contains one strong sigma (a) bond and one weak Pi (π) bond. The electrons of the π bond are delocalised and is thus a source of electrons. Any electrophile can come and attack it. That is why alkenes undergo electrophilic addition reactions.

The double bond is shorter in bond length (134 pm) than the C-C single bond (154 pm), π bond is a weaker bond due to poor overlapping between the two 2p orbitals. The strength of the double bond (bond enthalpy 681 kJ mol^-1) is greater than that of a C—C single bond (bond enthalpy 348 kJ mol^-1) in ethane.

Orbital picture of ethene depicting bonds only


Original picture of ethene showing formation of (a) π-bond. (b) π-cloud and (c) bond angles and bond lengths

Nomenclature of Alkenes:
In the IUPAC system, the longest chain of carbon atoms containing the double bond is ‘selected’. The numbering of the chain is done from the end which is nearer to the double bond. The suffix ‘ene’ replace ‘ane’ of alkanes.

Put n = 2 in C_n H_2n; C₂H₄ or H₂C = CH₂ is ethylene (common name) and ethene in IUPAC system.

Isomerism in Alkenes
Alkenes show both structural isomerism and geometrical isomerism.
(a) Chain isomerism

(b) Position isomerism
CH₃ – CH₂ – CH = CH₂ But -1-ene
and CH₃ – CH = CH – CH₃ But-2-ene
are position isomers as they differ in the position of the functional group.

(c) Geometrical isomerism
C_xy = C_xy and C_xy type of alkenes show geometrical isomerism
e.g. But-2-ene CH₃ – CH = CH – CH₃ exists in two forms- called geometrical or cis-trans isomers as shown below.

When the identical atoms or groups lie on the same side of the double bond it is called cis-isomer
When the identical atoms or groups lie on the opposite side of the double bond it is called trans-isomer.

The restricted rotation of atoms or groups around the doubly bonded carbon atoms gives rise to different geometries to such compounds. The stereoisomers of this type are called geometrical isomers.

Due to different Spatial arrangements of atoms or groups, geometrical isomers differ in their properties like m.p., b.p., dipole moment, solubility etc.

Cis-form of but-2-ene is more polar than the transform. (Dipole moment) p of cis-form is 0.35 Debye whereas p of transform is almost zero, or trans-2-butene is non-polar. In the transform, two methyl groups being in opposite directions cancel polarities due to each C – CH₃ bond.

In the case of solids, it is found that the trans isomer has higher m.p. than cis form. This is due to the better symmetry of the trans-isomers. Trans solids fit well into the crystal lattice.

Preparation 0f Alkenes
1. From Alkynes: Alkynes on partial reduction with a calculated amount of dihydrogen in the presence of partially deactivated palletised charcoal called Lindlar’s Catalyst to give cis-alkenes. However, alkynes on reduction with sodium in liquid ammonia form trans-alkenes.

2. From Alkyl Halides: Alkyl halides on heating with alcoholic potash (potassium hydroxide dissolved in alcohol) undergo dehydrohalogenation to give alkenes. This is an example of a β-Elimination reaction.

For halogens, the rate of reaction is Iodine > bromine > chlorine while for alkyl groups, it is tert > sec > prim.

3. From vicinal dihalides: Vicinal (on two adjacent C atoms) dihalides on treatment with zinc undergo dehalogenation to give alkenes.

4. From the acidic dehydration of alcohols: Alcohols on heating with conc. H₂SO₄ lose a molecule of H₂O (β-elimination reaction) to form an alkene.

Properties of Alkenes:
Physical properties:

1. The first three members are gases, the next 14 are liquids and the higher ones are solids.
2. Except for ethene, which has a pleasant smell, all alkenes are odourless and colourless.
3. They are insoluble in water but fairly soluble in non-polar solvents like benzene, petroleum, ether etc.
4. They show a regular increase in b.p. with an increase in size [For every — CH₂— group added b.p. increases by 20—30 K] Like alkanes, straight-chain alkenes have higher b.p. than isomeric branched ‘ alkenes.
5. Like alkanes, alkenes are generally non-polar but certain, alkenes are weakly polar due to their unsymmetrical geometry.

Chemical Properties:
(a) Addition reactions:
1. Addition of H₂ (catalytical hydrogenation)

2. Addition of halogens: [Electrophilic addition] Br₂ is a reddish-orange liquid that adds to the unsaturated site to give a colourless product. This reaction is used as a test of unsaturation.

3. Addition of hydrogen halides
The order of reactivity is HI > HBr > HCl
CH₂ = CH₂ + H – Br → CH₃ – CH₂Br
Markovnikov Rule. [Addition of HX to unsymmetric alkenes] “The negative part of addendum (the molecule to be added) goes to that carbon atom of the unsymmetrical alkene which is attached to lesser number of carbon atoms”.

The modern version of Markovnikov Rule. The product is formed from the more stable carbocation.

The more stable carbocation [which predominates because it is former faster] reacts with Br- to form the product.

Anti-Markovnikov Addition or Peroxide/Kharash Effect

In the presence of peroxide, the addition of HBr to unsymmetrical alkenes like propene takes place contrary to the Markovnikov rule. This happens only with HBr but not with HCl and HI. This addition reaction was observed by M.S. Kharash and F.R. Mayo in 1933 at the University of Chicago. This reaction is known as peroxide or Kharash effect or addition effect or addition reaction anti to Markovnikov rule.

Mechanism: Peroxide effect proceeds via free radical chain mechanism as given below:


The secondary free radical obtained in the above mechanism (iii) is more stable than the primary. This explains the formation of 1 — bromopropane as the major product. It may be noted that the peroxide effect is not observed in addition to HCl and HI.

This may be due to the fact that the H – Cl bond being stronger (430.5 kJ mol^-1) than H — Br bond (363.7 kJ mol^-1), is not cleaved by the free radical, whereas the H – I bond is weaker (296.8 kJ mol^-1) and iodine free radicals combine to form iodine molecules instead of an addition to the double bond.

4. Addition of sulphuric acid. Cold, concentrated sulphuric acid adds to alkenes in accordance with the Markovnikov rule as a result of electrophilic addition to form alkyl hydrogen sulphate.

5. Addition of water. In the presence of a few drops of the cone. H₂SO₄, alkenes undergo hydration with water in accordance with the Markovinkov rule to form alcohols.

6. Oxidation: (a) Alkenes on reaction with cold, dilute, 1 % alkaline potassium permanganate (KMnO₄) solution called Baeyer’s Reagent produce vicinal glycols. The colour of KMnO₄ is discharged, It is also used as a test of unsaturation.

(b) Acidic KMn04 or acidic K₂Cr₂O₇ oxidizes alkenes to ketones and/or acids depending upon the nature of the alkene and the experimental conditions

7. Ozonolysis. It involves the addition of O₃ molecules to the alkene to form ozonide followed by cleavage by Zn/H₂O to form aldehydes and ketones.


8. Polymerisation. When a large number of ethene molecules combine at high temperature, high pressure in the presence of a catalyst, Polythene is obtained.

These polymers are of great use in the manufacture of plastic bags, squeeze bottles, toys, pipes radio and TV cabinets, milk crates, plastic buckets and other moulded articles.

Alkynes: Alkynes are unsaturated hydrocarbons containing at least one — C = C — triple bond.
General formula: C_nH_2n-2
Common & I.U.P.A.C. names of Alkynes
n = 2 C₂H₂ H – C ≡ C – H Acetylene Ethyne

n = 3 C₃H₄ CH₃ — C ≡ CH MethylacetylenePropyne

n = 4 C₄H₆

• CH₃CH₂C = CH Ethylacetylene But-l-yne
• CH₃ – C = C-CH₃ Dimethylacetylene But-2-yne

n = 5 C₅H₈

Structures (i) and (iii) are position isomers
Structures (i) and (ii) and (iii) are chain isomers
Structure of Triple Bond.

Each carbon atom of ethyne has two sp hybridized orbitals. Carbons-carbon sigma (a) bond is obtained by the head-on overlapping of the two sp hybridised orbitals of the two carbon atoms. The remaining sp hybridised orbitals of the two carbon atoms. The remaining sp hybridized orbital of each carbon atom undergoes overlapping along the internuclear axis with the Is orbital of each of the two hydrogen atoms forming two C — H sigma bonds. H – C—C bond angle is 180°.

Each carbon has two unhybridised p orbitals which are perpendicular to each other as well as to the plane of the C – C sigma bond. The 2p orbitals of one carbon atom are parallel to the 2p orbitals of the other carbon atom, which undergo lateral or sideways overlapping to form two pi (p) bonds between two carbon atoms. Thus ethyne molecule consists of one C — C s bond, two C >- Hs bonds and two C — C p bonds.

The strength of the C = C bond (bond enthalpy 823 kJ mol^-1) is more than those of the C = C bond (bond enthalpy 681 kJ mol^-1) and C – C bond (bond enthalpy 48 kJ mol^-1). The C = C bond length is shorter (120 pm) than those of C = C (134 pm) and C – C) (154 pm). The electron cloud between two carbon atoms is cylindrically symmetrical about the internuclear axis. Thus ethyne is a linear molecule.

Orbital picture of ethyne showing (a) sigma overlaps (b) pi overlaps bond angles and bond lengths.

Preparation of Acetylene (Ethyne). Commercially, it is prepared by the action of water on calcium carbide.
CaC₂ + 2H₂O → Ca(OH)₂ + C₂H₂ (Ethyne)

2. From Vicinal Dihalidies. Vicinal dihalides on treatment with alcoholic potassium hydroxide undergo dehydrohalogenation.

Properties of Alkynes:
Physical properties.

1. First, three members are gases, the next eight are liquids and the higher ones are solids.
2. All alkynes are colourless.
3. Except for enthene which has a characteristic odour, others are odourless.
4. Alkynes are weakly polar in nature.
5. They are lighter than water and immiscible with water but soluble in organic solvents like ethers, benzene etc.
6. Their m, p., b, p, and density increases with an increase in molar mass.

Chemical Properties: Alkynes show usual addition reactions, acidic reactions and polymerisation reaction.
A. Acidic character of alkynes: Unlike alkenes, ethyne shows acidic reactions.

Alkanes, alkenes, and alkynes follow the following trend in their acidic behaviour.

1. H – C ≡ C – H > CH₂ = CH₂ > CH₃ – CH₃
2. H – C ≡ C – H > CH₃ – C ≡ CH > > CH₃ – C ≡ C – CH₃

B. Addition reactions,
1. Addition of dihydrogen.
Alkynes contain a triple bond. Therefore, they add up two molecules of H₂.

2. Addition of Halogens. When Br2 is added to alkynes, the reddish-orange colour of Br₂ disappear. It is a test of unsaturation.

3. Addition of hydrogen halides [HCl, HBr, HI]
Two molecules get added to alkynes to form gem dihalides.

4. Addition of water

5. Polymerisation

Aromatic Hydrocarbons or Arenes: Aromatic compounds containing benzene ring are known as Benzenoids and those not containing a benzene ring are called Non-Benzenoids. Some of the arenas are given below.


Where o = ortho (1,2)
m = meta (1, 3)
p = para (1, 4)

Structure of Benzene:

1. Molecular formula C₆H₆ indicates that benzene is an unsaturated hydrocarbon.
2. The unusual stability of benzene and no change of orange-red colour of Br₂ in addition to benzene ruled out the open chain structure of benzene.
3. It forms a triozonide which indicates the presence of three double bonds.
4. Benzene produces one and only one monosubstituted derivative which indicates that all the six-carbon and six hydrogen atoms of benzene are identical.
5. A. Kekule’ in 1865 proposed the cyclic structure for benzene with alternate single and double bonds in carbon atoms with each C atom carrying one hydrogen.

Kekule’ suggested the oscillating nature of double bonds.

Resonance and Stability of Benzene
Benzene is a resonance hybrid of various, resonating structures. The two structures are given above by Kekule’ are the main contributing st; lectures. The hybrid structure is (c) is given below.

The circle represents the six electrons that are delocalised between the six carbon atoms of the benzene ring.

Orbital Picture of Benzene.
All the six carbon atoms of benzene are sp² hybridised. Two of these three sp² hybrid orbitals of each C atom overlap with sp² hybrid orbitals of adjacent C atoms to form six C – C single bonds which are in the hexagonal plane. The remaining sp² orbital of each C atom overlaps with the s-orbital of each hydrogen atom to form six C — H single sigma bonds. Each C atom is now left with one unhybridised p- orbital perpendicular to the plane of the ring as shown on the next page.

The unhybridised p orbital of carbon atoms is close enough to form π (Pi) bond by sidewise overlap. These overlaps can be of overlaps of p-orbitals of C₁ — C₂, C₃ — C₄, C₅ – C₆ or C₃, C₄ — C₅, C₆ – C₁ respectively as shown in the following figures.

(a)

(b)
X-ray diffraction data reveals that benzene is a planar molecule. The six n electrons are delocalised and spread on the whole of the molecule: one half of the electron cloud above and the other half below the plane of the benzene ring. The presence of delocalised n electrons in benzene makes it more stable than the imaginary cyclohexatriene.

or

(Electron cloud)

If there were three single C – C bonds and three alternate C = C bonds present in benzene, the bond lengths should have been 154 pm and 134 pm respectively. In benzene there are neither C – C double bonds present as all the six C — C bonds in benzene are exactly alike and have a bond length of 139 pm. Thus the absence of pure double bonds in benzene accounts for the hesitation on the part of benzene to take part in additional reactions. Due to its extra stability, it prefers to show substitution reactions.

Aromaticity: Benzene is considered a parent aromatic compound. Now the name is applied to all the ring systems whether or not having benzene ring, possessing the following characteristics.

1. It should be planar
2. Complete delocalisation of % electrons in the ring.
3. Presence of (4n + 2) n electrons in the ring where n is an integer (n = 0, 1, 2). This is called Huckel Rule.

Preparation of Benzene
Benzene is commercially isolated from the ‘Light oil fraction’ of coal tar. However, it may be prepared in the laboratory by the following methods.
1. From ethyne

2. Decarboxylation of the aromatic acids Sodium salt of benzoic acid on heating with soda lime gives benzene.

3. Reduction of Phenol in the presence of zinc dust gives benzene

Properties of Benzene (Aromatic hydrocarbons)
Physical properties

1. Aromatic hydrocarbons are non-polar.
2. They are colourless liquids or solids with a characteristic aroma.
3. Aromatic hydrocarbons are immiscible with water but are readily miscible with organic solvents.
4. They burn with a sooty flame.

Chemical Properties
Arenes undergo electrophilic substitution reactions. However, under special conditions, they undergo addition and oxidation reactions.

Electrophilic Substitution Reactions of arenes are nitration, halogenations, sulphonation, Friedel Craft’s reactions.
In all these reactions, the attacking reagent is an electrophile E®.
1. Nitration.

2. Halogenation: Arenes react with halogen in the presence of Lewis acids like FeCl₃, FeBr₃, or AlCl₃ to yield halo arenes.

Order of reactivity of halogens is Cl₂ > Br₂ > I₂.

3. Sulphonation: Here H of the benzene ring is replaced by sulphonic group (— SO₂ OH). It is carried out by heating benzene with fuming sulphuric acid (oleum).

4. Friedel Craft’s Reaction:
(a) Alkylation: On reacting benzene with an alkyl halide in the presence of anhydrous Aluminium chloride, alkyl benzene is formed.

(B) Acylation: On treating benzene with an acyl chloride in the presence of Lewis acids (AlCl₃) gives acyl benzene.

Mechanism of electrophilic substitution reactions: It involves three steps:

1. Generation of an electrophile.
2. Formation of a resonance-stabilised carbocation intermediate.
3. Removal of proton H⁺ to form the product.

1. Generation of Eelecrophile (E⁺): In the above reactions electrophiles like Cl⁺ (chloronium ion) is generated during chlorination by reacting with any. AlCl₃.

In the case of Nitration, NO₂ (nitronium ion) is generated.

2. Formation of carbocation (arenium ion) results with one of the carbon getting sp³ hybridised on the attack of the electrophile (E)⁺

Sigma complex (arenium ion)
The intermediate arenium ion gets stabilised by resonance.

3. Removal of a proton (H⁺)


2. Addition reactions: Under drastic conditions of high temperature and or pressure in the presence in the presence of nickel catalyst, dihydrogen gets added to the benzene.

In the presence of ultraviolet light, three molecules of Cl₂ get added to benzene to form Benzene hexachloride [BHC] C₆H₆C₁₆ also called Gammaxene.

3. Oxidation by combustion: When heated in air, benzene burns with a sooty flame producing CO₂ and H₂O.
C₆H₆ + O₂ → 6 CO₂ + 3H₂O
General combustion reaction for any hydrocarbon is
C_xH_y +(x + y/4)O₂ → xCO₂ + y/2H₂O.

Directive Influence of a Functional Group in Monosubstituted Benzene
Ortho and para directing groups: The groups which direct the incoming group to ortho & para positions are called ortho & para directing groups. In phenol, for example, — OH (hydroxy) group attacked to benzene directs the new (or coming group) to ortho para positions as explained below:

From the above structures, it is clear that electron density is more at ortho & para positions (structure II, III & IV) to the – OH group. Hence the coming electrophile will prefer to attack ortho & para position rather than meta. However due to the — I effect exerted by the — OH group, electron density at o—&p—position is slightly reduced. But overall, there is an increase of electron density ato-Scp- position. Hence the substituent at o—&p — positions to the -OH group.

Therefore, the -OH group is an activating group, as it activates the benzene ring for the attack of an electrophile. Other activating groups are NH₂, -NHR, NHCOCH₃, -OCH₃, -CH₃, -C₂H₅ etc.

Halogens are a class among themselves. They are deactivating and at the same time o—Scp — directing. Because of the, I effect, the overall electron density on benzene decreases. It makes further substitution difficult. However, due to resonance, the electron density on the o—& p — position is greater than at the meta position. Hence they are also o— & p — directing.

Meta-directing groups. The groups which when present in the benzene ring direct the incoming groups to meta position are called meta-directing groups. Some of the meta-directing groups are

(nitro group), for example, reduces the electron density in the benzene ring due to its — I effect. Nitrobenzene is a resonance hybrid of the following five canonical structures.

In this case, the electron density on the benzene ring decreases making further substitution difficult. Therefore these groups are called deactivating groups. The electron density on the o – and p – position is comparatively less than that at the meta position. Hence, the electrophile attacks on comparatively electron-rich meta position, resulting in meta-substitution.

Carcinogenicity and Toxicity:
Benzene and polynuclear hydrocarbons containing more than two fused benzene rings are toxic and said to possess cancer-producing (carcinogenic) property. They enter into the human body and undergo various biochemical reactions and finally damage DNA and cause cancer. Some of the carcinogenic hydrocarbons are given below. Such polynuclear hydrocarbons are formed on incomplete combustion of organic materials like tobacco coal and petroleum.

By going through these CBSE Class 11 Chemistry Notes Chapter 12 Organic Chemistry Some Basic Principles and Techniques, students can recall all the concepts quickly.

Organic Chemistry Some Basic Principles and Techniques Notes Class 11 Chemistry Chapter 12

→ Carbon: Tetra valency of carbon, shape of organic compounds & characteristic features of π-bond.

→ Structural representation of organic compounds: Complete, condensed & bond line structural formulae.

→ A 3-dimensional representation of organic molecules & classification of organic compounds.

→ Acyclic or open chain compounds & Alicyclic or closed chain compounds or ring compounds & functional groups.

→ Homologous series, Nomenclature of organic compounds & I.U.P.A.C. nomenclature of alkanes.

→ Nomenclature of organic compounds having a functional group or groups & nomenclature of substituted benzene compounds.

→ Isomerism: Structural, chain, position, functional group isomerism, metamerism & stereoisomerism.

→ Fundamental concepts in organic reaction mechanism: Fission of a covalent bond, Nucleophiles & Electrophiles, Electron movement in organic reactions. Electron displacement effects in covalent bonds. Inductive effect, resonance structure & resonance effect.

→ Electromeric effect (E-effect), Hyperconjugation, types of organic reactions & mechanisms.

→ Methods of purification of organic compounds: Sublimation, crystallization, distillation, differential extraction, chromatography.

→ Qualitative analysis of organic compounds detection of C & H, N, S, halogens & for PO₄^3-.

→ Quantitative analysis of organic compounds: Elemental detection in the form of a percentage, C, H, N-(Dumas method, Kjeldahl’s method), halogens, sulfur, phosphorus & oxygen.

→ Orbital hybridization concept: The nature of the covalent bonding in organic compounds can be described in terms of the orbitals hybridization concept.

→ Three-dimensional representation of organic compounds: Three-dimensional representation of organic compounds on paper can be drawn by wedge & dash formula.

→ Functional group: A functional group is an atom or group of atoms bonded together in a unique fashion & which determines the physical & chemical properties of compounds.

→ I.U.P.A.C.: International union of pure & applied chemistry.

→ Organic reaction mechanism: Organic reaction mechanism concepts are based on the structure of the substrate molecule. Fission of a covalent bond, the attacking reagents, the electron displacement effects & the conditions of the reaction.

→ Cleavage of covalent bond: A covalent bond may be cleaved in a heterolytic or homolytic fashion. A Heterolytic cleavage yields carbocations or carbanions & a homolytic cleavage gives free radicals as reactive intermediates.

→ Nucleophile & Electrophile:

• Nucleophile – Electron pair donor
• Electrophile – Electron pair acceptor.

→ Organic reactions:

1. Substitution reactions
2. Addition reactions
3. Elimination reactions
4. Re-arrangement reactions

→ Methods of purifications of organic compounds:

1. Sublimation
2. Distillation &
3. Differential extraction

→ Chromatography is a useful technique of separation, identification & purification of compounds. It is classified into two categories adsorption & partition chromatography. Lassaigne’s Test: N, S, halogens & phosphorus are detected by Lassaigne’s test.

→ Estimation of C & H: Carbon & hydrogen are estimated by determining the amounts of CO₂ & water produced. Estimation of Nitrogen: Nitrogen is estimated by Duma’s or Kjeldahl’s method.

→ Halogens Estimation: Halogens are estimated by various methods. Estimation of S & Phosphorus: S & P are estimated by oxidizing them to sulphuric & phosphoric acid respectively.

→ The percentage of oxygen: The percentage of oxygen is usually determined by subtracted (the sum of percentages of all other elements present in the compound) out of 100.

→ Retardation factor:
Rf = \(\frac{\text { Distance moved by the substance from base line }}{\text { Distance moved by solvent from base line }}\)
(a) Percentage of carbon = \(\frac{12 \times m_{1} \times 100}{44 \times m}\)
m₁ = mass of CO₂
m = mass of organic compound

(b) Percentage of Hydrogen = \(\frac{2 \times m_{1} \times 100}{18 \times m}\)
m₁ = mass of H₂O
m = mass of organic compound

(c) Percentage of nitrogen by Dumas method = \(\frac{28 \times V \times 100}{22400 \times m}\)
V = Volume of nitrogen m mass of organic compound

(d) Percentage of nitrogen by kJeldahl’s method = \(\frac{1.4 \times \mathrm{M} \times 2\left(\mathrm{~V}-\mathrm{V}_{1} / 2\right)}{m}\)
m mass of organic compound
M = Molarity of H₂SO₄ taken
V = Volume of H₂SO₄ of molarity-M
V₁ = Volume of NaOH of molarity-M used for titration of excess of H2S04

(e) Percentage of halogens = \(\frac{\text { Atomic mass of }(\mathrm{X}) X m_{1} g \times 100}{\text { molecular mass of }(\mathrm{AgX}) X m}\)
m = mass of organic compound
m₁ = mass of AgX formed
X = halogen atom

(f) Percentage of sulphur = \(\frac{32 \times m_{1} \times 100}{233 \times m}\)
m = mass of organic compound
m₁ = mass of BaSO₄ formed

(g) Percentage of Phosphorus:
If Phosphorus is estimated as Mg₂P₂O₇
= \(\frac{62 \times m_{1} \times 100}{222 \times m}\)
m = mass of organic compound
m₁ = mass of Mg₂P₂O₇
222 = molar mass of Mg₂P₂O₇

If Phosphorus is estimated as (NH₄)₃ PO₄.12MoO₃ then percentage of Phosphorus = \(\frac{31 \times m_{1} \times 100}{1877 \times m}\)
here m₁ = mass of (NH₄)₃ PO₄.12MoO₃
1877 = molar mass of (NH₄)₃PO₄.12MoO₃

(h) Percentage of Oxygen:
= \(\frac{32 \times m_{1} \times 100}{44 \times m}\)
m = mass of organic compound
m₁ = mass of carbondioxide

Chapter In Brief:
Berzelius, A Swedish chemist proposed that a Vital Force was responsible for the formation of organic compounds. F. Wohler gave a death blow to the Vital Force theory when he synthesized organic compound urea from an inorganic compound ammonium cyanate.

Tetravalence of Carbon: Shapes of organic compounds: The formation of CH₄, C₂H₆ is due to sp³ hybridization of C; formation of CH₄ is on the basis of sp² hybridization of C, and formation of C₂H₂ is on the basis of sp hybridization of C. The presence of double bond in H₂C=CH₂ and triple bond in HC = CH is due to the presence of one π and two π bonds respectively in them.

In H₂C=CH₂, rotation about C-C bond is hindered due to the presence of π bond between the two C atoms, sp³ hybridization gives rise to tetrahedral shape t

o CH₄, sp² hybridization gives rise to a trigonal planar arrangement to C₂H_4, and sp hybridization gives linear shape to C₂H₂. An sp³ hybrid orbital can overlap with Is orbital of hydrogen to give a C—H bond (sigma a single bond). Overlap of an sp² orbital of one carbon with an sp² orbital of another results in the formation of a carbon-carbon bond.

The unhybridized p-orbitals on two adjacent carbons can undergo lateral (side-by-side) overlap to give a pi (π) bond. Organic compounds can be represented by various structural formulas. The three-dimensional representation of organic compounds on paper can be drawn by the wedge and dash formula.

Organic compounds can be classified on the basis of their structure or the functional groups they contain. A functional group is an atom or group of atoms bonded together in a unique fashion which determines the physical and chemical properties of the compounds. The naming of the organic compounds is carried out by following a set of rules laid down by the International Union of Pure and Applied Chemistry (IUPAC). In IUPAC nomenclature, the names are correlated with the structure in such a way that the reader can deduce the structure from the name.

Structural Representations Of Organic Compounds:
Complete, condensed, and Bond-line structural formulae

stand for the complete structural formulae of ethane, ethene ethyne, and methanol whereas CH₃—CH3₃ (or C₂H₆), H₂C=CH₂ (or C₂H₄), HC ≡ CH (or C₂H₂), and CH₃OH stand for their condensed structural formulae respectively.

Bond-line structural representation of 1,3 butadiene is

and that of 3-methyl octane is

(its condensed formula is CH₃CH₂CH(CH₃) (CH₂)₄CH₃

The bond-line structure of chlorocyclohexane is

Three Dimensional Representation Of Organic Molecules:
The three-dimensional (3-D) structure of organic molecules can be represented on paper by using certain conventions.

For example by using solid

and dashed

wedge formula, the 3-D image of a molecule from a two-dimensional picture can be perceived. The solid wedge projects towards the observer and the dashed wedge projects away from the observer. The bonds lying in the plane of the paper are depicted by using a normal line (—)

The 3-D representation of CH₄ is

Classification of Organic Compounds:

1. Acyclic or Open Chain Compounds: These compounds are also called aliphatic compounds and consist of straight or branched chain compounds.


is acetaldehyde and

is acetic acid.

2. Alicyclic or Closed Chain or Ring Compounds:
Some of the examples of alicyclic /closed chain or ring compounds are as follows:

Aromatic Compounds: Benzenoid Aromatic Compounds:

Non-Benzenoid Compounds

Hetero Cyclic Aromatic Compounds

→ Functional Group:
The functional group may be defined as an atom or group of atoms joined in a specific manner that is responsible for the characteristic chemical properties of the organic compounds. The examples are hydroxyl group (-OH), aldehyde group (-CHO) and carboxylic acid group (—COOH), etc.

→ Homologous Series:
A group or a series of organic compounds each containing a characteristic functional group forms a homologous series and the member of the series are called homologs. The members of a homologous series can be represented by general molecular formula and the successive members differ from each other in the molecular formula by a — CH₂ unit. There are a number of homologous series of organic compounds. Some of these are alkanes, alkenes, alkynes, alkyl halides, alkanols, alkanols, alkenones, alkanoic acids, amines, etc.

e.g. The general formula of alkanols is C_nH_2n+1-OH. Individual members of a homologous series are called Homologues.
CH₃OH, C₂H₅OH, C₃H₇OH are homologs of the alkanol family.

→ Nomenclature of Organic Compounds:
Earlier organic compounds were known by their common or trivial names. For example, HCOOH was called formic acid, CH₃ CHO was called acetaldehyde, and so on.

→ The I.U.P.A.C System Of Nomenclature:
To systematize the naming of millions of organic compounds IUPAC (International Union Of Pure And Applied Chemistry) pattern of naming is adopted.

The I.U.P.A.C. System Of Nomenclature


Common or Trivial names of some organic compounds

Alkyl, Radicals (R) C_nH_2n+1

Table: Some functional Groups and classes of organic compounds:


Note: Students are advised to follow different rules and conventions as per the IUPAC system as given in the Textbook. In the case of polyfunctional compounds, one of the functional groups is chosen as the principal functional group and the compound is named on that basis.

The remaining functional groups which are subordinate functional groups are named as substituents using the appropriate prefixes. The choice of the principal functional group is made on the basis of the order of preference. The order of decreasing priority for the same functional groups is:
-COOH, -SO₃H, -COOR (R = alkyl group)
-COCl, -CONH₂, -C ≡ N, -CHO, > C = O, -OH, -NH₂,

The R, C₆H₅, halogens (F, Cl, Br, I), NO₂, alkoxy (OR), etc. are always prefixed substituents.

For example:
(i) HOCH₂(CH₂)₃CH₂COCH₃ will be named as 7 hydroxyheptan- 2-one
(ii) Br CH₂CH = CH₂ is named as 3-Bromoprop-l-ene.
(iii) CH₂ = CH-CH = CH₂ is Buta-1, 3-diene.

Problem:
Derive the structure of
1. 2-chioropentane
Answer:

2. Pent-4-en-2-ol
Answer:

3. 3-Nitrocyclohexene
Answer:

4. Cyclohex-2-en-l-ol
Answer:

5. 6-Hydroxyheptanal
Answer:

Nomenclature of Substituted Benzene Compounds:




Problem: Write the structural formula of
(a) o-Ethyl anisole
Answer:

(b) p-Nitroaniline
Answer:

(c) 2, 3-dibromo-l-phenyl pentane
Answer:

(d) 4-Ethyl-l-fluoro-2-nitrobenzene
Answer:

Isomerism: The phenomenon of the existence of two or more compounds possessing the same molecular

formula but different properties is known as isomerism. Such compounds are called isomers. The above flow chart shows different types of isomerism.

Types of structural isomerism:
1. Chain isomerism: This type of isomerism is due to the difference in the nature of the carbon chain (i.e., straight or branched) which forms the nucleus of the molecule, e.g.,

2. Position isomerism: It is due to the difference in the position of the substituent atom or group or an unsaturated linkage in the same carbon chain. Examples are


3. Functional isomerism: Two or more compounds having the same molecular formula but different functional groups are called functional isomers and this phenomenon is termed functional group isomerism. For example, the molecular formula C₃H₆O represents an aldehyde and a ketone.

and C₃H₆O represents an ether and alcohol.

4. Metamerism: It is due to the difference in nature of the alkyl group attached to the same functional group. This type of isomerism is shown by compounds of the same homologous series.
For example.

II. Stereoisomers: Stereoisomers are compounds that have the same constitution and sequence of covalent bonds but differ in the relative positions of their atoms or groups in space.

5. Geometrical isomerism: The isomers which possess the same structural formula but differ in the spatial arrangement of the groups around the double bond are known as geometrical isomers and the phenomenon is known as geometrical isomerism. This Isomerism is shown by alkenes or their derivatives. When the similar groups lie on the same side, it is the cis-isomer, while when the similar groups lie on opposite sides, the isomer is trans. For example

Fundamental Concepts In Organic Reaction Mechanism:
In an organic reaction, the organic molecules (substrate) reacts with an appropriate attacking reagent and leads to the formation of one or more intermediates and finally product (s)

The general reaction is depicted as follows:

The substrate is that reactant which supplies carbon to the new bond and the other reactant is called reagent. A sequential account of each step, details of electron movement, energetics during bond breaking and bond formation, and the details of timing, when a reactant is transformed into the product are referred to as Reaction Mechanism.

Fission of a Covalent Bond: It occurs in two ways.
(A) Homolytical Fission/Cleavage or Homolysis
In such fission, each atom gets one electron of the shared pair of electrons.

Alkyl radicals are classified as primary secondary or tertiary. Alkyl radical stability increases as we proceed from primary to tertiary. Organic reactions, which proceed by homolytic fission are called free radical or homopolar, or non-polar reactions.

→ Heterolytical Fission/Cleavage or Heterolysis: The covalent bond breaks in such a way that the shared pair of electrons remains with one of the fragments

The species that has a sextet at the carbon and is positively charged is called a Carbocation (or carbonium ion)
The shape of methyl carbocation C is sp2 hybridized and its shape is Trigonal Planar

The observed order of carbocation stability is

The fission can occur, the other way.

: -CH₃ is called a Carbanion. Such a carbon species carrying a negative charge is called a Carbanion. Their stability decreases as follows:

Nucleophiles & Electrophiles:
(A) Nucleophile (Nu:): A reagent that is electron-rich and is in search of a relatively positive center is called a nucleophile. Example of nucleophiles are

Negatively charged reagents:

(B) Electrophiles: A recent which is electron-deficient and is in search of electron-rick site is called an electrophile Positively charged electrophiles are: H⁺, H₃O⁺, NO₂⁺, R⁺, Br₄
Neutral particles: BF₃, AlCl₃, SO₃

Inductive Effect (I effect):
It is the process of displacement of electrons along the chain of carbon atoms due to the presence of a polar covalent bond at one end of the chain. This is a permanent effect. It is of two types:
(A) -I effect: When the atom or group of atoms of the polar covalent bond is more electronegative than C, it is said to show the -I effect.

It is practically over after C₂
The —I effect of some of the atoms or groups of atoms in decreasing order is
-NO₂ > -CN > -COOH > -F > -Cl > -Br > -I

(B) + I effect: If the substituent attached to the end of the carbon chain is electron-donating, the effect is called + I effect.

The + I effect of some of the atoms or groups of atoms in the decreasing order is

→ Electromeric Effect (E-effect): It involves the complete- transfer of electrons of multiple bonds (double or triple bond) to one of the bonded atoms (usually more electronegative) at the call of the attacking reagent. It vanishes the moment the attacking reagent is removed. It is a temporary effect.

It is also of two types – E and + E effect.
If the electrons of the bond are transferred to that atom of the double bond to which the reagent finally gets attached the effect is called the + E effect.

If the electrons of the double bond are transferred to an atom of the double bond other than the one to which the reagent gets finally attached, the effect is called the — E effect.

→ Resonance Or Mesomerism: The phenomenon of resonance is said to occur whenever for a molecule we can write two or more Lewis structures that differ in the positions of electrons but not in the relative position of atoms. The various Lewis structures are called responding/canonical/contributing structures. The actual structure of the molecule is not represented by any of the resonance structures but is a resonance hybrid of all these canonical structures.

The various resonance structures are separated by a double-headed arrow ↔ Benzene is a resonance hybrid of the two Kekule structures.

Any of the two structures cannot explain all the properties of benzene. But the resonance hybrid which cannot be drawn on the paper and which is the actual structure of benzene will explain all the properties of benzene. For example, there are 3 double bonds and 3 single bonds (3 C = C and 3 C — C) in benzene corresponding to bond lengths of 1.34 Å and 1.54 Å respectively.

But as X-ray diffraction studies point out there are no single or double bonds in benzene and all the C—C bonds are having a bond length of 1.39 Å and are exactly equivalent. The resonance hybrid of benzene is generally shown by III.

Another example of resonance is provided by CH₃NO₂ (nitromethane).

Hyper Conjugation:
It is regarded as no bond resonance. Hyperconjugation is a general stabilizing interaction. It involves delocalization of an electron of C-H bond of an alkyl group directly attached to an atom of the unsaturated system; or to an atom with an unshared p orbital. The electrons of C—H a bond of the alkyl group enter into partial conjugation with the attached unsaturated system or with the unshared p orbital. Hyperconjugation is a permanent effect.

To understand the hyperconjugation effect, let us take an example of CH₃⁺CH₂ (ethyl cation) in which the positively charged carbon atom has an empty n orbital. One of the C-H CT bonds of the methyl group can align in the plane of this empty n orbital and the electrons constituting the C—H bond in-plane with this π orbital can then be delocalized into the empty π orbital as depicted in Fig.

Orbital diagram showing hyperconjugation in ethyl cation

This type of overlap stabilizes the carbocation because electron density from the adjacent bond helps in dispersing the positive charge. In general, the greater the number of alkyl groups attached to a positively charged carbon atom, the greater is the hyperconjugation interaction and stabilization of the cation. Thus, we have the following relative stability of carbocations:

Hyperconjugation is also possible in alkenes and alkyl arenes. Delocalisation of electrons by hyperconjugation in the case of an alkene can be depicted as in Fig.(b)

Orbital diagram showing hyperconjugation in propene

There are various ways of looking at the hyperconjugation effect. One of the ways is to regard the C-H bond as possessing partial ionic character due to resonance.

Problem: Explain why (CH₃)₃C⁺ is more stable than CH₃ CH₂⁺ and CH₃⁺ is the least stable cation.
Answer: Hyperconjugation interaction in (CH₃)₃C⁺ is greater than in CH₃CH₂⁺ as the (CH₃)₃C⁺ has nine C—H bonds. In CH₃, vacant n orbital is perpendicular to the plane in which C-H bonds lie, hence cannot overlap with it. Thus CH₃⁺ lacks hyper conjugative stability.

Types Of Organic Reactions:
1. Substitution Reactions

2. Addition reactions
H₂C = CH₂ + HBr → CH₃ – CH₂Br

3. Elimination reactions
CH₃—CHBr—CH₃ + KOH → CH₃-CH = CH₂ + KBr + H₂O

4. Rearrangement Reactions

Methods Of Purification Of Organic Compounds:
The common techniques used for the purification of organic compounds are based on their nature and the impurity present in them. The methods are as follows.

1. Sublimation
2. Crystallization
3. Distillation
4. Differential extraction
5. Chromatography

Finally, the purity of a compound is ascertained by determining its melting point or boiling point. Most of the pure compounds have sharp melting points and boiling points.

1. Sublimation: Some solid substances like camphor, naphthalene, etc. on heating change from solid to vapor state without passing through the liquid phase. The purification technique based on the above principle is known as sublimation and is used to separate sublimable compounds like benzoic acid from non-sublimable compounds like sodium chloride.

2. Crystallisation: It is based on the difference in the solubilities of the compound and the impurities in a suitable solvent. The impure compound is dissolved in a solvent in which it is sparingly soluble at room temperature but appreciably soluble at a higher temperature. The solution is concentrated by heating to get a nearly saturated solution. On cooling, crystals of the pure substance are removed by filtration.

3. Distillation: The process of distillation is carried out to separate

• volatile liquids from non-volatile impurities and
• liquids having sufficient differences in their boiling points.

Liquids having different boiling points vaporize at different temperatures. The vapors are cooled and get condensed into liquids. They are collected separately. CHCl₃ (b.p. 334 K) and aniline (b.p. 457 K) are easily separated by this method.

Fractional Distillation:
It is resorted to when the difference in boiling points of two liquids is not much. It is carried out through an a.fractionating column fitted over the mouth of the round bottom flask.

One of the technological applications of fractional distillation is to separate different fractions of crude oil in the petroleum industry.

Fractional. Distillation

The vapors of the less volatile liquid condense into the liquid which returns to the flask. The more volatile fraction passes over to the other side, condenses in the water condenser, and is collected in the receiver. When one fraction is completely separated the temperature is raised and the receiver is changed. Now, the second less volatile fraction distills over. Thus the more volatile liquid distills afterward. This is highly successful if the difference in b.p. of two liquids is less than 10-15 K.

Fractional. Distillation

Distillation Under Reduced Pressure:
This method is applicable to purify liquids having very high boiling points and those, which decompose at or below their boiling points. Such liquid is are made to boil at a temperature lower than their normal boiling points by reducing the pressure on their surface. A liquid boils at a temperature at which its vapor pressure becomes equal to the external pressure. The flowsheet diagram for distillation under reduced pressure is shown below.

Distillation under reduced pressure. A liquid boils at a temperature below its vapor pressure by reducing the pressure.

Steam Distillation:
This technique is applied to separate substances, which are steam, volatile, and are immiscible with water. In steam distillation, the steam generator is passed through a heated flask containing the liquid to be distilled. The mixture of steam and the volatile liquid is condensed and collected.

In steam distillation the liquid boils when the sum of the vapor pressures due to the organic liquid (p₁) and that due to water (p₂) become equal to the atmospheric pressure (p) i.e., p = p₁ + p₂ since p1 is lower than p, the organic liquid vapourised at a lower temperature than its b. pt. Thus if one of the substances in water and the other a water-insoluble substance such a mixture will boil close to but below 373 K. Aniline is separated from the aniline-water mixture by this method.

Steam distillation. Steam volatile component volatilizes, the vapors condense in the condenser and the liquid collects in a conical flask.

Differential Extraction:
When an organic compound is present in an aqueous medium, it is separated by shaking with an organic solvent in which is more soluble than water. The organic solvent and the aqueous solution should be immiscible with each other so that they form two distinct layers which can be separated by the separatory funnel. The organic solvent is later removed by distillation or by evaporation to get back the compound.

(a) Differential extraction. Extraction of the compound takes place based on the difference in solubility

Chromatography:
Chromatography is an important technique extensively used to separate mixtures into their components, purify compounds, and also test the purity of compounds. In this technique, the mixture of substances is applied into a stationary phase, which may be a solid or a liquid.

A pure solvent, a mixture of solvents, or a gas is allowed to move slowly over the stationary phase. The components of the mixture get gradually separated from one another. The moving phase is called the mobile phase.

Based on the principle involved, chromatography is classified into different categories. Two of these are:
(a) Adsorption chromatography and
(b) Partition chromatography,

(a) Adsorption Chromatography: Adsorption chromatography is based on the fact that different compounds are adsorbed on an adsorbent to different degrees. Commonly used adsorbents are silica gel and alumina. When a mobile phase is allowed to move over a stationary phase (adsorbent), the components of the mixture move by varying distances over the stationary phase.

Following are two main types of chromatography techniques based on the principle of differentials adsorption.
(a) Column chromatography and
(b) Thin layer chromatography

(a) Column Chromatography: Column chromatography involves the separation of a mixture over a column of adsorbent (stationary phase) packed in a glass tube. The mixture adsorbed on the adsorbent is placed on top of the adsorbent in the column. An appropriate element which is a liquid or a mixture of liquid is allowed to flow down the column slowly. Depending upon the degree to which the compounds are adsorbed, complete separation takes place. The most readily adsorbed substances are retained near the top and the others come down to various distances in the column.

Column chromatography. Different stages of separation of components of a mixture

(b) Thin Layer Chromatography (TLC): Another type of adsorption chromatography, which involves the separation, of substances of a mixture over a thin layer of an adsorbent coated on a glass plate. A thin layer of an adsorbent (silica or alumina SiO₂ or Al₂O₃ gel) is spread over a glass plate. The solution of the mixture to be separated is applied as a small spot about 2 cm above one end of the TLC plate. The glass plate is then placed in a closed jar containing the eluent (see fig. below).

As the eluent rises up the plate the components of the mixture move up along with the eluent to different distances depending on their degree of adsorption and separation takes place. The relative adsorption of each component is expressed in terms of its Retention Factor, i.e., R_f Value
R_f = \(\frac{\text { Distance moved by the substance from baseline }(x)}{\text { Distance moved by the solvent from baseline }(y)}\)

(a) Thin layer chromatography,
(b) Developed chromatogram. Chromatogram being developed

The spots of colored compounds are visible on the TLC plate due to their original color. Fig. (b) on the previous page.

Partition Chromatography:
Paper chromatography is a type of partition chromatography. Water trapped in chromatography paper acts as a stationary phase. It is based on the principle of continuous differential partitioning of components of a mixture between stationary and mobile phases. A strip of paper spotted at the base with the solution of the mixture is suspended in a suitable solvent.

The solvent acts as the mobile phase rise up due to capillary action and flows over the spot. The paper selectively retains different components according to their differing partition in two phases. The paper strip so developed is called Chromatogram. The spots of the separated colored compounds are visible at different heights from the position of the initial spot on the chromatogram. The spots may be observed under U. V. light as discussed in TLC.

Paper chromatography. Chromatography paper in two different shapes.

Qualitative Analysis Of Organic Compounds Detection Of Elements:
The elements present in the organic compound can be detected as follows:
1. Carbon and hydrogen: The given organic compound is mixed with about double the amount of pure and dry copper oxide. The mixture is heated in a hard glass tube. The CO₂ and H₂O produced due to combustion are tested by lime water and anhydrous copper sulfate. The lime water will turn milky and copper sulfate will turn blue.


2. Nitrogen, can be detected as
1. Soda-lime test: When the organic compound is heated with soda lime in a test tube, the evolution of ammonia indicates nitrogen.

2. Lassaigne’s test: A small piece of dry sodium is heated gently in a fusion tube till it melts to a shining globule. Then a small amount of organic substance is added and the tube is heated to red hot. The red hot tube is plunged into distilled water contained in a china dish. The contents of the dish are boiled, cooled, and filtered. The filtrate is known as sodium extract or Lassaigne’s extract.

For the nitrogen test, the sodium extract is made alkaline with a few drops of dil. NaOH. Freshly prepared FeSO₄ solution is added and the contents are warmed. Then a few drops of FeCl₃ are added followed by acidification with cone. HCl or H₂SO₄. The appearance of bluish-green coloration indicates nitrogen.

If the organic compound contains N and S together, sodium thiocyanate (Na CNS) may be formed with the sodium extract which gives blood-red coloration due to the formation of Fe(CNS)₃,

Sulfur: To the sodium extract.
1. Add lead acetate: The formation of black precipitate confirms sulfur.

2. To the other part of sodium extract add a few drops of sodium nitroprusside solution. The appearance of purple color indicates sulfur.

Halogens:
This test can also be done with sodium extract. The extract is boiled with a cone. HNO₃ to expel the gases. It is then cooled and treated with silver nitrate solution. The formation of different colored precipitates confirms halogens.

Quantitative Analysis:
1. Estimation of Carbon and Hydrogen: A known weight of the organic compound is heated with dry cupric oxide in the dry atmosphere free from CO₂. The carbon and hydrogen present, in the organic compound, are oxidized to CO₂ and water. The CO₂ is absorbed in potash bulbs and water is absorbed in CaCl₂ tubes. From the weights of CO₂, and H₂O form, the percentage of C and H are calculated as:

2. Estimation Of Nitrogen: Nitrogen can be estimated by one of the following two methods.
1. Duma’s Method: A known weight of the given organic compound is heated with dry cupric oxide in a current of CO₂. The N0 gas obtained is connected in Scliffs nitrometer at the prevailing temperature and pressure. Then, this volume of N, gas so collected is converted to volume at STP/NTP by using gas equation
P₁V₁/T₁ = P₂V₂/T₂
knowing 22.4 L of N₂ gas at STP weight = 28.0 gm.

Weight and percentage of Nitrogen can be calculated
% of N = \(\frac{28}{22400} \times \frac{\text { Volume of } \mathrm{N} \text { at } \mathrm{STP} \text { in } \mathrm{mL}}{\text { Mass of compound }}\) × 100

2. Kjeldahl’s Method: This is a more convenient method for the estimation of N particularly in foods, fertilizers, drugs etc. This method is, however not applicable to compounds containing nitrogen in the ring (Pyridine, quinoline, etc.) and compounds containing N directly linked to an oxygen atom (eg. NO₂) or another N atom. e.g. A Z O (—N = N—) compounds.

In this method, the given organic compound is treated with a cone. H₂SO₄ to couvert N into (NH₄)₂ SO₄ the ammonium sulfate [(NH₄)₂SO₄ is treated with 40% NaOH solution and the ammonia evolved is neutralized with an excess of a standard acid [known volume V of the acid taken. The excess of the residual acid is titrated with a standard solution of the alkali and the volume of the acid left unneutralized by ammonia (v ml) is noted.

∴ Volume of the acid neutraL ised by ammonia = (V – u) ml.
∴ % of N = \(\frac{1.4 \times N(V+v)}{W}\)
where N = Normality of acid taken
W = wt. of the organic compound.

Estimation Of Halogens: The given organic compound containing halogens is treated in Carius Method with fuming nitric acid in a long-necked Carius Tube and silver nitrate. The halogen present is converted into silver halide. From the weights of silver halide formed and the known weight of the organic compound taken, the percentage of halogen can be calculated.

% of halogen = \(\frac{\text { Atomic mass of halogen }}{108+\text { At. mass of halogen }}\) × \(\frac{\text { Mass of silver halide }}{\text { Mass of substance }}\) × 100

Estimation of Sulphur:
In the Carius method organic compound is treated with fuming HNO₃ and S is precipitated as BaSO₄ by the addition of BaCl₂. From the wt. of BaS04 formed, the percentage of S can be calculated.
% of sulphur = \(\frac{32}{233} \times \frac{\text { Mass of } \mathrm{BaSO}_{4}}{\text { Mass of substance }}\) × 100

Estimation of Phosphorus:
In Carius method, P is quantitatively oxidized to H₃PO₄ by fuming HNO₃ which is precipitated to Mg₂P₂O₇. knowing the wt. of Mg₂P₂O₇ and that of the organic compound, percentage of P can be determined
% of P = \(\frac{62}{222} \times \frac{\text { Mass of } \mathrm{Mg}_{2} \mathrm{P}_{2} \mathrm{O}_{7} \text { formed }}{\text { Mass of the substance }}\) × 100

Estimation of Oxygen:
There is no direct method available to estimate oxygen in the organic compound. The percentage of oxygen is usually found by subtracting the sum of the percentages of all elements present in the compound from 100.

By going through these CBSE Class 11 Chemistry Notes Chapter 11 The p-Block Elements, students can recall all the concepts quickly.

The p-Block Elements Notes Class 11 Chemistry Chapter 11

→ General trends in the chemistry of p-block elements.

→ Group-13 Elements: The Boron family-Electronic configuration, atomic radii, ionization, enthalpy, electro-negativity, physical & chemical properties.

→ Important trends & anomalous properties of boron

→ Important compounds of boron: Borax, orthoboric acid, diborane, uses of boron & aluminium & their compounds.

→ Group-14 Elements: The carbon family. Electronic configuration, covalent radius, ionization enthalpy, electronegativity, physical & chemical properties.

→ Important trends & anomalous behaviour of carbon.

→ Allotropes of carbon: Diamond, graphite & fullerenes & uses of carbon.

→ Important compounds of carbon & silicon: Carbon monoxide, carbon dioxide, silicon dioxide, silicones, silicates & zeolites.

→ P-BIock Elements: p-block of elements of the periodic table is unique in terms of having all types of elements-metals, non¬metals & metalloids. Group numbers ranging from 13-18.

→ Valence shell electronic configuration ns².np^1-6(Except for He).

→ pπ-pπ bonds and dπ – pπ or dπ-dπ bonds: The combined effect of size & availability of d-orbitals considerably influences the ability of these elements to form π-bonds. While the lighter elements form pπ-pπ bonds. The heavier ones form dπ-dπ bonds.

→ Electron deficiency in boron compounds: The availability of 3-valence electrons for covalent bond formation using four orbitals (2S, 2P_x, 2P_y & 2P_z.) leads to the so-called electron deficiency in boron compounds.

→ Boranes: Boron forms covalent molecular compounds with di-hydrogen as boranes. The simplest is diborane (B₂H₆).

→ Inert pair effect: Aluminium exhibits + 3 oxidation state. With heavier elements, the +1 oxidation state gets progressively stabilised on going down the group. This is a consequence of the so-called inert pair effect.

→ Catenation: The ability to form chains or rings not only with C – C single bonds but also with multiple bonds
(C = C or C ≡ C).

→ Allotropes of carbon: Three important allotropes of carbon are diamond, graphite & fullerenes.

→ Carbon monoxide: Carbon monoxide having lone pair of electrons on C forms metal carbonyls. It is deadly poisonous due to the higher stability of its haemoglobin complex as compared to that of the oxy-haemoglobin complex.

→ Carbon dioxide: Increased content of CO₂ in the atmosphere due to combustion of fossil fuels & decomposition of limestone is feared to cause an increase in the greenhouse effect. This in turn raises the temperature of the atmosphere & causes serious complications.

→ Compounds of silicon: Silica, silicates & silicones are important compounds & find applications in industry & technology.

Chapter in Brief:
In the case of elements of p-block, the last electron enters a p-orbital. As p-subshell can hold a maximum of 6 electrons in p_x, p_y and p_z atomic orbitals, p-block has 6 groups namely 13, 14, 15, 16, 17 and 18th groups. The valence shell electronic configuration is ns² np^1-6. In the case of the boron family (group 13), carbon family (group 14) and nitrogen family (group 15), the group oxidation states (the most stable oxidation states) are +3, +4 and +5 respectively for the lighter element an in the respective groups.

However, the oxidation state two until less than the group oxidation state becomes increasingly more stable for the heavier elements in each group. The occurrence of oxidation state two units less than the group’s oxidation state is due to the Inert Pair Effect.

General Electronic Configuration And Oxidation States Of P-Block Elements

Non-metals and metalloids exist only in the p-block. The non-metallic character of elements decreases down a particular group. In fact, the heaviest element in each group of the p-block is the most metallic in nature.

In general non-metals have higher ionisation enthalpies and higher electronegativities than metals. Hence in contrast to metals which readily form cations, non-metals readily form anions. The compounds formed by highly reactive non-metals like halogens with highly reactive metals like alkali metals are generally Ionic due to the large difference in their electronegativities.

On the other hand, compounds formed by non-metals themselves are largely covalent because of the small differences in their electronegativities. The change of non-metallic to metallic character can be best illustrated by the nature of oxides formed by them. The non-metallic oxides like CO₂ and SiO₂ are acidic or neutral whereas metallic oxides like CaO Na₂O are basic.

The first member of the groups of p-block differs from the remaining members of their corresponding group in two major respects. First is the size and all other properties which depend on size. Thus, the lightest p-block elements show the same kind of differences as the lightest s-block elements, lithium and beryllium. The second important difference, which applies only to the p-block elements, arises from the effect of d-orbitals in the valence shell of heavier- elements (starting from the third period onwards) and their lack in second-period elements.

The second-period elements starting from boron are restricted to a maximum covalence of four (using 2s and three 2p orbitals). In contrast, the third-period element of a p-group with the electronic configuration 3s²3pⁿ has the vacant 3d orbitals lying between the 3p and the 4s levels of energy. Using these d-orbitals the third-period elements can expand their covalence above four. For example, while boron form only [BF₄]^–, aluminium gives [AlF₆]^3- ion. The presence of these d-orbitals influences the chemistry of the heavier elements in a number of other ways.

The combined effect of size and availability of d orbitals considerably influences the ability of these elements to form their bonds. The first member of a group differs from the heavier members in their ability to form pπ-pπ multiple bonds to itself (e.g., C = C, C ≡ C, N ≡ N) and to other second-row elements (e.g., C = 0, C = N, C ≡ N, N = 0). This type of π-bonding is not particularly strong for the heavier p-block elements. The heavier elements do form, n bonds but this involves d orbitals (dπ-pπ or dπ—dπ).

As the d orbitals are of higher energy than the p-orbitals, they contribute less to the overall stability of molecules than does pπ -pπ bonding of the second-row elements. However, the coordination number in species of heavier elements may be higher than for the first element in the same oxidation state. For example, in the +5 oxidation state both N and P form oxoanions:

NO^3- (with π-bonding involving one nitrogen p-orbital ) and PO₄^3- (four-coordination involving s,p and d orbitals contributing to the π-bonding).

Group 13 elements: The boron family
Boron, Aluminium, Gallium, Indium and Thallium are the elements present in group 13. Boron (B) is a typical non-metal. Aluminium is a metal. Gallium, indium and thallium are almost exclusively metallic in character.

Atomic & Physical Properties of Group 13 Elements


^aMetallic radius, ^b6-coordination, ^cPauling scale,

For M³⁺ (aq) + 3e^– → M(s)
^eFor M⁺ (aq) + e^– → M(s).

1. Electronic Configuration: The outer electronic configuration of these elements is ns²np¹

2. Atomic Radii: Generally atomic radii increase in going down the group. However atomic radius of Ga is less than that of Al, due to the poor screening effect of the inner d-electrons for the valence electrons from the increased nuclear charge in gallium.

3. Ionisation Enthalpy: IE of Al is less than that of B due to the increased size of Al.

4. Electronegativity: Electronegativity first decreases from B to Al and then increases marginally.

5. Physical Properties: Boron is non-metallic, extremely hard and black coloured solid. It exists in many allotropic forms. It has unusually high M.Pt. The rest of the members are soft metals with low M.Pt. and high electrical conductivity Gallium with M.Pt. of 303 K is a liquid during summer. The density of elements increases down the group.

6. Chemical Properties: Due to its small size the sum of its first three enthalpies is very high. Therefore B does not form +3 cations and forms only covalent bonds. Al due to its low I.E. forms Al³⁺ ions. In the heavier metals due to the inert pair effect, they exhibit an oxidation state of +1.

BF₃ is an electron-deficient compound and acts as a Lewis acid by accepting a pair of electrons.

AlCl₃ achieves stability by forming a dimer.

Trivalent covalent state compounds are hydrolysed by water to form tetrahedral [M(OH)₄]^– species, the hybridisation state of M is sp³. AlCl₃ in acidified aqueous state forms octahedral [Al(H₂O)₆]³⁺ ion. Al is in d²sp³ hybridisation.

1. Reactivity towards air: Boron is unreactive in crystalline form. A1 forms a very thin oxide layer on the surface which protects the metal from further attack. On heating B₂O₃ and Al₂O₃ are formed. With N₂, they form nitrides at a higher temperature.

2. Reactivity towards acids and bases: B does not react. Al dissolves in dilute HCl and liberates H₂ gas
2Al(s) + 6HCl(aq) → 2Al³⁺ (aq) + 6Cl^–(aq) + 3H₂(g)
Cone. HNO₃ renders A1 passive by forming a protective oxide layer on the surface.

Al reacts with aq. alkalies and liberates H2 gas.

3. Reactivity towards halogens.
2E(s) + 3X₂(g) → 2EX₃(S) (X = F, Cl, Br, I)
E = B, Al, Ga, In.

Important Trends and Anomalous Properties of Boron:
1. The trihalides of all these elements are covalent in nature and hydrolysed by water
EX₃ + 3H₂O → E(OH)₃ + 3HX

2. Monomeric trihalides, being electron deficient are strong LEWIS ACIDS.

3. Maximum covalency shown by boron is 4 because it cannot expand its octet beyond 4 due to the absence of d-orbitals. Due to the availability of d-orbitals with other metals, the maximum covalent can be expected beyond 4.

AlCl₃ is dimerised to AlCl₆

Some Important Compounds of Boron:
1. Borax: It is a white crystalline solid of formula Na₂B₄O₇.10H₂O, more appropriately Na₂[B₄O₅(OH)₄].8H₂O. It dissolves in water to give an alkaline solution.

2. Orthoboric Acid: It is a white crystalline solid with soapy touch. Its formula is H₃BO₃. It is sparingly soluble in water but highly soluble in hot water.

Preparation:

1. Na₂B₄O₇ (Borax) + 2HCl + 5H₂O → 2NaCl + 4B(OH)₃ (Boric acid)
2. It is formed by hydrolysis with water of BCl3:
   BCl₃ + H₂O(aq) → H₃BO₃ + 3HCl.

Structure: It has a layer structure in which planar BO₃ units are joined by hydrogen bonds as shown in the figure below.

[Structure of boric apid H₃BO₃ dotted line represent hydrogen bonds.]

Properties of Boric Acid (H₃BO₃)

1. It is a weak monobasic acid.
2. It is not a protonic acid but acts as Lewis-acid by accepting electrons from a hydroxyl ion
   B(OH)₃ + 2HOH → [B(OH)₄]- + H₃O⁺
3. On heating above 370K, metaboric acid (HBO₂) is formed which on further heating yields boric oxide (B₂O₃).
   
   Diborane B₂H₆: It is the simplest of boron hydrides.

Preparation:


(iii) Industrially it is prepared by the reaction of BF₃ on sodium hydride.

Properties of Diborane:
1. If is a colourless, highly toxic gas with a B.Pt. of 180 K.

2. It catches fire spontaneously upon exposure to air. Enormous energy is released during the reaction.
B₂H₆ + 3O₂ → B₂O₃ + 3H₂O; Δ_CH° = -1976 kJ mol”1

3. Most of the higher boranes are highly flammable.

4. It is hydrolysed by water giving boric acid
B₂H₆(g) + 6H₂O(l) → 2B(OH)₃(aq) + 6H₂O

5. Diborane undergoes cleavage reactions with Lewis bases to give borane adduct
B₂H₆ + 2NMe₃ → 2BH₃ . NMe₃
B₂H₆ + 2CO → 2BH₃ . CO
B₂H₆ + 2NH₃ → B₂H₆.2NH₃
which is formulated as [BH₂(NH₃)₂]⁺ [NH₄]^– further heating gives [BH₂(NH₃)₂]⁺ [BH₄]^– , further heating gives Borazine or Borazole or Inorganic Benzene B₃N₃H₆


The structure of diborane is shown in Fig.(a) below. The four-terminal hydrogen atoms and the two boron atoms lie in one plane. Above and below this plane, there are two bridging hydrogen atoms. The four-terminal B—H bonds are regular two centre-two-electron bonds while the two bridge (B—H – B) bonds are different and can be described in terms of three centre-two electron bonds shown in Fig.(b)

(a) The strucwre of diborane, B₂H₆

Boron also forms a series of Hydridoborates; the most important one is the tetrahedral [BH₄]^– ion. Tetrahydridoborates of several metals is known. Lithium and sodium Tetrahydridoborates is also known as Borohydrides are prepared by the reaction of metal hydrides with B₂H₆ in diethyl ether.

(b) Bonding in diborane. Each B atom uses sp³ hybrids for bonding.

Out of the four sp³ Iribrids on each B atom, one is without an electron shown with broken lines. The terminal B-H bonds are normal 2 centre-2 electron bonds but lie two bridge bonds are 3 centre-2 electron bonds. The 3 centres 2 electron bridge bonds are also referred to as banana bonds.

2MH + B₂H₆ → 2M⁺[BH₄] ; M = Li or Na.
Both LiBH₄ and NaBH₄ are used as reducing agents in organic synthesis. They are starting materials for preparing other borohydrides.

Uses Of Boron & Aluminium And Their Compounds:
Boron is an extremely hard refractory solid of high melting point, low density and very low electrical conductivity find many applications. Boron fibres are used in making bullet-proof vest and light composite material for aircraft. The boron-10 (10B) isotope has a high ability to absorb neutrons and, therefore, metal borides are used in the nuclear industry as protective shields and control rods.

The main industrial application of borax and boric acid is in the manufacture of heat resistant glasses (e.g., Pyrex), glass-wool and fibreglass. Borax is also used as a flux for soldering metals, for heat, scratch and stain resistant glazed coating to earthenwares and as a constituent of medicinal soaps. An aqueous solution of orthoboric acid is generally used as a mild antiseptic.

Aluminium is a bright silvery-white metal, with high tensile strength. It has a high electrical and thermal conductivity. On a weight- to-weight basis, the electrical conductivity of aluminium is twice that of copper. Aluminium is used extensively in industry and everyday life.

It forms alloys with Cu, Mn, Mg, Si and Zn. Aluminium and its alloys can be given shapes of pipe, tubes, rods, wires, plates or foils and, therefore, find uses in packing, utensil making, construction, aeroplane and transportation industry. The use of aluminium and its compounds for domestic purposes is now reduced considerably because of its toxic nature.

Group 14 Elements: The Carbon Family
Carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb) are the members of group 14.
1. The valence shell electronic configuration of these elements is ns²np¹.

2. Covalent Radius: There is a considerable increase in covalent radius from C to Si, thereafter from Si to Pb, a small increase in radius is observed. This is due to the presence of completely filled d and f-orbitals in heavier members.

Atomic And Physical Properties Of Group 14 Elements:

^afor MIV oxidation state; ^b6-coordination, ^cPauling scale, ^d293 K; ^efor diamond; for graphite, density is 2.22; ^fβ-form (stable at room temperature)

3. Ionization Enthalpy: The first TE of group 14 members is higher than the corresponding members of group 13. It generally decreases from top to bottom. There is a small increase in the case of lead and it is due to the poor shielding effect of intervening d and f orbitals and the increase in the size of the atom.

4. Electronegativity: Due to the small size, the elements of this group are slightly more electronegative than group 13 elements. The electronegativity values for elements from Si to Pb are almost the same.

5. Physical Properties: All group 14 elements are solids, C and Si are non-metals, germanium (Ge) is a metalloid, whereas tin and lead are soft metals with low melting points. Melting points and boiling points of group 14 elements are much higher than those of the corresponding elements of the. group 13 elements.

6. Chemical Properties:
Oxidation states and trends in chemical reactivity: The common oxidation states shown by these elements are +4 and +2. Since the sum of four ionisation enthalpies is very high, compounds in the +4 oxidation state are generally covalent. The heavier members Ge, Sn and Pb, tendency to show an oxidation state of +2 increases due to the inert pair effect, i.e., the two electrons in ns² orbital prefer to remain paired if we go down the group and do not participate in bond formation.

C & Si mostly show an oxidation state of + 4.

Ge shows a + 4 states in stable compounds and only a few compounds in a + 2 oxidation state.
Sn forms compounds in both oxidation state + 4 and + 2 (Sn in + 2 states is a reducing agent)

Lead compounds in the + 2 state are stable and in the + 4 states are strong oxidising agents.
Being electron-precise molecules, they are neither electron- acceptors nor electron-donors.

Although C cannot expand its octet beyond 4 due to the non-availability of d-orbitals, other elements of the group can do so, because of the presence of d-orbitals in them. CCl₄ can’t undergo hydrolysis, whereas SiCl₄ can do so due to the same reason.

For examples, the species like SiF₅^–, SiF₆^2-, GeCl₆^2- and [Sn(OH)₆]^2- exist where the hybridisation of the central atom is sp²d³.
1. Reactivity towards oxygen: All members on heating in oxygen form oxides-MO and MO₂. Oxides in a higher oxidation state are more acidic than in a lower oxidation state. CO is neutral, CO is acidic. The dioxides-CO₂, SiO₂, GeO₂ are acidic, SnO₂ and PbO₂ are amphoteric. GeO is distinctly acidic, SnO and PbO are amphoteric.

2. Reactivity towards water:
C, Si, and Ge are not affected by water

Pb is not affected by water.

3. Reactivity towards halogens
M + X₂ → MX₂ M: Si, Ge, Sn, Pb
M + 2X₂ → MX₄ X: F, Cl, Br, I

Most MX₄ are covalent M shows sp3 hybridisation and MX₄ are tetrahedral in shape. SnF₄ & PbF₄ are ionic in nature. Pbl4 does not exist. Stability of MX₂ increases down the group.

GeX₄ is more stable than GeX₂, whereas PbX₂ is more stable than PbX₄. Sid₄ undergoes hydrolysis as shown below, but CCl₄ cannot undergo hydrolysis because carbon cannot expand its covalence beyond four due to the absence of d-orbitals

Important Trends and Anomalous Behaviour of C:
Carbon (C) the first member of group 14 differs from its congeners due to

1. Small size
2. Higher ionisation enthalpy and higher electronegativity.
3. Non-availability of d-orbitals.

In carbon, only s and p orbitals are available for bonding and, therefore, it can accommodate only four pairs of electrons around it. This would limit the maximum covalence to four whereas other members can expand their covalence due to the presence of d orbital.

Carbon also has a unique ability to form pπ-pπ multiple bonds with itself and with other atoms of small size and high electronegativity. Few examples of multiple bonding are: C = C, C ≡ C, C=0, C = S, and C = N. Heavier elements do not form pπ-pπ bonds because their atomic orbitals are too large and diffuse to have effective overlapping.

Carbon atoms have the tendency to link with one another through covalent bonds to form chain and rings. This property is called catenation. This is because C—C bonds are very strong. Down the group the size increases and electronegativity decreases, and, thereby, the tendency to show catenation decreases. This can be clearly seen from bond enthalpies values. The order of catenation is C >> Si > Ge = Sn. Lead does not show catenation.

Bond	Bond enthalpy/kJ mol-1
C-C	348
Si-Si	297
Ge-Ge	260
Sn-Sn	240

Due to the property of catenation and pπ-pπ bonds formation, carbon is able to show allotropic forms.

Allotropes of Carbon:
Carbon exists in crystalline and amorphous forms. Diamond and graphite are two well-known crystalline forms of carbon. In 1985, the third form of C known as Fullerenes was discovered:

The structure of diamond

Carbon in diamond is sp³ hybridised. Diamond has a crystal lattice. The C—C bond length is 154 pm. The structure is a rigid three-dimensional network of carbon atoms. In this structure shown on the side, directional covalent bonds are present throughout the lattice.

It is very difficult to break extended covalent bonding and therefore diamond is the hardest substance on the earth. It is used as an abrasive for sharpening hand tools, in making dies and in the manufacture of tungsten filaments for electric light bulbs.

Graphite:
Graphite has a layered structure. Layers are held by van der Waals forces and the distance between the two layers is 340 pm. Each layer is composed of planar hexagonal rings of C atoms. C—C bond length within a layer is 142 pm. Here C undergoes sp² hybridisation and makes three bonds with 3 neighbouring C atoms. The fourth electron forms a bond. The electrons are delocalised over the whole sheet.

These electrons in graphite are mobile and therefore, graphite conducts electricity- Graphite is very soft and is used as a dry lubricant in machines running at high temperature, where oil cannot be used as a lubricant.

Structure of graphite

Fullerenes:
Fullerenes are made by heating graphite in an electric arc in the presence of an inert gas such as helium or argon. The sooty material formed by the condensation of vapourised C₆₀ small molecules consists up mainly of a smaller quantity of C₇₀ and traces of fullerenes consisting of an even number of carbon atoms up to 350 or above. Fullerenes are the only pure forms of carbon because they have smooth structure without having “dangling” bonds.

Fullerenes are cage-like molecules. the molecule has a shape like a soccer ball and is called Buckminster Fullerene. It contains twenty six-membered rings and twelve five-membered rings. A six-membered ring is fused with six or five-membered rings but a five-membered ring can only fuse with six-membered rings.

All the carbons atoms are equal and they undergo sp³ hybridisation. Each carbon atom forms three sigma bonds with the other three carbon atoms. The remaining electron at each carbon atom is delocalised in molecular orbitals which give an aromatic character to the molecule.

This ball-shaped molecule has 60 vertices and each one is occupied by one C atom and it contains both single and double bonds with C-C distances of 143.5 pm and 138.3 pm respectively. Spherical fullerenes are also called Bucky Balls

The structure of C 60, Buckminster fullerene. Note that molecule has the shape of a soccer ball (football)

Graphite is a thermodynamically most stable allotrope of carbon and therefore Δ_fH° of graphite is taken as zero.

Uses of Carbon:

1. Graphite fibres embedded in plastic material form high strength, lightweight composites which find wide applications.
2. Being a good conductor, graphite is used as electrodes in batteries and in industrial electrolysis.
3. Crucibles made of graphite are inert to dilute acids and alkalies.
4. Graphite is used as a moderator in nuclear reactors to slow down the speed of fast-moving neutrons.
5. Being highly porous activated charcoal is used in absorbing poisonous gases. It is also used in water filters to remove organic contaminators and in the air conditioning system to control odour.
6. Carbon black is used as a black pigment in black ink and as filler in automobile tyres.
7. Coke is used as a fuel and largely as a reducing agent in metallurgy.
8. Diamond is a precious stone and used in jewellery. It is measured in carat (1 carat = 200 mg)

Some Important Compounds of Carbon and Silicon:

• Oxides of Carbon: Two important oxides of C are carbon monoxide CO and carbon dioxide CO₂.
• Carbon Monoxide (CO): Direct oxidation of carbon in a limited supply of air or oxygen yields CO.

1. Lab. method: On a small scale CO is prepared by dehydration of formic acid with cone. H₂SO₄ at 373K

2. Commercial-scale: It is prepared commercially by the passage of steam over hot coke. The mixture of CO and H₂ produced is called water-gas or synthesis gas

When air is used instead of steam a mixture of CO and N₂ produced which is called producer gas.

Properties:

1. It is a colourless odourless gas.
2. It is almost insoluble in water;
3. It- is a powerful reducing agent and reduces all metal oxides other than those of alkali and alkaline earth metals, aluminium and a few transition elements.
   
4.  In C ≡ O: there is one sigma and two n bonds between C and oxygen. Because of the presence of a lone pair of electrons on C, the CO molecule acts as a donor and reacts with certain metals when heated to form metal carbonyls.
   
5. Due to its highly poisonous nature, CO forms a complex with haemoglobin which is about 300 times more stable than the oxygen complex. This prevents haemoglobin in the red blood corpuscles from carrying oxygen around the body and ultimately results in death.

Carbon Dioxide:
Methods of Preparation.

1. Complete combustion of C and C containing fuels.
   
2. Lab. method
   CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g) + H₂O(l)
3. Commercially, it is prepared by heating lime stone,
   

Properties:

1. It is colourless and odourless gas.
2. It has low solubility in water. With water, it forms carbonic acid H₂CO₃ which is a weak dibasic acid.
   H₂CO₃^– + H₂O ⇌ HCO₃^– + H₃O⁺
   HCO₃^– + H₂O ⇌ CO₃^2- + H₃O⁺
3. 2NaOH + CO₂ → Na₂CO₃ + H₂O
4. Photosynthesis
   
5. Excess of CO₂ in the atmosphere leads to the greenhouse effect which will raise the temperature of the atmosphere.
6. CO₂ in the solid state is called Dry ice which is used as a refrigerant for ice cream and frozen food.

Structure of CO₂
C in CO₂ undergoes sp hybridisation. Two sp hybridised orbitals of carbon atom overlap with two p orbitals of oxygen atoms to make two sigma bonds while the other two electrons of the carbon atom are involved in pπ-pπ bonding with an oxygen atom. This results in its linear shape [with both C-O bonds of equal length (115 pm)] with no dipole moment. The resonance structures are shown below:

Resonance structures of carbon dioxide

Silicon Dioxide SiO₂
Silicon dioxide or silica along with silicates constitute 95 % of the earth’s crust. SiO₂ is a covalent three-dimensional network solid in which each silicon atom is covalently bonded in a tetrahedral manner to four oxygen atoms. Each oxygen atom in turn covalently bonded to another silicon atoms as shown.

Three-dimensional structure of SiO

Properties:

1. Silica in its normal state is almost non-reactive.
2. It is attacked by HF and NaOH.
   SiO₂ + 4HF → SiF₄ + 2H₂O
   SiO₂ + 2NaOH → Na₂SiO₃ (Sodium silicate) + H₂O

Uses: Silica gel is used as a drying agent, as a catalyst and in chromatography.

Silicones: They are a group of organosilicon polymers that have -R₂SiO- as a repeating unit. They are prepared as follows:

industries for cracking of hydrocarbons and isomerisation, e.g., ZSM-5 (A type of zeolite) used to convert alcohols directly into gasoline. Hydrated zeolites are used as ion exchangers in softening hard water.

By going through these CBSE Class 11 Chemistry Notes Chapter 10 The s-Block Elements, students can recall all the concepts quickly.

The s-Block Elements Notes Class 11 Chemistry Chapter 10

→ Gp. 1 Elements: Alkaline metals Electronic configuration, Atomic & Ionic radii Ionization enthalpy, hydration enthalpy.

→ Physical & Chemical properties.

→ Uses of Alkali metals.

→ General characteristics of the compounds of alkali metals-halides, salts of oxo-acids.

→ Anomalous properties of Lithium Points of difference between Li & other alkali metals

→ Points of Similarities between Lithium & Magnesium.

→ Important compounds of Sodium: Sodium carbonate, sodium chloride, sodium hydroxide & sodium hydrogen carbonates.

→ Biological importance of sodium & potassium

→ Gp. 2 Elements: Alkaline Earth metals

→ electronic configuration, Atomic & Ionic radii, Ionization enthalpy, hydration enthalpy

→ Physical & chemical properties & use of alkaline earth metals.

→ General characteristics of compounds of alkaline earth metals- oxides & hydroxides.

→ Halides, salts of oxo-acids & carbonates.

→ Anomalous behaviour of Beryllium-Diagonal relationship between Beryllium & Aluminium.

→ Some important compounds of calcium: Calcium oxide, calcium hydroxide & calcium carbonate, calcium sulphate & cement.

→ Biological importance of Mg & ca.

→ S-block Elements: Group-1 (Alkali metals) & Group-2 (Alkaline earth metals) Their oxides & hydroxides are alkaline in nature.

→ Ionization Enthalpy: Decreases down the group.

→ Atomic & Ionic sizes: Increases down the group.

→ Diagonal Relationship: Li in group-1 & Be in group-2 shows similarities in properties to the second member of the next group. Such similarities are termed a diagonal relationship.

→ Castner-Kellner process: Sodium hydroxides are manufactured by this process.

→ Solvay process: Sodium carbonate is prepared by this process.

→ Plaster of Paris: CaSO₄. \(\frac{1}{2}\) H₂O

→ Portland cement: It is an important constructional material. It is manufactured by heating a pulverised mixture of limestone & clay in a rotatory kiln.

→ Importance of Sodium, Potassium, Magnesium & Calcium: Monovalent Na, K ion & divalent Mg, Ca ions are found in large proportions in Biological fluids. These ions perform important biological functions such as maintenance of unbalance & nerve impulse conduction.

“The s-block, elements are called lighter metals because of their low density.

There are two groups (1 and 2) that belong to the s-block. In these two groups of elements, the last electron enters the s-subshell of the valence shell of their atoms. They are all highly reactive metals. The elements of group 1 are called alkali metals and consist up of elements: lithium, sodium, potassium, rubidium caesium and francium. These are so-called because these metals in reaction with water form hydroxides which are strongly alkaline in nature. Their general electronic configuration is ns type.

The elements of Group 2 include beryllium, magnesium calcium, strontium, barium and radium. These elements (except beryllium) are commonly known as alkaline earth metals. These are so .called because their oxides and hydroxides are alkaline in nature and these metal oxides are found in the earth’s crust. Their general electronic configuration is ns type.

Electronic Configuration Of Alkali Metals:

Francium is radioactive. Its largest-lived isotope 223 Fr has a half-life of only 21 minutes.

1. General characteristics of the alkali metals
(j) All the alkali metals have one valence electron ns1. This loosely held s-electron makes them the most electropositive metals which readily give M+ ions. Hence they are never found in a free state.
M → M+ + e-

Atomic and Ionic Radii
They have the largest sizes in a particular period in the periodic table. With the increase in atomic number, the atom becomes larger.

The monovalent ions (M⁺) are smaller than the parent atom, e.g.
Na⁺ → Na
K⁺ → K and so on

The atomic radii and ionic radii of alkali metals increase on moving down the group.
Li < Na < K < Rb < Cs and similarly
Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Ionisation Enthalpies
Due to large sizes, the ionisation enthalpies of alkali metals are considerably low and decrease down the group from Li so Cs, because the effect of increasing size outweighs the increasing unclear charge.

Hydration Energy: The hydration enthalpies of alkali metal ions decrease with an increase in ionic sizes.
Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺
Li⁺ ion has a maximum degree of hydration and for this reason, lithium salts are mostly hydrated e.g. LiCl.2H₂O.

Physical Properties:
1. Physical Appearance: Alkali metals are silvery-white, soft and light metals.

2. Density: Because of their large size, these elements have low densities, which increases down the group from Li to Cs. However, potassium is lighter than sodium.

3. Melting points & boiling points: The melting and boiling points of the alkali metals are low indicating weak metallic bonding.

4. Flame colouration: The alkali metals and their salts impart characteristic colour to an oxidizing flame.” This is due to energy imparted to the loosely bound electron as a result of which it gets excited and jumps to higher energy levels. When the excited electron comes back to the ground state, there is the emission of radiation in the visible region.

Alkali metals can therefore be detected by their flame tests.

Chemical Properties Of Alkali Metals:
The alkali metals are highly reactive due to their large size and low ionisation enthalpy and reactivity increases down the group.
1. Reactivity towards air: The alkali metals tarnish in dry air due to the formation of an oxide which in turn reacts with moisture to form hydroxides. They burn vigorously in oxygen forming oxides. Li forms monoxide, sodium forms peroxide, the other metals form superoxides.
2Li + O₂ → 2LiO (Oxide)
2Na + O₂ → Na₂O₂ (peroxide)
M + O₂ → MO₂ (superoxide)
[M = K, Rb, Cs]

Lithium (Li) shows exceptional behaviour in reacting with the nitrogen of air directly to form die nitride Li₃N as well.
6Li + N₂ → 2Li₃N (from the air)

Due to extreme reactivity, these metals are kept in kerosene oil.

2. Reactivity towards water:
2M + 2H₂O → 2M⁺ + 2OH^– + H₂↑
M = an alkali metal

Li reacts less vigorously with water. Other metals of the group react explosively with water. Reactivity increases down the group which is due to an increase in the electropositive character.

3. Reactivity towards hydrogen: Alkali metals react with hydrogen at about 673 K [Lithium at 1073 K] to form ionic hydrides which have high melting solids.
2M + H₂ → 2M⁺H^–

They also react with proton donors such as alcohol, gaseous ammonia and alkynes.
2C₂H₅OH + 2M → 2C₂H₅OM + H₂↑
CH = CH + Na → CH ≡ C^–Na⁺ + \(\frac{1}{2}\)H₂(g)

4. Reactivity towards halogens: The alkali metals react readily with halogens to form ionic halides M⁺X^–. However, lithium halide is somewhat covalent because of polarisation (The distortion of the electron cloud of the anion by the cation is called polarisation)
2M + X₂ → 2M⁺X^– Metallic (halide)

5. Solubility in liquid ammonia: All alkali metals are soluble in liquid ammonia. Dilute alkali metal-ammonia solution is blue in colour. With increasing concentration of metal in ammonia the blue colour starts changing to that of metallic copper after which a further amount of metal does not dissolve.

6. Reducing property (oxidation potentials): The tendency of an element to lose an electron is measured by its standard oxidation potential (E°), the more the value of E° of an element stronger will be its reducing character.
Since alkali metals have high values of E° these are powerful reducing agents and further lithium having the highest value is the strongest of them.

However, among the alkali metals, lithium although, has the highest ionization energy, yet is the strongest reducing agent. The greater reducing power of lithium is due to its larger heat of hydration which in turn is due to its small size.

7. Formation of alloys: The alkali metals form alloys amongst themselves as well as with other metals. The alkali metals dissolve readily in mercury forming amalgams. The process is highly exothermic.

General Characteristics Of The Compounds Of Alkali Metals:
(a) Oxides: Alkali metals when burnt in the air form oxides. The nature of oxides depends upon the nature of the alkali metal.

Under ordinary conditions, lithium forms the monoxide (Li₂O), sodium forms the peroxide (Na₂O₂) and the other alkali metals form mainly superoxides (MO₂) along with a small number of peroxides.

The increasing stability of the peroxide or superoxide, as the size of the metal ion increases, is due to the stabilization of large anions by larger cations through lattice energy effects. These oxides are easily hydrolysed by water to form the hydroxides according to the following reactions:
M₂O + H₂O → 2M⁺ + 2OH^–
M₂O₂ + 2H₂O → 2M⁺ + 2OH^– + O₂
2MO₂ + 2H₂O → 2M⁺ + 2OH^– + H₂O₂ + O₂

The oxides and the peroxides are colourless, but the superoxides are yellow or orange coloured. The superoxides are also paramagnetic. Sodium peroxide is widely used as an oxidizing agent in inorganic chemistry.

(b) Hydroxides: Alkali metal hydroxides, MOH are prepared, by dissolving the corresponding oxide in water. Their solubility in the water further increases as we move down the group due to a decrease in lattice energy.

Properties:

1. These are white crystalline solid, highly soluble in water and alcohols. Their solubility in the water further increases as we move down the group due to a decrease in lattice energy.
2. Since alkali metals are highly electropositive, their hydroxides form the strongest bases known. They dissolve in water with the evolution of much heat to give a strongly alkaline solution.
3. They melt without decomposition and are good conductors of electricity in the fused state.
4. These are stable to heat and do not lose water even at red heat. The thermal stability increases on moving from Li to Cs. However, they sublime at about 400°C and the vapours mainly consists of dimers. (MOH)₂.

(c) Halides: Alkali metal halides arc prepared by the direct combination of the element, M and halogens. They are normally represented by the formula MX and Cs and Rb, being of large size, also form Polyhalides, i.e. Csl₃

Properties:

1. All alkali halides except lithium fluoride are freely soluble in water (LiF is soluble in non-polar solvents).
2. They have high melting and boiling points.
3. Solubility of halides of alkaline metals: The solubility of alkali metal halides show a gradation. For example
   
4. They are good conductors of electricity infused state.
5. They have an ionic crystal structure. However, lithium halides have a partly covalent character due to polarising power of Li+ ions.

(d) Carbonates and bicarbonates: All alkali metals from carbonates of the type M₂CO₃. Due to the high electropositive nature of the alkali metals, their carbonates (and also the bicarbonates) are highly stable to heat (however, lithium carbonate decomposes easily by heat. Further, as the electropositive character increases in moving down the group, the stability of carbonates (and bicarbonates) increases in the same order.

Both carbonates and bicarbonates are quite soluble in water and their solubility increases as we move down the group from Li to Cs. Since carbonates are salts of a weak acid (carbonic acid H₂CO₃), they are hydrolysed in water to give a basic solution.
2M⁺ + CO₃^– + H – OH = 2M⁺ + HCO₃^2- + OH^–

Since the alkali metals are highly electropositive, these are the only elements that form stable solid carbonates. However, lithium due to its less electropositive nature does not form solid bicarbonate.

(e) Hydrides: Alkaline metals form hydrides of the type M⁺N^–. The presence of hydrogen as an anion in alkali metal hydrides is evidenced by the fact that on electrolysis hydrogen is liberated at the anode. The hydrides are not very stable. They react with water liberating hydrogen
LiH + H₂O → LiOH + H₂

These hydrides are, therefore, used as reducing agents. Lithium aluminium hydride, LiAlH₄ is even a stronger reducing agent and is used in organic chemistry.

2. Anomalous properties of Lithium:
1. Points of difference between lithium and other Alkali Metals:
(a) Lithium is much harder, its m.p. and b.p. are higher than the other alkali metals.

(b) Lithium is the least reacting but the strongest reducing agent among all the alkali metals. On combustion in air, it forms mainly monoxide Li₂O and the nitride, Li₃N, unlike other alkali metals.

(c) LiCl is deliquescent and crystallizes as a hydrate, LiCl.2H₂O whereas other alkali metal chlorides do not form hydrates. Lithium bicarbonate is not obtained in solid form while all other elements of this group form solid bicarbonate. Lithium unlike other alkali metals forms no acetylide on reaction with ethane.

(d) Lithium nitrate when heated gives lithium oxide Li₂O whereas other alkali metal nitrates decompose to give the corresponding nitrite.
4LiNO₃ → 2Li₂O + 4NO₂ + O₂
4NaNO₃ → 2NaNO₂ + O₂

(e) LiF and Li₂O are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

2. Points of similarities between Lithium and Magnesium
(a) Both lithium and magnesium are harder and lighter than other elements in their respective groups.
(b) Both Li and Mg react slowly with cold water. Their oxides and hydroxides are much less soluble and their hydroxides decompose on heating.
2LiOH → Li₂O + H₂O
Mg(OH)₂ → MgO + H₂O

(c) Both form nitrides by direct combination with N₂.
6Li + N₂ → 2Li₃N
3Mg + N₂ → Mg₃N₂

(d) Their oxides do not combine with an excess of O₂ to give peroxide or superoxide.

(e) The carbonates of both decompose on heating to give oxide and CO₂.

Solid bicarbonates are not formed by lithium and magnesium.

(f) Both LiCl and MgCL are soluble in ethanol.

(g) Both lithium perchlorate LiClO₄ and magnesium perchlorate Mg(ClO₄)₂ are extremely soluble in ethanol.

(h) Both LiCl and MgCl₂ are deliquescent and crystallise from aqueous solution as hydrates, LiCl.2H₂O and MgCl₂.8H₂O.

Some Important Compounds Of Sodium:
1. Sodium carbonate (washing soda) Na₂CO₃.10H₂O: Sodium carbonate is generally prepared by the Solvay process. In this process, the advantage is taken of the low solubility of sodium bicarbonate whereby it gets precipitated in the reaction of brine solution (sodium chloride) with ammonium bicarbonate. The latter is prepared by passing CO₂ to a concentrated solution of sodium chloride saturated with ammonia.

Ammonium carbonate first formed changes to ammonium bicarbonate.
1. 2NH₃ + H₂O + CO₂ → (NH₄)₂ CO₃
2. (NH₄)₂CO₃ + H₂O + CO₂ → 2NH4HCO₃
3. NH₄HCO₃ + NaCl → NH₄Cl + NaHCO₃

Sodium bicarbonate crystal separates. These are heated to give sodium carbonate.

In this process, NH₃ is recovered when the solution containing NH₄Cl is treated with Ca(OH)₂ Calcium chloride is obtained as a by-product.

5. 2NH₄Cl + Ca(OH)₂ → 2NH₃ + CaCl₂ + H₂O

Properties of sodium carbonate

1. It is a white crystalline solid which exists as decahydrate, Na₂CO₃10H₂O.
2. It is readily soluble in water.
3. On heating, the decahydrate loses its water of crystallisation to form monohydrate. Above 373 K, the monohydrate becomes completely anhydrous and changes to a white powder called soda ash.
   
4. It gets hydrolysed by water to form an alkaline solution
   CO₃^2- + H₂O → HCO₃^– + OH^–

Uses of sodium carbonate

1. It is used in water-softening, laundering and cleaning.
2. It is used in the manufacture of glass, soap, borax and caustic soda.
3. It is used in paper, paint and textile industries.
4. It is an important laboratory reagent both in qualitative and quantitative analysis.

Sodium Chloride NaCl:
Crude sodium chloride present in seawater (2.7 to 2.9% salt) is generally obtained by evaporation. It contains Na₂SO₄ CaSO₄, CaCl₂ and MgCl₂ as impurities. CaCl₂ and MgCl, are undesirable impurities because they are deliquescent (absorb moisture easily from the atmosphere).

To obtain pure NaCl, the crude salt is dissolved in a minimum amount of water and filtered to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas. Crystals of pure sodium chloride separate out. CaCl₂, and MgCl₂, being more soluble than NaCl remain in the solution.

Sodium chloride melts at 1081 K. It has a solubility of 36.Ogin 100g of water at 273K. The solubility does not increase appreciably with an increase in temperature.

Sodium Hydroxide (Caustic Soda) NaOH:
It is manufactured from the electrolysis of brine solution (an aqueous solution of NaCl) by Castner-Kellner cell. A mercury cathode and carbon anode are used.
Na⁺Cl^– (aq) → Na⁺(aq) + Cl^– (aq)
At cathode

At anode
Cl^– → \(\frac{1}{2}\) Cl₂ + e^–

The amalgam on treatment with water gives sodium hydroxide and H₂ gas.
2Na-amalgam + 2H₂O → 2NaOH + 2Hg + H₂

Properties

1. It is a white translucent solid.
2. Its M.Pt. is 591 K.
3. It gives a strongly alkaline solution in water.
4. Its crystals are deliquescent.
5. NaOH solution formed at the surface reacts with CO₂ from the atmosphere to form a crystal of Na₂CO₃.
   2NaOH + CO₂ → Na₂CO₃ + H₂O

Uses of sodium hydroxide:

1. It is used in the manufacture of sodium metal, soap, rayon, paper, dyes and drugs.
2. It is used in petroleum refining.
3. Sodium hydroxide is used for mercerizing cotton to make cloth unshrinkable.
4. It is used as a reagent in the laboratory.

Sodium Bicarbonate (Baking Soda) NaHCO₃:
Preparation
Na₂CO₃ + H₂O + CO₂ → 2NaHCO₃

Uses:

• Sodium bicarbonate is a mild antiseptic for skin infections.
• It is used in fire-extinguishers.
• It is known as baking soda because it decomposes on heating to generate bubbles of CO₂ (leaving holes in cakes or pastries and making them light and fluffy).

Biological Role Of Sodium & Potassium:
K+ ions and Na+ ions are present in the red blood cells. A 70 kg weighing man contains about 90g of Na and 170gof K. Sodium ions are found primarily on the outside of cells, is located in the blood plasma and in the interstitial fluid which surrounds the cells. These ions participate in the transmission of nerve signals in regulating the flow of water across cell membranes and in the transport of sugars and amino acids into cells.

Sodium and potassium which are chemically so alike, differ quantitatively in their ability to penetrate cell membranes, in their transport mechanisms and their efficiency to activate enzymes. Thus potassium ions are the most abundant cations within cell fluids, where they activate many enzymes, participate in the oxidation of glucose to produce ATP and with sodium, are responsible for the transmission of nerve; signals.

The ionic gradients of Na+ and K+ demonstrate that a discriminatory mechanism, called the sodium-potassium pump operate across the cell membranes which consumes more than one-third of the ATP used by a resting animal- about 15 kg per 21 h in a resting human.

Group-2 Elements: Alkaline Earth Metals: The group 2 elements comprise beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), Barium (Ba) and radium (Ra). They follow alkali metals in the periodic table. These (except Beryllium) are known as alkaline earth metals.
1. Atomic properties:
(a) Electronic configuration:

(b) Atomic and ionic sizes: The atomic and ionic radii of the alkaline earth metals are smaller than those of the alkaline metals in the corresponding periods. This is due to the increased nuclear charge in these elements.

(c) Ionization Enthalpies: The first ionization enthalpies of the alkaline earth metals are higher than those of Group 1 metals. The second ionization enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals.

2. Physical properties of the alkaline earth metals:
(a) Physical appearance: These metals in general are silvery-white, lustrous and relatively soft, but harder than the alkali metals. Beryllium and magnesium appear to be somewhat greyish.

(b) Melting and boiling points: The fairly higher melting and boiling points of the alkaline earth metals compared to those of the corresponding alkali metals and attributed to their smaller sizes and presence of two valence electrons. The trend is, however, not systematic.

(c) Flame colour: Chlorides of alkaline earth metals, except that of Be and Mg, produce the characteristic colour of flame due to easy excitation of electrons to higher energy levels. Beryllium and magnesium atoms due to their small size, bind their electrons more strongly, i.e., their ionisation energies are high. Hence these possess high excitation energy and not excited by the energy of the flame to a higher energy state with a result no colour is produced in the flame.

(d) Electrical and thermal conductivities: These properties are characteristics of typical metals.

3. Chemical Reactivity:
(a) Action of air: Their less reactivity than the alkali metals is evident by the fact that they are only slowly oxidised on exposure to air. However, when burnt in the air, they form ionic oxides of the type MO, except Ba and Ra which give peroxides. Thus, the tendency of the metal to form higher oxides like peroxide increases on moving down the group. On ignition powdered Be burns to give BeO & Be₃N₂. Mg also burns with dazzling brilliance to give MgO and Mg₃N₂.

(b) Action of water: These metals react slowly with water liberating hydrogen and forming metal hydroxides, e.g.
Ca + 2H₂O → Ca(OH)₂ + H₂
The reaction with water becomes increasingly vigorous on moving down the group.
Ba > Sr > Ca > Mg > Be (Reactivity with water)

The inertness of Be and Mg towards water is due to the formation of a protective thin layer of hydroxide on the surface of the metals.

(c) Action of hydrogen: All these elements, except beryllium, combine with hydrogen to form hydrides MH₂, BeH₂ is prepared indirectly.
2BeCl₂ + LiAlH₄ → 2BeH₂ + LiCl + AlCl₃

(d) Action of halogens: All these elements combine with halogens at elevated temperatures forming halides, MX₂. Beryllium halides are covalent, while the rest are ionic. The solubility of halides (except fluoride) decreases on moving down the group.

(e) Action with nitrogen: All these elements burn in nitrogen forming nitrides, M3N2 which react with water to liberate ammonia.
3Ca + N₂ → Ca₃N₂
Ca₃N₂ + 6H₂O → 3Ca(OH)₂ + 2NH₃
The ease of formation of nitrides decreases on moving down the group.

(f) Action with acids: On account of their high oxidation potentials, they readily liberate hydrogen from dilute acids. For example.
Mg + 2HCl → MgCl₂ + H₂

The reactivity of alkaline earth metals increases on moving down the group. This is due to an increase in electropositive character from Be to Ba. Thus beryllium reacts very slowly, Mg reacts very rapidly while Ca, Sr and Ba react explosively.

(g) Formation of amalgam and alloys: They form an amalgam with mercury and alloys with other metals.

(h) Complex formation: Beryllium, due to its small size, forms a number of stable complexes, e.g., [BeF₃]^–, [BeF₄]^2-, [Be(H₂O)]²⁺ etc.

(i) Reducing Character: They are strong reducing agents, Their reducing power is less than the corresponding alkali metals.

(j) Solubility in liquid ammonia: Alkaline earth metals dissolve in liquid ammonia giving coloured solutions. When the metal- ammonia solutions are evaporated, Hexammoniates M(NH3)6 are formed. The tendency for the formation of ammoniates decreases with an increase in the size of the metal atom, i.e., on moving down the group.
M + (x + y) NH₃ → M(NH₃)²⁺ + 2e^– (NH₃)y

4. General characteristics of compounds of the alkaline earth metals:
(a) Oxides and Hydroxides: The alkaline earth metal oxides, MO are prepared either by heating the metal in oxygen or better by calcination of carbonates.

These are extremely stable, white crystalline solids. Except for BeO, all the alkaline earth oxides are ionic, in which doubly charged ions are packed in a NaCl-type of lattice leading to their high crystal lattice energy and hence high stability. However, beryllium oxide is covalent due to its small size and relatively large charge on the beryllium ion. The high melting point of BeO is due to its polymeric nature.

The heavier metal oxides react with water to form soluble hydroxides which are strong bases.
MO + H₂O → M(OH)₂ + heat [where M = Ca²⁺, Ba²⁺ or Sr²⁺)

The solubility of hydroxides of alkaline earth metals in water increases on moving down the group. This is due to the fact that with the increase in the size of the cation (down a group), the lattice energy- decreases more than the decrease in hydration energy.

Halides:
They are obtained:

1. by heating the metal with halogens at high temperature or
2. by treating, metal carbonates with dilute halogen acids.

Beryllium halides are covalent compounds due to their small size and relatively high charge of Be²⁺ ion causing high polarising power. Due to the covalent bonding beryllium chloride, shows the following anomalous characteristics.

1. It has low melting and boiling points.
2. It does not conduct electricity in the fused state.
3. It is soluble in organic solvents such as ether.
4. It is hygroscopic and fumes in the air due to hydrolysis
   BaCl₂ + 2H₂O → Be(OH)₂ + 2HCl(g)
5. It is electron-deficient and behaves as Lewis acid.
   

The chlorides, fluorides, bromides and iodides of other alkaline earth metals are ionic solids and thus possess the following characteristics.

1. The melting and boiling points are high.
2. They conduct electricity in the molten state. Further, since the ionic character of the halides increases on moving down the group, the melting point and conductivity increase in the group from Mgd₂ to BaCl₂.
3. They are “hygroscopic and readily form hydrates, e.g., MgCl₂.6H₂O, CaCl₂.2H₂O, BaCl₂.2H₂O.
4. The halides (except fluorides) of the alkaline earth metals are soluble in water and their solubility decreases with an increasing atomic number of the metal due to a decrease in the hydration energy with the increasing size of the metal ion.

(c) Carbonates: The carbonates are invariable insoluble and therefore occur as solid rock materials in nature. However, the carbonates dissolve in water in the presence of carbon dioxide to give bicarbonates.

Most beryllium salts of strong oxo-acids crystallize as soluble hydrates. Beryllium carbonate is prone to hydrolysis and can be precipitated only in an atmosphere of carbon dioxide. The carbonates of magnesium and the other alkaline earth metals are all sparingly soluble in water, their thermal stability increases with increasing cationic size. Calcium carbonate finds use in the Solvay process for the manufacture of sodium carbonate in glassmaking and in cement manufacture.

(d) Sulphates: These can be prepared by dissolving the metal oxide in H₂SO₄.
MgO + H₂SO₄ → MgSO₄ + H₂O

The solubility of the sulphates of the alkaline earth metals decreases regularly on moving down the group. Thus beryllium sulphate is highly soluble in water, while barium and radium sulphates are practically insoluble.
The insolubility of barium sulphate is used for detecting an obstruction in the digestive system by the technique commonly known as barium meal.

The presence of BaSO₄ in the stomach helps in getting X-ray pictures because of the great scattering power of heavy Ba²⁺ ions. Barium sulphate is also used as a white pigment.

(e) Nitrates: The nitrates are made by the dissolution of the carbonates in dilute nitric acid. Magnesium nitrate crystallizes with six molecules of water. Barium nitrate crystallizes as an anhydrous salt. All of them decompose on heating giving the oxide.
2M(NO₃)₂ → 2MO + 4NO₂ + O₂ (M = Be, Mg, Ca, Sr or Ba)

Strontium and barium nitrates are used in pyrotechnics for giving red and green flames.

Anomalous behaviour of Beryllium
The anomalous behaviour of beryllium is mainly died to its very small size and partly due to its high electronegativity. These two factors increase the polarising power [Ionic charge/ (ionic radii)²] of Be²⁺ ions to such extent that it becomes significantly equal to the polarising power of Al³⁺ ions.

Hence the two elements resemble (diagonal relationship) very much.
1. Both of them have the same value of electronegativity (1.5).

2. The standard oxidation potential of Be and Al are of the same order (Be = 1.69 V, Al = 1.7 V)

3. In nature both occur together in beryl, 3BeO, Al₂O₃, 6SiO₂.

4. Due to its small size, beryllium has a high charge density and therefore, exhibits a strong tendency to form covalent compounds. Aluminium too has a strong tendency to form covalent compounds. Thus salts of both beryllium and aluminium have low m.p. are soluble in organic solvents and get hydrolysed by water.

Beryllium does show some tendency to form covalent compounds but other alkaline earth metals do not form covalent compounds.

5. Unlike other alkaline earth metals but like aluminium, beryllium is not easily affected by dry air.

6. Both (Be and Al) do not decompose water even on boiling; because of their weak electropositive character. Other alkaline earth » metals decompose even cold water evolving hydrogen.

7. Beryllium, like aluminium, reacts very slowly with dilute – mineral acids liberating hydrogen.
Be + 2HCl → BeCl₂ + H₂
2Al + 6HCl → 2AlCl₃ + 3H₂
Other alkaline earth metals react very readily with dilute acids.

8. The chlorides of both beryllium and aluminium have bridged chloride structures in the vapour phase.

9. Salts of these, metals form hydrated ions e.g., [Be(OH₂)₄]³⁺ and [Al(OH₂)₆]³⁺ in aqueous solutions.

10. Beryllium and aluminium both react with caustic alkalies to form beryllate and aluminate respectively. Other alkaline earth metals do not react with caustic alkalies.

Some Important Compounds Of Calcium:
1. Calcium oxide (Quick lime), CaO: Preparation: By heating limestone at 1273 K

(a) The reaction is reversible and thus in order to assure the complete decomposition of CaCO₃, carbon dioxide formed must be swept away by a current of air.

(b) Temperature should not be too high, because, at high temperature, clay (present as an impurity in limestone) will react with lime to form fusible silicates.

Properties:
1. Calcium oxide is a white amorphous substance.

2. When heated in an oxy-hydrogen flame, it gives an intense white light called limelight.

3. Action of water: On adding water, it gives a hissing sound and forms calcium hydroxide commonly known as slaked lime. The reaction is exothermic and known as slaking of lime.
CaO + H₂O → Ca(OH)₂ ΔH = – 64.5 kJ/mol

4. It reacts with SiO₂ and P₂O₅ at high temperature forming calcium silicate, CaSiO₃ and calcium phosphate; Ca₃(PO₄)₂ respectively
6CaO + 3P₂O₅ → 2Ca₃(PO₄)₂

5. With moist chlorine it forms bleaching powder, Ca(OCl)₂. With moist CO₂ it forms CaCO₃ and with moist SO₂ it forms CaSO₃ and with moist HCl gas, it forms CaCl₂. None of these gases will react when perfectly dried.

6. When heated with carbon at 2000°C, it forms calcium Carbide.

Uses of calcium oxide:
(a) It is used as a drying agent as such or as soda lime.
(b) Large quantities of quick-lime are used in the production of slaked lime.
(c) As a constituent of mortar, it is used on a very large scale in building constructions.

2. Calcium Hydroxide (Slaked lime), Ca(OH)₂:
Preparation:

1. By treating lime (quick lime) with water
   Ca O + H₂O → Ca(OH)₂
2. By the action of caustic alkalies on a soluble calcium salt.

Properties:
(a) It is a white amorphous powder, only sparingly soluble in water. Its solubility decreases with the increase in temperature.

(b) When dried and heated to redness, it loses a molecule of water and converted into calcium oxide (lime).

(c) Action of CO₂: Lime water is frequently used for the detection of C0₂ gas. C0₂ gas turns lime water milky due to the formation of CaC0₃.
Ca(OH)₂ + CO₂ → CaC0₃(s) + H₂O
However, the precipitate disappears on prolonged treatment with C0₂ because of the conversion of CaS0₃ (insoluble) to calcium bicarbonate (soluble).

The above solution, if heated again gives turbidity. This is due to the decomposition of calcium bicarbonate to calcium carbonate,

(d) Milk of lime reacts with chlorine to form hypochlorite, a constituent of bleaching powder
2Ca(OH)₂ + 2Cl₂ → CaCl₂ + Ca(ClO)₂ + 2H₂O

Uses of calcium hydroxide: Calcium hydroxide finds various uses:
(a) For absorbing acid gases
(b) For preparing ammonia from ammonium chloride
(c) In the production of mortar, a building material
(d) In glassmaking, tanning industry, for the preparation of bleaching powder and for purification of sugar.
(e) It is also used as a disinfectant
(f) As lime water in laboratories.

3. Plaster of Paris, CaSO₄. \(\frac{1}{2}\) H₂O
Preparation:
It is obtained when gypsum, CaSO₄.2H₂O is heated to 393 K
2(CaSO₄.2H₂O) → 2CaSO₄.H₂O + 3H₂O

Properties:
1. It is a white powder.

2. It has a very remarkable property of setting into a hard mass on wetting with water. So, when water is added to Plaster of Paris, it sets into a hard mass in about half an hour. The setting of Plaster of Paris is due to its hydration to form crystals of gypsum which set to form a hard solid mass.

The setting of plaster of Paris is accompanied by a slight expansion in volume due to which it is used in making castes for statues, toys, etc.

Uses of Plaster of Paris:
(a) It finds extensive use in surgical bandages, in casting and moulding.
(b) It is also employed in dentistry, in ornamental work and for taking castes of statues and busts.

Properties:
(a) It is a white powder and exists in two crystalline forms: Calcite and aragonite.
(b) It is insoluble in water but dissolves in the presence of CO2 due to the formation of calcium bicarbonate.
CaCO3 + H2O + CO2 → Ca(HCO3)2

Uses:
(a) Limestone is used:

• for the manufacture of the lime, element, washing soda and glass and
• as a flux, since CaO obtained from its decomposition combines with silica to form calcium silicate, CaSiO₃.

(b) Marble is used:

• for building purposes and
• in the laboratory for the production of CO₂ gas.

(c) Chalk is used:

• in paints (white ash) and distempers and
• in the production of CO₂ in the laboratory.

(d) Precipitated chalk is used:

• in toothpaste and powders
• in medicine for indigestion
• in adhesives and in cosmetic powders and
• to de-acidify wines.

4. Portland cement: It is made by heating a mixture of limestone (or chalk, shells etc.) with alumina silicates in carefully controlled amounts so as to give the approximate composition CaO 70%, SiO₂ 20%, Al₂O₃ 5%, FeCO₃ 3%. The new minerals are ground to pass 300-mesh sieves and then heated in a rotary kiln to 1773 K to give sintered clinker. This is ground to 325 mesh sieve and mixed with 2-5% gypsum. An average-sized kiln can produce 1000-2000 tonnes of cement per day.

When mixed with water, the setting of cement takes place. Chemically, it is the hydration of the molecules of the constitutions and their rearrangement.

The adhesion of other particles to each other and to the embedded aggregates is responsible for the strength of the cement which is due, ultimately, to the formation of Si-O-Si-O bonds.

The purposes of adding gypsum are only to slow down the process of setting the cement so that it gets sufficiently hardened.

Concrete: It is a mixture of cement, sand, gravel (small pieces of stone) and the appropriate amount of water. When the cement concrete is filled in and around a wire-netting or skeleton of iron rods and allowed to set, the resulting structure is known as reinforced concrete (RCC).

Uses of cement: It is used in concrete and reinforces concrete, in plastering and in the construction of bridges, dams and buildings.

Biological Importance Of Magnesium & Calcium:
An adult body contains about 25g of Magnesium and 1200 g of calcium as compared with only 5g of iron and 0.06g of copper. The daily requirement in the human body has been estimated to be 200-300 mg.

All enzymes that utilize ATP in phosphate transfer require magnesium as the cofactor.

The main pigment for the absorption of light in plants for photosynthesis is green coloured chlorophyll which contains magnesium.

About 99% of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function, interneuronal transmission, cell membrane integrity and blood coagulation. The calcium concentration in plasma is regulated at about 100 mg L^-1. It is maintained by two hormones.

Calcitonin and parathyroid hormone. Bone is not an inert and unchanging substance but is continuously being solubilized and redeposited to the extent of 400 mg per day in man. All the calcium passes through the plasma.

By going through these CBSE Class 11 Chemistry Notes Chapter 9 Hydrogen, students can recall all the concepts quickly.

Hydrogen Notes Class 11 Chemistry Chapter 9

→ Hydrogen: Hydrogen is the lightest atom with only one electron. Loss of this electron results in an elementary particle, the proton.

→ Isotopes of hydrogen: Protium (¹₁H), Deuterium (D or ²₁H) & Tritium (T or ³₁H). Tritium is radioactive.

→ Water-Gas Shift Reaction: Dihydrogen is obtained on an industrial scale by this reaction.

→ Bond dissociation Enthalpy: Bond dissociation enthalpy of dihydrogen is (435.88 KJ mol^-1) is highest for a single bond between two atoms of any elements.

→ Hydrides: Dihydrogen combines with almost all the elements under appropriate conditions to form hydrides. Three types of hydrides as Ionic or saline, covalent or molecular hydrides & metallic or non-stoichiometric hydrides.

→ Hydrogen Economy: The basic principle of a hydrogen economy is the transportation & storage of energy in the form of liquid or gaseous dihydrogen. In fact, it has promising potential for use as a non-polluting fuel of the near future.

→ Water: It is of great chemical & biological significance. Water molecule is highly polar in nature due to its bent structure. This property leads to hydrogen bonding which is maximum in ice & least in water vapor. Its property to dissolve many salts, particularly in large quantity makes it hard & hazardous for industrial use.

Temporary & permanent hardness can be removed by the use of, zeolites & synthetic ion exchangers.

→ Heavy water: D₂O is manufactured by the electrolysis of normal water. It is essentially used as a moderator in nuclear reactors.

→ Hydrogen peroxide: H₂O₂ has an interesting non-polar Structure & widely used as an industrial bleach & in pharmaceutical & pollution control treatment of industrial & domestic effluents.

→ Dihydrogen: Isotopes -Protium, Deuterium & Tritium, No. of neutrons-NIL, One, & Two Tritium is radioactive

→ Water-Gas: Mixture of CO & H₂.

→ Synthesis Gas or Syn Gas: When water gas is used for the synthesis of methanol & a no. of hydrocarbons. It is also called synthesis gas or syngas.

→ Coal Gasification: The process of producing syngas from coal is called coal gasification.

→ Uses of dihydrogen: In ammonia synthesis & in nitrogenous fertilizers & for preparing Vanaspati Ghee, manufacturing of organic chemical, particularly methanol; HCl & hydrides. Used as rocket fuel in space research. It does not produce any pollution & releases greater energy per unit mass of fuel in comparison to Gasoline & other fuels.

→ Hydrides: Three types of hydrides:

1. Ionic or Saline
2. Covalent or molecular
3. metallic or interstitial

→ Molecular hydrides:

1. Electron deficient
2. Electron precise
3. Electron rich

→ Water: Colourless & tasteless liquid. The unusual properties of water in the condensed phase (liquid & solid states) are due to the presence of extensive hydrogen bonding between water molecules.
Str. of the water molecule

→ Hard & Soft Water: The presence of Magnesium & Calcium salts in the form of hydrogen bicarbonates chloride & sulfate in water make water hard. Hard water does not give lather with soap. Water-free from soluble salts of calcium & magnesium is called Soft Water. It gives lather with soap easily.

→ Removal of Hardness of Water: Temporary hardness by boiling of water & by dark’s method. Permanent hardness by treatment with washing soda, Calgon’s method, ion exchange method & synthetic resins method.

→ Hydrogen-peroxide: Hydrogen peroxide is an important chemical used in the pollution control treatment of domestic & industrial effluents.

→ Storage of H₂O₂: H₂O₂ decomposes slowly on exposure to light. In presence of metal surfaces or traces of alkali (present in glass- containers), the following reaction is catalyzed.
2H₂O₂ (l) → 2H₂O(l) + O₂(g)

It is, therefore, stored in wax-lined glass or plastic vessels in dark. Urea can be added as a stabilizer. It is kept away from dust because dust can induce explosive decomposition of the compound.

→ Uses of H₂O₂:

1. As a hair bleach & a mild disinfectant. It is sold in the market as Perhydrol.
2. In manufacturing chemicals that are used in high-quality detergents.
3. In the synthesis of hydroquinone, tartaric acid & incertain food products & in pharmaceuticals (cephalosporin), etc.
4. It is also used in environmentally green chemistry.

→ Heavy Water: D₂O uses as a moderator in nuclear reactors & can be prepared by exhaustive electrolysis of water.

→ Hydrogen Economy: Hydrogen economy is an alternative. The basic principle of a hydrogen economy is the transportation & storage of energy in the form of liquid or gaseous dihydrogen. Nowadays, it is also used in fuel cells for the generation of electric power. Position of Hydrogen in the periodic table

Hydrogen is the first element in the periodic table. Its atom has only one proton and one electron. In the elemental form, it exists as H₂ and is called dihydrogen. The electronic configuration of hydrogen is Is1. Alkali metals have an electronic configuration of ns¹ which is similar. On the other hand like halogens (electronic configuration ns²np⁵, of 17th group) it is short by one electron than the corresponding noble gas configuration of He-1s². Hydrogen thus resembles both alkali metals as well as halogens.

But hydrogen also differs from alkali metals and halogens in some other respects. Thus it is unique in its behavior and is, therefore, best placed separately in the periodic table.

Occurrence of Dihydrogen (H₂):
It is the most abundant element in the universe (70% of the total mass of the universe) and is the principal element in the solar atmosphere. However, due to its light nature, it is much less abundant (0.15% by mass) in the earth’s atmosphere. In the combined form, it constitutes 15.4% of the earth’s crust and the oceans. In the combined form, besides in water, it occurs in plants, and animal tissues, carbohydrates, proteins, hydrides including hydrocarbons.

Isotopes of Hydrogen:
Hydrogen has three isotopes: protium ¹₁H, deuterium, ²₁H or D, and tritium, ³₁H or T.

They differ from one another in the number of neutrons. Ordinary hydrogen, protium has no neutrons, deuterium has one and tritium has two neutrons in the nucleus.

Tritium is radioactive and is present as 1 atom per 10¹⁸ atoms of protium. ²₁H or D is also known as heavy hydrogen.

Table: Physical properties of Dihydrogen and Dideuterium

Preparation of Dihydrogen

Laboratory Preparation of Dihydrogen H₂
(a) By the action of acids on metals: Metals (like Li. Na, Ba, Mg, Al, Zn, Fe, etc.) placed above hydrogen in the electrochemical series; when reacted with acids like HCl or dil. H₂SO₄ evolves hydrogen gas. Reaction with Li, K, Na, Ba, and Ca is violent while reaction with Zn, Fe, Al, and Mg is smooth.
Zn + H₂SO₄ → ZnSO₄ + H₂ (lab. method)
Fe + 2HCl → FeCl₂ + H₂

(b) By the action of alkalies on amphoteric metals Zn, Al, Pb, Sn, As, Sb, etc.)
Zn + 2NaOH → Na₂ZnO₂ (Sodium zincate ) + H₂
2Al + 2NaOH + 2H₂O → 2NaAlO₂ (Sodium aluminate) + 3H₂
Sn + 2KOH + H₂O → K₂SnO₃ (Potassium stannate) + 2H₂

(c) By the action of water on active metals (metals placed above electrochemical series)
1. Active metals like Na, K react at room temperature.
2Na + 2H₂O (cold) → 2NaOH + H₂ (violent)
Ca + 2H₂O (cold) → Ca(OH)₂ + H₂ (smooth)

2. Less active metals like Zn, Mg, Al liberate hydrogen only on heating.
Mg + 2H₂O (hot) → Mg(OH)₂ + H₂

3. Metals like Fe, Co, Ni, Sn can react only by passing steam.
3Fe (red hot) + 4H₂O (steam) → Fe₃O₄ + 4H₂

(d) By the action of water on a metal hydride:
LiH + H₂O → LiOH + H₂
CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂

Commercial production of Dihydrogen: The commonly used processes are:

Electrolysis of acidified water, using platinum electrodes, is employed for the bulk preparation of dihydrogen.

(a) Hydrogen of high purity (> 99.95%) is obtained by electrolysis warm aqueous barium hydroxide between nickel electrodes.

(b) Reaction of steam on hydrocarbons or coke at high temperatures in the presence of catalyst yields hydrogen gas.


CO is converted to CO₂ bypassing the gas’s steam over an iron oxide or cobalt oxide catalyst at 673K resulting in the generation of more H₂.

This is called the water-gas shift reaction.

(c) Relatively smaller quantities of dihydrogen (1-17 m³ h^-1) are obtained by passing a 1.1 molar mixture of vaporized methanol and water over a “base-metal chromite” type catalyst at 673 K. The mixture of hydrogen and carbon monoxide obtained is made to react with steam to give CO₂ and more hydrogen.

(d) It is also produced as a by-product of the brine electrolysis process for the manufacture of chlorine and sodium hydroxide. Presently-77% of the industrial hydrogen produced is from petrochemicals, 18% from coal, 4% from the electrolysis of aqueous solution, and 1 % from other sources.

Properties of Dihydrogen:
(a) Physical Properties:

1. Hydrogen is colorless, odorless, and tasteless gas.
2. It is the lightest element and also the lightest gas.
3. It is sparingly soluble in water.
4. Its critical temperature is very low (-236.9°C) at or below which can be liquefied by the application of suitable pressure. At -258.8°C it can be liquefied.
5. Its molecule is diatomic, indicated by the ratio of its specific heats at constant pressure and constant volume (C_p/C_v = 1.40).
6. It is adsorbed (occluded) by certain metals like Fe, Au, Pt, and Pd.

(b) Chemical properties:
1. Dihydrogen H₂ combines with halogens (X₂) to give hydrogen halides (HX). While the reaction with fluorine takes place even in the dark, with iodine a catalyst is required.
H₂(g) + X₂(g) → 2HX(g) (X = F, Cl, Br, I)

2. With dioxygen, dihydrogen forms water. The reaction is strongly exothermic
2H₂(g) + O₂(g) → 2H₂O(1) ΔH° = -285.8 kJ mol^-1

3. Reaction with dinitrogen it forms ammonia (NH₃)

ΔH° = -92.6 kJ mol^-1

4. Reaction with metals: With many metals, it combines at high temperatures to yield the corresponding hydrides.
H₂(g) + 2M(g) → 2MH(s)
Where M is an alkali metal

5. Reaction with metal ions and metal oxides: It reduces
some metal ions in aqueous solution and oxides of metals (less active than iron) into corresponding metals.
H₂(g) + Pd²⁺ (aq) → Pd(s) + 2H⁺(aq)
H₂(g) + Cu²⁺ (aq) → Cu(s) + 2H⁺(aq)
in general: YH₂(g) + MxOy(s) → xM(s) + yH₂O(l)

Reaction with organic compounds: It reacts with many organic compounds in presence of catalysts to give useful hydrogenated products of commercial importance. For example;

1. Hydrogenation of vegetable oils using nickel as catalyst gives edible fats (margarine and vanaspati ghee)
2. Hydroformylation of olefins yields aldehydes which further undergo reduction to give alcohol.
   H₂ + CO + RCH = CH₂ → RCH₂CH₂CHO
   H₂ + RCH₂CH₂CHO → RCH₂CH₂CH₂OH

1. The largest single use of dihydrogen is in the synthesis of ammonia which is used in the manufacture of nitric acid and nitrogenous fertilizers.

2. Dihydrogen is used in the manufacture of vanaspati fat by the hydrogenation of polyunsaturated vegetable oils like soybean, cotton seeds, etc.

3. It is used in the manufacture of bulk organic chemicals, particularly methanol.

4. It is widely used for the manufacture of metal hydrides.

5. It is used for the preparation of hydrogen chloride, a highly useful chemical.

6. In metallurgical processes, it is used to reduce heavy metal oxides to metals.

7. Atomic hydrogen and oxy-hydrogen torches find a use for cutting and welding purposes. Atomic hydrogen atoms (produced by dissociation of dihydrogen with the help of an electric arc) are allowed to recombine on the surface to be welded to generate a temperature of 4000 K.

8. It is used as rocket fuel in space research.

9. Dihydrogen is used in fuel cells for generating electrical energy. It has many advantages over conventional fossil fuels and electric power. It does not produce any pollution and releases greater energy per unit mass of fuel in comparison to gasoline and other fuels.

Hydrides:
Dihydrogen under certain reaction conditions combines with almost all elements except noble gases to form binary compounds called hydrides expressed as EH [like MgH₂] or EmHn (like B₂H₆).

They are of 3 types:

1. Ionic or Saline or Salt-Like Hydrides
2. Covalent or Molecular Hydrides
3. Metallic or Non-Stoichiometric Hydrides

1. Ionic or Saline or Salt-Like Hydrides: Lighter metal hydrides like LiH, BeH, and MgH, have significant covalent character. Ionic hydrides like K⁺H. Na⁺H are crystalline, non-volatile, and non-conducting in solid-state. However, their melts conduct electricity and in electrolysis liberate dihydrogen gas at the anode which confirms the existence of H⁺ ions

They are generally formed by s-block elements which are highly electropositive in character. These hydrides are Stoichiometric. Saline hydrides react violently with water producing dihydrogen gas.
NaH(s) + H₂O(aq) → NaOH(aq) + H₂(g)

Lithium hydride is rather unreactive at a moderate temperature with O₂ or Cl₂. It is, therefore, used in the synthesis of other useful hydrides, e.g.
8 LiH + Al₂Cl₆ → 2LiAlH₄ + 6 LiCl
2 LiH + B₂H₆ → 2LiBH₄

2. Molecular hydrides/[Covalent Hydrides]: These are formed by elements of highly electronegative elements (viz non-metals) which share electron(s) with hydrogen. In most cases, bonds are covalent in character, although in some cases (eg HF) bond is partly ionic in character. These have molecular lattices. The molecules are held together by weak van der Waal’s forces. These hydrides are soft, have low m.p. and b.p. They have low electrical conductivity.
The stability decreases progressively down a group, e.g.
NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃

In a period the stability increases with increasing electronegativity of the element forming the hydride.
e.g., CH₄ < NH₃ < H₂O < HF

These become increasingly acidic in character on moving from left to right along a given row in the periodic table. Thus, while NH₃ is a weak base, H₂O is neutral and HF is acidic. Similarly, in the next row, while PH₃ is a weak base, H₂S is a weak acid and HCl is highly acidic.

These are used as reducing agents.
Molecular hydrides are further classified according to the relative numbers of electrons and bonds in their Lewis structure into:

1. Electron-deficient,
2. Electron-precise, and
3. Electron-rich hydrides.

An electron-deficient hydride, as the name suggests, has too few electrons for writing its conventional Lewis structure. Diborane (B₂H₆) is an example. In fact, all elements of group 13 will form electron-deficient compounds. They act as Lewis acids i.e., electron acceptors.

Electron-precise compounds have the required number of electrons to write their conventional Lewis structures. All elements of group 14 form such compounds (e.g., CH₄) which are tetrahedral in geometry.

Electron-rich hydrides have excess electrons which are present as lone pairs. Elements of groups 15-17 form such compounds. (NH₃ has 1-one pair, H₂O-2, and HF-3 lone pairs). They will behave as Lewis bases i.e., electron donors. The presence of lone pairs on highly electronegative atoms like N, O, and F in hydrides results in hydrogen bond formation between the molecules. This leads to the association of molecules.

3. Metallic or non-stoichiometric (or Interstitial) Hydrides:
These are formed by many (d-block or f-block elements. However, metals of groups 7, 8, and 9 do not form hydrides. These hydrides conduct heat and electricity. They are non-stoichiometric. Hydrogen atoms occupy interstitial places in the lattices of metals. They are reducing’ agents and give out hydrogen easily. Hydrogen in them is present in atomic form.

Water:
A major part of all living organisms is made up of water. The human body has about 85 % and some plants have as much as 95 % water. It is a crucial compound for the survival of all life forms.

Physical properties of water:
It is a tasteless and colorless liquid.
The molecular mass of H₂O is = 18.0151 g mol^-1
Melting point = 273.0 K
Boiling point = 373.0 K
enthalpy of formation = – 285.9 kJ mol^-1
Enthalpy of fusion = 6.01 kJ mol⁺
Enthalpy of vaporisation (373 K) = 40.66 kJ mol^-1
Density (at 298 K) = 1.00 g cm^–3.

The unusual properties of water in the condensed phase (liquid and solid states) are due to the presence of extensive hydrogen bonding between water molecules. It boils at a higher temperature than H₂S or H₂Se only because of hydrogen bonding.

Structure of Water:
In the gas phase, water is a bent molecule with a bond angle of 104.5° and an O-H bond length of 95.7 pm as shown in Fig (a). It is a highly polar molecule, (Fig.(b)). Its orbital overlap picture is shown in Fig. (c) In the liquid phase water molecules are associated together by hydrogen bonds.

(a) The bent structure of water;
(b) the water molecule as a dipole and
(C) the orbital overlap picture In water molecule.

Structure of Ice:
Ice has a highly ordered three-dimensional hydrogen-bonded structure as shown in Fig. Examination of ice crystals with x-rays shows that each oxygen atom is surrounded tetrahedrally by four other oxygen atoms at a distance of 276 pm.

The structure of Ice

Hydrogen bonding gives the ice a rather open type structure with wide holes. These holes can hold some other molecules of appropriate size interstitially.

Chemical Properties Of Water:
1. Amphoteric Nature: It has the ability to act as an acid as well as a base, i.e., it behaves as an amphoteric substance. In the Bronsted sense, it acts as an acid with NH₃ and a base with H₂S.
H₂O(l) + NH₃(aq) ⇌ NH⁺₄ (aq) + OH^–(aq)
H₂O(l) + H₂S(aq) ⇌ H₃O⁺ (aq) + HS^– (aq)

The auto-protolysis (self-ionization) of water takes place as follows:

2. Reduction Reaction:
2H₂O(l) + 2Na(s) → 2NaOH(aq) + H₂(g)

3. Oxidation Reaction:
2F₂(g) + 2H₂O(aq) → 4H⁺(aq) + 4F^–(aq) + O₂

4. Hydrolysis Reaction:
P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq)
SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(aq)
N^3-(s) + 3H₂O(l) → NH₃(g) + 3OH^–(aq)

5. Hydrates Formation: From aqueous solutions, many salts can be crystallized as hydrated salts.

1. Coordinated water e.g. [Cr(H₂O)6]³⁺ 3Cl^–
2. Interstitial water e.g. BaCl₂.2H₂O.
3. Hydrogen bonded water e.g., [Cu(H₂O)₄]²⁺ SO₄ H₂O in CuSO₄. 5H₂O.

Hard and Soft Water:
The presence of calcium and magnesium salts in the form of hydrogen carbonate, chloride, and sulfate in water makes water Hard. Hard Water does not give lather with soap. Water-free from soluble salts of calcium and magnesium is called Soft Water. It gives lather with soap easily.

Hard water forms scum/precipitate with soap.

Disadvantages of using hard water

1. It is unsuitable for laundry.
2. It is harmful to boilers due to the deposition of salts as a scale on the walls of boilers. This reduces the efficiency of boilers.

There are two types of hardness:
(A) Temporary Hardness: It is due to the presence of the bicarbonates of calcium and magnesium, viz., Ca(HCO₃)₂ and Mg(HCO₃)₂

Temporary hardness can be removed by
1. Boiling

These precipitates are removed by filtration. The filtrate obtained is soft water.

2. Clark’s method: Lime water is added to hard water in Clark’s process to remove the precipitates of CaCO₃ formed.
Ca(HCO₃)₂ + Ca(OH)₂ → 2CaCO₃↓ + 2H₂O
Mg(HCO₃)₂ + 2Ca(OH)₂ → 2CaCO₃↓ + Mg(OH)₂↓ + 2H₂O

(B) Permanent Hardness: It is due to the presence of soluble salts of Mg and Ca in the form of chlorides and sulfates:
MgCl₂, CaCl₂, MgSO₄, CaSO₄. It can be removed by the following methods:
1. Treatment with washing soda (Na₂CO₃)
MCl₂ + Na₂CO₃ → 4 MCO₃↓ + 2NaCl
MSO₄ + Na₂CO₃ → MCO₃↓ + Na₂SO₄
M = Mg, Ca.

2. Calgon’s Method: Sodium hexametaphosphate Na6P60lg is commercially called Calgon. When added to hard water, the following reactions take place.
Na₆P₆O₁₈ → 2Na⁺ + Na₄P₆O₁₈^2-[M = Mg, Ca]
M²⁺ + Na₄P₆ O₁₈^2- → [Na₂MP₆O₁₈]^2-
The complex anion is not harmful.

3. Permutit Method: Permutit is an artificial zeolite- chemically it is sodium orthosilicate (Na₂Al₂Si₂O₈. xH₂O). Permutit removes cations like Ca²⁺, Mg^2+, and Fe²⁺ and releases an equivalent number of Na⁺ ions. For simplicity, it can be written as Na Z. When added to water the following reaction takes place
2NaZ + M²⁺(aq) → MZ₂(s) + 2Na⁺(aq); M = Mg, Ca

Permutit/zeolite is said to be exhausted when all the sodium in it is used up. It is regenerated when treating it with an aqueous sodium chloride solution.
MZ₂(S) + 2NaCl(aq) → 2NaZ + MCl₂(aq)

4. By the use of ion exchange resins (synthetic resins): This method removes all cations and anions present in water by means of ion-exchange resins. Water is first passed through cation exchange resins (giant organic molecules with-SO₃H or —COOH groups), which remove the cations like Na⁺, Ca²⁺, Mg²⁺ and others by exchange with H⁺. The resulting water is now passed through anion exchange resins (giant organic molecules with — NH₂ group) which remove the anions like Cl^–, SO₄^– and NO₃ by exchange with OH^–.

Hydrogen Peroxide (H₂O₂):
It is an important chemical used in the pollution control treatment of domestic and industrial effluents.

Preparation:
1. By the reaction of sulphuric acid or phosphoric acid on hydrated barium peroxide (BaO₂).
(a) BaO₂.8H₂O + H₂SO4 → BaSO₄(g) + H₂O₂ + 8H₂O

Anhydrous barium peroxide does not react readily with sulphuric
acid because a coating of insoluble barium sulfate is formed on its surface which stops further action of the acid. Hence hydrated barium peroxide, BaO2,.8H2O must be used.

(b) 3BaO₂ + 2H₃PO₄ → Ba₃(PO4)₂ + 3H₂O₂
Ba₃(PO₄)₂ + 3H₂SO₄ → 2BaSO₄(s) + 2H₃PO₄

Treatment with phosphoric acid is preferred to H₂SO₄ because soluble impurities like barium persulphate (from BaO₂.8H₂O+ H₂SO4) tend to decompose H₂O₂ while H₃PO₄ acts as a preservative (negative catalyst for H₂O₂). Moreover, excess barium peroxide should be avoided as it tends to decompose H₂O₂.
BaO₂ + H₂O₂ → BaO + H₂O + O₂

In both cases, BaSO₄ is removed by filtration and hence more or less a fuse H₂O₂ solution is obtained by this method.
1. By adding the calculated quantity of sodium peroxide to a 20 % ice-cold sulphuric acid solution (Merck’s process):
Na₂O₂ + H₂SO₄ → Na₂SO₄ + H₂O₂

Sodium sulfate is removed by cooling when crystals of Na₂SO₄ 10H₂O separate out.
In this method, sulphuric acid can be replaced by NaH₂PO₄

Manufacture of Hydrogen peroxide:
1. By electrolysis of 50 % sulphuric acid to give Perdisulphuric acid (H₂S₂O₈) which on distillation yields 30% solution of hydrogen peroxide.
2H₂SO₄ → 2H+ + 2HSO₄^–

At cathode (Cu coil):
2H+ + 2e^– → 2H + H₂ ↑

At anode (Pt)
2HSO₄^– → 2HSO₄ + 2e^–
2HSO₄ → H₂S₂O₈ (Persulphuric acid)
H₂S₂O₈ + 2H₂O → 2H₂SO₄ + H₂O₂

Alternatively, electrolysis may be done with ammonium hydrogen sulfate (ammonium sulfate + H₂SO₄).
(NH₄)₂SO₄ + H₂SO₄ → 2NH₄HSO₄
NH₄HSO₄ → H⁺ + NH₄SO₄

At cathode 2H⁺ + 2e^– → H₂
At anode 2NH₄SO₄ → (NH4)₂S₂O₈ + 2e^–

The ammonium persulphate formed is removed and quickly distilled with dil. H₂SO₄ under reduced pressure to give hydrogen peroxide.

2. By the auto-oxidation of 2-ethyl anthraquinone: In this process, the air is passed through a 10% solution of 2-ethyl anthraquinone in a mixture of benzene and higher alcohol.

The resulting 2-ethyl anthraquinone is then reduced by hydrogen in presence of palladium as a catalyst. Thus the continuity of the process is maintained and the process needs only H₂, atmosphere 0_2, and water as the major raw materials.

Physical Properties of hydrogen peroxide:

1. Pure hydrogen peroxide is a pale blue syrupy liquid.
2. It is an unstable liquid and decomposes into water and oxygen either on standing or on heating.
3. Hydrogen peroxide is diamagnetic.
4. In the pure state, its dielectric constant is 93.7 which increases with dilution.
5. It is more highly associated with hydrogen bonding than water.
6. Pure hydrogen peroxide is weakly acidic in nature while its aqueous solution is neutral.

Chemical properties:
It acts as an oxidizing as well as a reducing agent in both acidic and alkaline media. Simple reactions are described below.
1. Oxidising action in acidic medium
2Fe²⁺(aq) + 2H⁺(aq) + H₂O₂(aq) → 2Fe³⁺(ag) + 2H₂O(l)
PbS(s) + 4H₂O₂(aq) → PbSO₄(s) + 4H₂O(l)

2. Reducing action in acidic medium
2MnO₄ + 6H⁺ + 5H₂O₂ → 2Mn²⁺ + 8H₂O + 5O₂
HOCl + H₂O₂ → H₃O⁺ + Cl^– + O₂

3. Oxidising action in basic medium
2Fe²⁺ + H₂O₂ → 2Fe³⁺ + 2OH^–
Mn²⁺ + H₂O₂ → Mn⁴⁺ + 2OH^–

4. Reducing action in basic medium
I₂ + H₂O₂ + 2OH^– → 2I^– + 2H₂O + O₂
2MnO₄^– + 3H₂O₂ → 2MnO₂ + 3O₂ + 2H₂O + 2OH^–

Storage of H₂O₂
H₂O₂ decomposes slowly on exposure to light.
2H₂O₂(l) → 2H₂O(l) + O₂(g)

In the presence of metal surfaces or traces of alkali (present in glass containers), the above reaction is catalyzed. It is, therefore, stored in wax-lined glass or plastic vessels in dark. Urea can be added as a stabilizer. It is kept away from dust because dust can induce explosive decomposition of a compound.

Structure of H₂O₂
Hydrogen peroxide molecule has a non-polar structure. The molecular dimensions in the gas phase and chemical phase are shown in Fig.

(a) H₂O₂ structure (gas phase) Dihedral angle 111.5° (b) H,0, (so, id phase at 110 K. The dihedral angle is reduced to 90.2°.

Uses Of Hydrogen Peroxide:

1. In daily life, it is used as hair bleach and as a mild disinfectant. As an antiseptic, it is sold in the market as per hydro.
2. It is used to manufacture chemicals like sodium perborate and per-carbonate, which are used in high-quality detergents.
3. It is used in the synthesis of hydroquinone, tartaric acid, and certain food products and Pharmaceuticals (cephalosporin), etc.
4. It is employed in the industries as a bleaching agent for textiles, paper pulp, leather, oils, fats, etc.
5. Nowadays it is also used in Environmental (Green) Chemistry. For example, in pollution control treatment of domestic and industrial effluents, oxidation of cyanides, restoration of aerobic conditions to sewage wastes.

Heavy Water (D₂O):
It is extensively used as a moderator in nuclear reactors and in exchange reactions for the study of reaction mechanisms. It was first prepared by Urey by the exhaustive electrolysis of water.

It is used for the preparation of other deuterium compounds. For example,

Volume Strength Of Hydrogen Peroxide:
H₂O₂ is miscible with water in all proportions and forms a hydrate H₂O₂.H₂O (mp 221 K). A 30% solution of H₂O₂ is marketed as “100 Volume” hydrogen peroxide. It means that one milliliter of 30% H₂O₂ solution will give 100V of oxygen at STP. Commercially, it is marketed as 10V. It means it contains 3% H₂O₂.

Problem:
Calculate the strength of a 10 volume solution of hydrogen peroxide.
Answer:
10 volume solution of H₂O₂ means that 1L of this H₂O₂ will give 10L of oxygen at STP
2H₂O₂ (l) → O₂(g) + H₂O(l)
2 × 34 = 68g 22.4 L at STP
22.4 L of 02 at STP is produced from H₂O₂ = 68g

10 L of O₂ at STP is produced from H₂O₂ = \(\frac{68 \times 10}{22.4}\)g
= 30.36g
Therefore, the strength of H₂O₂ in 10 volume H₂O₂ = 30.36g L^-1.

Dihydrogen as a Fuel:
It releases large quantities of heat on combustion. On mass for mass basis H₂(g) can release, more energy than petrol (about three times). Moreover, pollutants in the combustion of dihydrogen will be less than petrol. The only pollutant will be oxides of dinitrogen (due to the presence of dinitrogen as an impurity with dihydrogen).

This, of course, can be minimized by injecting a small amount of water into the cylinder to lower the temperature so that reaction between dinitrogen and dioxygen may not take place.

However, the mass of the containers in which dihydrogen will be kept must be taken into consideration, A cylinder of compressed dihydrogen weighs about 30 times as much as a tank of petrol containing the same amount of energy. Also, dihydrogen gas is converted into a liquid state by cooling to 20K.

This would require expensive insulated tanks. Tanks of metal alloy like NaNi₅, Ti-TiH₂, Mg-MgH₂, etc. are in use of storage of dihydrogen in small quantities. These limitations have prompted researchers to search for alternative techniques to use dihydrogen in an efficient way.

In this view Hydrogen Economy is an alternative. The basic principle of a hydrogen economy is the transportation and storage of energy in the form of liquid or gaseous dihydrogen. The advantage of a hydrogen economy is that energy is transmitted in the form of dihydrogen and not as electric power.

It is for the first time in the history of India that a pilot project using dihydrogen as fuel was launched in Oct 2005 for running automobiles. Initially, 5 % dihydrogen has been mixed in CNG for use in four-wheeler vehicles. The percentage of dihydrogen would be gradually increased to reach the optimum level. Nowadays, it is also used in fuel cells for the generation of electric power.

By going through these CBSE Class 11 Chemistry Notes Chapter 8 Redox Reactions, students can recall all the concepts quickly.

Redox Reactions Notes Class 11 Chemistry Chapter 8

→ Reactions taking place in an electrochemical cell are redox reactions in nature.

→ In an electrochemical cell loss of free energy appears as electrical energy.

→ The reaction in an electrochemical cell is spontaneous in nature.

→ A salt bridge maintains the electrical neutrality of the two electrolytes in their half cells.

→ The e.m.f. of an electrochemical cell is E°_cathode — E°_anode cathode anode

→ According to the electronic concept, the loss of electron is oxidation, and the gain of the electron is reduced.

→ The oxidation number of free elements homo atomic molecules and also of the neutral molecule is zero.

→ Electrolysis is the migration of the ions of the electrolyte towards the oppositely charged electrode when the current is passed.

→ In an electrolytic cell, the redox reaction is non-spontaneous in nature.

→ The chemical energy of the redox reaction occurring in the galvanic cell is converted into electrical energy.

→ Electrons flow from anode to cathode in the external circuit while current flow from cathode to anode.

→ 95600 c of charge represents one Faraday.

→ Oxidation: Oxidation is a process in which an atom or ion loses an electron(s).

→ Reduction: Reduction is a process in which an atom or ion gains an electron(s).

→ Oxidizing agent: (Oxidant) is a species that can readily accept one or more electrons.

→ Reducing agent: (Reductant) is a species that readily lose one or more electrons.

→ Redox Reaction: Redox reaction is a chemical reaction in which oxidation and reduction occur simultaneously.

→ Electrochemical cell: Electrochemical cell is a device in which oxidation and reduction half-reactions are carried indirectly and the loss of chemical energy during the reaction appears as electrical energy.

→ Electrolytic cell: Electrolytic cell is a device in which electrical energy is supplied from an external source to bring about a chemical reaction.

→ Anode: Anode is an electrode where the electrons are released or where oxidation takes place.

→ Cathode: A cathode is an electrode where the electrons are accepted or where reduction takes place.

→ Half cell: A half cell is a portion of an electrochemical cell in which either oxidation or reduction takes place.

→ Standard hydrogen electrode: Standard hydrogen electrode is an electrode that is used to calculate the reduction potential of another electrode. Its own reduction potential is taken as zero.

→ Standard reduction potential: Standard reduction potential of an electrode is its reduction potential as compared to that of a standard hydrogen electrode which is taken as zero.

→ Salt bridge: Salt bridge is an inverted U-shaped glass tube that contains a suitable electrolyte and connects the two-half cells in an electrochemical cell.

→ E.m.f. of a cell: E.m.f. of a cell is the difference between the reduction potential of electrodes when the cell is not sending the current.

→ Potential difference: Potential difference is the difference of potential between two electrodes when the cell is sending currents.

→ Electro-chemical series: Electrochemical series is the series obtained by arranging the electrode in order of increasing standard reduction potential values.

→ Electrolyte: Electrolyte is a substance that is capable of conducting electricity either in a molten state or when dissolved in an aqueous solution.

→ Electrolysis: Electrolysis is the process of the decomposition of an electrolyte on passing electric current.

→ Oxidation Number: The oxidation number of an element is the residual charge which its atom appears to have when all other atoms present in its combination are removed as ions.

→ Disproportionation Reaction: In this reaction, an element in one oxidation state is simultaneously oxidized and reduced.

→ Redox couple: A redox couple consists of the oxidized and reduced forms of the same substance taking part in an oxidation and reduction half-reaction.

By going through these CBSE Class 11 Chemistry Notes Chapter 6 Thermodynamics, students can recall all the concepts quickly.

Thermodynamics Notes Class 11 Chemistry Chapter 6

Thermodynamic Terms: The laws of thermodynamics deal with energy changes at a macroscopic system involving a large number of molecules rather than a microscopic system containing a few molecules

→ System: The part of the universe which is under thermodynamic study/scrutiny is called a system.

→ Surroundings: The remaining part of the universe with which the system can exchange both matter and energy is called the surrounding. System and surrounding taken together constitute the universe.
The universe = The system + the surroundings.

Types of System
1. Open System: An open system is one in which there is an exchange of matter and energy between the system and surroundings, e.g., a reaction occurring in a test tube.

2. Closed System: A closed system is one in which there is only a transfer of energy and not matter with the surroundings, e.g., the presence of reactants in a closed vessel.

3. Isolated System: An isolated system is one in which no exchange of energy and matter is possible between the system and the surroundings. The presence of reactants in a thermos flask is an example of an isolated system.

→ The State of the System: When the properties of a system like its pressure, volume, temperature, and composition are specified quantitatively, it is said to have stated. If any of these observable (measurable) properties undergo a change, the system is said to have another state.

→ State Variables: Such properties of the system, by changing any one of which, the system changes its state, are called State Variables.

→ State Function: A physical quantity is said to be a state function if its value depends upon the initial state and final state of the system and not the path followed by the system. Internal energy, enthalpy, free energy, entropy are all state functions. Heat and work are not state functions, because they depend upon the path followed.

→ Extensive Properties: Extensive properties of a system are those properties that depend upon the amount of substance contained in the system. Mass, volume, heat capacity, internal energy, enthalpy, Gibbs free energy are all extensive properties.

→ Intensive Properties: Intensive properties of a system are those which are independent of the quantity of matter contained in the system. They depend upon only the nature of the substance. Temperature, pressure, viscosity, density, surface tension, refractive index, specific heat, freezing point, boiling point are all intensive properties.

There are four types of thermodynamic processes:

1. Isothermal Process: A process in which temperature remains constant.
2. Adiabatic Process: A process in which no heat exchange takes place,(i.e., q = 0) between the system and surroundings.
3. Isochoric Process: A process in which volume remains constant.
4. Isobaric Process: A process in which pressure remains constant.

→ Internal Energy ‘U’: Its absolute value cannot be determined.
The energy stored within a substance (or system) is called its internal energy. It is a state function.
ΔU denotes a change in internal energy.
ΔU = U₂ – U₁ = U_p – U_r
The internal energy of an ideal gas depends upon only its temperature.

Hence in an isothermal process involving an ideal gas ΔU = 0.

→ Reversible Process: When a process is carried out infinitesimally slowly so that it can be retraced at each step of the change, it is called a reversible process. In such a process system and surroundings are always in equilibrium. .

→ Irreversible Process: A process that is carried out so rapidly that the system does not get a chance to attain equilibrium It cannot be retracted to each step of the change. A reversible process may take infinite time for completion, whereas an irreversible process takes a finite time for completion.

• If w_reversible = work obtained in a reversible process
• If w_irreversible = work obtained in a reversible process is maximum.

→ Cyclic process: If a system after having undergone a series of changes returns back to its initial state, the process is called a cyclic process. The path of such a process is called a cycle.
In cyclic process ΔU = 0.

→ The first law of thermodynamics: States: “Energy can neither be created nor destroyed. However, it can be transformed from one form into another.”

In other words, the first law of thermodynamics can be stated as
“The total energy of the universe (system + surroundings) remains constant”.

The mathematical formulation of the first law of thermodynamics: Let us consider a system in its initial state having internal energy. U₁ If heat equal to q is supplied to the system, and work equal to w is done on the system, then the internal energy of the system is the final state (U₂) is given by
U₂ = U₁ + q + w
U₂ – U₁ = q + w
ΔU = q + w
or change in internal energy = Heat given to the system + work done on the syštem.

This relationship between internal energy, work, and heat is the mathematical statement of the first law of thermodynamics.

Applications: Heat, Work, and Internal Energy:
(a) Heat: The energy exchange between a system ad the surroundings when their temperatures are different is commonly known as heat.

If a system is at a temperature higher than the temperature of its surroundings, then it loses temperatures are different is commonly known as heat.

If a system is at a temperature higher than the temperature of its surroundings, then it loses heat energy to the surroundings. This transfer of heat lowers the temperature of the system and raises the temperature of the surroundings

(b) Work (w): Work is said to be done if the point of application of a force moves through a certain distance. For example, if a gas, enclosed in a cylinder fitted with an airtight piston has higher pressure than the external pressure, then the piston will move outwards. The piston continues to move until the pressure inside and outside are equal. In this, the gas inside the cylinder is pushing the piston outwards. Thus, the gas (system) is doing work.

(c) Internal Energy: the total energy contained in a system is called its internal energy or intrinsic energy. This is denoted by a symbol E or U.

The internal energy of a system depends upon the state of the system and not upon how the system attains that state. The internal energy thus is a state function.

Each substance or system possesses a definite amount of internal energy under a given set of conditions. For example, the internal energy of a system in state A be U_A and in state B, it is U_B. Then, change in the internal energy (ΔU) of the system is going from state A to B is
ΔU = U_final – U_initial = U_B – U_A = w_ad.
wad is positive when work is done on the system. wad is negative if the work is done by the system.

This change in the internal energy depends only upon the initial and the final state and not on how this change is brought about.

Heat capacity and specific heat capacity:
1. Heat capacity: The heat capacity (C) of a sample of a substance is defined as the quantity of heat energy required to raise its temperature by 1 K (or 1°C). From this definition, we can write
Q = C. Δt
where Q is the quantity of heat given to the sample, At is the rise in the temperature of the substance. The unit of heat capacity is Joules per degree Celsius (J/°C or J°C^-1) or Joules per Kelvin (J/K or JK^-1).

2. Specific heat capacity: The specific heat capacity of a substance is equal to the quantity of heat required to raise the temperature of 1 unit mass of capacity of the substance by 1 °C, or 1 K.
S.I. unit of specific heat is joules per kilogram per degree (J/Kg°CorJ/kg.K).

3. Molar heat capacity: The molar heat capacity (c_m) of a substance is defined as the quantity of heat required to raise the temperature of one mole of a substance by 1 K or (or 1°C). Thus

Molar heat capacity = Specific heat capacity × Molar mass
c_m = C × M
The unit of molar heat capacity is J mol^-1 K^-1.

There are two types of Heat Capacities.
1. Heat capacity at a constant volume (C_v): The heat supplied to a system to raise its temperature through 1°C keeping the volume of the system constant is called heat capacity at constant volume.

2. Heat capacity at constant pressure (C_p): The heat supplied to a system to raise its temperature through 1°C keeping external pressure constant is called heat capacity at constant pressure.
C_v = \(\frac{d \mathrm{E}}{d \mathrm{~T}}\)
C_p = \(\frac{d \mathrm{H}}{d \mathrm{~T}}\)

Relationship between (C_p) and (C_v) C_p – C_v = R

→ Enthalpy (H): The sum of internal energy (E) and the product of pressure and volume is called enthalpy.
H = E + PV
P = Pressure,
V = Volume like E, Enthalpy is also a state function.

Its absolute value like E cannot be determined.
Nor is there any necessity for it. What is required is ΔH:
Change in enthalpy.

The value of ΔE and ΔH can be measured with the help of a Bomb Calorimeter. Calorimetry is the technique to measure energy changes associated with chemical or physical processes.
ΔH = H_products – H_reactants

The enthalpy change of a reaction: (ΔrH) is equal to the heat absorbed or evolved during a reaction at constant temperature and pressure.

∴ ΔH = q_p where q_p: Heat absorbed or evolved at constant pressure.
ΔH = ΔE + P ΔV … (1)

When there is no change in volume (i.e. ΔV = 0)
ΔE = q_v where q_v = Heat change at constant volume.

Equation (1) becomes
q_P = q_v + PΔV
q_P = q_v + Δn_gRT
where Δn_g is the change in the no. of gaseous moles of the products minus that of the reactants.

Sign Convention for Heat and Work

• Heat absorbed by the system = q positive
• The heat evolved by the system = q negative
• Work done on the system = w positive
• Work done by the system = w negative

Enthalpies or Heats of Reactions (Δ_rH):
The enthalpy change (amount of heat evolved or absorbed) during a chemical reaction, the moles of reactants and products being the same as indicated by the balanced chemical equation is called Enthalpy of the reaction (Δ_rH).

Exothermic Reaction: These are the reactions in which heat is evolved. Here sum of heat contents of the reactants (H_R) > Sum of heat contents of the products (H_p).
or
H_R > H_p .
or
ΔH = H_p – H_R < 0
or
ΔH < 0
or
ΔH is negative

Examples of exothermic reactions:

1. C(s) + O₂(g) → CO₂(g) + 393.5 kJ
2. H₂(g) + \(\frac{1}{2}\)O₂(g) → H₂O(1) + 285.8 kJ
3. N₂(g) + 3H₂(g) → 2NH₃(g) + 92.4 kJ

or in terms of ΔH

1. C(s) + O₂(g) → CO₂(g); ΔH = – 393.5 kJ
2. H₂(g) + \(\frac{1}{2}\)O₂(g) → H₂O(l); ΔH = – 285.8 kJ
3. N₂(g) + 3H₂(g) → 2NH₃(g); ΔH = – 92.4 kJ

→ Endothermic Reactions: These reactions which proceed with an absorption of heat are called endothermic reactions.
In endothermic reactions sum of the enthalpies of the product > Sum of the enthalpies of reactants
or
ΣH_p > ΣH_R
or
ΔH = H_p – H_R > 0
or
ΔH >0
or
ΔH is positive

Examples of endothermic reactions

1. N₂(g) + O₂(g) → 2NO(g) – 180.7 kJ
2. C(s) + H₂O(g) → CO(g) + H₂(g) – 131.4 kJ
3. C(s) + 2S → CS₂(g) – 92.4 kJ

or writing the above thermochemical equation in terms of AH

1. N₂(g) + O₂(g) → 2NO(g); ΔH = + 180.7 kJ
2. C(s) + H₂O(g) → CO(g) + H₂(g); ΔH = + 131.4 kJ
3. C(s) + 2S → CS₂(g); ΔH = + 92.4 kJ

→ Thermochemical Equation: When a balanced chemical equation not only indicates the quantities of different reactants and products but also indicates the amount of heat evolved or absorbed, it is called a thermochemical equation.
thus N₂(g) + O₂(g) → 2NO(g); ΔH = + 180.7 kJ is a thermochemical equation

Factors on which the heat of the reaction depends

• Quantities of the reactants involved.
• The physical state of the reactants and products.
• Allotropic modifications
• The concentration of the solutions.
• Temperature
• Conditions of constant pressure or constant volume.

→ Standard Enthalpy Change: A substance in its most stable form at 25°C or 298 K under one atmospheric pressure is said to be in its standard form.

The enthalpy change of reaction when all the reactants and products are in their standard states i.e., at 25°C or 298 K and under a pressure of one bar is called standard enthalpy change.

It is usually represented by Δ_rH° or Δ_rH²⁹⁸.

Common Types of Enthalpy or Heat Changes
Enthalpy or Heat of formation: The enthalpy of formation is the enthalpy change when one mole a of the compound is formed from its elements. It is deno’.ed by ΔH_f.

For example, enthalpies of formation of carbon dioxide and methane (CH₄) may be expressed as
C(s) + O₂(g) → CO₂(g); ΔH = ΔH_f=- 395 kJ
C(s) + 2H₂(g) → CH₄(g); ΔH = ΔH_f = -74.8 kJ

If all the species of the chemical reactions are in their standard state, (i.e.. at 298 K and one atmospheric pressure) the enthalpy of formation is called standard enthalpy of formation, expressed as ΔH_f.

The enthalpy of every element in its standard state is arbitrarily assumed to be zero.

Thus, the enthalpy change accompanying the formation of one mole of a compound from its elements, all the substances being in their standard states (1 atm pressure and 298 K) is called standard enthalpy of formation.

The enthalpies of the formation of different substances can be used to calculate the enthalpy change of the reaction.

Standard enthalpies of
ΔH_f° = formation of all products
or
Standard enthalpies of formation of all reactants
Δ_fH° = ΣΔH_f° (products) – ΣΔH_f° (reactants)

→ Eptnalpy or Heat of combustion: The enthalpy change when 1 mole of a substance is completely burnt in excess of oxygen or air is called enthalpy of combustion. It is expressed as Δ_cH°
For example,
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l); Δ_fH° = – 890.3 kJ
C₆H₆(l) + \(\frac{15}{2}\) O₂(g) → 6CO₂(g) + 3H₂O(l); Δ_fH° = – 3268 kJ

→ Enthalpy or Heat of solution: The enthalpy change when one mole of a substance is dissolved in a specified quantity of a solvent at a given temperature is called enthalpy of solution.
NH₄NO₃(s) + H₂O(1) → NH₄+ NO₃^– (aq. 1.o m); ΔH = 26.0 kJ

To avoid the amount of solvent, the heat of the solution is usually defined as an infinitely dilute solution. Thus, the heat of solution at infinity dilution is the heat change when one mole of a substance is dissolved in such a large quantity of solvent so that further dilution does not give any further heat change.
NaCl(s) + aq → NaCl(aq); ΔH = 5.0 kJ

→ Enthalpy or Heat of neutralization: The heat change when one gram equivalent of an acid is completely neutralized by a base or vice versa in dilute solution, is called heat of neutralization.
For example,
1. Neutralisation of HCl with NaOH
HCl(aq) + NaOH(aq) → NaCI(aq) + H₂O(l); ΔH = – 57.1 kJ

2. Neutralisation of CH3COOH with NaOH
CH₃COOH(aq) + NaOH(aq) → CH₃COONa + H₂O(l); ΔH = – 55.9 kj.

→ Calorific Values of Foods and Fuels: Different fuels and foods produce different amounts of heat on combustion. They are usually expressed in terms of their calorific value. “The calorific value of a fuel or food is the amount of heat in calories (or joules) produced from the complete combustion of 1 gram of the fuel or the food.
e.g., C₆H₁₂O₆(S) + 6O₂(g) → 6CO₂(g) + 6H₂O(g) + 2840 kJ mol^-1
1 Mole = 180 g.
∴ Calorific value of glucose = \(\frac{2840}{180}\) = 15.78 kJ g^-1.

→ Measurement of Enthalpy of Combustion: At constant volume (q_v or q_v = AE) is done in a Bomb calorimeter. The known weight of the compound whose enthalpy of combustion is to determined is burnt in oxygen in a bomb calorimeter
M
ΔE = Q × Δf × \(\frac{M}{m}\)
where Q = Heat capacity of the calorimeter
Δt = rise in temperature of water in the calorimeter
m = mass of the substance taken
M = Molecular mass of the substance

Enthalpy Changes during Phase Transitions
1. Enthalpy of Fusion: It is the heat change accompanying the transformation of one mole of a solid substance into its liquid form at its melting point.
H₂O(S) (ice) → H₂O(l) (water); ΔH_fus = + 60 kJ mol^-1

2. Enthalpy of vaporization: It is the heat change accompanying the conversion of 1 mole of a liquid into its gaseous state at its boiling point.
H₂O(l) (water) → H₂O(g) (steam); Δ_vap H° = + 407 kJ mol^-1

3. Enthalpy of sublimation: It is the heat change that takes place when 1 mol of a solid changes directly into its vapor phase at a given temperature below its melting point.
I₂(s) → I₂(g); Δ_sub H° = + 62.39 kJ mol^-1
ΔH_sublimation = ΔH_fusion + ΔH_vaporisation

Hess’s Law of Constant Heat Summation
“The total amount of heat evolved or absorbed in a reaction is the same whether the reaction takes place in a single step or in a number of steps, provided the temperature is kept constant.
C(s) + O₂(g) → CO₂(g); ΔH = – 393.5 kJ mol^-1

It can be made to proceed in two steps

1. C(s) + \(\frac{1}{2}\)O₂(g) → CO(g); ΔH₁ = – 110.5 kJ mol^-1
2. CO(g) + \(\frac{1}{2}\)O₂(g) → CO₂(g); ΔH₂ = – 283.0 kJ mol^-1

Thus, the.total heat evolved in steps (1) and (2) is
ΔH₁ + ΔH₂ = – 110.5 + (- 283.0) = – 393.5 kJ mol^-1
which is the same when the reaction takes place directly in a single step.

From the above
Heat change is going from A → D via I route = Q
Heat change in going from A → B → C → D via route II = q₁ + q₂ + q₃
∴ Q = q₁ + q₂ + by Hess’s Law.

It is a corollary from the 1st Law of Thermodynamics.

Applications of Hess’s Law:

1. It helps to calculate the enthalpies of the formation of many compounds which cannot be determined experimentally.
2. It helps to calculate the enthalpy of allotropic transformations.
3. It helps to calculate the enthalpies of hydration.
4. It helps to calculate the enthalpies of different reactions.

→ Bond Enthalpy or Bond Energy: It is defined as the amount of energy required to dissociate one mole of bonds presents between the atoms in the gaseous molecules.
H₂(g) + 436 kJ mol^-1 → 2 H(g)
For diatomic molecules like H₂, bond energy is its bond dissociation energy.

→ Spontaneous Process: A process that can take place of its own or has an urge or tendency to take place is called a spontaneous process. It is simply a process that is feasible.
1. Examples of processes that take place by themselves

1. Dissolution of common salt in water
2. Evaporation of water in an open vessel
3. The flow of water downhill
4. The flow of heat from the hot end to the cold end.

2. Example of processes that take place on initiation

1. Lighting a candle (initiation by ignition)
2. Heating of CaCO₃ to give CaO and CO₂.

→ Non-Spontaneous Process: A process that can neither take place of its own nor on initiation is called a non-spontaneous process.

Examples of non-spontaneous processes.

1. The lighting of a candle (initiation by ignition)
2. Heating of CaCO₃ to give CaO and CO₂.

→ Non-Spontaneous Process: A process that can neither takes place of its own nor on initiation is called a non-spontaneous process.

Examples of non-spontaneous processes.

1. The flow of water uphill.
2. The flow of heat from a cold body to a hot body.

→ The Driving Force for a Spontaneous Process: The force which is responsible for the spontaneity of a process is called the driving force. It is based upon.

1. The tendency towards minimum energy: To acquire maximum stability the system tends to acquire minimum energy.
2. The tendency towards maximum randomness: When a system proceeds from a state of orderliness to disorderliness as the mixing up of two gases, it tries to have maximum randomness or disorder. An increase in disorder leads to the spontaneity of a process.

→ Entropy: It is a measure of disorder or randomness of the system. Like internal energy, enthalpy, entropy is also a state function and it is an extensive property. ΔS is called a change of entropy.
ΔS = S₂ – S₁ = ΣS_products – ΣS_reactants
= \(\frac{q_{\text {reversible }}}{\mathrm{T}}\) under isothermal conditions.

Thus Entropy change (ΔS) during a process is defined as the amount of heat (q) absorbed isothermally and reversible divided by the absolute temperature (T) at which the heat is absorbed.

→ Units of Entropy: As ΔS = \(\frac{q}{T}\) = jK^-1 mol^-1 (S.I. unit)
= Calories K^-1 mol^-1 (C.G.S. unit)

Entropy Change During Phase Transformations
1. Entropy of fusion (ΔS_fusion): The entropy of fusion is the change in entropy when 1 mole of a solid substance changes into its liquid form at its melting point.

Mathematically, ΔS_fus = S_liq – S_solid = \(\frac{\Delta \mathrm{H}_{\text {fusion }}}{\mathrm{T}_{m}}\)
where T_m = Melting point of the solid (K)
ΔH_fusion = Enthalpy of fusion per mole.

2. Entropy of Vaporization: It is defined as the entropy change when 1 mole of a liquid changes into its vapor at its boiling point.
Mathematically, ΔSvap = Svap – Sliq = \(\frac{\Delta \mathrm{H}_{\text {vap}}}{\mathrm{T}_{b}}\)
where, ΔH_vap = Enthalpy of vaporisation per mole
ΔS_vap = entropy of vaporisation
S_vap = Molar entropy of vapour
S_liq = Molar entropy of liquid
T_b = Boiling point of the liquid in Kelvin

3. Entropy as a state function: Let us examine entropy as a state function. Consider a cylinder containing a gas and is fitted with a frictionless and weightless piston which is in contact with a large heaf reservoir.

During isothermal and reversible expansion of the gas from volume V₁ to V₂, the substance will absorb heat, q at temperature T
∴ Change in entropy of the system ΔS_sys = \(\frac{q_{\mathrm{rev}}}{\mathrm{T}}\)

Since an equivalent amount of heat will be lost by the reservoir, change in entropy of reservoir will be ΔS_res = \(\frac{-q_{\mathrm{rev}}}{\mathrm{T}}\)

Total change in entropy
ΔS_t = ΔS_sys + ΔS_res
= \(\frac{q_{\mathrm{rev}}}{\mathrm{T}}+\left(\frac{-q_{\mathrm{rev}}}{\mathrm{T}}\right)\) = 0

If we compress the gas isothermally from a volume V₂ to V₁, heat given by the system is, – q_rev
ΔS_sys = \(\frac{q_{\mathrm{rev}}}{\mathrm{T}}\)
and ΔS_res = \(\frac{-q_{\mathrm{rev}}}{\mathrm{T}}\)

Total change in entropy
ΔS₁ + ΔS₂ = 0

At the end of the cycle, the entropy of the system is the same as it had initially. Therefore, entropy is a state function.

The total entropy change (AStota]) system and surrounding of a spontaneous process is given by
ΔS_total = ΔS_system + ΔS_surr > 0 When a system is in equilibrium, the entropy is maximum, and the change in the entropy, ΔS = 0.

We can say that entropy for a spontaneous process increases till it reaches maximum but at equilibrium change in entropy is zero. Since entropy is state property, we can calculate the change in – entropy of a reversible process by
ΔS_sys = \(\frac{q_{\text {sys rev}}}{\mathrm{T}}\)

We find that both for reversible and irreversible expansion for an ideal gas, under isothermal conditions, ΔU = 0, but ΔStotal i.e., ΔS_sys + ΔS_surr is not zero for an irreversible process. Thus ΔH does not discriminate between reversible and irreversible processes, whereas ΔS does.

→ Gibbs energy and spontaneity: We have seen that for a system, it is the total entropy change, ΔStotal which decides the spontaneity of the process But most of the chemical reactions fall into the category of either closed systems or open systems. Therefore, for most of the chemical reactions, there are changes in both enthalpy and entropy. It is clear that neither decrease in enthalpy nor an increase in entropy alone can determine the direction of spontaneous change for these systems.

For this purpose, we define a new thermodynamic function the Gibbs energy or Gibbs function, G, as
G = H -TS
Gibbs function, G is an extensive property and a state function.

The change in Gibbs energy for the system, ΔGsys can be written as
ΔG_sys. = ΔH_sys – TΔS_sys – S_sysΔT

At constant temperature,
ΔT = 0 ΔG_sys = ΔH_sys – TΔS

Usually the subscript ‘system’ is dropped and we simply write this equation as
ΔG = ΔH – TΔS

Thus, Gibbs energy change = enthalpy change – temperature x entropy change, and is referred to as the Gibbs equation, one of the most important equations in chemistry. Here, we have considered both terms together for spontaneity: energy (in terms of ΔH) and entropy (ΔS, a measure of disorder) Dimensionally if we analyzed, we find that ΔG has units of energy because, both ΔH and the TΔS are energy terms (Since TΔS = (K) J/(K mol) = J/mol or J mol^-1).

Now let us consider how ΔG is related to reaction spontaneity.
We know, ΔS_total = ΔS_sys + ΔS_surr

If the system is in thermal equilibrium with the surroundings, then the temperature of the surrounding is the same as that of the system.

Therefore, an increase in entropy of the surrounding is equal to a decrease in the entropy of the system.
Therefore, entropy change of surroundings,
ΔS_sur = – \(\frac{\Delta \mathrm{H}_{\text {surr}}}{\mathrm{T}}\) = – \(\frac{\Delta \mathrm{H}_{\text {sys}}}{\mathrm{T}}\)

ΔS_sur = ΔS_sys + (- \(\frac{\Delta \mathrm{H}_{\text {sys}}}{\mathrm{T}}\))

Rearranging the above equation:
TΔS_total = TΔS_sys -ΔH_sys

For spontaneous process,
ΔS_total > 0, so TΔS_sys – ΔH_sys > 0

Using the equation, the above equation can be written as
ΔG = ΔH – TΔS < 0

ΔH_sys is the enthalpy change of a reaction, TΔS_sys is the energy that is not available to do useful work. So ΔG is the net energy available to do useful work and is thus a measure of the ‘free energy. For this reason, it is also known as the free energy of the reaction.

ΔG gives a criterion of spontaneity at constant pressure and temperature.

• If ΔG is negative (< 0), the process is spontaneous.
• If ΔG is positive (> 0), the process is nonspontaneous.

Note. If a reaction has a positive enthalpy change and positive entropy change, it can be spontaneous when TΔS is large enough to outweigh ΔH.

This can happen in two ways;
(a) The positive entropy change of the system can be ‘small’ in which case T must be large.
(b) The positive entropy change of the system can be ‘large’, in which case T may be small. The former is one of the reasons why reactions are often carried out at high temperatures.

→ Gibbs Energy Change and Equilibrium: “We have seen how a knowledge of the sign and magnitude of the free energy change of a chemical reaction allows:

1. Prediction of the spontaneity of the chemical reaction.
2. Prediction of*the useful work that could be extracted from it.

So far we have considered free energy changes in irreversible reactions. Let us now examine the free energy changes in reversible reactions.

‘Reversible’ under strict thermodynamic sense is a special way of carrying out a process such that the system is at all times in perfect equilibrium with its surroundings. When applied to a chemical reaction, the term ‘reversible’ indicates that a given reaction can proceed in either direction simultaneously, so that dynamic directions should proceed with a decrease in free energy, which seems impossible.

Therefore, we say that at equilibrium total free energy of the system is a minimum. If it is not, the system would spontaneously change to the configuration of lower free energy.

So, the criterion for equilibrium
A + B ⇌ C + D; is
Δ_rG = 0

Free energy is an extensive property and its value for a given substance will depend on the concentration of the substance. So we have to set its value at an equilibrium concentration as we set temperature and pressure. If the concentrations vary from the equilibrium values, Gibbs energy will also change, we can write
Δ_rG = Δ_rG° + RT ln Q … (1)

Here, ΔrG° is the difference in standard Gibbs energies of formation of the products and reactants, both in their standard, states. Similarly, Δ_rG is the Gibbs energy change at a definite, fixed composition of the reaction mixture. Q is the reaction quotient. R- Gas constant = 8.314 JK^-1 mol^-1. If the species are gases, these concentrations are expressed in partial pressure, and the reaction quotient will be Q_p and if species are in solution, the reaction quotient will be expressed in terms of molar concentration as Q.

At equilibrium, Q = K called equilibrium constant, and Δ_rG = 0,
0 = Δ_rG° + RT ln K
or
Δ_rG° = – 2.303 RT log K

We also know that
Δ_rG° = Δ_rH° – TΔ_rS° = – RT ln K …(2)

For strongly endothermic reactions, the value of ArH° may be large and positive. In such a case, the value of K will be much smaller than 1 and the reaction is unlikely to form much product. In the case of exothermic reactions, Δ_rH° is large and negative, and Δ_rG° is likely to be a large and negative tool. In such cases, K will be much larger than 1. We may expect strongly exothermic reactions to having a large K and hence can go to near completion. Δ_rG° also depends upon

Δ_rS°, if the changes in the entropy of reaction are also taken into account, the value of K or extent of chemical reaction will also be affected, depending upon whether Δ_rS° is positive or negative.

Using equation (2)

1. It is possible to obtain an estimate of ΔG° from the measurement of ΔH° and ΔS° and then calculate K at any temperature for economic yields of the products.
2. If K is measured directly in the laboratory, the value of ΔG° at any other temperature can be calculated.

Effect of Temperature on Spontaneity of Reactions

The terms low temperature and high temperature are relative. For a particular reaction, the high temperature could even mean room temperature.

→ System: That part of the universe in which observations are made surroundings. The remaining part of the universe with which the system can exchange both matter and energy is called the surroundings.

There are three types of system:

1. Open system
2. Closed system
3. Isolated system.

→ The state of the system: When the properties of a system like its pressure, volume, temperature, and composition are specified quantitatively, it is said to have a state.

→ State functions or state variables: The average measurable properties like temperature (T), volume (V), pressure (P), and amount (n) of the system are called state functions because their values depend only on the state of the system and not on how it is reached.

Internal Energy: The total energy stored in a system is called its internal energy (U). It may be chemical, electrical, mechanical, or any other type of energy.
(a) Work: It is a form of energy. Work is said to be done if the point of application of a force moves through a certain distance.

Adiabatic process: It is a process in which there is no transfer of heat between the system and surroundings.

Adiabatic system: The system that does not allow the exchange of heat between itself and its surroundings through its boundary is called an adiabatic system.

(b) Heat: The change of internal energy of a system by transfer of heat from the surroundings to the system or vice-versa without the expenditure of work is called heat (q).

Note:

1. q is positive when heat is transferred from the surroundings to the system.
2. Similarly, w is positive when work is done on the system, w is negative when work is done by the system.

First Law of Thermodynamics:
The energy of an isolated system is constant.
or
Energy can neither be created nor destroyed.

Mathematically, the first law of thermodynamics is
ΔU = q + w
where ΔU is the change of internal energy.

q is the heat absorbed/evolved and iv is the work done
Work done on gas

= – p_exΔV, where p_ex is the external pressure and ΔV is the change in volume from initial volume V_i to final volume V_f.

→ Reversible process: A process is reversible if it is brought about in such a way that it could at any moment, be reversed by an infinitesimal chance. It proceeds infinitely slowly by a series of 1 equilibrium states such that the system and surroundings are always in equilibrium with each other.

→ Irreversible process: A process or change which is brought, about so rapidly that the system does not get a chance to attain equilibrium. It cannot be retraced at each step of the change.

Under reversible conditions

where a pint is the internal pressure of the system. Since dp × dV is very small, we can write

Writing as (the pressure of the gas inside) p and for n moles of an ideal gas
pV = nRT
p = \(\frac{nRT}{V}\)

∴ Under isothermal conditions (i.e., at constant temperature)

If the gas expands in a vacuum (p_ext = 0), it is called the free expansion of the gas.
w_rev = 0

No work is said to be done during free expansion of a gas whether the process is reversible or irreversible.
Isothermal and free expansion of an ideal gas T = Constant (Isothermal expansion). In vacuum p_ext = 0
∴ w = 0, q = 0
∴ from ΔU = q + w
ΔU = 0
1. For isothermal irreversible change
q = – w = P_ext(V_f – V_i)

2. For isothermal reversible change
q = – w = nRT ln \(\frac{\mathrm{V}_{f}}{\mathrm{~V}_{i}}\)
= 2.303 nRT log \(\frac{\mathrm{V}_{f}}{\mathrm{~V}_{i}}\)

3. For adiabatic change, q = 0
ΔU = w ad

Enthalpy: Sum of internal energy U and the product of pressure and volume (PV) is called enthalpy (H)
Mathematically H = U + pV

Enthalpy change (ΔH) can be written as
ΔH = ΔU + pΔV ………..(1)
ΔH = q_p

Thus enthalpy change is the heat absorbed by the system at constant pressure.

• ΔH is negative for exothermic reactions
• ΔH is positive for endothermic reactions

At constant volume
ΔH = ΔU – q_v

Thus internal energy change is the heat change at constant volume.
pΔV = Δn_g RT
where Δng refers to the number of moles of gaseous products minus the number of moles of gaseous reactants.
Substituting the value of pΔV from (2) to (1)
ΔH = ΔU + Δn_gRT

→ Extensive Property: An extensive property is one whose value depends on the quantity or size of matter present in the system. For example, mass, volume, internal energy, enthalpy/heat capacity, etc. are extensive properties.

→ Intensive Property: The property which does not depend on the quantity or size of matter present in the system is called Intensive property. For example, temperature, density, pressure is intensive properties.

Heat Capacity (C): It is defined as the quantity of heat energy required to raise the temperature of a substance by 1 K (or 1°C).
q = C × ΔT
where ΔT is the. temperature change
At constant volume the heat capacity C is denoted by C_v
At constant pressure the heat capacity C is denoted by C_p

The relationship between C_p and C_v for an ideal gas
q_v = C_vΔT = ΔU …….(3)
q_p = C_pΔT = ΔH ………(4)

For 1 mole of an ideal gas
ΔH = ΔU + Δ(pV)
= ΔU + Δ(RT) [As pV = RT for 1 mole]
or
ΔH = ΔU + RΔT …….(5)

For equations (3) and (4), equation (5) becomes
C_pΔT = C_vΔT + RΔT
C_p = C_v + R
C_p – C_v = R

Measurement of ΔU and ΔH is done with the help of Calorimeter: Calorimetry is an experimental technique to measure energy changes associated with chemical or physical processes.

The instrument used for it is called a calorimeter.

Enthalpy change Δ_rH of a reaction.
Δ_rH = (Sum of enthalpies of products) – (sum of enthalpies of reactants)

Here symbol X(sigma) is used for summation and ai and b{ are the stoichiometric coefficients of the products and reactants respectively in the balanced chemical equation.

Standard enthalpy of reactions: the standard enthalpy of reaction is the enthalpy change for a reaction where all the participating substances are in their standard states.

The standard state of a substance at a specified temperature is its pure form at 1 bar.
Standard enthalpy change is denoted by ΔH°.

Standard enthalpy of fusion or Molar enthalpy of fusion Δ_fusH°: The enthalpy change that accompanies melting of one mole of a solid substance in the standard state is called standard enthalpy of fusion or molar enthalpy of fusion Δ_fus H°.
H₂O(S) → H₂O(l); Δ_fusH° = 6.00 kJ mol^-1

→ Standard or Molar enthalpy of Vaporization (Δ_vap H°): It is the amount of heat required to vaporize one mole of a liquid at constant. temperature and under standard pressure (1 bar).
H₂O(l) → H₂O(g); Δ_vapH° = 40.79 kJ mol^-1

→ Standard enthalpy of sublimation (Δ_subH°): It is the enthalpy change when one mole of a solid substance sublimes at a constant temperature and under standard pressure (1 bar).

Solid CO₂ or dry ice sublimes at 195 K with Δ_subH° = 25.2 kJ mol^-1

→ Standard enthalpy of formation: The standard enthalpy change for the formation of one mole of a compound from its elements in their most stable states of aggregation (also known as reference states) is called Standard Molar Enthalpy of Formation (Δ_rH°).
H₂(g) + \(\frac{1}{2}\)O₂(g) → H₂O(l); Δ_fH° = – 285,8 kJ mol^-1

→ Thermochemical equation: A balanced chemical equation together with the value of its Δ_rH is called a thermochemical equation.
C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l); Δ_rH° = – 1367 kJ mol^-1

Hess’s Law of Constant Heat
If a reaction takes place in several steps, then its standard reaction enthalpy is the sum of the standard enthalpies of the intermediate reactions into whiter the overall reaction may be divided at the same temperature.
Or
The enthalpy change for a reaction is the same whether it occurs in one step or through a series of steps at the same, temperature.
Standard enthalpy of combustion: It is defined as the enthalpy change when one mole of a substance in its standard state is burnt completely in excess of oxygen or air. (Symbol: Δ_CH°).

→ Enthalpy of atomization (Δ_aH°): It is the enthalpy change on breaking one mole of bonds completely to obtain atoms in the gas phase.
CH₄(g) → C(g) + 4H(g); Δ_aH° = 1665 kJ mol^-1

Bond Enthalpy (Δbond H°): Two different terms are used
1. Bond dissociation enthalpy: It is the change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken to form products in the gas phase.
Cl₂(g) → 2Cl(g); Δ_Cl-Cl H° = 242 kJ mol^-1

It is for diatomic molecules like H₂, Cl₂, O₂ in gas phase.
For polyatomic molecules the term used is the mean bond enthalpy. Consider the reaction:
CH₄(g) → C(g) + 4H(g); Δ_aH° = 1665 kJ mol^-1,
Δ_C-H° = (Δ_aH°) = \(\frac{1}{4}\)(1665 kJ mol^-1)
= 416 kJ mol^-1

Thus the mean C-H bond enthalpy in methane is 416 kJ mol^-1.
→ Enthalpy of Solution (Δ_Sol H°): It is the enthalpy change when one mole of it dissolves in a specified amount of solvent.

→ Enthalpy of Solution at infinite dilution: It is the enthalpy change observed on dissolving the substance in an infinite amount of the solvent when the interactions between the ions (or solute molecules) are negligible.

→ Lattice Enthalpy: The lattice enthalpy of an ionic compound is the enthalpy change that occurs when one mole of an ionic compound dissociates into its ions in a gaseous state.
Na⁺Cl^–(g) → Na+(g) + Cb(g); Δ_latticeH° = + 788 kJ mol^-1

→ Bom-Haber Cycle: It is used to determine the lattice enthalpies. Let us calculate the lattice enthalpy of Na+Cl-(s) by the following steps given below:

1. Na(s) → Na(g); Δ_subH° = 108.4 kJ mol^-1
2. Na(g) → Na⁺(g) + e^–(g); Δ_iH° = 496 kJ mol^-1
3. Cl₂(g) → Cl(g); Δ_bondH° – 121 kJ mol^-1
4. Cl(g) + e^–(g) → Cl^–(g); Δ H° = – 348.6 kJ mol^-1
5. Na+(g) + Cl^–(g) → Na⁺Cl^–(s); Δ_latticeH° = U

Heat of formation of.NaCl(s) is foirnd to be 411.2 kJ mol^-1
Na(s) + \(\frac{1}{2}\) Cl₂(g) → NaCl(s); Δ_FH° = – 411.2 kJ mol^-1

Applying Hess’s Law
-411.2 = 108.4 + 121 + 496 – 348.6 + U
or U = – 788 kJ

Thus, lattice energy for NaCl(s) has a large negative value. This explains why the compound NaCl(s) is highly stable.

→ Spontaneous Process: A spontaneous process is an irreversible process and may only be reversed by some external agency.

A decrease in enthalpy is one of the contributing factors for the spontaneity of a process, but it is not true for all cases.

→ Entropy: It is a measure of disorder or randomness of the system like internal energy (U) and enthalpy (H), entropy (S) is a state function.
Change in entropy ΔS = S₂ – S₁
= ΣS_Products – ΣS_reactants
ΔS is related with q and T for a reversible reaction as
ΔS = \(\frac{q_{rev}}{T}\)

For a spontaneous process: The total entropy change (ΔS_total) for the system and surroundings is > 0, i.e.,
ΔS_total = ΔS_system + ΔS_surf > 0

For a system in equilibrium, the entropy is maximum and change in entropy, ΔS = 0

→ Important Note: For both reversible and irreversible expansion of an ideal gas under isothermal conditions (T = constant) ΔU – 0, but ΔS_total, i.e., ΔS_sys+ ΔS_surr is 0 for a reversible process, but not zero for an irreversible process. Thus ΔU does not discriminate between reversible and irreversible processes whereas ΔS does.

→ Gibbs Energy (or Gibbs function): Gibbs energy (G) is an extensive property and it is a state function.
It is the thermodynamic quantity of a system, the decrease in whose value during a process is equal to the maximum possible useful work that can be obtained from the system.

Mathematically G = H – TS
where H is the heat content, T is the absolute temperature and S is the entropy of the system
ΔG = ΔH – TΔS
where ΔG is the change in Gibbs energy, ΔH is the enthalpy change, ΔS is the change in entropy. The above equation is called Gibbs-Helmholtz Equation.
-ΔG = W_{non-expansion} = W_useful
or
– ΔG = W_max

1. If ΔG is negative, the process will be spontaneous.
2. If ΔG = 0, the process is in. equilibrium.
3. If ΔG is positive, the direct process is non-spontaneous; the reverse process may be spontaneous.

Standard Gibbs energy change is related to the equilibrium constant by
Δ_rG° = – RT ln K, where K is equilibrium constant
or
Δ_rG° = – 2.303 RT log K.

By going through these CBSE Class 11 Chemistry Notes Chapter 5 States of Matter, students can recall all the concepts quickly.

States of Matter Notes Class 11 Chemistry Chapter 5

Most of the observable characteristics of chemical systems represent the bulk properties of matter. These are the properties associated with a collection of a large number of atoms, ions or molecules. For example, an individual molecule of a liquid does not boil, but the bulk boils.

Intermolecular Forces: Intermolecular forces are the forces of attraction and repulsion between interacting particles (atoms and molecules). This term does not include electrostatic forces that exist between the two oppositely charged ions and the forces that hold the atoms together in a molecule, i.e., a covalent bond.

Attractive intermolecular forces are known as van der Waals forces. They vary considerably in magnitude and include:

1. Dispersion or London forces
2. Dipole-dipole forces
3. Dipole-induced-dipole forces
4. A particularly strong type of dipole-dipole interaction is hydrogen bonding. But only a few elements (F, O, N) can participate in hydrogen bond formation.

Attractive forces between an ion and a dipole are known as ion-dipole forces and these are not van der Waals forces.

Dispersion Forces or London Forces: Atoms and non-polar molecules are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed. But a dipole may develop momentarily even in such atoms and molecules. Suppose we have two atoms A and B close to each other [Fig. (a)]. It may so happen that momentarily electronic charge distribution in one of the atom A becomes unsymmetrical [Fig. (b) and (c)]. This results in the development of instantaneous dipole on atom A, for a very short time. This distorts the electron results a dipole is included in the atom.

Dispersion forces of London forces between atoms

The temporary dipole of atom ‘A’ and ‘B’ attract each other. Similarly, temporary dipoles are induced in molecules also. This force of attraction was first calculated by the German physicist Fritz London. For this reason force of attraction between two temporary dipoles is known as the London forces. Another name for this force is dispersion force.

These forces are always attractive and interaction energy is inversely proportional to the sixth power of the distance between two interacting particles (i.e., \(\frac{1}{r^{6}}\) where r is the distance between two particles). These forces are important only at short distances (~ 500 pm) and the magnitude of the force depends on the polarisability of the particle. Dispersion forces are present among all particles.

Dipole-Dipole Forces:
Dipole-dipole forces act between molecules possessing permanent dipole. Ends of the dipole possess “partial charges” and these charges are shown by the Greek letter delta (5) Partial charges are always less than the unit charge (1.6 × 10^-19 C) because of the electron sharing effect.

The polar molecules interact with neighbouring molecules. This interaction is weak compared to ion-ion interaction because only partial charges are involved. The interaction energy decreases with the increase of distance between the dipoles. It is proportional to \(\frac{1}{r^{6}}\) where r is the distance between polar molecules. Besides dipole-dipole interaction, polar molecules can interact by London forces also. Thus the cumulative effect is that total of intermolecular forces in polar molecules increase. Fig. shows electron cloud distribution in the dipole of hydrogen chloride.

(a) Distribution of electron cloud in HCl a polar molecule
(b) Dipole-dipole interaction between two HCl molecules

Dipole-Induced Dipole Forces
This type of attractive forces operates between polar molecules having permanent dipole and molecules lacking permanent dipole. The permanent dipole of the polar molecule induces dipole of the electrically neutral molecule by deforming its electronic cloud.

Thus induced dipole is developed in the other molecule. In this case, also interaction energy is proportional to -4 where r is the distance between two molecules Induced dipole moment depends UpOn the dipole moment present in the permanent dipole and the polarizability of the electrically neutral molecule.

Dipole-induced dipole interaction between a permanent dipole and induced dipole

In this case, the cumulative effect of dispersion forces and dipole- included dipole interactions exists.

Hydrogen Bond
The bond formed between the hydrogen atom of one molecule with the more electronegative atom (like N, O or F) of another molecule is called a hydrogen bond Such a molecule is highly polar. The electronegative atom of the covalent bond possesses lone pair of electrons. When two such molecules containing these type of bonds come close to each other, the hydrogen of one molecule is attracted towards the electronegative atom of the other molecule This interaction is represented by a dotted line and is called a hydrogen bond.

The energy of hydrogen bond varies between 10 to 100 kJ mol^-1. This is a very significant amount of energy, therefore hydrogen bonds are a powerful force in determining the structure and properties of many compounds. The strength of the hydrogen bond is determined by the Coulomb interaction between the lone-pair electrons of the electronegative atom and the hydrogen atom. The following diagram shows the formation of the hydrogen bond.

The intermolecular forces discussed so far are all attractive. Molecules also exert repulsive forces on one another. When two molecules are brought into close contact with each other, the repulsion between the electrons and between the nuclei in the molecules come into play. The magnitude of the repulsion rises very rapidly as the distance separating the molecules decreases. This is the reason that liquids and solids are hard to compress. In these states molecules are already in close contact, therefore they resist further compression.

→ Thermal Energy: Thermal energy is the energy of a body arising from the motion of its atoms or molecules. It is directly proportional to the temperature of the substance. It is the measure of the average kinetic energy of the particles of the matter and is thus responsible for the movement of particles. The movement of particles is called thermal motion.

Intermolecular Forces Vs Thermal Interactions
Intermolecular forces tend to keep molecules together but the thermal energy of the molecules tend to keep them apart. Three states of matter are the result of a balance between intermolecular forces and the thermal energy of the molecules.

The predominance of thermal energy and the molecular interaction energy of a substance in three states is depicted as follows:

The Gaseous State: Only 11 elements in the periodic table exist in the form of gases under normal conditions:

Gases are classified by the following properties.

1. Gases are the most compressible out of all the 3 states of matter.
2. Gases exert pressure equally in all directions.
3. Gases have a much lower density than solids and liquids.
   
   elements that exist as gases
4. They assume the volume and shape of the container in which they are kept.
5. Gases mix evenly, and completely without any mechanical help.

Gas Law: Boyle’s Law (Pressure-Volume Relationship)
“Volume of a given mass of a gas is inversely proportional to its pressure provided the temperature is kept constant.”
V ∝ \(\frac{1}{P}\) (mass and temperature kept constant)
or
PV = constant = k
or
P₁V₁ = P₂V₂, where P₁V₁ are the initial pressure and volume of the gas. P₂ and V₂ are the final values.

The value of k depends upon the amount of the gas, temperature of the gas, its nature and the units in which P and V are expressed.
Graphical Representation of Boyle’s Law

(a) Graph of Pressure Vs Volume (b) Graph of pressure p Vs \(\frac{1}{V}\)

Each curve in (a) above corresponds to a different constant temperature and is known as an Isotherm

(c) Graph of PV against P

1. A plot of P versus V at constant temperature for a fixed mass of gas would be a rectangular hyperbola (a)
2. A plot of P versus \(\frac{1}{v}\) at constant temperature for a fixed mass of gas would be a straight line passing through the origin (b)
3. A plot of P versus PV at a constant temperature. In a fixed mass of a gas is a straight line parallel to the pressure axis:

Relationship between density and pressure

Charle’s Law (Temp. Volume Relationship): The law states, “At constant pressure, the volume of a given mass of a gas increases or decreases by \(\frac{1}{273.15}\) of its volume at 0°C and 1°C rise or fall in temperature.” Thus if the volume of the gas at 0°C and at t°C is V_o and V_t. Then

At this stage, we define a new scale of temperature such that f°C on the new scale is given by T = 273.15 + t and 0°C will be given by T_o = 273.15. This new temperature scale is called the Kelvin temperature scale or Absolute temperature scale. It can be represented as follows:

Relationship between Kelvin Scale and Celsius Scale

One degree Celsius (1 °C) is equal to one Kelvin (K) in magnitude (Note that degree sign is not used while writing the temperature in absolute temperature scale, i.e. Kelvin scale). Only the position of zero has been shifted. Kelvin scale of temperature is also called the Thermodynamic scale of temperature and is used in all scientific works.

Thus we add 273 (more precisely 273.15) to Celsius temperature to obtain temperature at Kelvin scale.

A graph of volume Vs temperature at Kelvin scale will look like as shown.

If we write Tf = 273.15 + t and T_o = 273.15 in the equation we obtain the relationship
V_t = V_o(\(\frac{\mathrm{T}_{t}}{\mathrm{~T}_{0}}\))
⇒ \(\frac{V_{t}}{V_{0}}=\frac{T_{t}}{T_{0}}\)

Thus we can write a general equation as follows:
\(\frac{V_{2}}{V_{1}}=\frac{T_{2}}{T_{1}}\)
or
\(\frac{V}{T}\) = constant = k

Each line of V vs T graph is called Isobar

At constant pressure for a given mass of a gas
V ∝ T
or
V = kT (where k is a constant, the value of k depends upon nature and amount of the gas)
or
\(\frac{V}{T}\) = constant

If V₁ = initial volume of the gas
T₁ = initial temperature (absolute)
V₂ = Final volume
T₂ = Final temperature .
Then \(\frac{V_{1}}{T_{1}}=\frac{V_{2}}{T_{2}}\) (mass and pressure kept constant)

Relationship between density and temperature
V ∝ T
or
\(\frac{1}{d}\) ∝ T
or
dT = constant
or
d₁T₁ = d₂T₂

Gay-Lussac’s Law (Pressure-Temperature Relationship): It states that at constant volume, the pressure of a given mass of a gas varies directly with temperature on the Kelvin scale
Mathematically P ∝ T
It is also called Amonton’s Law
or
\(\frac{P}{T}\) = constant

Pressure Vs Temperature (K) graph (Isometrics) of a gas

Avogadro’s Law (Volume-Amount Relationship)
It states, “Equal volumes of all gases under similar conditions of temperature and pressure contain an equal number of molecules.”

Mathematically v ∝ n where n = no. of moles of the gas if T and P are constant.
n = \(\frac{m}{M}\); where n = number of moles of gas
m = mass of the gas;
M = Molar mass of the gas

where d is the density of the gas.

Thus the density of a gas is directly proportional to its formula mass.

Avogadro’s number or Avogadro’s constant is the number of particles present in 1 mole of a substance. This number is 6.023 × 10²³.

One mole of each gas will have volume = 22.4 L at STP.

A gas that follows Boyle’s Law, Charle’s and Avogadro’s law strictly is called an ideal gas. Such gas is a hypothetical gas. It is assumed that intermolecular forces are not present between the molecules of an ideal gas. Real gases follow these laws only under certain specific conditions when forces of interaction are practically negligible. In all other situations, these deviate from ideal behaviour.

Ideal Gas Equation

1. At constant T and n, V ∝ \(\frac{1}{P}\) Boyle’s Law
2. At constant P and n, V ∝ T Charle’s Law
3. At constant T and P, V ∝ n Avogadro’s Law
   Thus
   
   On rearrangingPV= nRT … (1)

R is called Gas Constant. It is the same for all gases. Therefore, it is called Universal Gas Constant, equation (1) above is called Ideal Gas Equation.

The value of R for 1 Mole of an ideal gas can be calculated ai 273.15 K and 1 atm (1.013 × 10⁵ Pascal) [Most of the real gases behave like ideal gases under these conditions) as follows.
R = \(\frac{\left(1.0325 \times 10^{5} \mathrm{~Pa}\right) \times\left(22.41383 \times 10^{-3} \mathrm{~m}^{3} / \mathrm{mol}\right)}{298.15 \mathrm{~K}}\)
= 8.31441 Pa m³ k^-1 mol^-1
= 8.3144 JK^-1 mol^-1

[For one mole V = 22.413832 for all gases at STP] For all practical purposes approximate value of R, i.e., 8.314 JK^-1 mol^-1 can be used.

In the equation PV = nRT

• P and T are Intensive properties. Intensive properties are independent of the ‘quantity’ or ‘bulk’ of the substance.
• n and V are Extensive properties. Extensive properties are dependent upon the ‘quantity’ or ‘bulk’ of the substance.

If temperature, volume and pressure of a fixed amount of the gas vary from T₁, V₁ and P₁ to T₂, V₂ and P₂.
Then \(\frac{\mathrm{P}_{1} \mathrm{~V}_{1}}{\mathrm{~T}_{1}}\) = nR and \(\frac{\mathrm{P}_{2} \mathrm{~V}_{2}}{\mathrm{~T}_{2}}\) = nR
or
\(\frac{P_{1} V_{1}}{T_{1}}=\frac{P_{2} V_{2}}{T_{2}}\)

It is called General Gas Equation or Combined gas law. Density and Molar Mass of a Gaseous substance

Dalton’s Law of Partial Pressures: The law states, “Total pressure exerted by a mixture of non-reacting gases is equal to the sum of partial pressures exerted by the individual gases-all measurement being made at the same temperature.” Thus at constant temperature and volume:

P_Total = p₁ + p₂ + p₃ + ………; P_Total = Total pressure.
p₁, p₂, …. are the partial pressures of the different non-reacting gases. ,

Gases are generally collected over water. The pressure of dry gas can be calculated by subtracting the vapour pressure of water from the moist gas which contains water vapours also. The pressure exerted by saturated water vapour is called aqueous tension.
P_{dry gas} = P_Total – Aqueous tension

Partial pressure in terms of mole fraction: Suppose three gases, enclosed in the volume V, exert partial pressure p₁, p₂ and p₃ respectively. Then
p₁ = \(\frac{n_{1} \mathrm{RT}}{\mathrm{V}}\)
p₂ = \(\frac{n_{2} \mathrm{RT}}{\mathrm{V}}\)
p₃ = \(\frac{n_{3} \mathrm{RT}}{\mathrm{V}}\)
where n1, n2 and n3 are number of moles of these gases. Thus expression for total pressure will be
P_Total = P₁ + P₂ + P₃
= n₁\(\frac{\mathrm{RT}}{\mathrm{V}}\) + n₂\(\frac{\mathrm{RT}}{\mathrm{V}}\) + n₃\(\frac{\mathrm{RT}}{\mathrm{V}}\)

= (n₁ + n₂ + n₃)\(\frac{\mathrm{RT}}{\mathrm{V}}\)
On dividing p₁ by P_Total we get
\(\frac{p_{1}}{P_{\text {Total }}}=\left(\frac{n}{n_{1}+n_{2}+n_{3}}\right) \frac{\mathrm{RTV}}{\mathrm{RTV}}=\frac{n_{1}}{n_{1}+n_{2}+n_{3}}\) = x₁

x₁ is called mole fraction of first gas
Thus, p₁ = x₁ P_Total.

Similarly, for the other two gases, we can write
P₂ = x₂P_Total and P₃ = x₃P_Total

Thus a general equation can be written as
p_i = x_i P_Total

where p_ix_i, and P_Totalare partial pressure and mole fraction of with gas and total pressure of the gas mixture respectively. If the total pressure of the gases is known, the equation can be used to find out the pressure exerted by the individual gas.

Molar Volume of the gas under different conditions
(a) Standard temperature and pressure (STP) conditions are 0° C or 273,15 K and one atmospheric pressure. Under these conditions, 1 mole of the gas occupies a volume of 22.413996 L = 22.4 L or 22400 mL.

When STP conditions are taken as 0°C and 1 bar pressure [as 1 bar < 1 atom and 1 bar = 0.987 atm) molar volume is slightly higher and = 22.71098 L mol^-1 = 22.7 L mol^-1 or 22700 mL.

→ Standard Ambient Temperature and Pressure (SATP): Conditions are also used in some scientific works. SATP conditions mean 298.15 K and 1 bar (i.e. exactly 10s Pa). At SATP (1 bar and 298.15 K), the molar volume of an ideal gas is 22.789 mol^-1 or 22800 mL.

→ Graham’s Law of Effusion/Diffusion: It states, “Under similar conditions of temperature and pressure, rate of diffusion of a gas is inversely proportional to the square root of the density of a gas.”

Mathematically R_d ∝ \(\sqrt{\frac{1}{d}}\)
R_d = Rate of diffusion, d = density of the gas
Rate of diffusion R_d = R\(\sqrt{\frac{1}{d}}\)

If two gases A and B diffuse under similar conditions of temperature and pressure rate r_A and r_B respectively and their densities are dA and dB respectively then
r_A = k\(\sqrt{\frac{1}{d_{\mathrm{A}}}}\)
r_B = k\(\sqrt{\frac{1}{d_{\mathrm{B}}}}\)

On dividing r_A and r_B we obtain
\(\frac{r_{\mathrm{A}}}{r_{\mathrm{B}}}=\sqrt{\frac{d_{\mathrm{B}}}{d_{\mathrm{A}}}}=\sqrt{\frac{d_{\mathrm{B}} \mathrm{V}}{d_{\mathrm{A}} \mathrm{V}}}=\sqrt{\frac{\mathrm{M}_{\mathrm{B}}}{\mathrm{M}_{\mathrm{A}}}}\)
where V is the volume of gas A and gas B which undergoes diffusion and MA and MB are molecular masses of gas A and gas B respectively. Rate of diffusion should be proportional to the average speed of the molecules of a gas. If average speed of molecules of gas A and gas B are n_A and n_B then
\(\frac{u_{\mathrm{A}}}{u_{\mathrm{B}}}=\sqrt{\frac{\mathrm{M}_{\mathrm{B}}}{\mathrm{M}_{\mathrm{A}}}}\)

Above equation can be used to determine the molar mass of gases.

→ Student’s Note: Though Graham’s Law of Diffusion is not included in the CBSE syllabus, but students are advised to go through it for various competitive examinations.

Kinetic Molecular Theory of Gases Postulates. Assumption of Kinetic Theory of Gases
1. Every gas is made up of a large number of extremely small particles called molecules. All the molecules of a particular gas are identical in mass and size and differ, in these from gas to gas.

2. The molecules of a gas are separated from each other by large distances so that the actual volume of the molecules is negligible as compared to the total volume of the gas.

3. The distances of separation between the molecules are so large that the forces of attraction or repulsion between them are negligible.

4. The force of gravitation on the molecules is also supposed to be negligible.

5. The molecules are supposed to be moving continuously in different directions with different velocities. Hence they keep on colliding with one another (called molecular collisions) as well as on the walls of the containing vessel.

6. The pressure exerted on the walls of the containing vessel is due to the bombardment of the molecules on the walls of the containing vessel.

7. The molecules are supposed to be perfectly elastic hard spheres so that no energy is wasted when the molecules collide with one another or with the walls of the vessel. The energy may, however, be transferred from some molecules to the other on collision.

8. Since the molecules are moving with different velocities, they possess different kinetic energies. However, the average kinetic energy of the molecules of a gas is directly proportional to the absolute temperature of the gas.

The behaviour of Real Gases
Deviation from ideal gas behaviour: A gas that obeys the ideal gas equation. (PV = nRT) at all temperatures and pressure is called Ideal Gas or Perfect Gas. Actually, none of the known gases obeys the ideal gas equation under all conditions of temperature and pressure. Such gases which do not obey the gas equation are called Real Gases.

However, most of the gases follow the ideal gas equation at low pressures and high temperatures. Appreciable deviation from the ideal gas behaviour is observed at low temperature and high pressure.

The temperature at which a real gas obeys the ideal gas law over an appreciable range of pressure is called Boyle Temperature or Boyle Point. Boyle point of a gas depends upon its nature.

A gas deviates from ideal behaviour when the product of the observed pressure and volume (P × V) is lower or higher than that expected from the ideal gas equation. Thus for a gas showing deviations from ideal behaviour (PV) observed < nRT
or
(PV)observed > nRT

Thus, to describe the behaviour of real gases, the ideal gas equation should be modified. This is conveniently done by inviting the ideal gas equation ‘in’ the form PV = Z(nRT).
or
Z = \(\frac{P V}{n R T}\)

The plot of PV Vs Pressure for real and ideal gas

The plot of Pressure Vs Volume for real gas and ideal gas

Where Z can have a value of one, less than one or more than one. The factor Z is called Compressibility Factor.
(a) When Z = 1, then PV = nRT, i.e., the gas shows an ideal gas behaviour. For an ideal gas Z = 1 under all conditions of temperature and pressure.
(b) When Z < 1, the observed PV value is less than the value for an ideal gas. Thus the gas shows a negative deviation when Z < 1.
(c) When Z > 1, the observed PV value is higher than that for an ideal gas. So when Z > 1, the gas shows positive deviations.

Equation of state for Real Gases (Van der Waal’s equation)
The van der Waal’s equation for one mole of a real gas is written
(P + \(\frac{a}{v^{2}}\))(v-b) = RT
If n moles of a real gas is taken

The equation becomes
(P + \(\frac{a n^{2}}{v^{2}}\))(v – b) = nRT
Here ‘a’ and ‘b’ are called van der Waals parameters or constants which vary from gas to gas. Their value depends upon the gas nature.

The cause of the deviation is actually due to two faulty assumptions of the kinetic theory
They are

• There is no force of attraction and repulsion between the molecules of a gas.
• The volume of the molecules of a gas is negligibly small as compared to the empty space between them.

Derivation of van der Waal’s equation
1. Correction for volume: Suppose the volume occupied by the gas molecules is v. When the molecules are moving, their effective volume is four times the actual volume i.e., 4v. Let us called it b i.e. b = 4v (called excluded volume of co-volume). Thus the free volume available to the gas molecules for movement i.e.

Corrected volume = (V – b) for one mole
= (V – rib) for n moles

2. Correction for pressure. A molecule (A) lying within the vessel is attracted equally by other molecules on all sides but a molecule near the wall (B) is attracted (pulled back) by the molecules inside (Fig.) Hence it exerts less pressure. In other words, the observed pressure is less than the ideal pressure. Hence

Corrected pressure = P + p

Backward pull on molecule B by other molecules

Evidently, the correction term p is proportional to density of the gas near the wall and the density of the gas inside i.e.
p ∝ (density)²
or
p ∝ d²
But d ∝ \(\frac{1}{V}\) for one mole
or
d ∝ \(\frac{1}{\mathrm{~V}^{2}}\) for 1 mole
or
p = \(\frac{a}{\mathrm{~V}^{2}}\)
or
p ∝ \(\frac{n^{2}}{\mathrm{~V}^{2}}\) for n moles
or
p = \(\frac{a n^{2}}{V^{2}}\)

∴ Corrected pressure = (P + \(\frac{a}{\mathrm{~V}^{2}}\)) for 1 mole
= (P + \(\frac{a n^{2}}{V^{2}}\)) for n moles
where a is constant depending upon the nature of the gas.

Substituting the correct values of volume and pressure in the ideal gas equation we get
(P + \(\frac{a}{\mathrm{~V}^{2}}\))(V – b) = RT for 1 mole
or
(P + \(\frac{a n^{2}}{V^{2}}\))(V – nb) = nRT for n moles

Significance of van der Waal’s constants
1. van der Waal’s constant ‘a’: Its value is a measure of the magnitude of the attractive forces among the molecules of the gas. The greater the value of ‘a’, the larger is the intermolecular forces of attraction.

2. van der Waal’s constant ‘b’: Its value is a measure of the effective size of the gas molecules. Its value is equal to four times the actual volume of the gas molecules. It is called excluded volume or co-volume.

Units of van der Waal’s constant
1. Units of ‘a’. As p = \(\frac{a n^{2}}{V^{2}}\)
∴ a = \(\frac{p \times V^{2}}{n^{2}}\)
= atm L² mol² or bar dm⁶ mol^-2

2. Units of ‘b’. As volume correction
v = nb,
∴ b = \(\frac{v}{n}\) = K mol^-1dm³mol^-1.

Explanation of the behaviour of Real gases by van der Waal’s equation
1. At very low pressure, V is very large. Hence the correction term a.V² is so small that it can also be neglected in comparison to V. Thus van der Waal’s equation reduces to the form PV = RT. This explains why at very low pressures, the real gases behave like ideal gases.

2, At moderate pressures, V decreases. Hence a/V² increases and cannot be neglected. However, V is still large enough in comparison to ‘b’ so that ’b’ can be neglected. Thus van der Waal’s equation becomes

Thus compressibility factor is less than 1. As pressure is increased at a constant temperature, V decreases so that the factor n/RTV increases. This explains why initially a dip in the plot of Z versus P is observed.

3. AT high pressures, V is so small that ‘b’ cannot be neglected in comparison to V. The factor a/V² is no doubt large but as P is very high, a/V² can be neglected in comparison to P. Thus van der Waal’s equation reduces to the form
P(V -b) = RT
or
PV = RT + Pb ⇒ \(\frac{\mathrm{PV}}{\mathrm{RT}}\) = 1 + \(\frac{\mathrm{P} b}{\mathrm{RT}}\)
or
Z = 1 + \(\frac{\mathrm{P} b}{\mathrm{RT}}\)

Thus compressibility factor is greater than 1. As P is increased (at constant T), the factor Pb/RT increases. This explains why after minima in the curves, the compressibility fact of increases continuously with pressure.

4. At high temperatures, V is very large (at a given pressure) so that both the correction factors (a/V² and b) become negligible as in case (i). Hence at high temperature, real gases behave like an ideal gas.

→ Explanation of the exceptional behaviour of hydrogen and helium may be seen that for H₂ and He, the compressibility factor Z is always greater than l and increases with the increase of pressure: This is because H₂ and He being very small molecules, the intermolecular forces of attraction in them are negligible i.e., V is very very small so that a/V² is negligible
P(V -b) = RT
or
PV = RT + Pb
or
\(\frac{\mathrm{PV}}{\mathrm{RT}}\) = 1 + \(\frac{\mathrm{PB}}{\mathrm{RT}}\)

Thus, \(\frac{\mathrm{PV}}{\mathrm{RT}}\) i.e., Z > 1 and increases with increase in the value of P at constant.

(a) Z vs P for different gases

(b) Z as P to N₂ gas at different temperatures

Plots in Fig. (b) show that as the temperature increases, the minimum in the curve shifts upwards. Ultimately, a temperature is reached at which the value of Z remains close to 1 over an appreciable range of pressure. For example, in the case of N₂ at 323 K, the value of Z remains close to 1 up to nearly 100 atmospheres.

For a real gas Z = \(\frac{p V_{\text {real }}}{n \mathrm{RT}}\) ……….(1)
If the gas shows ideal behaviour then
V ideal = \(\frac{nRT}{p}\)
On putting the value of \(\frac{nRT}{p}\) in the above equation (1)
Z = \(\frac{\mathrm{V}_{\text {real }}}{\mathrm{V}_{\text {ideal }}}\)

Thus compressibility factor may also be defined as the ratio of actual molar volume of a gas to the molar volume of it if it were ideal at that temperature and pressure.

→ Liquefaction of Gases and Critical Point: At high pressures and lower temperatures deviation from ideal behaviour for gases is observed. At high pressures, molecules of the gas come closer and attractive forces start operating. As the temperature is lowered further, the attractive forces draw the molecules together to form a ‘liquid’. This temperature is called liquefaction temperature. It depends upon the nature of the gas and its pressure.

The critical temperature for gas is that temperature above which it is not possible to liquefy it, however large is the pressure applied on it. CO₂ gas remains a gas even when a pressure of 73 atmospheres is applied to it. However at this pressure when it is cooled to 30.98°C, it starts liquifying. So for CO₂ 30.98°C is its critical temperature. The volume of a gas at critical temperature is called its critical volume and pressure (73 atm. for CO₂) is critical pressure.

Critical pressure is the pressure required to liquefy the gas at the critical temperature. All three of them are collectively called critical constants of the gas and are represented by Tc, Pc and Vc. For example, critical constants of CO₂ are:
T_c = 31.1°C, P_c = 73.9 atm, V_c = 95.6 cm³ mol^-1

Isotherms of CO₂ (From Andrew’s experiment)

At the lowest temperature employed i.e., 13.1°C, at low pressure. CO₂ exists as a gas, as shown at point A. As the pressure is increased, the volume of the gas starts. Hence volume decreases rapidly along with BC because the liquid has much less volume than gas. At point C, liquefaction is complete. Now the increase in pressure has very little effect upon volume because liquids are very little compressible.

Hence a step curve CD is obtained. As the temperature is increased, the horizontal portion becomes smaller and smaller and at 31.1°C it is reduced at a point, P. This means that above 31.1°C, the gas cannot be liquefied at all, however, high pressure may be applied. Thus 31.1°C, is the critical temperature. The corresponding pressure to liquefy the gas at the critical temperature is its critical pressure, P (i.e., 73.9 atm). The volume occupied by 1 mole of the gas under these conditions is its critical volume, V. (i.e., 95.6 ml).

→ Liquid State: In liquids, intermolecular forces of attraction are much large than in gases. Unlike gases, liquids have a definite volume, but no definite shape. The molecules in a liquid are in constant random motion. The average kinetic energy of the molecules in a liquid is proportional to absolute temperature.

→ Vapour Pressure: Vapour pressure of a liquid at any temperature is the pressure exerted by the vapour present above the liquid in equilibrium with the liquid. The magnitude of vapour pressure depends upon the nature of, liquid and temperature. At equilibrium between the liquid and the vapour phase, the vapour pressure becomes steady or constant and at this stage, it is called Equilibrium Vapour Pressure or saturated vapour pressure.

The Boiling Point of a liquid is the temperature at which its vapour pressure becomes equal to external pressure. At 1 atm pressure boiling temperature is called the normal boiling point. If pressure is one bar, then the boiling point is called standard boiling point standard boiling point of a liquid is slightly lower than the normal boiling point (because 1 bar < 1 atm)

The normal boiling point of water is 100°C (373 K) and its standard boiling point is 99.6°C (373.6 K).

The Vapour pressure of a liquid increases with an increase in temperature, with more molecules of the liquid going into the vapour phase, the liquid becomes less dense and the density of the vapour phase increases. When the density of the liquid and vapours becomes the same the clear boundary between the liquid and the vapour disappears. This temperature is called the critical temperature of the liquid.

→ Surface Tension: The surface tension of a liquid is defined as the tangential force acting along the surface of a liquid at right angles along one unit length drew on the surface of the liquid.

It is the work done to increase the free surface area of any liquid by one unit at constant temperature and pressure.

The S.I. unit of surface tension is Nm^-1. The C.G.S. unit of surface tension is dyne^-1 cm^-1. The lowest energy state of the liquid will be when its surface area is minimum. Spherical shape satisfies this condition. Hence raindrops or mercury drops are spherical in shape. Liquids tend to rise in the capillary due to surface tension. Surface tension decreases as the temperature is raised. It has dimensions of kg s^-2. In terms of surface energy per unit area, its dimensions are Jm^-2. ,

→ Viscosity: Viscosity is a measure of resistance that arise due to the internal friction between layers of fluid as they slip past one another. It is the internal resistance to the flow of liquid. The liquids which flow rapidly have low internal resistance and hence are said to be less viscous, i.e., their viscosity is low. On the other hand, the liquids which flow slowly have high internal resistance and hence are said to be more viscous, i.e., their viscosity is high The viscosity of honey or glycerol is higher than, say, water.

The coefficient of viscosity is defined as the force applied per unit area which will maintain a unit relative velocity between the two layers of a liquid at unit distance from each other. It is represented by q. Liquids having low q values are called mobile while those having high q value are termed as viscous.

If f is the force required to maintain the flow of layers, then
f ∝ A [A is the area of contact]
f ∝ \(\frac{d u}{d Z}\)
or
f = ηA\(\frac{d u}{d Z}\); η is called coefficient of viscosity.

S.I. unit of viscosity coefficient is 1 newton second per square metre [Nsm^-2]

In the CGS system, its unit is poise
1 poise = 1 gm cm^-1 s^-1 = 10^-1 kg m^-1 s^-1

Viscosity decreases with an increase in temperature. Hydrogen bonding and van der Waal’s forces are strong enough to cause High viscosity.

→ van der Waals forces: Attractive intermolecular forces between interacting particles (atoms and molecules) are called van der Waals forces.

→ Dispersion Force/London Force: Atoms and non-polar molecules are electrically symmetrical and hence no dipole moment. But a dipole may develop momentarily even in such atoms and molecules. This force of attraction between two temporary dipoles is called the London force or dispersion force.

→ Dipole-Dipole Forces: Dipole-dipole forces act between the molecules possessing permanent dipole. The polar molecules interact with neighbouring molecules as in the case of HCl molecules.

→ Dipole-induced Dipole Forces: This type of attractive forces operate between the polar molecules having permanent dipole and the molecules lacking permanent dipole. The permanent dipole of the polar molecule induces dipole on the electrically neutral molecule by deforming its electronic cloud Thus an induced dipole is developed in the other molecule.

→ Hydrogen Bond: It is a special case of dipole-dipole interaction It is present in molecules having H-atom attached to a small-sized highly electronegative atom like F, O, N. The highly electronegative attracts H-atom of the neighbouring molecule through electrostatic force of attraction It is a weak bond and present in molecules like I IF, FRO, R-OH etc.

→ Thermal Energy: It is the energy of a body arising due to the motion of its atoms or molecules.

→ Boyle’s Law: At constant temperature, the volume of a given mass of a gas is inversely proportional to its pressure.

→ Charles’ Law: Pressure remaining constant, the volume of a given mass of a gas is directly proportional to its absolute temperature.

→ Isotherm: Each curve obtained on plotting different values of, pressure against volume at constant temperature is called Isotherm.

→ Isobar: Each line of the volume vs temperature graph at constant j pressure is called Isobar.

→ Kelvin temperature/Absolute temperature scale. It is obtained by adding 273.15 to the temperature on the Celsius scale.
Kelvin-Tem. = T = t°C + 273.15. It is also called the Thermodynamic ’ scale.

→ Absolute zero: The volume of all gases becomes zero at – 273.15°C. The lowest hypothetical or imaginary temperature at which gases cease to exist is called absolute zero.

→ Gay-Lussac’s Law: At constant volume, the pressure of a given j mass of a gas varies directly with the pressure.

→ Isochore: Each line of the graph obtained by plotting pressure vs Kelvin temperature at constant volume is called.Isochore.

→ Avogadro’s Law: Equal volumes of all gases under the similar condition of temperature and pressure contain an equal number of molecules.

→ Avogadro’s constant: The number of molecules in one mole of a gas is – 6.022 × 10²³ and the number 6.022 × 10²³ is called Avogadro’s constant.

→ Ideal Gas Equation.
PV = nRT
where P = Pressure,
V = volume,
R = Universal gas constant
n = no. of moles,
T = absolute temperature

For 1 mole of an ideal gas PV = RT.

Combined gas law \(\frac{\mathrm{P}_{1} \mathrm{~V}_{1}}{\mathrm{~T}_{1}}=\frac{\mathrm{P}_{2} \mathrm{~V}_{2}}{\mathrm{~T}_{2}}\)

→ Dalton’s Law of Partial Pressures: The total pressure exerted by the mixture of non-reactive gases is equal to the sum of partial pressures exerted by individual gases under similar conditions of temperature and volume.

→ Aqueous Tension: Pressure exerted by the saturated water vapour is called Aqueous Tension. ,

→ Compressibility Factor: The deviation of a gas from ideal behaviour can be measured in terms of compressibility factor Z which is the ratio of product pV and nRT.
Z = \(\frac{pV}{nRT}\)

→ Boyle Point or Boyle Temperature: The temperature at which a real gas obeys the ideal gas law is called Boyle point/temperature.

→ Critical Temperature: The temperature above which a gas cannot be liquified, however large may be the pressure applied on it.
Or
It is a temperature below which a gas can be liquified with the application of pressure is called its critical temperature (T_c) Critical Pressure: The pressure required to liquefy a gas at its critical temperature is called its critical pressure (P_c).

→ Critical Volume: The volume possessed by gas at its critical temperature is called its critical volume (V_c).

→ Boiling Point: It is a temperature at which the vapour pressure of a liquid becomes equal to external pressure.

→ Normal Boiling Point: At 1 atmospheric pressure, the boiling temperature of a liquid is called its normal boiling point.

→ Standard Boiling Point: If the pressure is 1 bar, the boiling temperature of the liquid is called standard boiling point, the standard boiling point of a liquid is slightly lower than the normal boiling point (as 1 bar is slightly < 1 atm).

→ Surface Tension: It is defined as the force acting per unit length perpendicular to the line drawn on the surface of the liquid.

→ Unit: Surface Tension (γ-Gamma) in S.I. scale = Nm^-1 V Its dimensions are kgs^-2.

→ Viscosity: It is a measure of the resistance of the flow of a liquid which arises due to internal friction between layers of it as they slip past one another while liquid flows.

Important Formulae
→ Boyle’s Law: At T constant
V ∝ \(\frac{1}{P}\)
or
PV = constant
If P₁, V₁ are the initial pressure, the volume of a gas at constant temperature and P₂, V₂ are its final pressure and volume respectively.
then P₁V₁ = P₂V₂

→ Charles’ Law: At constant pressure
V ∝ T, where T is the absolute temperature
or
\(\frac{V}{T}\) = Constant
or
\(\frac{\mathrm{V}_{1}}{\mathrm{~T}_{1}}=\frac{\mathrm{V}_{2}}{\mathrm{~T}_{2}}\)
or
\(\frac{\mathrm{V}_{1}}{\mathrm{~V}_{2}}=\frac{\mathrm{T}_{1}}{\mathrm{~T}_{2}}\)
Gay Lassac’s Law : P ∝ T if volume is kept constant
or
\(\frac{P}{T}\) = constant

→ Ideal Gas Equation:
PV = nRT

→ Combined Gas Law:
\(\frac{P_{1} V_{1}}{T_{1}}=\frac{P_{2} V_{2}}{T_{2}}\)

→ Coefficient of Viscosity : If the velocity of the layer at a distance dZ is changed by a value du, then velocity gradient is given by the amount \(\frac{d u}{d Z}\). The force required lo maintain the flow of layers is proportional to the area of contact of layers and velocity gradient,
i.e., f ∝ A [A is the area of contact]
f ∝ \(\frac{d u}{d Z}\)
∴ f ∝ A\(\frac{d u}{d Z}\)
or
f = ηA\(\frac{d u}{d Z}\)

where η is proportionality constant called coefficient of viscosity.
If A = 1 \(\frac{d u}{d Z}\) = 1
then f = η
Therefore, the Coefficient of viscosity is defined as the force when the velocity gradient is unity and the area of contact is unity.

SI unit of viscosity coefficient is 1 newton sec m^-2
= Ns m^-2 = pascal second
1 Pa s = 1 kg m^-1 s^-1

In CGS system coefficient of viscosity is poise
1 poise = 1 g cm^-1 s^-1 = 10^-1 kg m^-1 s^-1.

By going through these CBSE Class 11 Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure, students can recall all the concepts quickly.

Chemical Bonding and Molecular Structure Notes Class 11 Chemistry Chapter 4

Chemical Bond: The attractive force which holds various constituents (atoms, ions, etc.) together in different chemical species is called.a chemical bond.

There are various theories to explain the formation of a chemical bond. They are

1. Kossel-Lewis approach.
2. Valence-Shell Electron-Pair Repulsion (VSEPR) Theory
3. Valence Bond (VB) Theory
4. Molecular Orbital Theory.

Kossel-Lewis Approach to Chemical Bonding: All noble gases [except He] have 8 electrons in their valence shells. They are chemically inactive. In the case of all other elements, there are less than 3 electrons in the valence shells of their atoms, and hence they are chemically reactive.

“The atoms of different elements combine with each other in order to complete their respective octets (i.e8 electrons in their outermost shell) or duplet (i.e., the outermost shell having 2 electrons in case of H, Li, and Be) to attain stable inert gas configuration”.

Lewis Symbols: Writing symbols of elements with valence shell electrons represented by dots are called Lewis Symbols.

→ Significance: The number1 of valence electrons helps to calculate the common or Group Valence of the element. The group valence of the elements is generally either equal to the number of dots in Lewis Symbols or 8 minus the number of dots or valence electrons.

Facts for chemical bonding (Kossel)

1. The highly electronegative halogens and highly electropositive alkali metals are separated by noble gases in the periodic table.
2. The formation of a negative ion and a positive ion is due to the gain or loss of electrons by atoms.
3. The negative and positive ions thus formed attain the nearest noble gas configuration. The noble gases (with the exception of He which has a duplet of electrons) have a particularly stable outer shell configuration of eight (Octet) electrons, ns²np⁶.
4. The negative and positive ions are stabilized by electrostatic attraction.

Electrovalent or Ionic Bond: When a bond is formed by complete transference of electrons from one atom to another so as to complete their octets or duplets the bond is called an ionic or electrovalent bond. For example

→ Formation of Calcium chloride

→ Formation of MgO

The no. of electrons gained or lost during the formation of an ionic bond is called the electro valency of the elements.

Factors favoring the formation of an ionic bond.

1. Atom going to lose electrons must have low ionization enthalpy so that it can lose electrons readily. Elements of I and II. groups prefer to form, ionic bonds because they have low values of Ionisation enthalpy.
2. Elements that accept electrons must have a High Negative Value of electron gain enthalpy so that they can retain the electrons. Halogens have high values of E.G. enthalpy.
3. Lattice enthalpy should be high. This, in turn, depends upon.
   (a) Charge on the ion: Greater the charge on the ion, greater the lattice energy.
   (b) Size of the ions: Smaller the size of the ion, the greater the lattice energy.

→ Octet Rule: According to the electronic theory of chemical bonding as developed by Kossel and Lewis, atoms can combine either by transfer of valence electron’s from one atom to another (gain or loss) or by sharing valence electrons in order to have an octet in their valence shells. This is known as the Octet rule.

→ Covalent Bond: A covalent bond is formed as a result of the mutual sharing of electrons between the two atoms to complete their octet or duplet. No. of electrons contributed by each atom is called covalency.

If one pair of electrons (one electron from each atom) is mutually shared between the two atoms, a single covalent bond is formed.

If two pairs of electrons (two electrons from each atom) are mutually shared between the two atoms, a double covalent bond is formed.

If three pairs of electrons (three electrons from each atom) are mutually shared between the two atoms.


Lewis Representation of Simple Molecules (the Lewis Structures)
The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pairs of electrons and the octet rule. While such a picture may not explain the bonding and behavior of a molecule completely, it does help in understanding the formation and properties of the molecules to a large extent. Writing of Lewis dot structure m molecules is, therefore, very useful.

The Lewis-dot-structures can be written by adopting the following steps:
1. The total number of electrons required for writing the
structures is obtained by adding the valence electrons of the combining atoms. For example, in the CH₄ molecule, there are eight valence electrons available .for bonding (4 from carbon and 4 from the four hydrogen atoms).

2. For anions, each negative charge would mean the addition of one electron. For cations, each positive charge would result in the subtraction of one electron from the total number of valence electrons. For example, for the CO₃^2- ion, the 2- charges indicate that there are two additional electrons than those provided by the neutral atoms. For NH₄⁺ ion, 1+ charge indicates the loss of the electrons from the group of neutral atoms.

3. Knowing the chemical symbols of the combining atoms and
having knowledge of the skeletal structure of the compound (known or guessed intelligently), it is easy to distribute the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.

4. In general the least electronegative atom occupies the central position in the molecule/ion. For example in the NF^–₃ and CO₃^2-, nitrogen and carbon are the central atoms whereas fluorine and oxygen occupy the terminal positions.

5. After accounting for the shared pairs of electrons for single bonds, the remaining electron pairs are either utilized for multiple bonding or remain as the lone pairs. The basic requirement being that each bonded atom gets an octet of electrons.

→ Lewis representation of a few molecules/ions are given in the table below:

Table: The Lewis Representation of Some Molecules

Each H atom attains the configuration of helium (a duplet of electrons)

→ Formal Charge: Lewis dot structures, generally, do not represent the actual shape of the molecules.

The formal charge of an atom in a polyatomic molecule or ion is the difference between the number of valence electrons in that atom in an isolated state or free State and the no. of electrons assigned to that atom in the Lewis structure.

Formal charge (F.C) on an atom in a Lewis structure = [total no. of valence electrons in the free atom] – [total no. of non-bonding or lone pair of electrons] – \(\frac{1}{2}\) [total no. of bonding (shared) electrons]

The molecule of O₃ can be considered

The atoms have been numbered 1. 2 and 3

The formal charge on

1. the central atom O marked 1
   = 6 – 2 – \(\frac{1}{2}\) [6] = + 1
2. the end O atom marked 2 = 6 – 4 – \(\frac{1}{2}\)[4] = 0
3. the end O atom marked 3 = 6 – 6 – \(\frac{1}{2}\)(2) = – 1

Hence we represent O₃ along with formal charges as follows:

Generally, the lowest energy structure is the one with the smallest formal charges on the atoms. The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally by neighboring atoms.

Limitations of the Octet Rule: It is mainly useful to understand the structures of organic compounds and applies to the elements of the second period of the periodic table. But it is not universal.

1. The incomplete octet of the central atom: In some compounds, the number of electrons surrounding the central atom is less than eight. This especially is the case with elements having less than four valence electrons. Examples are LiCl, BeH₂ and BCl₃

Li, Be and B have 1, 2, and 3 valence electrons only. Other such compounds are AlCl₃ and BF₃.

2. Odd electrons molecules: In molecules with an odd number of electrons like nitric oxide, NO, and nitrogen oxide, NO₂, the octet rule is not satisfied for all the atoms.

3. The expanded octet Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals are available for bonding. In a number of compounds of these elements, there are more than eight valence electrons around the central atom. This is termed the expanded octet. The octet rule does not apply in such cases.

Some of the examples of such compounds and PF₃, SF₆, H₂SO_4, and a number of coordination compounds.

4. It is clear that the octet rule is based upon the chemical inertness of noble gases. However, some noble gases (Xenon and Krypton) also combine with oxygen and fluorine to form a number of compounds like XeF₂, XeF₄, SeOF₂, KIF_2, etc.

5. The theory does not account for the shape of molecules.

6. It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.

Ionic or Electrovalent Bond: An ionic (or electrovalent) bond is formed by a complete transfer of one or more electrons from the atom of metal to that of a non-metal.

As a result of this electron transfer, the following changes occur in the reacting atoms.
(a) Both the atoms acquire stable noble gas configuration.
(b) The atom that loses its electrons becomes positively charged ions called a cation, whereas the atoms which gain these electrons becomes negatively charged ion called an anion.
(c) The two oppositely charged ions, i.e., ‘he cation and the anion, are then held together by the Coulomb forces of attraction to form an ionic bond.

Thus, an ionic bond that may be defined as the Coulomb force of attraction which holds the oppositely charged ions together is called an ionic bond.
1. Lattice Enthalpy (Lattice energy): The Lattice enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. For example, the lattice enthalpy of NaCl is 788 kJ mol^-1. This means that 788 kJ of energy is required to separate to an infinite distance 1 mol of solid NaCl into 1 mol of Na⁺ (g) and mole of Cl^– (g).

For two atoms to form an ionic bond, the following factors are necessary

1. Low ionization enthalpy. The atom going to lose electrons must have a low value of ionization enthalpy.
2. High electron gain enthalpy. The atom going to accept electrons must have a high negative value of electron gain enthalpy.
3. High lattice enthalpy. The lattice energy of the compound should be high. The greater the charge on the ion and the smaller the size, the greater is the lattice enthalpy.

An ionic bond is formed if lattice energy and electron affinity took together are greater than ionization energy. The larger the negative value of lattice energy greater is than the stability of the ionic compound. The stability of an ionic compound is provided by its enthalpy of lattice formation and not simply by achieving an octet of electrons around the ionic species in a gaseous state.

Example of Ionic Bond

One Na⁺ is surrounded by C Cl^– ions and vice-versa Rock Salt Structure
1. Na⁺ Cl^– (Ionic Solid)


2. MgO

3. CaF2

Properties of Ionic solids:

1. Ionic compounds are solids, whereas covalent compounds can be solids, liquids, or gases.
2. Ionic compounds do not contain molecules but are made up of ions, whereas covalent compounds are molecular.
3. Ionic compounds have high melting points and boiling points due to strong electrostatic forces of attraction between ions. Covalent compounds, on the other hand, have low melting points and boiling points.
4. Ionic compounds are generally soluble in water or any polar
   solvent whereas covalent compounds are insoluble in water, but soluble in non-polar solvents.
5. Ionic compounds take part in ionic reactions whereas covalent compounds take part in molecular reactions.
6. Ionic bonds are non-directional in nature, whereas covalent bonds have directional characteristics.

Bond Parameters
1. Bond length: Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. Bond lengths are measured by spectroscopic, X-ray diffraction, and electron diffraction techniques. Each atom of the bonded pair contributes to the bond length (Fig.) In the case of a covalent bond, the contribution from each atom is called the covalent radius of that atom.

The bond length in a covalent molecule AB
R = r_A + r_B (R is the bond length and r_A and r_B are the covalent radii of atoms A and B respectively)

The covalent radius is measured approximately as the radius of an atom’s core which is in contact with the core of an adjacent atom in a bonded situation. The covalent radius is half of the distance between two similar atoms joined by a covalent bond in the same molecule.

The van der Waals radius represents the overall size of the atom which includes its valence shell in a nonbonded situation. Further, the van der Waals radius is half of the distance between two similar atoms in separate molecules in a solid. Covalent and van der Waals radii of chlorine and depicted in Fig.

Covalent and van der Waals radii in a chlorine molecule. The inner circles correspond to the size of the chlorine atom (r_vdw and r₀ are van der Waals and covalent radii respectively)

2. Bond angle: It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complexion. For example, the bond angle in water can be represented as under:

3. Bond enthalpy: It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.

The unit of bond enthalpy is kJ mol^-1 For example H-H bond enthalpy in hydrogen molecule is 435.8 kJ mol^-1
H₂ (g) → H (g) + H (g); ΔH = 435.8 kj mol^-1

4. Bond order: In the Lewis description of the covalent bond the bond order is given by the number of bonds between the two atoms in a molecule. The bond order in H₂ is one, in O₂ it is two and in N₂ it is three.

Isoelectronic molecules and ions have identical bond orders. For example F₂ and O₂^2- have bond order = 1.N₂, CO and NO⁺ have bond order 3.

The stabilities of molecules can be understood by the statement. With the increase in bond order, bond enthalpy increases and bond length decreases.

5. Resonance structures: According to the concept of resonance, whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, positions of nuclei, bonding, and non-bonding pairs of electrons are taken as the canonical structures of the resonance hybrid which describes the molecule accurately. Thus for Cl the two structures shown below constitute the canonical structures and their hybrid, (III, below) represents the structure of O, more accurately. Resonance is represented by a double-headed arrow.

Resonance in O₃ molecule. I and II represent the two canonical forms III is the resonance hybrid.

Some of the other examples of resonance structures are provided by the carbonate ion and the carbon dioxide molecule.
1. Carbonate ion (CO₃)

2. Carbon-dioxide

6. Electronegativity: The electronegativity is the ability of an atom of an element in a molecule to attract electrons towards itself in a shared pair of electrons.

Table: Difference between electronegativity and electron affinity (electron gain enthalpy)

Electronegativity	Electron affinity (electron gain enthalpy)
1. It is the tendency of an atom to attract shared pair of electrons.	1. It is the tendency of an atom to attract outside electrons.
2. It is the property of a bonded atom.	2. It is the property of an isolated atom.
3. The elements with symmetrical configurations have specific electronegativities.	3. The elements with symmetrical configurations have almost zero electron affinities.
4. It has no units.	4. It has units of kJ mol-1 of V atom-1

7. Polarity of Bonds: An ionic bond is formed due to the complete transfer of electrons from one atom to another. A covalent bond is formed due to the mutual sharing of electrons. A covalent bond between two atoms of different elements is called a polar covalent bond. A polar bond is partly covalent bond and partly ionic. The percentage of ionicity in a covalent bond is called the percentage ionic character in that bond. The ionic character in a bond is expressed in terms of bond dipole moment (μ).

The dipole moment of a bond depends upon the difference in the electronegativity of the two atoms held together by the chemical bond.

The dipole moment of two and opposite charges is given by the product of the charge and the distance separating them. Thus
Dipole moment, μ = charge (q) × Distance separation (r)
= q × r

In the HCl molecule, the chlorine is more electronegative. Therefore, the bonding electrons lie closer to the chlorine atom. As a result, H- atom develops a slight positive (+q) and Cl^– atom a slight negative (-q) charge, thereby generating a dipole in the HC1 molecule.

For non-polar bonds, there is no charge separation. Therefore, q = 0. As a result, the dipole moment of a non-polar bond is zero.

Dipole moment and the ionic character in bonds: Ionic character in a covalent bond depends upon the difference between the electronegativity of two atoms. The greater the electronegativity difference, the greater is the charge separation, and therefore, the greater is the ionic character in the bond.

The valence shell electron pair repulsion (VSEPR) theory: The main postulates of the VSEPR theory are:

1. Pairs of electrons in the valence shell of a central atom repel each other.
2. These pairs of electrons tend to occupy positions in space that minimize repulsion and thus maximize distances between them.
3. The valence shell is taken as a sphere with the electron pairs localizing on the spherical surface at a maximum distance away from one another.
4. Multiple bonds are treated as if it is a single electron pair and the two or the three electron pairs of multiple bonds are treated as single super pair.
5. Where two or more resonance structures can depict a molecule, the VSEPR model is applicable to any such structure.

Predicting the shape of molecules on the basis of VSEPR theory: According to the VSEPR theory/the geometry of a molecule is determined by the number of electron pairs around the central atom. So to predict the geometry of the molecules, no. of electron pairs (both shared and one pair) should be known.

The repulsive interactions of electron pairs decrease in the order l – l – p > l – pb – p > b – pb – p repulsions.

The geometry of Molecules in which the Central Atom has No Lone Pair of Electrons


Shapes (Geometries) of some simple molecules/ions with central Ions having One or More Lone Pairs of Electron (E)


Shapes of Molecules containing Bond Pair and Lone Pair



Valence Bond Theory: The main postulates of the valence bond theory are:

1. A covalent bond is formed due to the overlap of the outermost half-filled orbitals of the combining atoms. The strength of the bond is determined by the extent of overlap.
2. The two half-filled orbitals involved in the covalent bond formation should contain electrons with opposite spins. The two electrons then move under the influence of both nuclei.
3. The completely-filled orbitals (orbitals containing two paired electrons) do not take part in the bond formation.
4. An s-orbital does not show any preference for direction. The non-spherical orbitals such as p-and d-orbitals tend to form bonds in the direction of the maximum overlap, i.e., along the orbitals axis.
5. Between the two orbitals of the same energy; the orbital which is non-spherical (e.g., p- and d-orbitals forms stronger bonds than the orbitals which are spherically symmetrical, e.g., s-orbital.
6. The valence of an element is equal to the number of half-filled orbitals present in it.

In the valence bond model, the stability of a molecule is explained in terms of the following type of interactions:
(a) electron-nuclei attraction interactions, i.e., the electrons of one atom are attracted by the nucleus of the other atom also.
(b) electron-electron repulsive interactions, i.e., electrons of one atom are repelled by the electrons of the other atom.
(c) nucleus-nucleus repulsive interactions, i.e., the nucleus of one atom is repelled by the nucleus of the other atom.

Valence Bond Description of Hydrogen molecule: A hydrogen molecule is a stable molecule. That is why one mole of H₂ molecules requires energy equal to 433 kJ to dissociate into hydrogen atoms, viz.
H₂ (g) + 433 kJ → 2H (g)

The formation of a hydrogen molecule from two hydrogen atoms may be considered to take place through the following steps:
Step 1. Consider two hydrogen atoms H_A and H_B at large separation from each other, so that there is no interaction between them. Since, there is no interaction between the two H atoms, hence the total energy is equal to the sum of the energies of the two H atoms. (Fig. a)

(a) No interactions at large distances

Step 2. Now, let the two H atoms approach each other. When the two atoms come closer new attractive and repulsive forces begin to operate. The electrons of one atom are attracted by the nucleus of the other atom. At this stage, both the electrons are attracted by both nuclei. Since, attractive interactions lead to a decrease of energy, hence as the two

(b) Interactions start as atoms come closer.
atoms approach each other the energy of the system starts decreasing. (Fig. b, e, d).

Step 3. When two hydrogen atoms, H_A and H_B come still closer the electron-electron (e^–_A — e^–_B) and the nucleus- nucleus repulsive interactions start operating. The repulsive interaction tends to increase energy. As long as the attractive interactions are stronger than the repulsive interactions, the energy of the system continues to decrease.

At a certain distance between the two atoms, the attractive and repulsive interaction balance each other, and the energy of the system attains a minimum value. At the state, two H atoms have a fixed distance between them and form a stable H₂ molecule. The internuclear separation when the energy of the system is minimum is called bond length The H—H bond length in H₂(g) is 75 pm (0.74 A).

New forces of attraction and repulsion between two H-atoms approaching each other

Step 4. Now, if the two H- atoms in an H₂ molecule are formed to come closer than the equilibrium internuclear separation (bond length: 74 pm), the repulsive forces start predominating and as a result, the energy of the system increases very sharply (Fig. above)

(d) Potential energy diagram showing the variation of energy with the internuclear distance between two H-atoms

2. Overlap of atomic orbitals: Various types of atomic orbitals overlap leading to the formation of covalent bonds are:
(a) s-s overlap: In this type of overlap, half-filled s-orbital of the two combining atoms overlap each other.

(b) s-p overlap: Here a half-filled s-orbital of one atom overlaps with one of the p-orbitals having only one electron in it.

(c) p-p overlap along the orbital axis: This is called head-on, end-on, or end-to-end linear overlap. Here, the overlap of the two half-filled p-orbitals takes place along the line joining the two nuclei.

(d) p-p sideways overlap: This is also called lateral overlap, in this type of overlap, two p-orbitals overlap each other along a line perpendicular to the internuclear axis, i.e., the two overlapping p- orbitals are parallel to each other.

The p-p sideways overlap

Hybridization of Atomic Orbitals:
Hybridization: It may be defined as the phenomenon of mixing of orbitals of nearly the same energy so as to redistribute their energies and to give rise to new orbitals of equivalent energies and shapes.

The new orbitals that are formed are called hybridized or hybrid orbitals. As a general rule, the number of hybrid orbitals produced from hybridization is equal to the number of orbitals that are mixed together. For example, when an s orbital is mixed with a p orbital two sp hybrid orbitals are produced. The two orbitals lie in a straight line and make an angle of 180°. Examples of molecules with sp hybrid orbitals are BeF₂ and CH ≡ CH (acetylene).

The ground state configuration of the boron is 1 s²2s²2p¹_x. If one of the 2s electrons is promoted to an orbital of a little higher energy, say 2p¹_y orbital, then the configuration 1s² 2s¹ 2p¹ 2p¹_y would result. The three equivalent orbitals formed from one 2s and two 2p orbitals are called sp² hybrid orbitals, which are coplanar and directed at an angle of 120° to each other as shown. The orbitals form a stronger bond than the sp orbitals.

sp³ Hybridization: The ground state configuration of carbon = 6 = 1s², 2s², 2p_x¹ 2p¹_y


4sp³ hybrid orbitals of C

The four sp³ hybrid orbitals are directed towards the four corners of the tetrahedron. The bond angles are 109°28′.

Example of sp³ hybridization: Shape of CH₄
The shape of CH₄ Molecule

4sp³ hybrid orbitais of C

Sigma (σ) and Pi (π) bonds
1. Sigma (σ) Bond: It is formed as a result of s-s, s-p, or p-p overlap axially or on the internuclear axis. A sigma bond îs strong as a result of overlapping on the internuclear axis.
1. Combination of s-orbitais

2. Combina Lion of s- and p- orbitais bond

3. End to end combination of two p-orbitals

2. Pi (π) Bond: When a bond is formed by the lateral or sideways overlapping of atomic orbitals, it is known as a pi (π) bond. A π bond is weaker than a bond as the overlapping is not effective and is made up of two half-electron-clouds, 1/2 above and 1/2 below the internuclear axis. It is always present in addition to a_o bond in molecules containing a double bond or a triple bond. Between two atoms forming a single covalent bond, it has to be a bond. A Pi (π) bond has no primary effect on the direction of the bond. It however shortens the internuclear distance.

The relative bond strength.has the order p-p < s-p < s-s

In addition to the sp hybridization in BeF2 sp² hybridization in BF3 and sp³ hybridization in CH4, we have hybridization involving d-orbitals also.

The elements of the third period contain d-orbitals also in addition to s and p orbitals. The 3d orbitals are comparable in energy to 3s and 3p orbitals.

Due to the availability of d orbitals, valencies of 5, 6, and 7 are also expected.
1. sp³d hybridization. This involves the mixing of one s three p and one d-orbitals. These orbitals hybridize to form five sp³d hybrid orbitals which are directed towards the corners of a regular trigonal bipyramidal geometry. For example, PF₅ has this geometry.

2. sp³d² hybridization. This involves the mixing of one-s orbital, three p, and two d-orbitals. These six orbitals hybridize to form, six sp³d² hybrid orbitals which are directed towards the corners of an octahedral geometry. For example, the SF₆ molecule has octahedral geometry.

3. sp³d³ hybridization. This involves the mixing of one s, three p, and three-d orbitals forming seven sp³d³ hybrid orbitals. These adopt pentagonal bipyramidal geometry. For example, IF₇ has this geometry.

4. dsp² hybridization. In this case, one d(^{d_x2_{– y}2}), one, and two p orbitals get hybridized to form four dsp2 hybrid orbitals. These hybrid orbitals have square planar geometry. For example [Ni(CN)₄]^2- ion.

Molecular Orbital Theory: Molecular orbital theory was developed by Hund and Milliken in 1932. The basic idea of molecular orbitals theory is that atomic orbitals of individual atoms combine to form molecular orbitals. The electrons in molecules are present in these molecular orbitals which are associated with several nuclei. These molecular orbitals are filled in the same way as the atomic orbitals in atoms are filled.

The molecular orbitals are formed by the combination of atomic orbitals of the bonded atoms.

The salient features of the Molecular orbital theory are:

1. The electrons in a molecule are present in the various molecular orbitals as the electrons of atoms are present in the various atomic orbitals.
2. The atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals.
3. While an electron in an atomic orbital is influenced by one nucleus, in a molecular orbital, it is influenced by two or more nuclei depending upon the number of atoms in the molecule. Therefore, we can call- a molecule orbital Polycentric.
4. The no. of MOs formed is equal to the no. of combining AOs. When two atomic orbitals combine, two MOs are formed. One is called bonding molecular orbital and the other is called an antibonding molecular orbital.
5. The bonding MO has lower energy and hence greater stability than the corresponding antibonding MO.
6. Just as the electron probability distribution around a nucleus in an atom is given by an atomic orbital the electron probability distribution around a group of nuclei in a molecule is given by a molecular orbital.
7. The MOs obey the Aufbau principle, Pauli Exclusion principle, and Hund’s rule like AOs.

Formation of Molecular Orbitals/Linear Combination of Atomic Orbitals (ICAO)
If φ_A and φ_B represent the wave functions of atomic orbitals A and B of hydrogen atoms forming hydrogen molecule, then
φ_MO = φ_A ± φ_B
Therefore the two molecular orbitals a and a* are formed as follows:
σ = φ_A + φ_B
σ* = φ_A – φ_B
where σ is called bonding molecular orbital and σ* is called an antibonding molecular orbital.

For example, s-orbitals of two hydrogen atoms combine to form two MOs — σ_1s and σ*_1s (bonding and antibonding MOs).

Similarly, 2s atomic orbitals may combine to form two MOs: σ2s and σ*2s.
If the 2 axes is assumed to be the internuclear axis, then 2p atomic orbitals combine to form sigma orbitals. The 2p_z atomic orbitals combine to form aMOs (σ2p_z, σ*2p_z) while 2p_x and 2p_y orbitals combine to form nMOs (π2p_x, π*2p_x, π2p_y, π*2p_y).

Energy Level Diagrams for Molecular Orbitals: 1s, 2s and 2p orbitals of two atoms combine to form bonding MOs: σ1s, σ2s, σ2p_z, π2p_x, π2p_y. and antibonding MOs:

σ*l s,’ σ*2s, σ*2p_z, π*2p_x and π*2p_y. The energy levels of these molecular orbitals have been determined experimentally by spectroscopic methods. It is observed that for diatomic molecules of the second period Li2 to Ne2 there are two types of energy levels of MOs. For molecules Li₂, Be₂, B₂, C_2, and N₂ the sequence of energies of MOs may be written as
σ1s, σ*ls, σ2s, σ*2s, π2p_x = π2p_y, σ2p_z, π*2p_x π*2p_y, σ*2p_z

On the other hand, for the molecule 02, F2 and Ne2.the sequence of energies of MOs are:
σ1s, σ1s, σ2s, σ*2s, σ2p_z, π2p_x = π2p_y, π*2p_x = π*2p_y, σ*2p_z.

The main difference between the two types of sequences is that for molecules O₂, F₂ and Ne₂, the σ2p_z MO is lower in energy than π2p_x and π2p_y MOs while in the case of molecules Li₂, Be₂, B₂, C_2, and N₂, σ2p_z MOs higher energy than π2p_x arid π2p_y MOs.

Bond Order: It is defined as the number of covalent bonds in a molecule. Bond order can .be calculated from the number of electrons in bonding and antibonding molecular orbitals as:

Bond order = \(\frac{\text {No. of electrons in bonding MOs-No. of electrons in anit-bonding MOs }}{2}\)
or
= \(\frac{\mathrm{N}_{b}-\mathrm{N}_{a}}{2}\)

The bond orders of 1, 2, or 3 correspond to the single, double, or triple bonds. But bond order may be fractional also in some cases.

Information Conveyed by Bond Order.

1. If the value of bond order is positive; it indicates a stable molecule and if the value of bond order is negative or zero, it means that the molecule is unstable.
2. The stability of a molecule is measured by its bond
   dissociation energy. But the bond dissociation energy is directly proportional to the bond order.
3. Bond order is inversely proportional to the bond length. The higher the bond order value, the smaller is the bond length.
4. If all electrons are doubly occupied in different molecular orbitals the molecule is diamagnetic. If there are one or more unpaired electrons, it is paramagnetic like O₂.

Conditions for the combination of Atomic Orbitals:

1. The combining atomic orbitals must have the same or nearly the same energy. This means that Is orbital can combine with another 1s, but not with 2s orbitals because the energy of 2s is appreciably higher than that of Is.
2. The combining atomic orbitals must have the same symmetry as the molecular axis.
3. The combining atomic orbitals must overlap to the maximum extent. In this type, sigma (σ) molecular orbitals are symmetrical around the bond axis while pi (π) molecular orbitals are not symmetrical. If the internuclear axis is taken to be in the Z-direction it can be seen that a linear combination of 2p_z-orbitals of two atoms also produces two sigma molecular orbitals σ2pz and σ*2p_z.

Molecular orbitals obtained from 2p_x and 2p_y orbitals are not symmetrical around the bond axis because of the positive lobes above and negative lobes below the molecular plane. Such molecular orbitals are labeled as π and π*. A π bonding MO has a large electron density above and below the internuclear axis. The σ* antibonding MO has a node between the nuclei.

Contoursmnd energies of bonding and antibonding molecular orbitals formed through combinations of (a) Is atomic orbitals; (b) 2p_z atomic orbitals and (c) 2p_x atomic orbitals

Molecular Orbital Electronic Configurations: Molecular orbital electronic configurations of some molecules/ions are given below:

→ Hydrogen Bond: The attractive force which binds the hydrogen atom of one molecule with the electronegative atom (F, O, N) of another molecule is known as a hydrogen bond.

A hydrogen bond is weaker than a covalent bond. The strength of a hydrogen bond ranges from 10 – 40 kJ mol^-1 while that of a normal covalent bond is of the order of 400 kJ mol^-1

Cause of formation of the hydrogen bond. When hydrogen is bonded to a strongly electronegative element, A (such as F, O, or N) the electron pair shared between the two atoms lies far away from the hydrogen atom. As a result, the hydrogen atom becomes highly electropositive with respect to the other atom, A.

Since the electrons are displaced towards A, it acquires partial charge (δ⁺). In other words, the bond H-A becomes polar and may be represented as H^δ+ – A^δ–. The electrostatic force of attraction between positively charged hydrogen- atom of one molecule and negatively charged atom of neighboring molecule results in the formation of the hydrogen bond. This may be represented as:

Conditions for Hydrogen Bonding: The conditions for hydrogen bonding are:

1. The molecule must contain a hydrogen atom attached to the highly electronegative atom (F, O, or N).
2. The size of the electronegative atom should be quite small. (HF), water (H₂O), ammonia (NH,) alcohols (ROH), carboxylic acids (RCOOH), amines (RNH₂), etc. Due to hydrogen bonding, some of the properties of these molecules are influenced.

→ Types of Hydrogen Bonds: Hydrogen bonding may be classified into two types.
(a) Intermolecular hydrogen bond: It is a hydrogen bond formed between two different molecules of the sai^e or different substances. This type of linking results in the association of molecules and increases the melting point, boiling point, viscosity, solubility, etc.

(b) Intramolecular hydrogen bond: Intramolecular hydrogen bond is formed between the hydrogen atom and the highly electronegative atom (F, O, or N) present in the same molecule. An intramolecular hydrogen bond results in the cyclization of the molecules and prevents their association. Consequently, the effect of intramolecular hydrogen bonds on physical properties is negligible.

For example, intramolecular hydrogen bonds are present in molecules such as o-nitro phenol, o-nitro benzoic acid, etc.
Ortho nitro phenol Ortho nitrobephobic acid

→ Chemical Bond: The force which holds the constituent atoms together with a molecule is called a chemical bond.

→ Electrovalent or Ionic bond: A bond formed by the complete transference of electrons from one atom to another to acquire the stable electronic configuration of the nearest gas is called an Ionic bond.

→ Electrovalency: The number of electrons so transferred (gained or lost) is called the electrovalency of the element.

→ Covalent Bond: When two atoms share the electrons mutually in order to complete their octet or duplet, the bond so formed is called a covalent bond.

→ Coordinate Bond: When the electron pair shared between the two atoms is donated by one of the two atoms, it is called a coordinate bond.

→ Sigma (σ) bond: When a bond is formed between the two atoms by the overlap of their atomic orbitals along the internuclear axis (end to end of head-on overlap) the bond so formed is called sigma (σ) bond.
It is due to s-s, s-p, p-p overlap.

→ Pi (π) Bond: It is formed by the side-wise or lateral overlap of only p-orbitals in a direction perpendicular to the internuclear axis.

1. Pi (π) bond is weaker than a sigma (σ) bond.
2. Whenever a Pi (π) bond is formed, it is formed in addition to a sigma (σ) bond in multiple bonds (a double or triple bond)
3. A single bond between the 2 atoms is always a sigma (σ) bond.

→ Bond length: The average distance between the centers of the nuclei of the two bonded atoms is called it’s the bond length.

→ Bond Enthalpy (Bond Energy): The energy required to break one mole of a particular type of bond so as to separate them into gaseous atoms is called bond dissociation energy or bond energy.

→ Dipole Moment: The product of the magnitude of charge (q) and the distance (d) between the centers of charges is called a dipole moment. Mathematically
Dipole moment = µ = q × d

→ Hybridization: It is the phenomenon of mixing of the atomic orbitals belonging to the same atom but having slightly different energies so that redistribution of energy takes place between them resulting in the formation of new orbitals of equal energies and identical shapes.

→ Hydrogen bonding: It can be defined as the attractive force which binds the hydrogen atom of one molecule with the more electronegative atom (like F, O, or N) of another molecule.

1. Intermolecular hydrogen bond: It is formed between two different molecules of the same or different compounds like H-bonds in water (H₂O), alcohol (R-OH), and HF.
2. Intramolecular hydrogen bond: It is formed between the H atom and the more electronegative atom (F, O, or N) within the same molecule as in o-nitrophenol.

→ Ionic compound: It is a three-dimensional aggregation of positive and negative ions in an ordered arrangement called the crystal lattice.

→ Valence Shell Electron Pair Repulsion (VSEPR) Theory: According to this theory used to explain the geometrical shapes of molecules-molecular geometry is determined by repulsions between lone pairs and bond pairs as follows: The order of these repulsions is lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

→ Bond Angle: It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule or complexion.

The bond angle in the water of H-O-H is 104.5° as shown above.

→ Bond enthalpy: It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. The unit of bond enthalpy is kJ m-1.

→ Bond order: In the Lewis description of covalent bond, the Bond order is given by the number of bonds between the two atoms in a molecule, e.g., bond orders in H₂ (H-H), O₇(O = O), and N₂(N = N) are 1, 2 and 3 respectively.

→ Resonance: Whenever a single Lewis structure cannot explain all the observed properties of a molecule, a number of structures with almost similar energy, positions of the nuclei, bonding, and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately.

By going through these CBSE Class 11 Chemistry Notes Chapter 3 Classification of Elements and Periodicity in Properties, students can recall all the concepts quickly.

Classification of Elements and Periodicity in Properties Notes Class 11 Chemistry Chapter 3

Why do we need to Classify Elements?
At present 114 elements are known. Efforts to synthesise new elements are continuing. It is very difficult to study such a large number of elements individually and to study their properties separately. The necessity was felt to organise their knowledge in a systematic way by classifying the elements.

Genesis of Periodic Classification: Classification of elements into groups led to the development of Periodic Law and Periodic Table.

J. Dobereiner was the first who made several groups of three elements (Triads). In each case he noticed the middle element of each of the Triads had an atomic weight about halfway between the atomic weights of the other two elements. Also, the properties of the middle element were in between these of the other two elements. He referred to them as the Law of Triads. It worked for only a few elements.

Dobereiner’s Triads:

J.A. Newlands propounded the Law of Octaves. He arranged the elements in increasing order of their atomic weights and noted that every eighth element has properties similar to the first element like the octaves of music – sa, re, ga, ma, pa, dha, nee, sa Newland’s Law of Octaves seemed to be true only for elements up to calcium.

Newlands Octaves:

Russian chemist Mendeleev and German chemist Lothar Meyer working independently, let to the Periodic Law as we know it today. Mendeleev succeeded in arranging the elements in vertical columns called groups and horizontal rows called periods in his table based upon the increasing order of their atomic masses. He gave his well- known law:

”The properties of the elements are a periodic function of their atomic weights”. It was called Mendeleev law. He realised that .some of the elements did not fit in with his scheme of classification if the order of atomic weights was strictly followed. He ignored the order of atomic weights and placed the elements -with similar properties together. For example, iodine with a lower atomic weight than that of tellurium (Group VI) was placed in Group VII along with other halogens F, Cl, Br because of similarities in properties.

He left some gaps for elements that were yet undiscovered. For example, both gallium (Ga) and germanium (Ge) were unknown at his time: He left the gap below aluminium and a gap below silicon and called these elements Eka-Aluminium and Eka-Silicon. The boldness of Mendeleev’s quantitative predictions and their eventual success made him and his Periodic Table famous. Mendeleev’s periodic table published in 1905 is shown.

Characteristics of Mendeleev’s Periodic Table:
It consists up of

1. Eight vertical columns called groups. Except for the VIII group, each group is. further subdivided in A and B. This subdivision is made on the basis of difference in their properties.
2. Six horizontal row called periods.

Significance of Mendeleev’s Periodic Table

1. Instead of studying properties of elements separately, they can be studied in groups containing elements with the same properties. It led to the systematic study of the elements.
2. Prediction of new elements. At his time only 56 elements were known. He left blank spaces or groups for unknown elements.
3. Mendeleev’s periodic table corrected the doubtful atomic weights.

Defects in the Mendeleev’s Periodic Table

1. Hydrogen was placed in group IA. However, it resembles both groups IA elements (alkali metals) and group VII A (halogens). Therefore the position of hydrogen in the periodic table is Anomalous or Controversial.
2. Anomalous pairs of elements. Some elements with higher atomic weight like Argon (39.9) precede potassium (39.1) with lower atomic weight.
3. Based upon atomic weights, isotopes of an element could not be assigned to different groups. They have been assigned only one position in a group.
4. Some dissimilar elements are grouped together while some similar elements are placed in different groups. For example, alkali metals in group IA which are highly reactive are in the same group as coinage metals like Cu, Ag, Au of group IB. At the same time, certain chemically similar elements like Cu (group IB) and Hg (group IIB) have been placed in different groups.
5. Position of elements of group VIII. No proper place has been allotted to nine elements of group VIII which have been arranged in three triads without any justification.

Mendeleev’s Periodic Table published in 1905

Modern Periodic Law and the present form of the Periodic Table.
Modern Period Law: The physical and chemical properties of the elements are periodic functions of their atomic number.

According to the recommendations of IUPAC, the groups are numbered from 1 to 18 replacing the older notation of groups 0, 1A, II A

There are 7 periods. The first period contains 2 elements. The subsequent periods contain 8, 8, 18, 18 32 elements respectively. The 7th period is incomplete and like the 6th period would have a theoretical maximum of 32 elements. In this form of Periodic Table, the elements of both the sixth and seventh periods (lanthanoids and actinoids respectively) are placed in Separate panels at the bottom.

Long Form of Periodic Table
General characteristics of the long form of the Periodic table:
1. There are in all, 18 vertical columns or 18 groups in the long-form periodic table.

2. These groups are numbered from 1 to 18 starting from the left.

3. There are seven horizontal rows called periods in the long-form periodic table. Thus, there are seven periods in the long-form periodic table.

The first period contains	2 elements	Shortest period
The second & Third period contains	8 elements each	Short period
The fourth & Filth period contains	18 elements each	Long-period
Sixth periods contains	32 elements	Longest period
The seventh period contains	It is incomplete	Incomplete period

4.  The elements of group 1, 2 .and 13 to 17 are called the main group elements. These are also called typical or representative or normal elements.

5. The elements of group 3 to 12 are called transition elements.

6. Elements with atomic number 58 to 71 (Ce to Lu) occurring after lanthanum (La) are called Lanthanides. Elements with atomic numbers 90 to 103 (Th to Lr) are called Actinides. These elements are called f-block elements and also inner-transition elements.


Long-form of the Periodic Table of the EIenient with their atomic numbers and ground state outer
electronic configurations. The groups are numbered i-18 in accordance with the 1984 IUPAC recommendations. this notation replaces the old numbering scheme of IA-VIIA, 4dIII, IB—VIIB and I) for the elements.

Nomenclature of the Elements with Atomic No. > 100
Table: Nomenclature of elements with atomic number above 103

Notation for IUPAC Nomenclature of Elements

Electronic Configuration of the Elements and Periodic Table: In the long-form periodic table, the elements are arranged in the order of their atomic numbers. An atomic number of an element is equal to the number of protons inside the nucleus of its atom. In an atom, the number of electrons is equal to the number of protons (hence equal to the atomic number). As a result, there is a close connection between the electronic configurations of the elements and the long form of the periodic table.

Table: The relationship between the electronic configuration of electrons and their positions in the periodic

Electronic Configuration in Periods: The period indicates the value of n- for the outermost or valence shell. In other words, a successive period in the Periodic Table is associated with the filling of the next higher principal energy level (n = 1, n = 2, etc.). It can be readily seen that the number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled.

→ The first period (n = 1) starts with the filling of the lowest level (1s) and therefore has two elements hydrogen (1s¹) and helium (1s²) when the first shell (K) is completed. The second period (n = 2) starts with lithium and the third electron and has the electronic configuration 1s22s2. Starting from the next element boron, the 2p orbitals are filled with electrons when the L shell is complete at neon (2s²2p⁶).

→ Thus there are 8 elements in the second period. The third period (n = 3) begins at sodium, and the added electron enters a 3s orbital. Successive filling of 3s and 3p orbitals gives rise to the third period 8 elements from sodium to argon. The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s orbital.

→ Now you may note that before the 4p orbital is filled, filling up of 3d orbitals becomes energetically favourable and we come from scandium (Z = 21) which has the electronic configuration 3d¹, 4s². The fourth period ends at krypton with the filling up the ‘Ip orbitals. Altogether we have 18 elements in this fourth period. The fifth period (n = 5) beginning with rubidium is similar to the fourth period and contains the 4d transition Serietarting at yttrium (Z = 39). This period ends al xenon with the filling up of the Sp orbitals.

→ The sixth period (n = 6) contains 32 elements and successive electrons enter 6s, 4f, Sd and 6p orbitals, in the order-filling up of the 4f orbitals begins with caesium (Z = 58) and ends at lutetium (Z = 71) Lo give 4f-inner transition series which is called lanthanoid sêries. The seventh period (n = 7) is similar to the sixth period with the successive filling up of the 7s, 5f, 6d and 7p orbitals and includes most of the man-made radioactive elements.

→ This period will end at the elements with atomic number 118 which would belong to the noble gas family. Filling up of the 5f orbital after actinium (Z = 80) gives the 5f-inner transition series known as the actinoid series. The 4f- and 5f- transition series of elements are placed separately in the Periodic Table to maintain its structure and to preserve the principle classification by keeping elements with similar properties in a single column.

Thus it can be seen that the properties oían element have periodic dependence upon its atomic number and not on relative atomic mass. All the elements in the same group have the same number of electrons (ns¹ in alkali metals) in their valence shells and thus have, same properties.

Electronic Configurations and Types of Elements s-, p-, d-, f-Blocks:
The Aufbau (build-up) principle and the electronic configuration of atoms provide a theoretical foundation for periodic classification. The elements in a vertical column of the Periodic Table constitute a group or family and exhibit similar chemical behaviour. This similarity arises because these elements have the same number and same distribution of electrons in their outermost orbitals. We can classify the elements into four blocks viz., s-block, p-block, d- block and f-block depending on the type of atomic orbitals that are being filled with electrons. We notice two exceptions to this categorization.

Strictly, helium belongs to the s-block but its positioning in the p-block along with another group of 18 elements is justified because it has a completely filled valence shell (1s²) and as a result, exhibits properties characteristic of other noble gases. The other exception is hydrogen. It has alone s-electron and hence can be placed in group-1 (alkali metals). It can also gain an electron to achieve a noble gas arrangement and hence it can behave similarly to group 17 (halogen family) elements. Because it is a special case, hydrogen is placed separately at the top of the Periodic Table.

→ The s-Block Elements: The elements of group 1 (alkali metals) and group 2 (alkaline earth metals) which have ns1 and ns² outermost electronic configuration belong to the s-Block Elements. They are all reactive metals with low ionization enthalpies. They lose the outermost electron(s) readily to form + lions (in the case of alkali metals) or + 2ion (in the case of alkaline earth metals). The metallic character and the reactivity increase as we go down the group. Because of high reactivity, they are never found pure in nature. The compounds of the s-block elements, with the exception of those of lithium and beryllium, are predominantly ionic.

→ The p-Block Elements: The p-Block Elements comprise those belonging to groups 13 to 18 and these together with the s-block elements are called the Representative Elements or Main Group Elements. The outermost electronic configuration varies from nsTnpl to ns²np⁶ in each period. At the end of each period is noble gas elements with a closed valence shell ns2np6 configuration.

All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity Preceding the noble gas family are two chemically important groups of non-metals. They are the halogens (Group 17) and the chalcogens (Group 16).

These two groups of elements have high negative electron gain enthalpies and readily add one or two electrons respectively to attain the stable noble gas configuration. The non-metallic character increases as we move from left to right across a period and the metallic character increases as we go down the group.

→ The d-Block Elements (Transition Elements): These are the elements of Group 3 to 12 in the centre of the Periodic Table These are characterised by the filling of inner (n – 1) d orbitals by electrons and are therefore referred to as d-Block Elements. These elements have the outer electronic configuration (n – 1)d^1-10ns^1-2. They are all metals. They mostly form coloured ions, exhibit variable valence (oxidation .states), paramagnetism and are often used as catalysts.

However, Zn > Cd and Hg which have, the electronic configuration, (n – 1)d¹⁰ns² do not show most of the properties of transition elements. In a way, transition metals form a bridge between the chemically active metals of s-block elements and the less active metals of Groups 13 and 14 and thus take their familiar name “transition elements”.

→ The f-Block Elements (Inner-Transition Elements): The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce (Z = 58) – Lu (Z = 71) and Actinoids, Th (Z = 90) – Lr (Z = 103) are characterised by the outer, electronic configuration (n – 2)f^1-14(n – 1)d^0-1 ns². The last electron added to each element is an f-electron.

These two series of elements are hence called the inner- transition elements (f-Block Elements). They are all metals. Within each series, the properties of the elements are quite similar. The chemistry of the early -actinoids .is more complicated than the corresponding lanthanoids, due to a large number of oxidation states possible for these actinoid elements.

Actinoid elements are radioactive. Many of the actinoid elements have been made only in nanogram quantities or even less by nuclear reactions and their chemistry is not fully studied. The elements after uranium are called transuranium elements.

The Lanthnoids show predominantly an oxidation state of + 3 (with a few exceptions + 2 or + 4) whereas Actinoids show variable oxidation states.

Metals, Non-Metals and Metalloids: The elements can be divided into Metals and Non-metals. Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table. Metals are usually solids at room temperature (mercury is an exception; gallium and caesium also have very low melting point 300K and 302K respectively). Metals usually have high melting and boiling points. They are good conductors of heat and electricity. They are malleable (can be flattened into thin sheets by hammering) and ductile (can be drawn into wires). In contrast, non-metals are located at the top right-hand side of the Periodic Table.

In fact, in a horizontal row, the property of elements change from metallic on the left to non-metals on the right. Non-metals are usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions). They are poor conductors of heat and electricity. Most non-metallic solids are brittle and are neither malleable nor ductile. The elements become more metallic as we go down a group; the non-metallic character increases as one goes from left to right across the Periodic Table.

The change from metallic to non-metallic character is not abrupt. The elements (e.g.r silicon, germanium, arsenic, antimony and tellurium) bordering this line and running diagonally across the Periodic Table show properties that are characteristic of both metallic and non-metals. These elements are called Semi-Metals or Metalloids.

Periodic Trends in Properties of Elements:
1. Ionization Enthalpy: The amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state is called ionisation enthalpy.
X (g) → X(g) + e^–

The energy required to remove the first electron is known as the first ionization enthalpy and the second electron is the second ionization enthalpy and so on. The second ionization enthalpy is always greater than The first ionization enthalpy because once an electron is removed, it becomes a positive ion and its nucleus has increased attraction for electrons. This makes it more difficult to remove the second electron.

Ionization enthalpy decreases from top to bottom in a group and increases horn left to right in a period. Thus Cs has the lowest ionization enthalpy and fluorine has the highest ionization enthalpy. Thus IE₃ > IE₂ > IE₁. IE of Li > Na > K > Rb > Cs and IE of Li in 2nd Period is lowest, whereas that of Na is highest.

The ionization enthalpy depends on two factors:

1. The attraction of electrons to the nucleus.
2. The repulsion of electrons from each other.

Screening or Shielding effect: A valence electron in a multi-electron atom is pulled by the nucleus And repelled by the other electrons in the core (inner shells). Thus the effective pull on the electron will be the pull due to the positive nucleus, less the repulsion due to core electrons.

The effective repulsive effect to the core electrons is called the screening or shielding effect. For example, the 2s electron in lithium is shielded by the inner core of Is electrons. As a result valence electron experiences a net positive charge less than + 3. In general, shielding is effective when the orbitals in the inner shells are completely filled.

2. Electron Gain (Enthalpy ΔegH): It is defined as the enthalpy change when a neutral gaseous atom takes up extra electrons to form an anion.
X(g) + e^– → X^– (g)

The value of electron gain enthalpy depends upon the atomic size, nuclear charge etc. with an increase in size, the electron gain enthalpy decreases as the nuclear attraction decreases. Thus electron gain enthalpy generally decreases in going from top to bottom in a group (F < Cl > Br > I). In a period from left to right, the electron gain enthalpy generally increases due to an increase in nuclear charge.

Thus halogens have very high electron affinity. However, the electron affinity of fluorine is less than that of chlorine. This is because fluorine has a small atomic size (only two shells). Also, electron repulsion is more in the case of fluorine, because of mutual electron-electron repulsions.

Factors affecting electron gain enthalpy:

1. Atomic size: Smaller the atom, the greater is the magnitude of the electron gain enthalpy. This is because the added electron can go closer to the nucleus and as a result releases more energy.
2. Effective nuclear charge: Greater the nuclear charge, the larger is the magnitude of electron gain enthalpy. This is because the electron would experience stronger attraction by virtue of a higher nuclear charge.
3. Electronic configuration of the atom: An atom with a stable electronic configuration has a little or no tendency to add another electron. As a result, such elements have zero, or nearly zero electron gain enthalpy.

Periodic Variation of Electron gain enthalpy:
1. Variation of electron gain enthalpy in a group: The magnitude of the electron gain enthalpy of elements decreases in going from top to bottom in a group. However, the electron gain enthalpy of fluorine is lower than that of chlorine.

On moving from top to the bottom of a group,
(a) the-atomic size increases and
(b) the nuclear charge also increases.

The effect of these factors is opposite to each other. In group, the effect of an increase in the atomic size outweighs the effect of the increased nuclear charge. As a result, the tendency to accept an electron in its valency. shell and hence the magnitude of electron gain enthalpy decreases as we go down the group.

2. Variation of electron gain Enthalpy in a period: The magnitude of electron gain enthalpy increases across a period is going from left to right.

On moving from left to right in a period, the size of the atoms decreases, and the effective nuclear charge increases. Both these factors increase the force of attraction exerted by the nucleus on the electrons. As a result, the atom has a greater tendency to gain an extra electron from outside and therefore, the magnitude of electron gain enthalpy increases in going from left to right. However, some elements in each period show an exception to such periodicity. For example, Be, Mg, N, P and noble gases show very low, or even zero electron gain enthalpies.

Periodic Variation of Atomic and Ionic Radii: Atomic radius is one half of the distance between the nuclei of two identical atoms in a molecule bonded by a single bond. In a period, as, we move from left to right, the number of shells remain the same but as more electrons are added the nuclear charge increases.

This results in an increase in the number attraction for the electrons, which in turn brings about a decrease in the radius of the elements in a period. Thus, the radius decreases along a period (from left to right). However, in a group from top to bottom; the number of shells increases, therefore, the radius also increases, e.g., r_Li < r_Na < r_K < r_Rb < … etc.

→ Ionic Radius: Ionic radius may be defined as “the effective distance from the centre and*nucleus of an ion up to which it has an influence on its electron cloud”.

→ Variation of Ionic radii in a group: The ionic radii of the ions belonging to the same group of elements and having identical charges, increases in going from top to bottom in a group. For example, the radii of the monovalent alkali metal ions increase from Li⁺ to Cs⁺.

→ Variation of the radii of isoelectronic ions: Atoms/ions of different elements having the same number of electrons are called isoelectronic ions. In other words, the ions which have the same number of electrons but different nuclear charges are called isoelectronic ions. The size of such ions depends upon the nuclear charge (as the total number of electrons remains the same).
Thus N^3- > O^2- > F^– > Na⁺ > Mg²⁺ > Al³⁺
Thus for isoelectronic ions, the ionic size decreases as the nuclear charge increases.

Cations are always smaller than their parent atoms because of

1. the disappearance of the valence shell in some cases
2. increase in effective nuclear charge.
   Thus Na⁺ < Na and Mg²⁺ < Mg.

Anions are always bigger in size than their parent atom. With the addition of one or more electrons to an atom, the effective nuclear charge per electron decreases. The electron cloud thus increases leading to an increase in the size of the anion.
Thus, F^– > F; Cl^– > Cl; O^2- > O; S^2- > S and so on.

→ Electronegativity: The ability of an atom in a chemical compound to attract the shared pair of electrons towards itself is called Electronegativity. It is not a measurable quantity, unlike IE and electron gain enthalpy. L Pauling gave numerical values of electronegativities to elements. F has been given a value of 4.0 arbitrarily, the highest in the periodic table. Electronegativity values are not constant. Electronegativity generally decreases from top to bottom in a group and increases across a period from left to right.

Electronegativity values (Pauling scale) across a period

Within a Group

Group-1	Group-17
Li = 1.0	F = 4.0
Na = 0.9	Cl = 3.5
K = 0.8	Br = 2.8
Rb = 0.8 Cs = 0.7	I = 2.5 At = 2.2

Periodicity of Valence: Valence is the most characteristic property of the elements and is based upon electronic configurations. The valence of representative elements is usually (though not necessarily) equal to the number of valence electrons arid or equal to eight minus the number of outermost electrons.

→ Anomalous Behaviour of Second Period Elements:
The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups 13-17 (boron to fluorine) differs in many respects from the other, members of their respective group. For example, lithium unlike other alkali metals, .and beryllium unlike other alkaline earth metals, form compounds with pronounced covalent character, the other members of these groups predominantly form ionic compounds.

In fact, the behaviour of lithium and beryllium is more similar to the second element of the following group i.e., magnesium and aluminium, respectively. This sort of similarity is commonly referred to as a diagonal relationship in the periodic properties.

→ What are the reasons for the different chemical behaviour Of the first members of a group of elements in the s- and p- blocks compared to that of the subsequent members in the same group?

The anomalous behaviour is attributed to their small size) large charge/radius ratio and high electronegativity of the elements. In addition, the first member of the group has only four valence orbitals (2s and 2p) available for bonding, whereas the second member of the groups has nine valence orbitals (3s, 3p, 3d). As a consequence of this, the maximum covalency of the first member of each group is 4, whereas the other members of the groups can expand their valence shell to accommodate more than four pairs of electrons.

Furthermore, the first member of p-block elements displays greater ability to form p_π – p_π multiple bonds to itself (e.g., C = C, C ≡ G, N = N, N ≡ N) and to other second period elements (e.g., C = O, C = N, C ≡ N, N = O) compared to subsequent members of the same group.

1. Lothar-Meyer Arrangement of Elements: When plotting a graph between the atomic volumes (gram atomic weight divided by density) and atomic weights of the elements he observed that the elements with similar properties occupied similar positions on the curve.

2. Mendeleev’s Periodic Law: The physical and chemical properties of elements are a periodic function of their atomic weights.

3. Moseley’s/Modem Periodic Law: The physical and chemical properties of the elements are a periodic function of their atomic numbers.

4. Periodicity of Properties of Elements: According to Modem Periodic Law the properties of the elements are repeated after certain regular intervals when these elements are arranged in order of their increasing atomic numbers. These regular intervals 2, 8, 8, 18, 18, 32 are called Magic Numbers.

5. Cause of Periodicity: The cause of periodicity in properties is the repetition of similar electronic configurations of the valence shells after certain regular intervals.

6. Groups: The vertical columns of elements in the periodic table are called Groups.

7. Periods: The horizontal rows of elements in the periodic table are called Periods:

→ There are 18 groups and 7 periods in the Modern Periodic Table
1. s-block Elements: These are the elements in which the last electron enters the s-subshell of the valence orbit.

2. p-Block Elements: The p-block elements are those in which the last electron enters the p-subshell of the vaLence orbit. The elements of groups 13, 14, 15, 16, 17, 18 (excluding helium) in which p-orbitals are being progressively filled in are called p-block elements. Since each group has five elements, therefore, ¡ri all, there are 30 p-block elements in the periodic table.

The elements of the 18th group are called noble gases.
General outer shell electronic configuration of p-block elements: ns2 np1 6

3. d-Block Elements: Elements in which the last electron enters any one of the five d-orbitals of their respective penultimate, i.e., (n – 1)the shells are called d-bock elements.

Transition Elements: Since the properties of these elements are midway between those of s-block and p-block elements, they are called Transition Elements.

General outer shell electronic confìguration of d-block elements: (n—1)d^1-10 n².
Zn, Cd, Hg do not show most of the properties of transition elements.

4. f-Block Elements: Elements in which the last electron enters any onè o the seven f-orbitals of their respective ante-penultimate shells are called f-block elements. In all these elements, the s-orbital of the last shell (n.) is completely filled, the d-orbital of the penultimate (n – 1) shell invariably contains zero or one electron but the J-orbitals of the antepenultimate (n – 2) shell (being lower in energy than d – orbitals of the penultimate shell) gets progressively filled in. Hence

General outer shell electronic configuration.off-block elements: (n – 2)^0-14 (n – 1) d^0-1 ns²

→ They are called Lanthanides/Lanthanones/Lanthanoids. In their case, the antepenultimate 4/subshell is being filled up.

→ They are called Actinides/Actinones/Actinoids if their antepenultimate 5/subshell is being b lied up.

All the actinoids are radioactive elements.

→ Elements from neptunium to lawrencium (₉₃Np – ₁₀₃Lr) which have been prepared artificially through nuclear reactions are called Transuranic or Transuranium elements as (hey follow uranium in the periodic table.

→ Metalloids: The elements like silicon, germanium, arsenic, antimony and tellurium (Si, Ge, As, Sb, Te) which show the properties of both metals and non-metals are called Metalloids.

→ Ionization Enthalpy: The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom of an element is called it’s Ionization Energy/Ionization Potential/Ionization Enthalpy.
M (g) + energy → M+ (g) + e^– (g)

→ Units of I.E.: electron volts (eV) per atom
or
kilo calories per mole (k cal mol^-1)
or
kilo Joules per mole (kj mol^-1)

1 eV per atom = 23.06 kcal mol^-1
= 96.49 kj mol^-1.

→ Successive I.Es.
IE₃ > IE₂ > IE₁
where M (g) + IE₁ → M⁺ (g) + e^– (g)
M⁺ (g) + IE₂ → M²⁺ (g) + e^– (g)
M²⁺ (g) + IE₃ → M³⁺ (g) + e^– (g)

→ Variation of IE across a period: It generally increases from left to right in a period.

→ Variation of IE within a group: It generally decreases from top to bottom within a group.

→ Electron Gain Enthalpy of an element may be defined as the energy released when a neutral isolated gaseous atom accepts an extra electron from outside.
X(g) + e^– → X^– (g); ΔH = Δ_egH

→ Variation across a period: In general, it becomes more and more negative from left to right in a period.

→ Variation within a group: In general, it becomes less negative as we move down a group.

→ Atomic Radius: The average distance from the centre of the nucleus to the outermost shell containing electrons.

They are of three types:

1. Covalent radius
2. Vander Waal’s radius
3. Metallic radius

→ Variation across a period: In general it decreases in a period with an increase in atomic number from left to right.

→ Variation within a group: It increases with the increase in atomic number from top to bottom within a group.

→ Ionic Radius: It is defined as the effective distance from the centre of the nucleus of the ion up to which it exerts Its influence on the electronic cloud.

→ Variation within a group: The ionic radius increases as we move down a group.

The size of a cation is always less than its corresponding atom
Na⁺ < Na
Mg²⁺ < Mg

The size of an anion is always larger than its corresponding atom
Cl^– > Cl
O^2- > o

→ Isoelectronic Ions or species: Ions of different elements which have the same number of electrons but the different magnitude of the nuclear charge are called Isoelectronic ions.
Na⁺, F^–, Mg²⁺ etc. are iso electronic ions.

The ionic radii of isoelectronic ions decrease with the increase in nuclear charge.

→ Trends in groups and periods: Electronegativity. values for the representative elements increase along period and decrease down the group.

Summary of the Trends in the Periodic Properties of Elements in the Periodic Table

→ Valency: The electrons present in the outermost shell of an atom are called valence electrons and the number of these electrons determine the valence or the valency of the atom. It is because of this reason that the outermost shell is also called the valence shell of the atom and the orbitals present in the valence shell are called valence orbitals.

→ Variation along a period: As the no. of electrons increases from 1 to 8 in representative elements across a period, the valency first increases from 1 to 4 and then decreases to zero in the case of noble gases.

→ Variation within a group: As the no. of electrons remain the same within a group, therefore, all the elements in a group exhibit the same valency.

The resemblance of properties of Li with Mg (which is diagonally situated); Be with Al and B with Si is called Diagonal relationship.

It is due to similar polarising power.