Prasanna

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.