Chemical Changes A chemical change is a permanent change in which the original substance, with its own composition and properties, gives rise to one or more new substances with different composition and properties, gives rise to one or more new substances with different composition and properties. For example, when a piece of iron is left in moist air it slowly changes into a brittle, brown substance known as iron rust. Experiments have shown that iron rust is not iron. It is hydrated iron oxide. It is a new substance and entirely different from iron. This is chemical change. Iron is grey, hard solid with a metallic lustre. It is also a good conductor of heat and electricity. If hydrochloric acid is poured on iron, hydrogen is produced. On the other hand, iron rust is a dull brown brittle substance. It conducts heat very poorly, does not conduct electricity and liberates no hydrogen when hydrochloric acid is poured on it. Thus, rusted iron is a new substance and the formation of rust on the iron is an example of chemical change. Characteristics of a Chemical Change:

1. A chemical change is a permanent change. This means that it cannot be reversed back by ordinary means to get back the original substance. For Example, it is not possible to change back curd into milk after the milk has been converted into curd.

2. One or more new substances (called products) are formed. The chemical properties of these new substances are different from those of the original substance. For example, the properties of CO2 formed by burning carbon are entirely different from those of carbon.

3. A change in the mass of specific substances can take place. In chemical change, the mass of the substances may change, however, the total mass of the substances taken together remains unchanged. For example, the weight of iron rust is more than that of iron from which it is produced. But the weight of iron, oxygen and water taken together remains the same.

4. An exchange of 3energy is generally involved in a chemical change. for Example, when wax ( in the form of a candle) is burnt a chemical change takes place, and carbon dioxide and water vapour are formed. But heat is released during the reaction (exothermic reaction). Experiments to show chemical changes

Experiment 1: When wood is burnt in the air, it gets decomposed to give carbon dioxide and water vapour. Heat and light (energy) were also produced. Impurities which are non-volatile remain behind.

C6H10O5 + 6O2 → 6O2 + 5H2O + Heat

(wood) (from air) Experiment 2: Hold a piece of magnesium ribbon with a pair of tongs and take it to a

flame. It will burn with a dazzling light and change into a white powder. The white powder will have a different mass and it will not burn if put over a flame (See Figure).

Magnesium oxide weighs more than the magnesium from which it is made.

2Mg + O2 → 2MgO

However, if the mass of oxygen is also taken into account, the total mass before and after the reaction remains the same

Experiment 3: Keep some milk in a dish and leave it overnight in a warm place (at room temperature). It will curdle, i.e. change into curd. You cannot change back curd into milk.

Experiment 4: Pass electric current through acidulated water in a voltmeter. Two new substances (gases) are formed, i.e., hydrogen and oxygen. These two substances do not recombine to give back water when the flow or electric current is stopped.

2H2O → 2H2 + O2

Experiment 5: Take some raw bananas and place them on your dining table for a few days. By the end of the few days your bananas will have ripened. They become black. Chemical change has taken place. You cannot ‘unripe’ the bananas.

Experiment 6: Take some sugar crystals in a dry test tube and heat them. The sugar crystals will melt and then turn

brown, and finally change into amorphous black carbon. During this change, you will notice the formation of a colourless liquid (water) on the cooler parts of the test tube. Heat changes sugar into carbon and water. The two, however, do not combine to form sugar.

C2H22O11 → 12C + 11H2O

∆

Sugar → Carbon + Water

Some more examples of chemical change 1. Burning of a matchstick 2. Digestion of food 3. Bursting of a cracker 4. Photosynthesis in plants 5. Burning of paper 6. Preparing ‘ghee’ from butter 7. Decomposition of leaves 8. Slaking of lime 9. Burning of LPG gas in a stove 10. Change of plain bread into toast

Photosynthesis: It is the process of conversion of water and carbon dioxide into carbohydrates by plants in presence of chlorophyll and sunlight.

Sunlight

6CO2 + 6H2O → C6H12O6 + 6 CO2

(Chlorophyll) Examples of substances which undergo physical as well as chemical change

1. Burning of candle. Candles are made of wax and a cotton thread (wick of the candle). As the candle burns, (a) The wax melts (melting of wax is a physical change) and resolidifies later. (b) The melted wax undergoes combustion as the cotton athread gets burnt. The combustion produces carbon

dioxide and water vapour. (The burning of wax and the cotton thread is a chemical change.) Thus, we observe that the burning of a candle is both a chemical and physical change.

2. Sublimation of ammonium chloride. When ammonium chloride is heated following changes take place: (a) Solid ammonium vapourises. (Physical change, no new substance produced) (b) Partial dissociation of ammonium chloride occurs. (Chemical change, new products are formed) (c) Ammonia and hydrochloride recombine to give ammonium chloride (Chemical change, new products is

formed) (d) Ammonium chloride vapours soldify as a sublimate. (Physical change, no new substance produced)

NH4 ⇌ NH3 + HCl

However, the entire change is considered to be a physical change. 3. Zinc oxide shows both physical and chemical change when heated.

First, it decomposes into zinc oxide and water vapour thereby showing a chemical change. But zinc oxide, i.e., the solid residue, which is yellow when hot, turns white on cooling. This shows only a phyical change. ∆

Zn (OH)2 → ZnO + H2O (chemical change)

∆

ZnO ⇌ ZnO (physical change)

(white) (yellow)

4. On heating sodium nitrate, it first melts, i.e., a physical change takes place. It then decomposes to give two new products, sodium nitrate and oxygen. This is a chemical change. ∆

NaNO3 ⇌ NaNO3 (physical change)

∆

2NaNO3→2NaNO2 + O2 (chemical change)

Table 5.1 Differences between Physical Change and Chemical Change

Physical Change Chemical Change

1. No new substances are formed as the composition of original substance remains unaltered.

2. The change is temporary and reversible. 3. There is no change in the composition of of

the new substance formed. 4. There is no change in the energy except a

change in state. 5. There is no change in the total mass.

1. New substances are formed as the composition of the reactants changes.

2. The change is permanent and irreversible. 3. Composition of the new substances formed is

different. 4. Energy change is generally involved in a chemical

change. 5. The respective masses of the substances either

increase or decrease. However, the total mass remains unchanged.

Conditions necessary for a chemical reaction to take place A chemical reaction occurs when either one or more of these conditions are met:

1. When two substances are mixed in their solid state, a chemical reaction occurs in some cases. For example, if you mix iodine and phosphorus an explosive reaction occurs.

2. A chemical reaction occurs in some cases when you mix certain substances in solution form. For example, oxalic acid crystals and sodium carbonate react in water solution only.

3. Some chemical reactions occur only by heating. For example, copper carbonate decomposes on heating into copper oxide and carbon dioxide.

4. Some chemical reactions take place by the action of light. These reaction are called photochemical reactions. For example, plants form glucose from carbon dioxide and water in the presence of chlorophyll and sunlight.

5. Some chemical reactions take place when pressure is also applied. this (i.e., the application of pressure) results in a quick reaction. For example, a mixture of KCIO3 and sulphur explodes when rubbed under pressure. another common example of reaction under pressure is how fire crackers explode when thrown against a wall or hammered forcefully.

6. Some reactions take place when there is an increase in the concentration of reaction. In other words, the higher the concentration of the substances the quicker the reaction.

7. Some chemical reactions occur only when electricity is passed through the substance. For example, electrolysis of water occurs (in the presence of electricity) to give hydrogen and oxygen.

8. Some chemical reactions need a catalyst to accelerate or deaccelerate their rates. However, catalysts do not take part in the reaction. For example, potassium chlorate by itself reacts only at 700˚C and even then, the rate of release of oxygen is very slow. But when it is heated with manganese dioxide, decomposition begins at a much lower temperature, i.e., at 300°C but the manganese dioxide remains unaffected. Thus, in this reaction manganese dioxide acts as a catalyst.

MnO2

2KClO3 → 2KCl + 3O2

9. The rate of reaction also depends on the increase in concentration of the reactants, i.e., a higher concentration of either or4 all of the reactants means a faster reaction. This can be shown by the following experiment: Prepare dilute and concentrated solutions of sodium thiosulphate (Na2S2O3) and take 5 ml of each of these solutions in two separate test tubes. Now, add 1 ml of HCl to each of them. Shake the solutions and keep them for some time. You will notice that the milkiness (due to the formation of collidial sulphur) appears first in the test tube that contains the concentrated solution of Na2S2O3.

Types of Chemical Changes Any chemical change in matter whether by combination, decomposition, displacement or by rearrangement, which involves the transformation of a substance into one or more new substances is termed as chemical reaction. The substances that undergo a chemical change are called reactants and the substances produced as a result of the chemical change are called products. The properties of the products are altogether different from those of the reactants. A breaking of bond between the atom of the reactants and the formation of new bonds between the atoms of the products take place to produce new substances. 1. Direct Combination or Synthesis Reaction

A combination reaction is a reaction in which two or more elements or compounds combine together to form a single compound.

A synthesis reaction is a reaction is a reaction in which two elements combine together to form a new compound. Examples of combination reaction 1. Take some black lead sulphide in a test tube and heat it. The black sulphide will combine with oxygen to form

white lead sulphate. PbS + 2O2 → PbSO4

2. Put a mixture of fine iron powder and sulphur in a test tube and heat it strongly. A large amount of heat will be produced. In the test tube there will be signs of melting. Iron sulphide ( a greyish black solid substance) will be formed.

Fe + S → FeS 3. Hold a piece of magnesium ribbon over a flame. It will burn with a dazzling light forming magnesium oxide.

2Mg + O2 → 2MgO (synthesis) More examples of combination reaction: (a) Reaction of metals with oxygen to form oxides:

∆ 4K + O2 –––––>2K2O ∆ 2Pb + O2 –––––> 2PbO (yellow) ∆ 2Cu + O2 –––––> 2CuO (black) ∆ 2Zn + O2 –––––> 2ZnO (yellow when hot) ∆ 3Fe + 2O2 –––––> Fe3O4

(b) Reaction of non-metals with oxygen to form oxides: ∆ S + O2 –––––> SO2 ∆ P4 + 5O2 –––––> 2P2O5 electric spark

N2 + O2 –––––––––––––> 2NO

(c) Reaction of metals and non-metals to form chlorides: sunlight H2(g) + Cl2(g) –––––––––> 2HCl(g) ∆ 2Na + Cl2 ––––––––> 2NaCl ∆ Cu + Cl2 ––––––––> CuCl2 (blue)

(d) Reaction of elements to form sulphides: ∆ 2Na + S –––––––> Na2S ∆

Fe + S –––––––> FeS (black) ∆ Zn + S –––––––> ZnS (white)

(e) Reaction of elements to form nitrides: 450°C(Fe)

N2(g) + 3H2 (g) ⇌ 2NH3 + H

500 atm ∆ 3Ca + N2 –––––––> Ca3N2

∆ 2Al + N2 –––––––> 2AlN

2. Decomposition or Dissociation Reactions: This type of reaction is the breaking down of a compound into either elements or simpler compounds such that these products of decomposition do not recombine to form the original compound. A decomposition reaction may be brought about by the presence of either heat or light or the passage of electricity. Such type of reaction that is brought about by heat is known as thermal decomposition Examples of thermal decomposition reaction: (1) The decomposition of metal oxides by heat:

∆ 2HgO ––––––> 2Hg + O2 ∆

2Ag2O ––––––> 4Ag + O2 ∆ 2Pb3O4 ––––––> 6PbO+ O2

(2) The decomposition of metal hydroxides by heat: ∆ Ca(OH)2 –––––> CaO + H2O ∆

Cu(OH)2 –––––> CuO + H2O ∆ 4 Ag(OH) –––––> 4Ag + O2 + 2H2O

(3) The decomposition by heat of metal nitrates: Alkali metal nitrates decompose on heating to nitrites and oxygen. Other nitrites decompose on heating to the oxides of the metal liberating nitrogen dioxide and oxygen except the nitrites of noble metals which give metal, nitrogen dioxide and oxygen.

∆ 2KNO3 ––––––––> 2KNO2 + O2 ∆ 2Pb(NO3)2 –––––––> 2PbO + 4NO2 + O2 (colourless) (yellow) (Brown fumes) ∆ 2AgNO3 ––––––––> 2Ag + 2NO2 + O2 (4) Decomposition of carbonates by heat:

All carbonates of metals decompose except the carbonates of alkali metals, liberating carbon dioxide and the oxide of metal. Carbonates of mercury and silver give metal, nitrogen dioxide and oxygen. ∆ CaCO3 –––––––––> CaO + CO2 ∆ CuCO3 –––––––––> CuO + CO2

∆ 2Ag2CO3 –––––––> 4Ag + O2 + 2CO2

(5) Decomposition of sulphates by heat: Certain metal sulphates decompose on heating forming products depending on the nature of the metal. ∆ 2ZnSO4 ––––––> 2ZnO + 2SO2 + O2

∆ 2FeSO4 ––––––> Fe2O3 + SO2 + SO3 ∆ Al2(SO4)3–––––> Al2O3 + 3SO3 Experiments to show decomposition reaction: (a) Take some lead nitrate crystals in a test tube and heat them. The crystals will begin to melt. As you carry on heating the crystals further, they will give out both nitrogen dioxide, which is a reddish brown gas, and oxygen. A yellow solid (lead monoxide) will be left in the test tube. ∆ 2Pb(NO3)2––––––> 2PbO + 4NO2 + O2 (colourless) (yellow) (Brown fumes)

(b) Heat some potassium nitrate crystals in a test tube. The crystals will begin to melt and slowly give out oxygen which

will rekindle a glowing splinter.

2KNO3 △ 2KNO2 + O2

(c) Put some zinc carbonate in a test fitted with a cork and bent glass tube. On heating, carbon dioxide is given out

which will turn lime water milky. The residue, i.e., zinc oxide is yellow when hot but turns white on cooling.

ZnCO3 △ ZnO+CO2

(d) Heat ammonium dichromate in a test tube. Upon heating, it decomposes into chromium oxide, water vapour

and nitrogen.

(NH4)2Cr2O7 △ Cr2O3 + 4H2O+N2

3. Reversible Thermal Decomposition Reaction

These are reactions in which the products formed by hating the reactants recombine on cooling.

Experiments to show thermal decomposition (also called thermal dissociation) reactions.

(a) Heat some solid ammonium chloride in a test tube. Two colourless gases are produced which combine to

reform ammonium chloride(which appears as white sublimate on the cooler parts of the test tube).

heat

NH4CI ⇌ NH3↑ +HCI ↑

cool

(b) Nitrogen tetroxide, on heating, changes to nitrogen dioxide which is a reddish brown gas. On cooling, nitrogen

dioxide changes into the original compound, nitrogen tetraoxide.

heat

N2O4 ⇌ 2NO2

cool 4. Electrolytic Decomposition Reactions

In these reactions, electricity is used for decomposing the reactant molecules:

NaCI(molten) �������� �������

(������������)> Na(cathode)+

�

� CI2(at anode)

2H2O(acidified) �������� �������

(������������ ������������)> 2H2 (at cathode) + O2 (at anode)

AI2O3(molten)�������� �������

>2AI (at cathode) +�

� O2(at anode)

5. Photo Decomposition Reactions

In these reactions, the decomposition of the reactant molecules is brought about by light energy (sunlight).

I2(g) Sunlight 2I(g)

AgBr Ag (in the form of particles)+Br

light reflected form

the objet to be photographed

(coated on photographic plate)

6. Displacement (or Substitution) Reactions

A displacement reaction is a chemical change in which a more active element displaces the less active element

from a compound.

This reaction can be presented by the following general chemical equation.

C+AB CB+A

Here AB is a compound of A. The element C is more reactive than element A and hence displaces it from its

compound. AB, Thus, C is the displacing element and A is the element which is being displaced by C.

Examples of displacement reactions

1. Displacement of hydrogen from water or steam by more reactive metals. Sodium, potassium and calcium

displace hydrogen with cold water. Zinc, iron, etc., liberate hydrogen from steam.

2K +2H2O 2KOH+H2

Ca +2H2O Ca(OH)2 +H2

Zn + H2O ZnO +H2

3Fe +4H2O ⇌ Fe3O4 +4H2

2. Displacement of H2 from acids by more reactive metals. Potassium, sodium and calcium displace hydrogen

from dilute mineral acids with violent explosion. Magnesium, aluminium and zinc reacts vigorously and slowly.

2Na +2HCI 2NaCI +H2

Ca +2HCI CaCI2 +H2

Mg +2HCI MgCI2 +H2

Fe+ 2HCI FeCI2 +H2

3. Displacement of less reactive metal from its salt solution by a more reactive metal. More reactive metals

displace less reactive metals from the aqueous solution of its salt.

Fe +CuSO4 FeSO4 +Cu

Zn+ FeSO4 ZnSO4+Fe

4. Displacement of less reactive metal from its oxide by a more reactive metal. When an active metal is heated

with the oxide of a less active metal, it forms its own oxide.

Mg+Zno MgO+Zn

2AI+Fe2O3 AI2O3+2Fe

5. Displacement of the less reactive halogen (non-metallic elements) from its metallic halide by more reactive

halogen. Chlorine displaces bromine and iodine form bromides and iodides respectively. Bromine displaces iodine

from iodides.

CI2+2KBr 2KCI +Br2

CI2+2NaI 2NaCI +I2

Br2+2NaI 2NaBr+I2

Experiments to show displacement reactions

(a) Take a solution of copper sulphate in a beaker, add a few pieces of zinc and stir with a glass rod. The blue

colour of the solution will gradually fade and soon the solution will become colourless. At the same time, reddish

brown particles of copper will settle down in the beaker.

CuSO4+Zn ZnSO4+Cu

(b) Take some dilute sulphuric acid in a test tube and drop a small piece of magnesium ribbon. A very brisk

effervescence will take place an hydrogen will evolve. This will burn with a ‘pop’ sound when you bring a burning

match stick near the mouth of the test tube.

Mg+H2SO4 MgSO4+H2

(c) Bubble chlorine through a solution of potassium iodide. The solution will turn brown due to the displacement

of iodine.

2KI+CI2 2KCI + I2

7. Double Decomposition (also called Double displacement)

These are reactions in which two compounds in a solution react to form two other compounds by exchanging

their radicals. usually a solid is formed as result of the reaction

This reaction can be represented by the following general equation:

AB+CD AD+CB

Double decomposition reactions are of the following types:

(i) Precipitation reactions : In these reactions, the aqueous solutions of the two compounds react together by

exchanging their radicals to give one of the products as a precipitate.

NaCI(aq)+AgNO2 NaNO3(aq) +AgCI(s)↓

Na2SO4 +Pb(NO3)2(aq) 2NaNO3(aq) + PbSO4(s)↓

FeCI2(aq) + 2NaOH(aq) 2NaCI(aq) + FeOH2↓

(dirty green ppt.)

BaCI2 +H2SO4(aq) 2HCI +BaSO4(s) ↓

(white ppt.)

CuSO4(aq) + H2S(aq) H2SO4(aq)+CuS(s)↓

(black ppt.)

(ii) Neutralisation reactions: In these reactions, an acid and a base (or basic oxide) react together by exchanging

their radicals, forming salt and water. The reaction takes place because the hydrogen ion (H+) from the acid

combines with the hydroxyl ion (OH-)from the base to form water.

NaOH + HCI NaCI +H2O or Na+ OH- + H+CI- Na+OH-CI-+H+OH-

Cancelling the common ions,Na+ and CI- the only change is the combination of H+ and OH- ions to form un-

ionized water, i.e., H+ +OH- H2O

Acid Base of Basic Oxide Salt Water HCI + H2SO4 + 2HNO3 + 2HCI +

NaOH 2KOH Ca(OH)2 CuO

NaCI + K2SO4 + Ca(NO3)2 + CuCI2 +

H2O 2H2O 2H2O H2O

Experiments to show double displacement reactions:

(a) Take a solution of silver nitrate in a test tube and add dilute hydrochloric acid or a solution of sodium chloride.

A white curdy precipitate will be formed.

AgNO3 + HCI AgCI ↓ + HNO3

AgNO3 + NaCI AgCI ↓ + NaNO3

(b) Fill one third of a test tube with a solution of dilute sulphuric acid and add a solution of barium chloride. A

thick white precipitate will be formed immediately.

BaCI2 + H2SO4 BaSO4 ↓ + 2HCI

(c) Heat some sodium nitrate crystals with concentrated sulphuric acid in a test tube. Sharp and pungent vapour of

nitric acid will be formed immediately.

BaCI2 + H2SO4 NaHSO + HNO3

Did you know? 1. If you are stung by a bee it hurts because of the formic acid entering your skin from the

bee’s sting. Don’t panic, rub the spot with slaked lime or baking soda both of which are bases.

2. One of the reasons we often have stomach ache is because our stomach glands secrete excess HCI. One home remedy available to relieve this kind of pain is milk of magnesia or sodium hydrogen carbonate both of which are bases.

3. Acid that is accidently spilled on your clothes or body can be neutralized with ammonia solution.

4. If soil is somewhat acidic and thus unfavourable for the growth of certain types of crops, lime is added to it to neutralize the excess acid. These are some of the uses of neutralization (reaction) as a chemical process.

8. Hydrolysis

This is the proces sby which a salt and water react in a limited way to form an acid and a base. You must

remember, however, that not all salt undergo hydrolysis. If a salt has been formed by the action of a strong acid on

a strong base, it will not undergo hydrolysis (process of decomposition) For instance, sodium chloride, formed

from the strong base NaOH and a strong acid HCI, does not hydrolyse and its solution does not have any effect on

litmus. But ammonium acetate, formed from the weak base NH4OH and weak acid, acetic acid. does hydrolyse: but

its solution has no effect on litmus.

Whenthe acid is strong (HCI2 H2SO4 HNO3) and the base is weak, the hydrolysis of their salts takes place and

blue litmus turns red.

FeCI3 + 3H2O Fe(OH)3 + 3HCI

(weak base) (strong acid)

ZnSO4 + 2H2O Zn(OH)2 + H2SO4

(weak base) (strong acid)

Solutions of iron (III) chloride and zinc sulphate are acidic.

On the other hand, if the acid is weak and the base is strong then hydrolysis of their salt takes place and red litmus

changes to blue.

NaCO3 + 2H2O 2NaOH + H2CO3

(strong base) (weak acid)

K3PO4 + 3H2O 3KOH + H3PO4

(strong base) (weak acid)

Solutions of sodium carbonate and potassium phosphate are alkaline.

Note: Strong acids and strong bases are those that ionize to a very large extent in aqueous solution. Weak acids and

weak bases are those that ionize to only a small extent in aqueous solution.

9. Oxidation and Reduction Reactions

In a chemical reaction, where one substance is oxidised, the other substance must necessarily be reduced. This is

because the electrons lost during oxidation are simultaneously gained during reduction and vice versa. Thus, the

half fractions.

Zn Zn2+ +2e-

Cu2+ + 2e- Cu

cannot occur in isolation, they occur simultaneously as Cu2++Zn Zn2+ + Cu

Thus, oxidation and reduction always go side by side. Such reactions are called oxidation reduction reactions or, in

other words, redox reactions.

Oxidation number: This is the number of electrons lost or gained by an atom of an element or an ion.

Examples of oxidation reactions

1. Addition of oxygen: Elements on burning in the air add oxygen to form oxides.

�������� �� ������

���������

�������� �� ��

2Mg +O2 2MgO C+O2 CO2

2. Removal of hydrogen: Chlorine, when reacting with hydrogen sulphide, removes hydrogen thus oxidising

hydrogen sulphide to sulphur.

������� �� ��������

���������

H2S+CI2 2HCI +S

3. Addition of electronegative radical or element: Chlorine reacts with ferrous chloride to form ferric chloride. Thus,

ferrous chloride is oxidized to ferric chloride.

addition of electronegative element

oxidation

2FeCl2+Cl2 2FeCl3

4. Removal of electropositive radical or element: Potassium iodide, when treated with hydrogen peroxide, loses

electropositive radical potassium as such oxidized to iodine

removal of electropositive radical:

oxidation

2KI + H2O 2KOH + I2

5. Loss of electrons: When an atom or iron loses electrons it is oxidized as its oxidation number increases.

(a) Zn Zn2+ + 2e- (0 +2)

Oxidation

In this reaction, Zn atom loses two electrons and hence is said to have been oxidized to Zn2+ ion.

The oxidation number of Zn increases from 0 to +2

(b) Fe2+ Fe3+ + e- (+2 +3)

Oxidation

Fe2+ ion loses one electron and is oxidized to Fe3+ ion. The oxidation number of Fe increases from +2 to +3.

(c) I- I + e- ( -1 0)

Oxidation

I- Ion loses one electron and gets oxidized to I atom. The oxidation number of I- increases from -1 to 0.

Examples of reduction reaction

1. Addition of hydrogen: Bromine when treated with hydrogen sulphide forms hydrogen bromide and sulphur. Since

hydrogen adds to bromine it is reduced to hydrogen bromide.

addition of hydrogen

reduction

Br2 + H2S 2HBr + S

2. Removal of oxygen: When hydrogen is passed over red hot lead oxide it is reduced to lead as oxygen is removed

by hydrogen to form steam ( water vapour).

removal of oxygen

reduction

PbO + H2 Pb + H2O

3. Addition of electropositive radical or atom: When iron is treated with iron (III) chloride, the latter is reduced to

iron (II) chloride.

addition of electropositive Fe atom

reduction

2FeCl3 + Fe 3FeCl2

4. Removal of electronegative radical or atom: When hydrogen sulphide is passed over heated iron (III) chloride it is

reduced to iron (II) chloride due to loss of electronegative chlorine atom.

removal of electronegative Cl- atom

2FeCl3 + H2S 2FeCl2 + 2HCl + S

5. Gain of electrons: When an atom or ion gains electrons it is reduced as its oxidation number decreases.

reduction

(a) Fe3+ + e- Fe2+ (+3 +2)

Fe3+ ion gains an electron and is reduced to Fe2+ ion. The oxidation number of Fe3+ ion decreases from +3 to +2

reduction

(b) Cu2+ + 2e- Cu (+2 0)

Cu2+ ion gains two electrons and is reduced to Cu- metal. The oxidation number of Cu2+ decreases from +2 to 0.

reduction

(c) F + e F- (0 1)

F atom gains one electron and is reduced to F- ion. In this reaction, the oxidation number of F atom decreases from 0

to -1.

Remember: Oxidation and reduction reactions are the reverse of each other and take place simultaneously.

We have seen from the above examples that oxidation involves loss of electrons while reduction involves gain of

electrons. Thus, oxidation and reduction reactions are the reverse of each other, i.e., if one atom or iron loses electrons,

(i.e., oxidises) in a reaction, the other atom or ion gains an equal number of atoms (i.e., reduces). Thus, oxidation and

reduction reactions take place simultaneously. Oxidation and reduction reactions never occur singly and every oxidation

reaction is accompanied by reduction reaction and vice versa.

Table 5.2 Differences between Oxidation and Reduction

Oxidation Reduction Oxidation is a chemical reaction which involves:

1. Addition of oxygen or any other electronegative atom or ion to a substance.

2. Removal of hydrogen or any other electro-positive atom or ion from a substance.

3. Loss of electrons from an atom or an ion. 4. Increase in oxidation number of the effective

element.

Reduction is a chemical reaction which involves: 1. Addition of a hydrogen or any other

electropositive atom or ion to a substance. 2. Removal of oxygen or any other electronegative

atom or ion from a substance. 3. Gain of electrons by an atom or an ion. 4. Decrease in oxidation number of the effective

element.

10. Reversible and Irreversible Reactions

Reversible reactions

These are reactions in which the products formed in the reaction react together to form the original reactants.

In a reversible reaction, both the reactants and the products are present in equilibrium with each other. The

concentraton of the reactants and products in these reactions depend on the degree of reversibility of the reaction. A

reversible reaction is represented by two half-arrows pointing in opposite directions (⇌). the conditions for the

forward and backward reactions may be the same or different.

Examples:

(a) Red hot iron reacts with steam and forms Fe3O4 and H2. It is a reversible reaction in which the forward and the

backward reactions take place under the same circumstances.

red hot

3Fe + 4H2O (steam) ⇌ Fe3O4 + 4H2

red hot

(b) CaCO3 react with the aqueous solution of CO2 in cold (at low temperature) to give Ca(HCO3)2. Ca(H2O3)2, on

being heated, gives the original products viz., CaCO3, H2O and CO2.

In Cold

CaCO3 + H2O + CO2 ⇌ Ca(HCO3)2

Heat

(c) NH4Cl(s) ⇌ NH3 (g) + HCl(g)

Cold

Heat (140˚C)

(d) N2O4(g) ⇌ 2NO2

Cool (20˚C)

Heat

(e) PCL ⇌ PCI2 + CI2

Cool

(f) H2 + I2 ⇌ 2HI

(g) 2SO2((g) + O2 (g) ⇌ 2SO3(g)

Irreversible reactions

These are reactions in which the reactants are completely converted into products.

In an irreversible reaction, the reactions proceed only in the direction towards the formation of such products

Examples

(a) 2Mg (s) + O2(g) –––––––> 2MgO(a)

(b) NaOH(aq)+HCI(aq) –––––––> NaCI(aq)

(c) NaCI(aq)+AgNO3(aq) –––––––>AgCI(g)+NaNO3(aq)

(d) C(s)+ O2(g) –––––––> CO2(g)

ENERGY CHANGES IN CHEMICAL CHANGE

Endothermic and Exothermic Reactions:

1. Endothermic reactions

These are eactions that take place onloy when heat energy is supplied from an external source

To the reactant molecules.

Examples:

N2(g) + O2(g) + Heat –––––––> 2NO(g)

C(s) + H2O(g) + Heat –––––––> CO(g) + H2(g)

Thus, in an endothermic reaction, heat energy is absorbed by the reactants. Therefore, in order to show these

reactions, the word of term ”heat” is written on the reactant side

Reactions + heat –––––––> Product (endothermic reaction)

2. Exothermic reaction

These are reactions in which the heat energy is evolved (or,liberated, or released or generated).

Examples:

C(s) + O2 (g) –––––––> CO2 + Heat

N2(g) + 3H2(g) –––––––> 2NH3(g) + Heat

Thus, in an exothermic reaction, heat energy is either absorbed or liberated or released. Therefore, in order to show

these reactions, (also called thermochemical reactions), the word or term ‘heat’ is written on the product side

Reactions –––––––> Products + heat (exothermic reaction)

3. Photochemical reactions

Chemical reactrions that take place only when light energy is supplied to the reactant molecules are called

photochemical reactions. In other words, in photochemical reactions the reactant molecules absorb light.

The rate of a photochemical reaction that takes place depends on the energy associated with the light absorbed by

the reactant molecules.

Examples:

Visible light

(a) Phtodissociation of I2(g) 2I(g)

(=499.1 nm)

Diffused sunlight

(b) Formation of HCI(g) H2 + CI2 ––––––––––––––> 2HCI

(slow reaction)Sunlight)

(c) Photosynthesis 6CO2 + 6H2 O –––––––> C6H12O6 + 6O2

Chlorophyll

(From atmosphere) (Present in the soil Glucose

or from atmosphere) (a carbohydrate)

Light reflected form the object

(d) Photodecomposition of silver bromide (AgBr) AgBr –––––––––––––––––––––––> Ag+Br

( Coated on to be photographed (in the form

photographic plate) of particles)

Note: AgBr and other silver salts are, therefore, kept in dark coloured bottles.

UV radiation

(e) Convertion of O2(g) into O3(g) 3O2 (g) –––––––––––––> 2O3 (g)

light

(f) Phtodecomposition of AgNO3 2AgNO3 –––––––––––––> 2Ag+2NO2 +O2

4. Electrochemical reactions

We have seen, from some of the above reactions, that reactant molecules react together only when heat or

light energy is supplied to them. However, there are many other reactions which take place only when electric energy is

supplied to the reactant molecules. Such reactions , that depend on electric energy, are called electrochemical

reactions.

Electrochemical reactions are defined as reactions which proceed when electric energy is absorbed by the

reactant molecules.

Examples:

Electric current

(a) Fused potassium chloride breaks into its charges particles(ions) 2KCI –––––––––––––>K + CI

(b) Acidulated water also breaks into hydrogen and oxygen

electric Current

2H2O –––––––––––––> 2H2 O2

(c) Electolytic decomposition of molten sodium chloride

electric cuttent

NaCI (molten –––––––––––––> Na(at cathode+I/2CI2(at anode)

(electrolysis)

(d) Electrolytic dissociation of acidified water electric current

sH2O (acidified) –––––––––––––> 2H2 (at cathode) + O2 (at anode)

electolyic dissociation)

(e) Electrolytic decomposition of molten AI2 O3

electric current

AI2O3(molten) –––––––––––––> 2AI(atcathode) + 3/2 O2(at anode)

BURNING OR COMBUSTION

Burning or combustion is chemical change in which combustible substances combine with oxygen to product new

compounds, called oxides with the liberation of large amounts of energy in the form of heat and light.

Oxygen in the air is mainly responsible for burning,

What happens when a substance burns?

1. New compounds, such as the respective oxides, are formed

2. Heat is released.

3. Light is generated or evolved as well.

Examples of Combustion

1. Combustion of metals:

(a) K, Na and AI burn in air or O2 at ordinary temperature and form their corresponding oxides

(b) Ca, Mg, Zn, Fe and Pb burn in O2 or air on strong heat and form their corresponding oxides

2M(M = Mg, Pb + O2 –––––––––––––> 2MO

2. Combustion of non-metals: Non-metals like P,S and C Combine with O2 to give their oxides

P4 +5O2 ––––––> 2P2 O5 + Heat

3. Combustion of hydrocarbons: All hydrocarbons born I O2 of the air to form water vapour and CO2. Heat is also

generated in these reactions

CH4 + 2O2 -–––––––––––––> CO2 + 2H2O + Heat

(methane) (from air) (water vapour)

2C2H6 + 7O2 –––––––––––––> 4CO2 + 6H2O + Heat

( ethane)

2C2H2 + 5O2 –––––––––––––> 4CO2 + 2H2O + Heat

Table 5.3 Combustible and Non-combustible Substances

Examples

Substances 1. Combustible 2.Non-combustible (incombustible)

Definition A Substance which can burn in air or oxygen is called a combustible substance A Substance which cannot burn in air or oxygen is called a non-combustible substance.

Solids Liquids Gases Candle wax,paper, Alcohol, ether Carbon coal,wood, coke petrol,kerosene Monoxide sodium phosphorus etc hydrogen hydrogen sulphide etc Silicon dioxide, Water, nitric Nirtogen alumina, common gold, sulphuric carbon sale, gold etc acid etc dioxide, cholorine oxygen nitrogen dioxide

Table 5.4 Supporter and Non-Supporter of Combustion

Term 1. Supporter of combustion 2. Non-supporter of Combustion

Definition It is the gaseous environment of supports combustion by allowing oxidation It is the gaseous environment that does not support combustion or oxidation

Examples Nitrous oxide,oxygen and chlorine Nirtogen, hydrogen, carbon dioxide,hydrogen chloride etc

Ignition Temperature of a Substance

As ignition temperature of a substance is the lowest temperature at which the substance catches fire and starts

burning. In other words, the ignition temperature is the minimum temperature to which the substance must be heated

in order to make it burn.

The ignition temperature is also called the Ignition point or kindling temperature.

(Note: To kindly means to set on fire)

For Example, since the ignition temperature of white phosphorus is 350 C it must be heated atleast 35˚Cto make it

catch fire and burn. The temperature of 35˚C is very low and hence when the room temperature reaches 35˚ C white

phosphorus catches fire and burns.

Types of Combustion

Combustion may be of the following types:

1. Slow Combustion: The combustion reaction in which only a small amount of heat is produced is called Slow

Combustion. Respiration and rusting of iron are examples of slow combustion.

2. Rapid Combustion: A combustion reaction in which a large amount of heat and light are evolved in a short period

of time is called rapid combustion. The immediate burning of domestic gas in a burner to give heat and light Is an

example of rapid combustion.

3. Spontaneous Combustion: A Combustion reaction in which no external heat is required to start. It is called

spontaneous combustion. In this combustion the substance undergoes slow oxidation at room temperature. The heat

liberated in this process rates the temperature until the ignition temperature of the substance is attained. Then, at

this ignition temperature, the substance bursts into flames which is rapid combustion. When white phosphorus is

exposed to air for some time it burns on its own. The burning of phosphorus in this way is called spontaneous

combustion of phosphosrus.

Experiment to show that a air is required for burning to take place

Light two candles (A and B) and mount each of them on two tiles. Cover one of the candles (B) with an inverted

jar. Candle B on one of the tiles continues to burn for some time but soon. Its flame gradually becomes shorter and

shorter and smoky till it finally goes off since it has used up all the oxygen in the jar. The candle on the other tile

continues to burn brightly with a steady flame since ti get a free supply of air.

Burning of Substance in the absence of air

We have seen from this experiment that, in order to burn a substance, it should be placed in an environment

which has plenty of iar or O2. Air is essential for burning. However, there are certain substances which can burn

even in the absende of air or O2 after the initial heating.

Example: (i) When gun powder, which contains nitre (KNO3), sulphur and charcoal (carbon is heated in the

absence of air, it burns spontaneously. Since the thermal decomposition of KNO3 gives O2, KNO3 forms an

explosive mixture with sulphur and carbon.

(ii) KClO3 and K2 Cr2 O7 readily give O2 on heating and hence form explosive

Mixtures with combustible substances.

2KClO3 –––––––––––––> 2KCl + 30

4K2Cr2O7 –––––––––––––> 4K2CrO4+2CR2O3+3O2

Burning of substances in an atmosphere of a gas other than air or O2

(i) H2 gas burns in Cl2 to form HCl

H2+Cl2 –––––––––––––> 2HCl + Heat

(ii) Burning Mg continues to burn in atmosphere of CO2. This happens because the heat produced by

burning of Mg Splits CO2 into carbon (black specks) and O2. The O2 thus produced is used in the

combustion of Mg.

2Mg + CO2 –––––––––––––>2MgO+2 (black Specks) + Heat

Factors on which the rate of combustion of substances depend

Following are the important factors:

1. Surface are of the combustible substances. The smaller the size of the particles of the combustible

substance, the larges is the area needed for the reaction to occur and hence faster is the rate of combustion.

For example, in many industrial processes, coal dust is used instead of coal chunks of pieces, since coal dust

provides larger surface are for the reaction of occur.

2. The amount or volume of the supporter of combustion. The volume of O2 and N2 is not. It is due to less

volume of O2 that combustion (burning) of substances takes place at a moderate and safe rate. If the volume

of O2 and N2 in air was 79% and 21%, then, due to the greater volume of O2, the combustion of substances

would have occurred at an uncontrollably rapid rate. Large fires would be produced and would ultimately

destroy life and property the world over.

3. Nature of combustible substances. The substances which have low ignition temperature are volatile and

inflammable and hence burn rapidly. Alcohol and petrol are examples of such substances. These substances

readily burst into flames.