Functional Groups Definitions, Equations and Examples

Functional Groups

A functional group in an organic compound is an atom or a group of atoms bonded together in a unique fashion, which is usually the site of chemical reactivity in an organic molecule.

If in a hydrocarbon chain, one or more hydrogen atoms are replaced by atoms of other elements such as halogens, oxygen, nitrogen, sulphur etc, such that the valency of carbon remains satisfied, then the element replacing hydrogen is referred to as a heteroatom. These heteroatoms confer specific properties to the compound, regardless of the length and nature of the carbon chain, and hence are called functional groups.

Functional Groups Definitions, Equations and Examples

Important Functional Groups

Some important functional groups are given in the table below.
Free valency or vaLencies of the group are shown by the single line. The functional group is attached to the carbon chain through this valency by replacing one hydrogen atom or atoms.
Functional Groups Definitions, Equations and Examples 1
Functional Groups Definitions, Equations and Examples 2
The electron dot structure of two compounds propanal, an aldehyde having molecular formula C2H5CHO and propanone, a ketone, having molecular formula CH3COCH3 are shown below:

Functional Groups Definitions, Equations and Examples

Electron Dot Structure of Propanal

Functional Groups Definitions, Equations and Examples 3

Electron Dot Structure of Propanone

Functional Groups Definitions, Equations and Examples 4

Some More Examples of Functional Groups
Functional Groups Definitions, Equations and Examples 5
Functional Groups Definitions, Equations and Examples 6

Class 10 Science Notes

Hydrocarbons Definitions, Equations and Examples

Hydrocarbons

The compounds made up of hydrogen and carbon only are called hydrocarbons. These are the simplest organic compounds and all other compounds are considered to be derived from them by the replacement of one or more hydrogen atoms by other atoms or groups of atoms.

The most important natural source of hydrocarbons is petroleum. There are two types of hydrocarbons:

  1. Saturated hydrocarbons
  2. Unsaturated hydrocarbons

Saturated Hydrocarbons or Alkanes:

  1. The hydrocarbons in which the carbon atoms are connected by only single bonds are called saturated hydrocarbons or alkanes.
  2. The general formula of saturated hydrocarbons are alkanes is CnH2n+2, where n is the number of carbon atoms in one molecule. The first few alkanes are methane (CH4), ethane (C2H6), and propane (C3H8).
  3. The saturated hydrocarbons are not very reactive.
  4. The saturated hydrocarbons generally give a clean flame. This is because the percentage of carbon is comparatively low which gets oxidized completely on combustion.

Structure of Saturated Hydrocarbons

The first step is to link the carbon atoms together with a single bond and then use the hydrogen atoms to satisfy the remaining valencies of carbon as shown in fig below.
Hydrocarbons Definitions, Equations and Examples 1
Methane: The simplest alkane is methane (CH4).
Hydrogen has a valency of 1. As carbons has four valence electrons, carbon shares these electrons with four atoms of hydrogen in order to achieve noble gas
configuration.

Hydrocarbons Definitions, Equations and Examples

It is widely used as a fuel and is a major component of bio-gas and Compressed Natural Gas (CNG).
Hydrocarbons Definitions, Equations and Examples 2
Ethane: Ethane is an alkane having two carbon atoms. The molecular formula of ethane is C2H6. There are seven single covalent bonds present in one molecule of ethane – one covalent bond between the two carbon atoms and six single bonds between carbon and hydrogen atoms.
Hydrocarbons Definitions, Equations and Examples 3

Unsaturated Hydrocarbons (Alkenes and Alkynes)

  1. The hydrocarbons in which the two carbon atoms are connected by a double bond or a triple bond are called unsaturated hydrocarbons.
  2. Unsaturated hydrocarbons may be alkenes (CnH2n) or alkynes (CnH2n-2).
  3. The general formula of an alkene is CnH2n, where n is the number of carbon atoms in one molecule.
  4. The general formula of an alkyne is CnH2n-2, where n is the number of carbon atoms in one molecule.
  5. These are more reactive than saturated hydrocarbons due to the presence of double and triple bonds which are the sites of chemical reactivity.
  6. These give a yellow flame with lots of black smoke. This is because the percentage of carbon is comparatively higher than saturated hydrocarbons which does not oxidize completely on combustion.

Structure of Unsaturated Hydrocarbons

In the first step, two carbon atoms linked together by single bond. Each carbon atom combines with two hydrogen atoms. One valency per carbon atom remains unsatisfied which can be satisfied only if there is a double bond between the two carbon atoms. Ethene : The simplest alkene: Ethene is the simplest alkene having two carbon atoms and its molecular formula is C2H4. There is a double bond between the two carbon atoms and four single bonds between carbon and hydrogen atoms.
Hydrocarbons Definitions, Equations and Examples 4
Ethyne: The simplest alkyne: The simplest alkyne is ethyne having two carbon atoms and its molecular formula is C2H2. There is a triple bond between the two carbon atoms and two single bonds between carbon and hydrogen atoms.
Hydrocarbons Definitions, Equations and Examples 5
Chains, Branches and Rings
Carbon atoms can form Long chains’ containing tens of carbon atoms.

When carbon atoms combine, three types of chains can be formed:

  1. Straight chains
  2. Branched chains
  3. Closed chains or ring-type chains.

Hydrocarbons Definitions, Equations and Examples 6
Some compounds have carbon atoms arranged in the form of a ring as in the case of cyclohexane (C6H12). Its structure is shown below. Similarly, the structure of benzene (C6H6) is also shown alongside:
Hydrocarbons Definitions, Equations and Examples 7

Hydrocarbons Definitions, Equations and Examples

Example 1.
What wilt be the formula and electron dot structure of cyclopentane?
Answer:
The formula of cyclopentane is C5H10 and its electron dot structure is given below:

Structural formula:

Hydrocarbons Definitions, Equations and Examples 8

Electron dot Structure:

Hydrocarbons Definitions, Equations and Examples 9

Structural Isomerism:

Organic compounds having some molecular formula but different physical and chemical properties due to different structures or catted structural isomers and this property is called isomerism. Isomerism is possible only with hydrocarbons having 4 or more carbon atoms.

If we look at the structure of butane (C4H10), we find that two different skeletons’ are possible with four carbon atoms having a single covalent bond:
Hydrocarbons Definitions, Equations and Examples 10
We find that both these compounds have the same molecular formula C4H10 but different structures and hence they are called Give isomers.

Isomers of hexane:
Hydrocarbons Definitions, Equations and Examples 11

Hydrocarbons Definitions, Equations and Examples

Example 2.
How many structural isomers can you draw for pentane?
Answer:
The molecular formula for pentane is C5H12. There are three structural isomers of’ pentane as given below
Hydrocarbons Definitions, Equations and Examples 12

Class 10 Science Notes

Versatile Nature of Carbon Definitions, Equations and Examples

Versatile Nature of Carbon

It is estimated that there are about three million carbon compounds whose formulae are known to chemists which is much greater than the compounds formed by all the other elements put together.
The factors due to which this is possible in the case of carbon are:

Catenation

The property of carbon elements due to which its atoms conjoin or link with one another to form long carbon chains is called catenation. These compounds may have long chains of carbon, branched chains of carbon or even carbon atoms arranged in rings.

In addition, carbon atoms may be linked by single, double or triple bonds.

Versatile Nature of Carbon Definitions, Equations and Examples

Tetra valency

The atomic number of carbon is 6 and its electronic configuration is 2,4. has a valency of 4, it can bond with four other atoms of carbon or atoms of other monovalent elements. Carbon forms compounds with oxygen, hydrogen, nitrogen, sulphur, chlorine and many other elements and these compounds have specific properties which depend on the elements other than carbon present in the molecule.

Strong Bonds due to Small Atomic Size:

The bonds that carbon forms with most other elements are very strong due to its small size making these bonds very stable. This enables the nucleus to hold on to the shared pairs of electrons strongly. The bonds formed by elements having larger atoms are much weaker.

Class 10 Science Notes

Covalent Bonding in Carbon Definitions, Equations and Examples

Covalent Bonding in Carbon:

The amount of carbon present in the earth’s crust and in the atmosphere is very less. The earth’s crust has only 0.02% carbon in the form of minerals (like carbonates, hydrogencarbonates,coal and petroleum) and the atmosphere has 0.03% of carbondioxide. Yet, we find that a large number of things that we use in our daily life are made of carbon compounds. Food, clothes, medicines, books and many other things are some examples.

The presence of carbon in a material can be tested by burning the substance in air and passing the gas formed through lime water. If the lime water turns milky, then the given material contains carbon.

Most carbon compounds are poor conductors of electricity and have low boiling and melting points, from which it can be concluded that the forces of attraction between these molecules are not very strong. Since these compounds are largely non-conductors of electricity, we can conclude that the bonding in these compounds does not give rise to any ions.

In the case of carbon, it has four electrons in its outermost shell and in order to attain noble gas
configuration, it needs to either gain four electrons or lose four electrons.

  1. It could gain four electrons forming C4- anion. But as the nucleus contains only six protons, it would be difficult for the nucleus to hold on to ten electrons.
  2. It could lose four electrons forming a C4+ cation. But a large amount of energy would be required to remove four electrons leaving behind a carbon cation with six protons in its nucleus holding on to just two electrons.

Carbon overcomes this problem by sharing its valence electrons with other atoms of carbon or with atoms of other elements. Apart from carbon, there are many other elements also which form molecules by sharing electrons in this manner. The shared electrons ‘belong’ to the outer shells of both the atoms and lead to both atoms attaining the noble gas configuration.

Covalent Bond: The types of bonds that are formed by the sharing of an electron pair between two atoms are known as covalent bonds.

Depending upon the number of pairs of electrons shared between atoms, there can be single, double or triple covalent bonds.

Covalent Bonding in Carbon Definitions, Equations and Examples

Some Examples of Covalent Compounds
Covalent Bonding in Carbon Definitions, Equations and Examples 1
Covalent Bonding in Carbon Definitions, Equations and Examples 2
Covalent Bonding in Carbon Definitions, Equations and Examples 3

Example 1.
Draw the electron dot structures for
(A) Ethanoic acid.
(B) H2S
(C) Propanone.
(D) F2
Answer:
(A) Ethanoic acid (CH3COOH):
A number of valence electrons of carbon = 4.
hydrogen = 1. oxygen = 6
Structural formula and electron dot structure is given below:
Covalent Bonding in Carbon Definitions, Equations and Examples 4
Covalent Bonding in Carbon Definitions, Equations and Examples 5
(B) Hydrogen suLphide (H2S):
Number of valence electrons of sulphur = 6, hydrogen = 1
Structural formula and electron dot structure is given:
Covalent Bonding in Carbon Definitions, Equations and Examples 7
(C) Propanone (CH3COCH3):
Number of valence electrons of carbon = 4,
hydrogen = 1, oxygen = 6
Structural formula and electron dot structure is given as:
Covalent Bonding in Carbon Definitions, Equations and Examples 7
(D) FLuorine (F2):
Number of valence electrons of fluorine = 7
The structural formula and electron dot structure is given below:
Covalent Bonding in Carbon Definitions, Equations and Examples 8

Covalent Bonding in Carbon Definitions, Equations and Examples

Properties of Covalent Compounds

Property Description
1. Covalent compounds are usually liquids or gases 1. This is due to the weak forces of attraction between their molecules.
2. Covalent compounds have usually low melting and boiling points 2. Covalently bonded molecules are seen to have strong bonds within the molecule, but intermolecular forces are small
3. Usually insoluble in water but soluble in organic solvents 3. This is also due to the presence of strong bonds within the molecule and small intermolecular forces
4. Covalent compounds do not conduct electricity 4. Since the electrons are shared between atoms and no charged particles are formed, such covalent compounds are generally poor conductors of electricity.

Differences between Ionic Bond and Covalent Bond

Ionic Bond Covalent Bond
1. An ionic bond is a chemical bond between two dissimilar (i.e. a metal and a non-metal) atoms in which one atom gives up an electron to another. 1. In a covalent bond the two atoms come together to share the electron, instead of an atom taking an electron from another
2. An ionic bond is formed between a metal and a non-metal. 2. A covalent bond is formed between two non-metals that have similar electronegativities.
3. Molecules have no definite shapes, as they have lattice structures 3. Molecules have a definite shape.
4. Electrical and thermal conductivity is High 4. No electrical conductivity but Thermal conductivity is usually low
5. Usually High melting point 5. Lower melting point
6.Usually highly soluble in water 6. lower solubility
7. Usually solids at room temperature 7. Exists as solids, liquids, gases

Allotropes of Carbon:

The various physical forms in which an element can exist are called allotropes of the element. Carbon exists in three solid forms called allotropes. The three alLotropes of carbon are:

  1. Diamond
  2. Graphite
  3. Fullerenes

Covalent Bonding in Carbon Definitions, Equations and Examples

Diamond

  • Diamonds are colourless, transparent, sparkle and reflect light, which is why they are described as Lustrous.
  • It is extremely hard and has a high melting point.
  • It does not conduct electricity.

Structure of diamond: Diamond is one giant molecule of carbon atoms. Every atom in a diamond is bonded to its neighbours by four strong covalent bonds, leaving no free electrons and no ions. This explains why diamond does not conduct electricity.
Covalent Bonding in Carbon Definitions, Equations and Examples 9
The Structure of Diamond

Uses of diamond:

  • Diamond is used in cutting instruments like glass cutters and in rock drilling equipment, as it is extremely hard.
  • Diamonds are used for making jewellery.
  • Sharp-edged diamonds are used by eye surgeons as a tool to remove cataract.

Graphite

  • Graphite is black, shiny and opaque.
  • It is a very slippery material.
  • Graphite is insoluble in water.
  • It has a high melting point and is a good conductor of electricity, which makes it a suitable material for the electrodes needed in electrolysis.

Structure of graphite: Graphite contains layers of carbon atoms.ln graphite, each carbon atom is bonded to three other carbon atoms in the same plane giving a hexagonaL array. One of these bonds is a double-bond, and thus the valency of carbon is satisfied. Graphite structure is formed by the hexagonal arrays being placed in layers one above the other.
Covalent Bonding in Carbon Definitions, Equations and Examples 10
The Structure of Graphite

Covalent Bonding in Carbon Definitions, Equations and Examples

Uses of graphite:

  • Powdered graphite is used as a lubricant for the fast-moving parts of machinery.
  • It is used for making electrodes in dry cells and electric arcs as it is a very good conductor of electricity.
  • It is used for making the core of pencils called ‘pencil Leads’.

Fullerenes:

Fullerenes form another class of carbon allotropes. The first one to be identified was C-60 which has carbon atoms arranged in the shape of a football. Since this looked like the geodesic dome designed by the US architect Buckminster Fuller, the molecule was named fullerene.
Covalent Bonding in Carbon Definitions, Equations and Examples 11
Structure of C-60 Bacminster Fullerene

Class 10 Science Notes

Corrosion of Metals Definitions, Equations and Examples

Corrosion of Metals:

The phenomenon of a surface of metal being attacked by air, water or any other substance around it is known as corrosion and the metal being attacked is said to corrode.

Examples of corrosion:

  1. When iron is exposed to moist air for a long time, its surface acquires a coating of a brown, flaky substance called rust, which is mainly hydrated iron oxide (Fe203.xH20).
  2. The surface of copper acquires a green coating of basic copper carbonate in moist air.
  3. The surface of aluminium gets coated with a thin layer of aluminium oxide which prevents the metal underneath from further damage.
  4. Silver articles become black after sometime when exposed to air as silver reacts with sulphur present in the air to form a coating of siLver sulphide.
  5. Noble metals such as goLd and silver do not corrode readily. Both air and water are necessary for corrosion to take place.

Corrosion of Metals Definitions, Equations and Examples

Example 1.
You must have seen tarnished copper vessels being cleaned with lemon or tamarind juice. Explain why these sour substances are effective in cleaning the vessels.
Answer:
Copper vessels become tarnished when the copper metal gets corroded and forms a layer of basic copper carbonate by reacting with carbon dioxide present in air.

Sour substances such as lemonjuice ortamarind juice are acidic in nature as they contain citric acid and tartaric acid respectively. Being acidic, they neutralize the basic copper carbonate and dissolves the layer of copper carbonate formed.

Prevention of Corrosion

Rusting of iron can be prevented by painting, oiling and greasing, galvanizing (by coating iron objects with zinc), chrome plating etc.

Galvanization:

Galvanisation is the method of protecting steel and iron from rusting by coating them with a thin layer of zinc. The galvanized article remains protected from corrosion even if the zinc layer is broken.

Important
If the zinc coating is broken, the galvanised object remains protected against rusting because zinc is more reactive than iron and hence can be easily oxidised .Thus when zinc layer breaks down, the zinc continues to react and gets oxidized.

Alloying:

Alloying is a method for improving the properties of a metal thereby getting desired properties.
Example: Pure iron is very soft and stretches easily when hot. It is therefore never used in its pure state, but instead mixed with a small amount of carbon (0.05 %) due to which it becomes hard and strong.

Alloys: An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by first melting the main metal and then dissolving the other elements in it in a definite proportion.

Corrosion of Metals Definitions, Equations and Examples

Amalgam:

An alloy whose one of the metals is mercury is known as an amalgam.
Properties of alloys:

  1. The electrical conductivity of an alloy is less than that of pure metals.
  2. Sometimes an alloy has lower melting point than any of its constituents.

Some common alloys:
Corrosion of Metals Definitions, Equations and Examples 1
Corrosion of Metals Definitions, Equations and Examples 2

Class 10 Science Notes

Occurrence and Extraction of Metals Definitions, Equations and Examples

Occurrence And Extraction of Metals:

A major source of metals is the earth’s crust. Soluble salts of certain metals such as sodium chloride, magnesium chloride, etc., are also found in seawater. Minerals: The inorganic elements or compounds which occur naturally in the earth’s crust are known as minerals, for example, zinc blende, cinnabar, etc. Ores: Those minerals from which a metal can be profitably extracted are called ores.

Gangue: The impurities like sand and rocky materials which are present in the ore are called gangue.

Occurrence of Metals:

(A) The metals at the bottom of the activity series (Au, Ag, Pt etc) are the least reactive and are often found in a free state.
(B) The metals in the middle of the activity series (Zn, Fe, Pb, etc) are moderately reactive and are found in the earth’s crust mainly as oxides, sulphides or carbonates.
(C) Metals at the top of the activity series (K, Na, Ca, Mg, Al) are very reactive and are never found as free elements.

Important
The ores of many metals are oxides because oxygen is a very reactive element and is very abundant on the earth.

Extraction of Metals:

The extraction of metals and the refining of them for use is known as metallurgy. The process of metallurgical operations consists of mainly three steps:

  1. Enrichment of ores
  2. Reduction
  3. Refining.

Occurrence and Extraction of Metals Definitions, Equations and Examples

Steps involved in the extraction of pure metals from ores
Occurrence and Extraction of Metals Definitions, Equations and Examples 1
Enrichment of Ores
The removal of impurities like sand, soil etc from the ore prior to the extraction of the metals is called enrichment of ores. The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the ore and the gangue.

Reduction of Ore to Metal:

Reduction: The process of obtaining metals from their oxides is known as reduction. Metal sulphides and carbonates are first converted into metal oxides as it is easier to reduce metal oxides as compared to metal sulphides and carbonates.

  • Extracting Metals Low in the Activity Series: Metals which are low in the activity series are very unreactive. Their oxides can be reduced to metals by heating alone.
    Examples:
  • Extraction of Mercury: Cinnabar (HgS) is an ore of Mercury, which is first converted to the mercuric oxide by heating and then it is reduced to mercury on further heating:
    Occurrence and Extraction of Metals Definitions, Equations and Examples 2
  • Extraction of Copper: Copper which is found as Cu2S is obtained from its ore by heating in air:
    Occurrence and Extraction of Metals Definitions, Equations and Examples 3

2. Extracting Metals in the Middle of the Activity Series
Metals in the middle of the activity series such as iron, zinc, lead, copper etc are moderately reactive and are usually present as sulphides or carbonates in nature.

  • Roasting: The process of converting sulphide ores into oxides by strongly heating in presence of excess of air is known as roasting.
    2ZnS(s) + 3O2(g) → 2ZnO(S) + 2SO2(g)
  • Calcination: The process of converting carbonate ores into oxides by strongly heating in absence of excess of air is known as calcination.
    ZnCO3(s) → ZnO(s) + CO2(g)

Occurrence and Extraction of Metals Definitions, Equations and Examples

The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon or by using displacement reactions with a more reactive metal.

Methods of Reduction:

Metal Type Method of Reduction
Unreactive metals which are Low in the activity series, like mercury Oxides of these metals are reduced by heating alone. Cinnabar (HgS) is heated in air and converted to mercuric oxide (HgO): 2HgS(S) + 3O2(g) → 2HgO(s) + 2SO2(g)

Mercuric oxide is then reduced to mercury on further heating: 2HgO(s) → 2Hg(l) + O2(g)
Moderately reactive metals which are in the middle of the activity series, such as iron, zinc, lead, copper etc Oxides of these metals are reduced by heating with carbon. ZnO(s) + C(s) → Zn(s) + CO(g)
Very reactive

metals which are high up in the activity series, such as sodium, magnesium, calcium etc

 Sometimes, displacement reactions can also be used. The highly reactive metals such as sodium, calcium, aluminium, etc, are used as reducing agents.

3Mn02(S) + 4Al(s) → 3Mn(l) + 2Al2O3(S) + Heat
The oxides of these metals are reduced by electrolytic reduction, i.e., by the electrolysis of their fused chlorides. The metals are deposited at the cathode and chlorine is liberated at the anode.

At cathode: Na+ → e– + Na
At anode:    2CI– → Cl2 + 2e–

Thermite Reactions: The reactions of metal oxides with aluminium are highly exothermic and the amount of heat evolved is so large that the metals are produced in the molten state. The reaction of iron oxide (Fe203) with aluminium is used to join railway tracks or cracked machine parts.
Fe2O3(S) + 2Al(S) → 2Fe(l) + Al2O3(s)

3. Extracting Metals Towards the Top of the Activity Series
The highly reactive metals are extracted by the electrolysis of their molten chlorides or oxides. These oxides are very stable and cannot be reduced by the most common reducing agent carbon as these metals have more affnity for oxygen than carbon.

Extraction of Sodium Metal: Sodium metal is extracted by the electrolysis of molten sodium chloride. When electric current is passed through molten sodium chloride, it decomposes to form sodium metal and chlorine gas:
2NaCl → 2Na + Cl2
The reactions taking place at the two electrodes are given below:
At the Cathode: 2Na+ + 2e– → 2Na
At the Anode: 2Cl– + 2e– → Cl2

Occurrence and Extraction of Metals Definitions, Equations and Examples

Example 1.
Which chemical process is used for obtaining a metal from its oxide?
Answer:
The chemical process used for obtaining a metal from its oxide is reduction. Metal oxides can be reduced to metal either by heating alone (less reactive metals) or by heating with carbon (moderately reactive metals).

Example: Reduction of HgO (oxide of less reactive metal) by heating alone:
2HgO(s) → 2Hg(l) + O2(g)
Reduction of ZnO (oxide of moderately reactive metal) by heating with carbon:
ZnO(s) + C(s) → Zn(s) + CO(g)

Refining of Metals

The process of purification of impure metals is known as refining of metals and the method used for refining an impure metal depends upon:

  1. Nature of the element
  2. Nature of impurities present in it.

Electrolytic refining of copper: The metaLs such as copper, zinc, nickel, silver, gold, etc., are refined electrolytically.
The impure metal is made as anode and a thin strip of pure metal is made as cathode A solution of the metal salt is used as an electrolyte.
Occurrence and Extraction of Metals Definitions, Equations and Examples 4
Electrolytic refining of copper

On passing a current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode.
The solid impurities go into the solution and the insoluble impurities settle down at the bottom of the anode and are known as anode mud.
At anode: Cu → Cu2+ + 2e–
At cathode: Cu2+ + 2e– → Cu

Anode: A thin block of impure metal is connected to the positive terminal of the battery.
Cathode: A thin strip of pure metal is connected to the negative terminal of the battery.
Electrolyte: A water-soluble salt of the metal to be refined is taken as the electrolyte.

Class 10 Science Notes

Reaction Between Metals and Non-Metals Definitions, Equations and Examples

Electronic Configuration of Noble Elements:

Noble gases have a completely filled valence shell and hence are chemically inert. The usual number of valence electrons in the atom of noble gases is 8 except helium in which it is 2.

Cause of Chemical Reactivity:

The reactivity of an element can be explained as the tendency of an atom to attain a completely filled valence shell.
The atoms combine with one another to achieve the inert gas electron configuration and become more stable. Atoms therefore form chemical bonds to achieve stability by acquiring the inert gas electron configuration. An atom can achieve the inert gas electron configuration in the following ways:

  • By losing one or more electrons to another atom.
  • By gaining one or more electrons from another atom.
  • By sharing one or more electrons with another atom.

Ions:

An ion is an electrically charged atom or group of atoms. An ion is formed by the loss of electrons (sodium ion, Na+, magnesium ion, Mg2+) or gain of electrons (chloride ion, CT, oxide ion, O2-) by an atom. There are two types of ions:- cations and anions.

Cation:

A positively charged ion is known as cation. Sodium-ion, Na+and magnesium ion, Mg2+ are cations because they are positively charged ions since the number of protons in the nucleus is more-than the number of electrons due to which there is a net positive charge on the atom. A cation has less electrons than protons.

Reaction Between Metals and Non-Metals Definitions, Equations and Examples

Anion:

A negatively charged ion is known as anion. Chloride ion, Cl” and oxide ion, O2- c.e joined by the addition of electrons due to which t! ere is a net negative charge on the atom. An anion has more electrons than protons.

Ionic Compounds:

When metals react with non-metals, they form compounds by the transfer of electrons from a metal to a non-metal. Such compounds are known as ionic or electrovalent compounds.

Ionic Bond:

The chemical bond formed by the transfer of electrons from one atom to another is known as an ionic bond. An ionic bond is formed when a metal reacts with a non-metal and the transfer of electrons takes place from metal atoms to non-metal atoms.

The metal atom develops a positive charge after donating electrons and thus becomes a cation.
The non-metal atom, on the other hand, becomes a negatively charged ion, anion, after gaining or accepting electrons.

Formation of Sodium Chloride (NaCl):

Sodium donates one electron to a chlorine atom and forms a sodium ion, Na+.

Formation of Sodium-ion, Na+: The atomic number of sodium is 11. Its electronic configuration is: 2, 8, 1 Since it does not have an octet, it is not very stable. In order to become stable, the sodium atom donates its 1 outermost electron to some other atom.
Reaction Between Metals and Non-Metals Definitions, Equations and Examples 1
Formation of Chloride ion, Cl–: Chlorine atom takes one electron from the sodium atom and forms a negatively charged ion, CT.
Cl + e– → Cl–

Reaction Between Metals and Non-Metals Definitions, Equations and Examples

Sodium chloride is an ionic compound and contains ionic bonds. Both these ions, being oppositely charged, attract each other and are held by strong electrostatic forces of attraction to form an ionic compound. Ionic compounds do not exist as molecules but as aggregates of oppositely charged ions.
Reaction Between Metals and Non-Metals Definitions, Equations and Examples 2

Formation of Magnesium Chloride (MgCl2):

Magnesium donates two electrons to two chlorine atoms and forms a magnesium ion, Mg2+. Similarly, 2 chlorine atoms take one electron each from the magnesium atom and form 2 negativeLy charged ions, Cl”.
Reaction Between Metals and Non-Metals Definitions, Equations and Examples 3
Reaction Between Metals and Non-Metals Definitions, Equations and Examples 4
Formation of Some Common Ionic Compounds
Reaction Between Metals and Non-Metals Definitions, Equations and Examples 5

Reaction Between Metals and Non-Metals Definitions, Equations and Examples

Properties of Ionic Compounds

The general properties of ionic compounds are summarised below:

Property name Property Shown by Ionic Compounds
1. Physical nature Ionic compounds are generally crystalline solids and hard due to the strong force of attraction between the positive and negative ions. They are generally brittle.
2. Melting and boiling points They have high melting and boiling points as a large amount of energy is required to break the strong inter-ionic attraction.
3. Solubility These are generally soluble in water but insoluble in organic solvents Like ether, kerosene, petrol etc.
4. Conduction of electricity These conduct electricity in the molten state as the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat.

Moreover, they also conduct electricity when dissolved in water as its solution in water contains ions.

However, these do not conduct electricity in the solid state due to their rigid structure.

Class 10 Science Notes

Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples

Chemical Properties of Metals and Non-Metals:

Metals are electropositive in nature as they lose electrons and form positive ions.
K → K+ + e–
Ca → Ca2++ 2e–

Non-metals are electronegative in nature as they accept electrons and form negative ions.
Cl + e– → cr
S + 2e– → S2-

Most non-metals produce acidic oxides
Examples: SO2, CO2
Most metals produce basic oxides. Examples: Na2O, CaO

Chemical Properties of Metals:

Reaction with Oxygen:
Metals combine with oxygen to form basic or amphoteric oxides.
Metal + Oxygen → Metal Oxide
Examples: 2Cu + O2 → 2CuO (Copper oxide)
4Al3O2 → 2Al2O3
Although metal oxides are basic in nature some metals show both acidic and basic nature.

Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples

Amphoteric oxides:

The metal oxides which show both acidic and basic nature are known as amphoteric oxides.
For example, aluminium oxide and zinc oxide are amphoteric oxides.
The reaction of aluminium oxide with acids and bases is given below:
Al2O3 + 6HCl → 2AlCl3 + 3H2O
Al2O3 + 2NaOH → 2NaAlO2 + H2O

Alkalies:

The oxides of most metals are insoluble in water but some of these dissolve in water to form alkalies.
Examples: Sodium oxide and potassium oxide
Na2O(s) + H2O(l) → 2NaOH(aq)
K2O(s) + H2O(l) → 2KOH(aq)

The reaction of different metals with oxygen: Different metals show different reactivities towards oxygen.

  1. The reaction of alkali metals such as potassium and sodium with oxygen is so vigorous that they are kept immersed in kerosene oil to prevent accidental fires in case they are kept open.
  2. The surfaces of metals such as magnesium, aluminium, zinc, lead, etc., are covered with a thin layer of oxide at ordinary temperature, which prevents the metal from further oxidation and is, therefore, a protective layer.
  3. Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of a burner.
  4. Copper does not burn, but the hot metal is coated with a black coloured layer of copper (II) oxide.
  5. Silver and gold do not react with oxygen even at high temperatures.

Anodising:

Anodising is the process of forming a thick oxide layer of aluminium. Aluminium forms a thin oxide layer when exposed to air which makes it resistant to further corrosion. During anodising, a clean aluminium article is made the anode and is electrolysed with dil. H2SO4. The oxygen gas evolved at the anode reacts with aluminium to form a thicker protective oxide layer.

Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples

Example 1.
An element reacts with oxygen to give a compound a high melting point. This compound is also soluble in water. The element is likely to be:
(a) Calcium
(b) Carbon
(c) Silicon
(d) Iron
Answer:
(a) Calcium

Explanation: As the compound has a high melting point, it is an ionic compound. Since metal oxides are ionic in nature, therefore, the element could be calcium or iron.
However, it is also given that the compound is also soluble in water and as calcium oxide is soluble in water, whereas iron oxide is not soluble, the given element is calcium.
2Ca(s) + O2(g) → 2CaO(s)
CaO(s) + H2O(l) → 2Ca(OH)2(aq) + Heat

Reaction with Water:

Metals react with water and produce a metal hydroxide or oxide and hydrogen gas is evolved.
Metal + Water → Metal oxide + H2

Metal oxides that are soluble in water dissolve in it to further form metal hydroxide.
Metal oxide + Water → Metal hydroxide

1. Metals like K and Na react violently with cold water.
2K(S) + 2H2O(l) → 2KOH(aq) + H2(g) + Heat
2Na(S) + 2H2O(l) → 2NaOH(aq) + H2(g) + Heat

2. The reaction of calcium with water is less violent.
Ca + 2H2O → Ca(OH)2 + H2
Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.

3. Mg does not react with cold water but reacts with hot water to form Mg(OH)2– It also starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.

4. Metals like Al, Zn and Fe react with steam to form a metal oxide.
2Al(S) + 3H2O(g) → Al2O3(s) + 3H2(g)
3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)

5. Metals like Pb, Cu, Ag and Au do not react with water at all.

Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples

Example 2.
Case-Based:
Collect the samples of the following metals: aluminium, copper, iron, lead, magnesium, zinc and sodium. Put small pieces of the samples separately in beakers half-filled with cold water. Put the metals that did not react with cold water in beakers half-filled with hot water. For the metals that did not react with hot water, arrange the apparatus as shown in Figure and observe their reaction with steam.
Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples 1
(A) Select the metals which react with cold water:
(I) Al
(II) Mg
(III) Na
(IV) K
(a) Both (I) and (II)
(b) Both (I) and (III)
(c) Both (III) and (IV)
(d) (I), (II) and (IV)
Answer:
(c) Both (III) and (IV)
Explanation: Metals like Kand Na react violently with cold water.
2K(S) + 2H2O((l) → 2KOH(aq) + H2(g) + Heat
2Na(s)+ 2H2O(l) → 2NaOH(aq) + H2(g) + Heat
Mg does not react with cold water but reacts with hot water to form Mg(OH)2.
Al reacts with steam to form a metal oxide.
2Al(s) + 3H20(g) → Al2O3(s) + 3H2(g)

(B) A metal X when treated with cold water gives the metal hydroxide Y and starts floating.
The metal X and compound Y could be:

Metal X Compound Y
(I) Mg Mg(OH)2
(II) Ca Ca(OH) 2
(III) Na NaOH
(iv) K KOH

Select the correct option:
(a) Only (I)
(b) Only (II)
(c) Both (I) and (III)
(d) Both (II) and (IV)
Answer:
(b) Only (II)
Explanation: Out of the metals Mg, Ca, Na and K, both Na and K react so violently with water that the hydrogen gas evolved immediately catches fire.
The reaction of Ca with cold water is less violent and calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.
Ca + 2H2O → Ca(OH)2 + H2
Mg does not react with cold water but reacts with hot water to form Mg(OH)2 and hydrogen.

(C) Name the metals which do not react either with cold or hot water.
Answer:
The metals Al, Zn and Fe do not react either with cold or hot water but they react with steam to form the metal oxide and hydrogen gas.
2Al(S) + 3H2O(g) → Al2O3(s) + 3 H2(g)
3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)

(D) Arrange the metals aluminium, copper, iron, lead, magnesium, zinc and sodium in decreasing order of their reactivity with water.
Answer:
The decreasing order of reactivity of metals with water is:
Na > Mg > Al > Zn > Fe > Pb > Cu

(E) Assertion (A): Sodium catches fire when it reacts with hot water.
Reason (R): The reaction between sodium and cold water is highly exothermic.
(a) Both (A) and (R) are true and (R) is the correct explanation of the (A).
(b) Both (A) and (R) are true, but (R) is not the correct explanation of the (A).
(c) (A) is true, but (R) is false.
(d) (A) is false, but (R) is true.
Answer:
(d) (A) is false, but (R) is true.
Explanation: Sodium reacts violently even with cold water and the reaction is such a highly exothermic reaction that the hydrogen gas evolved immediately catches fire.

Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples

Example 3.
Give reasons why copper is used to making hot water tanks and not steel (an alloy of iron).
Answer:
Steel is an alloy of iron and iron reacts with steam to form Fe3O4.
3Fe(S) + 4H2O(g) → Fe3O4(s) + 4H2(g)
However, copper is less reactive and does not react either with cold water, hot water or steam. Therefore, copper is used to making hot water tanks being less reactive than iron.

Reaction with Acids:

1. When a metal reacts with dil HCl and dil H2SO4, a salt is formed and H2 gas is evolved.
Metal + Dil. Acid → Salt + Hydrogen
Zn(s) + ZnSO4(aq) → ZnSO4(aq) + H2

2. H2 gas is not evolved when a metal reacts with HNO3 (nitric acid) as HNO3 is a strong oxidizing agent and oxidizes the H2 produced to water and is itself reduced to any of the oxides of nitrogen. Exceptions: Magnesium and Manganese.

3. The reactivity of all metals towards acids is not the same and can be found out from the rate at which bubbles of hydrogen gas are formed and the amount of heat generated. Amongst Al, Mg, Fe and Zn, the reactivities towards acids is: Mg > Al > Zn > Fe

4. Aqua regia: It is a freshly prepared mixture of cones. HCl and cone. HNO3 in the ratio of 3:1. It is a highly corrosive, fuming liquid which can even dissolve gold and platinum, even though neither of these acids can do so alone.

Important
Metals such as copper, mercury, silver and gold, which are placed below hydrogen in the reactivity series of metals, do not react with dilute acids.

Example 4.
A man went door-to-door posing as a goldsmith. He promised to bring back the glitter of old and dull gold ornaments. An unsuspecting lady gave a set of gold bangles to him which he dipped in a particular solution. The bangles sparkled like new but their weight was reduced drastically. The lady was upset but after a futile argument, the man beat a hasty retreat. Can you play the detective to find out the nature of the solution he had used?
Answer:
The man used aqua regia which is a mixture of cone. HCl and cone. HNO3 acids in the ratio 3:1 and has the property of dissolving gold. When the unsuspecting Lady gave a set of gold bangles to the man, he dipped them in aqua regia solution due to which some gold dissolved in the solution bringing back the sparkle but at the same time the weight of gold bangles was reduced as it went into the solution.

Example 5.
Name two metals which will displace hydrogen from dilute acids, and two metals which will not.
Answer:
Two metals which will displace hydrogen from dilute acids are magnesium and aluminium, and two metals which will not displace hydrogen are copper and silver.

The reaction of Metals with Solutions of Other Metal Salts:

1. Reactive metals displace less reactive metals from their compounds in solution or molten form. Displacement reactions give better evidence about the reactivity of metals. If metal A displaces metal B from its solution, it is more reactive than B.
Metal A + Salt solution of B → Salt solution of A + Metal B

2. If we put an iron nail in a solution of copper sulphate, we observe that the blue colour of copper sulphate solution in the test tube fades gradually and red-brown copper metal is formed. The reaction taking place is:
CuSO4(aq) + Fe(s) → FeSO4(aq) + Cu(s)

However, no reaction takes place if a strip of copper metal is placed in iron sulphate solution because copper is less reactive than iron and hence cannot displace iron from iron sulphate solution.

Important
Some examples of reactions of metals with solutions of other metal salts.

Example of Reaction of No. Metal with a solution of Metal Salt Chemical Equation
1. Reaction of zinc with copper sulphate solution CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s)
2. Reaction of copper with a silver nitrate solution 2AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2Ag

If zinc oxide, magnesium oxide and copper oxide were heated, turn by turn, with zinc, magnesium and copper metals, we can note down the displacement reactions taking place as shown:
Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples 2

Activity Series (The Relative Reactivities of Metals):

The metals have been arranged in order of their decreasing reactivity by using displacement reactions since all metals are not equally reactive. The activity series is shown below:
Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples 3

Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples

Example 6.
Samples of four metals A, B, C and D were taken and added to the following solution one by one. The results obtained have been tabulated as follows:
Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples 4
Use the table above to answer the following questions about metals A, B, C and D.
(A) Which is the most reactive metal?
Answer:
We observe that A undergoes displacement reaction with Copper(ll) sulphate. Therefore, A > Cu.
Metal B displaces iron from iron(ll) sulphate. Therefore, B > Fe.
Metal C displaces silver from silver nitrate. Therefore, C > Ag.
Metal D does not undergo a reaction with any of the given salt solutions. Therefore D is the least reactive metal.
As reactivity of Fe, Cu, Zn and Ag can be written as: Zn > Fe > Cu > Ag, and B displaces Fe, therefore B is the most reactive metal.

(B) What would you observe if B is added to a solution of Copper(ll) sulphate?
Answer:
If B is added to a solution of CuSO4, it will displace copper from the solution as B is more reactive than copper.

(C) Arrange the metals A, B, C and D in the order of decreasing reactivity.
Answer:
Metals A, B, C and D arranged in decreasing order of reactivity: B > A > C > D

Chemical Properties of Non-Metals

Reaction with Oxygen
Non-metals combine with oxygen to form acidic or neutral oxides which are covalent compounds.
C(s) + O2(g) → CO2(g)
CO2(g) + H2O(l) → H2CO3(aq)

Examples of neutral oxides are carbon monoxide (CO), nitrous oxide (N2O).
Non-metals do not displace hydrogen from dilute acids as it cannot supply electrons to H+ ions.

Reaction with Hydrogen

A few active metals Like Na, K and Ca force the hydrogen atom to accept electrons to form salts called hydrides.
2Na(s) + H2(g) → 2NaH(s)
Ca(s) + H2(g) → CaH2(s)

Non-metals combine with hydrogen to form covaLent hydrides.
N2(g) + 3H2(g) → 2N2H3(g)

Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples

Example 7.
Pratyush took a sulphur powder on a spatula and heated it. He collected the gas evolved by inverting a test tube over it, as shown in Figure below.
Chemical Properties of Metals and Non-Metals Definitions, Equations and Examples 5
Collection of gas

(A) What will be the action of gas on:
(i) dry litmus paper?
(ii) moist litmus paper?
Answer:
The gas evolved when Pratyush heated sulphur powder in a spatula is sulphur dioxide, SO2.
(i) There will be no change by the action of this gas (SO2) on dry litmus paper.
(ii) When this gas acts on moist blue litmus paper, it changes the colour of litmus from blue to red as sulphur dioxide is acidic in nature since oxides of non-metals are acidic in nature but in presence of water.

(B) Write a balanced chemical equation for the reaction taking place.
Answer:
Balanced chemical equation for the reaction taking place is:
S(s) + O2(g) → SO2(g)
SO2(g) + H2O(l) → H2SO3(aq)

Class 10 Science Notes

Physical Properties of Metals And Non-Metals Definitions, Equations and Examples

Physical Properties of Metals And Non-Metals:

On the basis of their properties, all the elements can be divided into two main groups: metals and non-metals, Apart from metals and non-metals, there are some elements that show characteristics of both metals and non-metals and known as metalloids, for example, Si, Ge, As, Sb and Te.

Metals are opaque, lustrous elements that are good conductors of heat and electricity. Examples: Copper, Iron, Silver, Aluminium, etc.

Non-metals are elements that have properties that are different from those of metals. Examples: Carbon, Nitrogen, Oxygen, Sulphur etc.

Physical Properties of Metals:

  1. State: Metals are generally solids at room temperature except for mercury which is a liquid.
  2. Lustre: In pure state metals have a shining surface. Al and Mg are white whereas Au is yellow and Cu is reddish-brown in colour.
  3. Hardness: Metals are generally hard but hardness varies from metal to metal.
  4. Malleability: Some metals can be beaten into thin sheets and this property is called malleability. Gold and silver are the most malleable metals.
  5. Ductility: Some metals can be drawn into very thin wires and this property is called ductility. Gold is the most ductile metal.
  6. Thermal and electrical conductivity: Metals are generally good conductors of heat and electricity. Silver is the best conductor of heat followed by copper and aluminium. Silver and copper are the best conductors of electricity.
  7. Melting and Boiling points: Metals generally have high melting and boiling points. The boiling and melting points of aluminium are 933 K and 2792 K respectively.
  8. Sonorous: Metals that produce a sound on striking a hard surface are said to be sonorous.

Physical Properties of Metals And Non-Metals Definitions, Equations and Examples

Example 1.
Case-Based:
A student took an aluminium or copper wire and clamped this wire on a stand, as shown in Fig.
Physical Properties of Metals And Non-Metals Definitions, Equations and Examples 1
He then fixed a pin to the free end of the wire using wax and heated the wire with a spirit lamp, candle or a burner near the place where it is clamped.

Another student set up an electric circuit as shown in the Figure below to test the electrical conductivity of metals. He placed the metal to be tested in the circuit between terminals A and B as shown.
Physical Properties of Metals And Non-Metals Definitions, Equations and Examples 2
(A) The activity performed by the first student (Figure) shows that:
(a) Metals are good conductors of heat and have low melting points.
(b) Metals are poor conductors of heat and have high melting points.
(c) Metals are good conductors of heat and have high melting points.
(d) Metals are poor conductors of heat and have low melting points.
Answer:
(c) Metals are good conductors of heat and have high melting points.

Explanation: Metals have a high melting point and are generally good conductors of heat. Most metals are also good conductors of electricity.

(B) The metals which are poor conductors of heat are:
(a) Lead and zinc
(b) Mercury and zinc
(c) Lead and tin
(d) Lead and mercury
Answer:
(d) Lead and mercury
Explanation: The metals lead and mercury are comparatively poor conductors of heat. Whereas the best conductors of heat are silver and copper.

(C) What conclusion can be drawn from the activity performed by the second student (Figure)?
Answer:
The conclusion that can be drawn from the activity performed by the second student (Figure) is that the bulb glows when a metal is inserted between A and B showing that most of the metals conduct electricity.

(D) Name the two best conductors of electricity.
Answer:
Silver and copper are the best conductors of electricity.

(E) Assertion (A): Metal such as aluminium is used for making cooking utensils.
Reason (R): Aluminium is a highly ductile metal.
(a) Both (A) and (R) are true and (R) is the correct explanation of the (A).
(b) Both (A) and (R) are true, but (R) is not the correct explanation of the (A).
(c) (A) is true, but (R) is false.
(d) (A) is false, but (R) is true.
Answer:
(b) Both (A) and (R) are true, but (R) is not the correct explanation of the (A). Explanation: MetaL such as aluminium is used for making cooking utensils as it is a good conductor of heat and has high melting point. The property of metals used for making electrical wires is ductility and good electrical conductivity.

Physical Properties of Non-Metals:

  1. State: Non-metals are either solids or gases except bromine which is a liquid.
  2. Lustre: Non-metals do not have any lustre.
  3. Hardness: Non-Metals are generally soft. Exceptions: Diamond is an allotrope of carbon and is the hardest natural substance
  4. Malleability: Non-metals are not malleable. Actually, they are brittle.
  5. Ductility: Non-metals are not ductile also.
  6. Thermal and electrical Graphite, which is an allotrope of Carbon conducts electricity. Therefore, this is an exception. do not conduct heat and electricity.
  7. Melting and boiling points: Non-metals have low melting and boiling points.

Physical Properties of Metals And Non-Metals Definitions, Equations and Examples

Example 2.
You are given a hammer, a battery, a bulb, wires and a switch.
(A) How could you use them to distinguish between samples of metals and non-metals?
Answer:
Metals and non-metals have different physical properties. If we take the given samples and strike them with a hammer, metals will be converted into sheets, whereas non-metals will not. This is because metals are malleable.
Also, metals are sonorous and produce sound when hit with a hammer, whereas non-metals are not sonorous.
The battery, wires, bulb and switch can be arranged in the form of a circuit for testing the electrical conductivity of samples (refer to Figure).

When samples of metals are inserted between A and B, the bulb glows which shows that metals are good conductors of electricity, whereas the bulb does not glow when samples of non¬metals are inserted which shows that non-metals are poor conductors of electricity.

(B) Assess the usefulness of these tests in distinguishing between metals and non-metals.
Answer:
These tests are useful in distinguishing between metals and non-metals based on the differences in properties of metals and non-metals as metals are malleable, sonorous and good conductors of electricity whereas non-metaLs do not possess any of these properties.

Exceptions in Physical Properties:

Physical Property

 Exception
1. Physical State Mercury is a metal that is a liquid at room temperature. Bromine is a non-metal that is liquid at room temperature.
2. Melting Point and Boiling point Sodium, potassium, caesium and gallium are metals that have low melting points. Di­amond is a non-metal which has a high melting point.
3. Density Alkali metals (lithium, sodi­um, potassium) are so soft that they can be cut with a knife.
4.Electrical Conductivity Mercury, a metal, offers very high resistance to the passage of current.
5. Lustre Iodine is a non-metal which is lustrous.

Class 10 Science Notes

Salts Definitions, Equations and Examples

Salts:

Salts are produced when an acid reacts with a base and such reactions are known as neutralization reactions.

Family of Salts:

Salts having the same positive or negative radicals are said to belong to the same family.
Examples:

  1. NaCl and Na2SO4 belong to the family of sodium salts as they have the common positive radical, namely, sodium.
  2. KCl and CaCl2 belong to the family of chloride salts as they both have the common negative radical, namely, chloride.

Salts Definitions, Equations and Examples

pH of Salts:

Salts may be neutral, acidic or basic depending upon the strength of the acid and base from which it has been derived or formed.

Neutral Salts:

Salts of a strong acid and a strong base are neutral with pH value of 7.
Examples:

  1. NaCl: Obtained from strong acid, HCl and strong base NaOH.
  2. K2SO4: Obtained from strong acid, H2SO4 and strong base KOH.

Acidic Salts:

Salts of a strong acid and weak base are acidic with pH value less than 7.
Examples:

  1. NH4Cl: Obtained from strong acid, HCl and weak base NH4OH.
  2. (NH4)2SO4: Obtained from a strong acid, H2SO4 and weak base NH4OH.

Basic Salts:

Salts of a strong base and weak acid are basic in nature, with pH value of more than 7.
Examples:

  1. CaCO2: Obtained from a weak acid, H2CO3, and strong base, Ca(OH)2–
  2. CH3COONa: Obtained from a weak acid, CH3COOH, and strong base, NaOH.

Salts Definitions, Equations and Examples

Common Salt:

The salt formed by the combination of hydrochloric acid and sodium hydroxide solution is called sodium chloride and is the salt that is used in food.

Preparing Common Salt

  1. Seawater contains many salts dissolved in it. Sodium chloride is separated from these salts by the process of evaporation.
  2. Deposits of solid salt are also found in several parts of the world. These large crystals are often brown due to impurities. This is called rock salt. Beds of rock salt were formed when seas of bygone ages dried up. Rock salt is mined like coal.

Chemicals Obtained from Common Salt:

Common salt is an important raw material for various materials of daily use, such as sodium hydroxide, baking soda, washing soda, bleaching powder and many more.

Sodium Hydroxide:

When electricity is passed through an aqueous solution of sodium chloride (called brine), it decomposes to form sodium hydroxide. The process is called the chlor-alkali process because of the products formed- chlor for chlorine and alkali for sodium hydroxide.

2NaCl(aq) + 2H2O(l) → 2NaOH(aq) + Cl2(g) + H2(g)

Chlorine gas is given off at the anode, and hydrogen gas at the cathode.
Sodium hydroxide solution is formed near the cathode.

Salts Definitions, Equations and Examples

Uses of sodium hydroxide:

  • It is used for making soaps and detergents.
  • It is used for making artificial textile fibres such as rayon.
  • It is used in the manufacture of paper.
  • It is used in purifying bauxite ore from which aluminium metal is extracted.
  • It is used in degreasing metals, oil refining, and making dyes and bleaches.

Uses of chlorine:

  • It is used to sterilize the drinking water supply and the water in swimming pools as it is a disinfectant.
  • It is used in the production of bleaching powder.
  • It is used in the production of hydrochloric acid.
  • It is used to make plastics such as PVC (polyvinyl chloride), pesticides, CFCs, chloroform, paints and dyestuffs.
  • It is used for making solvents for dry cleaning (such as trichloroethane).

Uses of hydrogen:

  • It is used in the hydrogenation of oils to obtain solid fats called vegetable ghee.
  • It is used in the production of hydrochloric acid.
  • It is used to make ammonia for fertilizers.
  • It is used to make methanol.
  • Liquid hydrogen is used as a fuel for rockets.

Uses of hydrochloric acid:

  • It is used for cleaning iron sheets before tin plating or galvanisation.
  • It is used in the preparation of chlorides such as ammonium chloride which is used in dry cells.
  • It is used in medicines and cosmetics.
  • It is used in textile, dyeing and tanning industries.
  • It is used in making plastics like PVC.

Salts Definitions, Equations and Examples

Bleaching Powder:

Bleaching powder is produced by the action of chlorine on dry slaked lime [Ca(OH)2]. Bleaching powder is represented as CaOCl2, though the actual composition is quite complex.
Ca(OH)2 + Cl2 → CaOCl2 + H2O

Properties of bleaching powder

  1. It is a white powder that gives a strong smell of chlorine.
  2. It is soluble in cold water.
  3. It reacts with dilute acids to produce chlorine.

Uses of bleaching powder:

  1. It is used for bleaching cotton and linen in the textile industry, for bleaching wood pulp in paper factories and for bleaching washed clothes in the laundry.
  2. It is used as an oxidising agent in many chemical industries; and
  3. It is used for disinfecting drinking water to make it free of germs.

Salts Definitions, Equations and Examples

Example 1.
Name the substance which on treatment with chlorine yields bleaching powder.
Answer:
When dry slaked lime or calcium hydroxide is treated with chlorine gas, we get bleaching powder.
Ca(OH)2 + Cl2 → CaOCl2 + H2O

Baking Soda:

The chemical name of the compound is sodium hydrogen carbonate (NaHCO3). It is produced using sodium chloride as one of the raw materials.
NaCl + H2O + CO2 + NH3 → NH4Cl + NaHCO3

Properties of baking soda:

  1. It consists of white crystals which are sparingly soluble in water.
  2. It is a mild, non-corrosive base.
  3. Its solution in water is mildly alkaline.
  4. When heated, it decomposes to give sodium carbonate with the evolution of carbon dioxide gas.
  5. The following reaction takes place when it is heated during cooking:
    2NaHCO3 → Na2CO3 + H2O + CO2

Uses of sodium hydrogen carbonate:

1. For making baking powder, which is a mixture of baking soda (sodium hydrogen carbonate) and a mild edible acid such as tartaric acid. When baking powder is heated or mixed in water, the following reaction takes place:

NaFICO(aq) + FI + (from tartaric acid) → CO2(g) + H2O(l) + Sodium salt of acid
Carbon dioxide produced during the reaction causes bread or cake to rise to make them soft and spongy.

2. Sodium hydrogen carbonate is also an ingredient in antacids. Being alkaline, it neutralises excess acid in the stomach and provides relief.

3. It is also used in soda-acid fire extinguishers.

Salts Definitions, Equations and Examples

Example 2.
What will happen if a solution of sodium hydro carbonate is heated? Give the equation of the reaction involved.
Answer:
When a solution of sodium hydrogen carbonate is heated, carbon dioxide is liberated along with the formation of sodium carbonate and water. The equation of the reaction involved is given below:
Salts 1
Washing soda
Sodium carbonate is obtained by heating baking soda and recrystallisation of sodium carbonate gives washing soda (Na2CO310H2O).
Na2CO3 + 10H2O → Na2CO3.10H2O

Properties of Washing Soda:

  1. It Is a transparent crystalline solid.
  2. It is soluble in water.
  3. The solution of washing soda in water is alkaline which turns red litmus blue.
  4. It has cleansing properties due to which it is used in detergents.

Uses of washing soda:

  1. Sodium carbonate (washing soda) is used in glass, soap and paper industries.
  2. It is used in the manufacture of sodium compounds such as borax.
  3. Sodium carbonate can be used as a cleaning agent for domestic purposes.
  4. It is used for removing the permanent hardness of the water.

Salts Definitions, Equations and Examples

Example 3.
Name the sodium compound which is used for softening hard water.
Answer:
The sodium compound which is used for softening hard water is sodium carbonate (Na2CO3 10H2O).

Plaster of Paris:

On heating gypsum at 373 K, it loses water molecules and
becomes calcium sulphate hemihydrate (CaSO4 \(\frac{1}{2}\) H2O).
Salts 2
This is called Plaster of Paris, the substance which doctors use as plaster for supporting fractured bones in the right position.

Properties of plaster of Paris:

  1. It is a white powder.
  2. It has a property of setting into a hard mass on wetting with water which is due to its hydration to form crystals of gypsum which set to form a hard solid mass. Plaster of Paris should therefore be stored in moisture-proof containers.
    CaSO4. 2 H2O + 1 \(\frac{1}{2}\) H2O → CaSO4.2H2O

Note: Only half a water molecule is shown to be attached as water of crystallisation. It is written in this form because two formula units of CaSO4 share one molecule of water.

Uses of Plaster of Paris:

  1. Plaster of Paris is used for making toys, materials for decoration and for making surfaces smooth.
  2. It is used for setting fractured bones in the right position by doctors.
  3. It is used as a fire-proofing material.
  4. It is used in chemical laboratories for sealing air gaps in apparatus.

Salts Definitions, Equations and Examples

Example 4.
Plaster of Paris should be stored in a moisture-proof container. Explain why?
Answer:
The plastic of Paris should be stored in moisture-proof containers because if it comes in contact with water, it sets into a hard solid mass, Gypsum.

The equation for the reaction taking place is:
CaSO4 \(\frac{1}{2}\) H2O + 1 \(\frac{1}{2}\) H2O → CaSO4.2H2O

Water of Crystallization:

The water of crystallisation is the fixed number of water molecules present in one formula unit of salt. Five water molecules are present in one formula unit of copper sulphate. The chemical formula for hydrated copper sulphate is CuSO4.5H20.

Copper sulphate crystals that seem to be dry contain water of crystallisation. When we heat the crystals, this water is removed and the salt turns white.
CUSO4.5H2O → CuSO4 + 5H2O

The dehydration of copper sulphate crystals is a reversible process. When water is added to anhydrous copper sulphate, it gets hydrated and turns blue:
CuSO4 + 5H2O → CuSO4

Some examples of hydrated salts:

  1. Copper sulphate – CuSO4.5H2O
  2. Sodium carbonate (washing soda) – Na2CO3.10H2O
  3. Calcium sulphate (Gypsum) – CaSO4.2H2O
  4. Iron sulphate – FeSO4.H2O

Class 10 Science Notes

Strength of Acid and Bases Definitions, Equations and Examples

Strength Of Acid And Bases

The strength of acids and bases depends on the number of H+ ions and OH– ions produced, respectively.
More the number of H+ ions produced in solution, stronger is the acid and more the number of OH– ions produced in solution, stronger is the base.

If we take hydrochloric acid and acetic acid of the same concentration, say one molar, then these produce different amounts of hydrogen ions.

The pH Scale:

It is a scale for measuring hydrogen ion concentration in a solution. The p in pH stands for ‘potenz’ in German, meaning power. On the pH scale we can measure pH from O (very acidic) to 14 (very aLkaLine).

Higher the hydronium ion concentration, the lower is the pH value. The pH of a neutraL solution is 7.

Acidic nature increasling Neutral Basic nature increasing
Strength of Acid and Bases Definitions, Equations and Examples 1
Important
The solution having Lower pH will have more hydrogen ion concentration.

pH of some Common Substances
Strength of Acid and Bases Definitions, Equations and Examples 2

Strength of Acid and Bases Definitions, Equations and Examples

Example 1.
Do basic solutions also have H+(aq) ions? If yes, then why are these basic?
Answer:
Yes, basic solutions also have hydrogen ions. They are basic because in such solutions, the concentration of hydroxide ions is much greater than the concentration of hydrogen ions.
Higher the concentration of hydroxide ions, more basic is the solution.

Example 2.
Five solutions A, B, C, D and E when tested with universal indicator showed pH as 4, 1, 11, 7 and 9, respectively. Which solution is
(A) neutral?
Answer:
Solution D is neutral as it has pH value equal to 7.

(B) strongly alkaline?
Answer:
Solution C is strongly acidic as it has pH value greater than 7 and closer to 14

(C) strongly acidic?
Answer:
Solution B is strongly acidic as it has pH value less than 7 and closer to 0.

(D) weakly acidic?
Answer:
Solution A is weakly acidic as it has pH value less than 7 and closer to 7.

(E) weakly alkaline?
Arrange the pH in increasing order of hydrogen-ion concentration.
Answer:
We can arrange the solutions alon,g with their pH values as:
(A) 4
(B) 1
(C) 11
(D) 7
(E) 9
Solution E is weakly basic as it has pH value greater than 7 and closer to 7.

As pH of a solution varies inversely as the H+ ion concentration, the pH value in increasing order of H+ ion concentration is: C < E < D < A < B

Universal Indicator:

A universal indicator is a mixture of several indicators or dyes which gives different colours at different concentrations of hydrogen ions in a solution.
The colours produced by the universal indicator at various values of pH are given below:
Strength of Acid and Bases Definitions, Equations and Examples 3

Importance of pH in Everyday Life:

Plants and animals are pH-sensitive: Our body works within the pH range of 7.0 to 7.8. Living organisms can survive only in a narrow range of pH change.

Acid Rain: When pH of rainwater is less than 5.6, it is called acid rain. When acid rain flows into the rivers, it lowers the pH of the river water which makes the survival of aquatic life in such rivers very difficult. pH of soil Plants require a specific pH range for their healthy growth. Farmers often treat their soil with basic substances such as quick lime (calcium oxide) or slaked lime (calcium hydroxide) if the soil is found to be acidic.

pH in our digestive system and acidity: Our stomach produces hydrochloric acid which helps in the digestion of food without harming the stomach. During indigestion the stomach produces too much acid which causes pain and irritation. People use bases called antacids to get rid of this pain by neutralizing the excess acid.

Example of antacid: Magnesium hydroxide (Milk of magnesia)

pH of milk:

The pH of fresh milk is 6, which is slightly acidic. The milkman, therefore, adds a small amount of sodium hydrogen carbonate, which is basic in nature, so as to shift the pH of fresh milk to make it slightly alkaline so as to preserve the fresh milk for longer time.

Strength of Acid and Bases Definitions, Equations and Examples

Example 3.
Fresh milk has a pH of 6. How do you think the pH will change as it turns into curd? Explain your answer.
Answer:
The pH of fresh milk will decrease from 6 as it turns into curd as lactic acid is produced during curd formation, which lowers the pH of fresh milk.

Example 4.
A milkman adds a very small amount of baking soda to fresh milk.
(A) Why does he shift the pH of the fresh milk from 6 to slightly alkaline?
Answer:
As the fresh milk has a pH of 6, it is slightly acidic. The milkman therefore adds a very small amount of a basic substance such as baking soda to fresh milk to shift the pH from 6 to slightly alkaline so that milk does not turn sour easily.

(B) Why does this milk take a long time to set as curd?
Answer:
This milk is slightly alkaline due to which it will take longer to form curd and become sour since its pH value is now greater than 7.

pH change as the cause of tooth decay: Tooth decay starts when the pH of the mouth is lower than 5.5. Tooth enamel, made up of calcium phosphate is the hardest substance in the body. Bacteria present in the mouth produce acids by degradation of sugar and food particles remaining in the mouth after eating. Using toothpastes, which are generally basic, for cleaning the teeth can neutralise the excess acid and prevent tooth decay.

Self-defence by animals and plants through chemical warfare: Bee-sting leaves an acid which causes pain and irritation. Use of a mild base like baking soda on the stung area gives relief. Stinging hair of nettle leaves inject methanoic acid causing burning pain.

pH of plants: The leaves of Nettle, which is a herbaceous plant that grows in the wild, have stinging hair, which cause painful stings when touched accidentallydue to the methanoic acid secreted by them. However, nature provides neutralization options in the form of dock plant, which often grows beside the nettle in the wild, whose leaf is rubbed on the stung area to provide relief from pain and irritation.

Example 5.
Under what soil condition do you think a farmer would treat the soil of his fields with quick lime (calcium oxide) or slaked lime (calcium hydroxide) or chalk (calcium carbonate)?
Answer:
A farmer would treat the soil of his fields with quicklime or slaked lime or chalk if the soil has become acidic as all these fertilizers are basic. This will neutralize the acidic nature of the soil.

Class 10 Science Notes