Chemical Reaction and Equation Class 10 Notes Science

Chemical Reaction and Equation Class 10 Notes CBSE

Introduction

                 Change is the law of nature. We observe various types of changes around us. Plants grow into trees, a child grows into an adult, fruits ripen, water evaporates, water freezes in refrigerator, mercury rises in a thermometer on a hot day, iron articles rust in moist air, milk changes into curd, oil burns in stoves, sometimes a glass tumbler breaks, etc. Scientists classify these changes as physical and chemical changes.

Physical Change

                A physical change is one in which the substance undergoing the change is not destroyed. That is, the substance does not lose its identity. In a physical change, no new element or no new compound is formed. When the source responsible for the physical change is removed, the substance regains its original state.

Examples of Physical Changes

                i               Evaporation of a liquid (L \rightleftarrows V)                          

                ii             Sublimation (S \rightleftarrows V)

                iii            Melting of a solid (S \rightleftarrows L)                     

                iv            Dissolution of sugar or salt in water

                v              Powdering of sugar                                             

                vi            Mixing of iron particles and sand

Comments

                In the various example mentioned above, the substance undergoing change is not destroyed and no new substance is formed. Hence, each one of these changes is physical.

Chemical Change

                A chemical change is one in which the identity of the original substance changes and a new substance or new substances are formed. In a chemical change, the properties of the substances before and after the change are entirely different. The original properties disappear and new properties are observed. A chemical change has taken place during chemical reaction.

Examples of Chemical Changes

                Ripening of fruits, digestion of food, souring of milk, rusting of iron, burning of a candle, combustion of glucose in metabolic process, preparation of water from hydrogen and oxygen, photosynthesis by plants, functioning of cells and batteries, growth of a plant, etc.

Why are the above changes called chemical changes?

                In the example of chemical changes mentioned above, the original properties disappear. For example :

                (i)       The raw and ripe fruits have different colours and different tastes.

                (ii)      Before rusting, iron is a grey solid and is attracted by a magnet. After rusting, iron loses its properties. The rust formed is a brown-brittle powder and nonmagnetic.

Chemical Reaction

                Chemical reaction is a process in which some known substances are changed into new substance or new substances. The properties of new substances are different from the properties of the starting substances.

                In the language of chemistry,

                The starting substances are called reactants.

                The new substances formed are called products.

                A chemical reaction is represented by shorthand notation called chemical equation as

Reactants → Products

                The arrow (→) indicates the direction of the reaction that reactants are changed into products.

Examples of Some Chemical Reactions

Hydrogen and oxygen combine so that water is formed under suitable experimental conditions.

\mathop {Hydrogen}\limits_{({\mathop{\rm Re}\nolimits} ac\tan t)} \,\, + \,\,\mathop {Oxygen}\limits_{({\mathop{\rm Re}\nolimits} ac\tan t)} \,\, \to \mathop {Water}\limits_{(\Pr oduct)} \,\,\, + \,\,Heat

When zinc is placed in sulphuric acid, hydrogen gas is liberated and zinc sulphate is formed.

\mathop {Zinc}\limits_{({\mathop{\rm Re}\nolimits} ac\tan t)} \,\, + \,\,\mathop {Sulphuric{\rm{ acid}}}\limits_{({\mathop{\rm Re}\nolimits} ac\tan t)} \,\, \to \mathop {Hydrogen}\limits_{(\Pr oduct)} \,\,\, + \,\,\mathop {Zinc\;sulphate}\limits_{(\Pr oduct)} \,\, + \,\,Heat

Methane gas burns in oxygen so that carbon dioxide and water are produced and heat energy is generated.

Methane  +  Oxygen  → Carbon dioxide  + Water + Heat energy

When a solution of lead nitrate is mixed with a solution of potassium iodide, then yellow precipitate of lead iodide is formed.

\mathop {Lead\,\,nitrate}\limits_{(Colourless{\rm{ solution)}}} \,\, + \,\,\mathop {Potassium {\rm {iodide}}}\limits_{{\rm{(Colourless solution}})} \,\, \to \mathop {Lead\,\,iodide}\limits_{(Yellow{\rm{ solid)}}} \,\,\, + \,\,\mathop {Potassium\,\,nitrate}\limits_{(Colourless\,\,\,solution)}

When magnesium is burnt, it combines with the oxygen of the air so that magnesium oxide is formed.

                Magnesium   +  Oxygen  → Magnesium oxide

Identification of Chemical Reactions

               A change is called a chemical reaction if it shows all or some of the following characteristics :

                (i)    Formation of new substance or substances

                (ii)   Production of heat or light or heat and light both

                (iii) Change in the colour

                (iv)  Change in temperature

Activity   

Main aim of this activity

                To demonstrate that burning of magnesium in air is a chemical reaction.

Experimental Steps

                (a)   Clean the surface of magnesium ribbon (strip of magnesium metal) by rubbing it with a sand paper and remove the dust.

                (b)   Light a burner or spirit lamp.

                (c)   Hold the magnesium strip carefully with a holder or with a pair of tongs.

                (d)   Bring one end of the magnesium strip in contact with flame of the burner. (Be careful, it burns with dazzling light which is harmful for eyes).

                (e)   Collect and cool the white ash (white powder) in a china dish.

                (f)    Put a small portion of the white ash on moist red litmus paper.

                (g)   Record your observations.

Figure : Burning of magnesium in air is an example of chemical reaction.

Observations

                (i)    Magnesium burns with dazzling light.

                (ii)   As a result of burning of magnesium, a white ash is formed.

                (iii) The white ash turns moist red litmus paper blue.

Conclusions

Magnesium combines with the oxygen of the air and forms a new substance.

Burning of magnesium is a chemical reaction and is represented as

Magnesium  +  Oxygen  →  Magnesium oxide

\mathop {2Mg(s)}\limits_{Greyish\,\,solid} \,\,\,\,\, + \,\,\,\,\mathop {{O_2}(g)}\limits_{Colourless\,\,gas} \,\, \to \,\,\,\,\mathop {2\,MgO(s)}\limits_{White\,\,powder}

The new substance has basic nature because MgO reacts with water and forms a base Mg (OH)₂

\mathop {MgO}\limits_{Magnesium\,\,\,Oxide} \,\, + \,\,\,\,\mathop {{H_2}O}\limits_{Water} \,\, \to \mathop {Mg{{(OH)}_2}}\limits_{Magnesium\,\,\,hydroxide}

Chemical Equation

                The chemical reactions are carried to produce new type of substances in the laboratory and industry. For example,

                In the laboratory, hydrogen gas is prepared by the reaction between zinc metal and dilute sulphuric acid.

Zinc + Dilute sulphuric acid → Zinc sulphate + Hydrogen gas

                In industry, ammonia gas is manufactured from hydrogen and nitrogen by Haber’s process.

Hydrogen  + Nitrogen → Ammonia

                The above type of  statement of any chemical reaction can be described in terms of symbols and formulae of the reactants and products involved in that reaction. This type of description of a chemical reaction is called chemical equation.

Definition of Chemical Equation

                A chemical equation is defined as a shorthand notation of an actual chemical reaction in terms of the symbols and formulae along with the number of atoms and molecules of its reactants and products.

Illustration of Chemical Equation

                Let us illustrate a chemical equation by taking an example of a useful reaction. Methane gas is used as fuel. This gas reacts with oxygen and burns. Due to this reaction, carbon dioxide and water are formed, and large quantity of heat is liberated. This reaction may be represented in the form of word equation as :

Methane + Oxygen → Carbon dioxide + Water +  Heat       … (i)

                In the form of chemical equation, the reaction is represented as :

CH₄(g) + 2O₂(g) →  CO₂ (g)  +  2H₂O (I) + Heat       … (ii)

                Equation (ii) represents a chemical reaction in terms of the formula of reactants and products. It also tells us about the number of atoms and molecules of various substances of the reaction.

Balanced Chemical Equations

                A chemical equation is said to be balanced when the number of atoms of each element on the reactant side is equal to the number of atoms of corresponding element on the product side.

Illustration and Verification of a Balanced Chemical Equation

                Let us consider the chemical equation of decomposition of ferric chloride

2FeC{l_3}\,\xrightarrow{{}}\,2FeC{l_2}\,\, + \,\,C{l_2}

The number of atoms of various types of elements on the reactant side and product side are given below :

From the above table, it is clear that the number of atoms of each element on the reactant side is equal to the number of atoms of the corresponding element on the product side. Therefore, the chemical equation is balanced.

Methods of Balancing A Chemical Equation

                There are following main methods employed to balance a chemical equation :.

                (a)   Hit and success method (hit and trial method)

                (b)   Partial equation method (half equation method)

Balancing Chemical Equation By Hit and Trial Method

                The following steps are to be followed :

                (a)   The symbol or formula of each one of the reactants and products is written in the form of expression (skeleton equation).

                (b)   The number of each type of atoms on the two sides of the skeleton equation is counted and tabulated.

                (c)   The number of each type of atoms on left side are made equal to the number of corresponding atoms on the right side of equation by using coefficients, if required.

                Example 1.

                Test whether the following is a balanced chemical equation? If not, balance it by hit and trial method.

Fe +  H₂O → Fe₃O₄ + H₂

                Solution

                Let us count and tabulate the number of various types of atoms on the two sides of the expression.

Fe + H₂O → Fe₃O₄ + H₂   … (i)

Atom	Reactant side	Product side
Number of Fe atoms	1	3
Number of O atoms	1	4
Number of H atoms	2	2

    In the above table,  it is seen that the number of atoms of each type of elements on the reactant side is not equal to the number of atoms of the corresponding element on the product side. Therefore, expression (i) is not a balanced chemical equation.

                Balancing Fe atoms

                There is one Fe atom on left side while there are three Fe atoms on right side. Therefore, a suitable coefficient of Fe on left side is 3 as 3Fe. Thus :

3Fe + H₂O → Fe₃O₄ + H₂   … (ii)

                Balancing O atoms

                There is one O atom (in H₂O) on left side and there are four O atoms on the right side of (ii). Therefore, a proper coefficient of H₂O is 4 as 4 H₂O.  Thus

3Fe +  4H₂O → Fe₃O₄ + H₂   … (iii)

                Balancing H atoms

                There are eight H atoms (in 4H₂O) on left side, but only two H atoms (in H₂) on the right side. Therefore, an appropriate coefficient of H₂ is 4 as 4 H₂. Thus

3Fe + 4H₂O → Fe₃O₄ + 4H₂      … (iv)

                Equation (iv) is a balanced chemical equation because the number of atoms of each type of element are conserved. How? These are shown below :

Atom	Reactant side	Product side
Number of Fe atoms	3	3
Number of O atoms	4	4
Number of H atoms	8	8

               Example 2.

                Count and tabulate each type of atom on both the side of the following equation, and explain whether it is a balanced chemical equation or not.

2KClO₃  → 2KCl  + 3O₂   … (i)

Atom	Reactant side	Product side
Number of K atoms	2	2
Number of Cl atoms	2	2
Number of O atoms	3 × 2 = 6	3 × 2 = 6

     The equation is a balanced chemical equation since the atoms are conserved on the two sides.

Balancing Chemical Equations By Half Equation Method

            More complex equations and ionic equations are balanced by half equation method. In this method, the main reaction is supposed to take place in several steps.

            Separate equation is written for each step. These are called half equations. Each half equation is balanced by hit and trial method. The balanced half equations are multiplied by suitable numbers (if necessary) to cancel the common substances which are not present in the final equation. These half equations are finally added to get the balanced overall equation. We will discuss this method in balancing of ionic equations.

Physical States of Reactants and Products

            Rate of reaction depends on the physical states of the reactants. For example, a reaction between solids is slow, a reaction in solution is fast and a reaction in gaseous state is very fast. Therefore, it is very important to mention the physical states of the substances involved in the reaction. As per IUPAC recommendation, the physical state is indicated by its symbol in a small bracket in front of the substance in a chemical equation. For example :

            The gaseous state is indicated by (g) as CO₂ (g)

            The liquid state is indicated by (l) as H₂O (l)

            The solid state is indicated by (s) as Fe (s)

            The solution of a substance in water is indicated by (aq) as NaOH (aq)

Balancing of Ionic Equations

            The chemical equations which involve ions along with atoms and molecules are called ionic equations. Usually such equations are used to represent oxidation-reduction reaction. For example, the reaction between zinc metal and copper sulphate solution is written as :

 Zn (s)         +        CuSO₄          →       ZnSO₄     +      Cu(s)

            Since sulphate ions SO^2–₄   are unaffected in the reaction, these are called expectators. Now, the above oxidation-reduction reaction is represented as :

Zn(s)          +      Cu²⁺ (aq)        →       Zn²⁺ (aq)    +      Cu(s)

            This is a balanced ionic equation. Why? Because the atoms and charge both are conserved. We shall confirm it by writing half reactions :

                Zn          →  Zn²⁺         +        2e^–      … (i) loss of electrons is oxidation

                Cu²⁺            +       2e^–      →    Cu       … (ii) gain of electrons is reduction

Add     

          Zn         +      Cu²⁺      →    Zn²⁺    +     Cu        … (iii)    … Oxidation-reduction  

            Equation (i) and (ii) represent half reaction while equation (iii) represent complete reaction.

Steps of Balancing lonic Equation

            (a)  The species oxidized and reduced are recognised.

            (b)  Half reaction equation for the loss of electrons is written.

            (c)  Half reaction equation for the gain of electrons is written.

            (d)  Each half reaction is separately balanced by hit and trial method with respect to number of atoms and number of each type of charge.

            (e)  Balanced half reaction equations are added and equal and opposite terms are cancelled.

            Example-3

            Balance the ionic equation Al + H⁺ → Al³⁺  +  H₂

            Solution

            In the given reaction, Al loses 3 electrons and it is oxidized to Al³⁺. A molecule of H₂ is formed from H⁺. That is, 2H⁺ ions gain 2e^– to give one molecule of H₂. Thus, the half reactions are written as :

Al  → Al³⁺  +  3e^–                    … (i)                 … loss of 3e^–

2H⁺ + 2e^– → H₂                       … (ii)                … gain of 2e^–

            The above two half equations are balanced with respect to the number of atoms of each element but not with respect to charge. Therefore, we shall multiply equation (i) by 2 and equation (ii) by 3 to get balanced half reactions.       

             2Al → 2Al³⁺  +  6e^–                 

              6H⁺ + 6e → 3H₂                     

Add         

          2Al + 6H⁺  +  6e^– → 2Al³⁺ + 6e^– + 3 H₂

On cancellation of equal number of electrons on the two sides, we get the balanced equation :

2 Al        +             6H⁺  →      2Al³⁺     +          3H₂

Information Conveyed by Chemical Equation

            We get the following information from a balanced chemical equation :

1. The formula, symbol, names and physical states of the reactants and products.
2. The relative number of atoms and molecules of the reactants that take part in the reaction.
3. The relative number of atoms and molecules of the products formed in the reaction.
4. The ratio of the moles of the reactants and products.
5. The ratio of the masses of the reactants and products.
6. The ratio of the volumes of the gaseous reactants and products.

Advantages of the Uses of a Chemical Equation.

            (a)  The representation of actual chemical reaction becomes easy by using equation. It saves time and space in writing.

            (b)  To prepare a known amount of the product we can calculate the amount of the reactant to be taken.

            (c)  The effect of temperature, pressure and concentration on the rate of a reaction can be investigated with the help of a perfectly balanced chemical equation.

            (d)  The chemical equation can be understood by the chemists belonging to any country of the world irrespective of the language they speak. For example, the chemical equation N₂ + 3H₂ → 2NH₃, can be understood by chemist as : nitrogen and hydrogen combine in 1 : 3 mole ratio to give 2 mole ratio of ammonia.

Limitations of a Chemical Equation

            (a)  A chemical equation does not tell us about the rate of the reaction and the time taken for the completion of the reaction.

            (b)  Some reactions may be even explosive. This is not revealed by a chemical equation.

            (c)  As such chemical equation does not tell about the actual quantity of the reactants consumed or products formed. It gives only ratio. Nevertheless, chemical equation is very useful to calculate the amount of all other substances consumed or formed if the amount of any one is known.

Characteristics of A Chemical Reaction

            Chemical reactions show one or more of the following characteristics :

Evolution of Gas

            Some chemical reactions take place with the evolution of a gas.

            Some reactions in which a gas is evolved are described below :

            Reaction between a metal such as zinc, magnesium, or iron and dilute

            sulphuric acid produces hydrogen gas. 

Figure : Evolution of gas

Zn (s)         +     H₂SO₄ (dil)        →     ZnSO₄ (aq)             +    H₂ (g)        

zinc               sulphuric acid                   zinc sulphate            hydrogen gas

Reaction between iron sulphide and dilute sulphuric acid produces hydrogen sulphide gas.

 FeS (s)        +     H₂SO4(dil)      →     FeSO₄ (aq)             +   H₂S (g)

iron sulphide        sulphuric acid       ferrous sulphate        hydrogen sulphide gas    

Heating a mixture of potassium chlorate (KClO3) and manganese dioxide (MnO2) gives oxygen gas.

2KClO₃ (s)            2KCl (s)     +        3O₂(g)

  Potassium chlorate             potassium chloride  oxygen gas

MnO₂ is used as a catalyst in this reaction.

            This reaction is used for the preparation of oxygen in the laboratory.      

Change of Colour

            There are some reactions in which there is a colour change.

            For example :

            When red lead oxide is heated, yellow lead monoxide is formed.

  2Pb₃O₄ (s)           6 PbO(s)              +      O₂ (g)

red lead oxide                 lead monoxide

(red colour)                    (yellow colour)

            When hydrogen sulphide gas is passed through a blue coloured solution of copper sulphate, black coloured copper sulphide is formed.

CuSO₄ (aq)      +      H₂S (g)      →        CuS (s)       + H₂SO₄ (aq)

copper sulphate   hydrogen sulphide   copper sulphide     sulphuric acid

                       (Blue Colour)                                            (Black Colour)

Formation of Precipitate

            In certain reactions, when solution of two reagents are mixed, one of the products formed gets precipitated immediately. Colour of the precipitate depends upon the reagents used in the reaction.

            When silver nitrate solution is mixed with a solution of sodium chloride, white precipitate of silver chloride is formed.

AgNO₃ (aq)      +    NaCl (aq)    →  AgCl (s)         +          NaNO₃ (aq)

Silver nitrate      sodium chloride       silver chloride       sodium nitrate

                                                                                    (white precipitate)         

            When potassium iodide solution is added to solution of lead nitrate, yellow precipitate of lead iodide is formed.

Pb (NO₃)₂ (aq)   +     2KI (aq)      →   PbI₂ (s)          +          2KNO₃ (aq)

 lead nitrate       potassium iodide         lead iodide             potassium nitrate

                                                                                   (yellow precipitate)

 When a few drops of ammonium hydroxide are added to a solution of ferric chloride, reddish-brown precipitate is formed.

 FeCl₃ (aq)       +     3NH₄ OH    →           Fe (OH)₃ (s)       +          3NH₄Cl (aq)

ferric chloride         ammonium            ferric hydroxide                ammonium

                             hydroxide                        (reddish-brown precipitate)               chloride

Energy Changes

            During a chemical change, energy is either evolved or absorbed. The energy evolved or absorbed may be in the form of heat, light, electricity, sound etc.

            When coal (or carbon) is burn, heat and light are produced.

C (s)           +       O₂ (s)        →     CO₂ (g)     +   Heat     +   Light

  carbon/coal             Oxygen                     Carbon dioxide

(from the air)

            In this reaction, heat and light are evolved.

            Heat and light are evolved in the following reactions also :

            (i)   During the burning of a candle.

            (ii)  During the burning of LPG.

            When a small quantity of water is added to quicklime, a large amount of heat is evolved.

 CaO (s)        +      H₂O (l)      →  Ca (OH)₂ (s)     +          Heat    

quicklime                water                      slaked lime

            This reaction takes place when lime is added to water for preparing the lime suspension which is used for whitewashing.

            Limestone (CaCO₃) is burnt to obtain lime (or quicklime). In this reaction, heat is absorbed.

CaCO₃ (s)       +        Heat        →      CaO (s)         +          CO₂ (g)

limestone                                        quicklime                    carbon dioxide

                                                      (or lime)

Change of Physical State

            In certain reactions, the physical state of products is different from that of the reactants. That is, there is a change of state during a chemical reaction.

            When a mixture of hydrogen and oxygen is ignited with an electric spark at room temperature, liquid water is formed.

2H₂ (g)         +       O₂ (g)          2H₂O (l) 

hydrogen gas        oxygen gas       water (liquid)

            When ammonia (NH₃) gas is allowed to come in contact with hydrogen chloride (HCl) gas, solid ammonium chloride is obtained.

 NH₃ (g)         +      HCl (g)      →                 NH₄ Cl (s)       

 ammonia        hydrogen chloride             ammonium chloride

                                      (white powder)

Types of Chemical Reactions

Types of Chemical Reactions

 

Direct combination                Decomposition             Displacement         

Exothermic &  Endothermic        Redox Reactions Reactions                                

                                                                                                             

Element + Element                            Thermal                                 Single        Double                                 

Oxidation         Reduction

Displaement               Displacement         

Compound + Compound        Electrical

Element + Compound             Photo

Combination Reactions

            Those chemical reactions which involve the combination of two or more substances to form a single new substance are called combination reactions. Combination reactions may involve either.

            (i) Combination of two elements or,

            (ii) Combination of an element and a compound or,

            (ii) Combination of two compounds.

            Let us now discuss all these types of combination reactions one by one.

(a)        Combination reactions involving two elements

            Some examples of combination reactions involving two elements are :

(i)         Magnesium ribbon burns in oxygen with a brilliant flame and a white residue of magnesium oxide is formed.

2 Mg (s)       +        O₂ (g)    →    2MgO (s)

Magnesium              Oxygen               Magnesium oxide

                                              (White residue)

(ii)         Hydrogen combine with chlorine in the presence of sunlight to give hydrogen chloride.

H₂ (g)         +        Cl₂ (g)                2HCl (g)

Hydrogen               Chlorine              Hydrogen chloride        

            The hydrogen chloride gas produced above on dissolving in water forms hydrochloric acid. Therefore, this reaction is used in the industry for the manufacture of hydrochloric acid.

            In all the above reactions, two elements combine to form a single new compound and are, therefore, combination reactions. These reactions are also called synthesis reactions.

(b)        Combination reactions involving an element and a compound

            Some examples of combination reactions involving an element and a compound are :

(i)         Carbon monoxide burns in oxygen to form carbon dioxide

2CO (g)        +        O₂ (g)    →        2CO₂ (g)    

Carbon monoxide          Oxygen         Carbon dioxide

(ii)         Nitric oxide combines with oxygen at room temperature to form brown fumes of nitrogen dioxide (NO2)

2NO          +           O₂        →         2NO₂

 Nitric oxide              Oxygen               Nitrogen dioxide

            In the above combination reactions, compounds like carbon monoxide (CO) nitric oxide (NO) or sulphur dioxide (SO₂) combine with an element oxygen to form new compounds like carbon dioxide (CO₂), nitrogen dioxide (NO₂) and sulphur trioxide (SO₃) respectively. So, all the above reactions are combination reactions.

(c)        Combination reactions involving two compounds

            Some examples of such reactions are :

Activity

            Take 2g of quick lime i.e. calcium oxide in a test tube. Add water to it slowly and observe. You will observe that calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide) and the test tube becomes hot. This is due to large amount of heat released during the reaction

CaO (s)        +        H₂O (l)    →   Ca (OH)₂ (aq)

 Calcium oxide                           Calcium hydroxide

            Do you know that a solution of slaked lime as produced by the above reaction is used for white washing purpose. Calcium hydroxide reacts slowly with carbon dioxide in the air to form a thin layer of calcium carbonate on the walls. The formation of calcium carbonate takes place in two to three days after white washing which gives a shine to the walls. It is also interesting to note that the chemical formula for marble is also CaCO₃.

Ca(OH)₂ (aq)    +        CO₂(g)    →       CaCO₃ (s)                +          H₂O (l)

Calcium hydroxide                    (From air)                       Calcium carbonate

                                                       (Shine on walls)

Sulphur trioxide combines with water to form sulphuric acid.

SO₃ (g)        +        H₂O (l)    →    H₂SO₄ (aq)

Sulphur trioxide            Water                   Sulphuric acid

            In the above reactions, two different compound combine with one another to form a new compound, therefore, all these reactions are known as combination reactions.

Decomposition Reactions

            The conversion of a single compound into two or more simple substances is called decomposition. For example:

            AB → A + B

            Here AB is the original compound and A and B are simple substances. A and B may be elements or simpler compounds into which the original compound AB decomposes. We can see that decomposition is reverse of combination.

Types of Decomposition Reactions

            Decomposition reactions are promoted by heat, electricity or light. Therefore, these reactions are classified as follows :

            (a)  Thermal decomposition (caused by heat)

            (b)  Electrolytic decomposition (caused by electricity)

            (c)  Photo decomposition (caused by light)

Thermal Decomposition

            Heat is called thermal energy. On heating, some substances undergo decomposition and new substances are formed. This process is called thermal decomposition or thermal dissociation. For example :

            Mercury (II) oxide decomposes on heating to form mercury and oxygen.

2HgO(s)              2Hg (l)         +              O₂(g)

Mercury (II) oxide         Mercury                      Oxygen

Thermal Decomposition of Metal Carbonates

            In general, heating causes the decomposition of metal carbonates so that carbon dioxide and respective metal oxides are formed.

Metal carbonate + Heat → Metal oxide + Carbon dioxide

CaCO₃(s)              CaO(s)        +             CO₂(g)

Calcium carbonate                Calcium oxide                    Carbon dioxide

                   (White)                  (White)

Electrolytic Decomposition Reaction

            The decomposition reactions caused by electric current is called electrical decomposition reaction or electrolysis. For example :

(a)        On passing electricity through molten sodium chloride, it decomposes into sodium metal and chlorine gas.

2NaCl                       2Na (s)         +         Cl₂ (g) … electrolysis

Sodium chloride                                                       Sodium                 Chlorine

                              (Molten)

(b)        On passing electricity through acidified water, it decomposes into hydrogen and oxygen.

2H₂O (l)                          2H₂(g)          +          O₂(g) …electrolysis

Photodecomposition Reactions

            A decomposition reaction caused by light is called photodecomposition reaction.

            For example:

            In the presence of sunlight, hypochlorous acid decomposes to give hydrochloric acid and oxygen

2HOCl (aq)          2HCl (aq)              +            O₂(g)          

Hypochlorous acid                   Hydrochloric acid           Oxygen

Displacement Reactions

            The chemical reactions in which an atom or a group of atoms in the molecule is replaced by another atom or a group of atoms are called displacement reactions. Let AL be any molecule and X be the displacing group,  then the displacement reaction may be written as :

X + A – L → A – X + L

            Here L is the leaving group.

            These reactions are generally found to occur in the solution. The elements involved may be metals or non-metals, i.e. a more active metal may displace a less active metal or a more active non-metal may displace a less active non-metal from its compound.

            Before we take up examples of displacement reactions, it is important to understand the relative reactivities of metals and the relative reactivities of non-metals

            Relative Reactivities of Metals (Reactivity series or Activity series of metals). We know that different metals posses different reactivities. Some metals are highly reactive whereas some other metals are less reactive. Based on the study of their reactions, the metals have been arranged in order of their reactivity. This arrangement of metals in a vertical column in order of their decreasing reactivity from top to bottom is called reactivity series or activity series of metals. Thus, more active metals lie above whereas less reactive metals lie below in the series. Potassium lying at the top is the most reactive metal whereas platinum lying at the bottom is the least reactive metal. Thus, any metal can displace the metal  lying below it from its salt solution. It is important to note that hydrogen (through non-metal) is included in the series because like metals, hydrogen can lose electron  to form  positive ion (H⁺). As metals lying above hydrogen are more reactive, they can displace hydrogen from acids or water, i.e. they can react with an acid to give out hydrogen gas.

Figure : Activity series of metals

            On the other hand, metals lying below hydrogen in the series, being less reactive than hydrogen, cannot displace hydrogen from acids, i.e., they do not react with the acid to give out hydrogen gas.

            Relative Reactivities of Non-metals For example, the relative reactivities of halogens is in the order :

F >  Cl  > Br  > I

            Thus, fluorine is most reactive while iodine is least reactive. Fluorine (F₂) can displace chlorine from NaCl, bromine from NaBr and iodine from NaI. Similarly, chlorine can displace bromine from KBr and iodine from KI and bromine can displace iodine from KI.

Types of Displacement Reactions

            The displacing group X may have many types of electronic configuration and hence many types of substitution reactions are possible. In general :

            (a)   A more active metal will displace a less active metal from compound.

            (b)   Some active nonmetals will displace less active nonmetals.

            (c)   Some metals will displace a nonmetal.

Displacement of A Less Active Metal By A More Active Metal

            Reaction 1.

            When an iron nail is dipped in a copper sulphate solution, it gets coated with copper.

Fe(s)       +        CuSO₄ (aq)  →     FeSO₄ (aq)     +          Cu (s)

Iron              Copper sulphate              Iron sulphate              Copper

            In this reaction, Fe has taken the place of Cu in the compound CuSO₄. In other words, we say that Fe has displaced Cu from the compound CuSO₄.

            Conclusion

            From this reaction we conclude that Fe (iron) is more reactive metal than Cu (Copper).

Displacement of A Less Active Non metal By A More Active Non metal

            Reaction 1.

            When chlorine gas is passed through sodium bromide solution, sodium chloride and bromine are formed.

2NaBr(aq)      +         Cl₂(g)          →       2NaCl (aq)             +    Br₂(g)

Sodium bromide         Chlorine                    Sodium chloride         Bromine       

            Conclusion.

            In this reaction, chlorine has displaced bromine from NaBr. Therefore, chlorine is more reactive than bromine.

            Reaction 2.

            When chlorine is bubbled through sodium iodide solution, sodium chloride and iodine are formed

2NaI(aq)       +        Cl₂ (g)          →        2NaCl (aq)             +     I₂ (g)

Sodium iodide           Chlorine         Sodium chloride      Iodine

            Conclusion

            In this reaction Cl has displaced I₂ from sodium iodide (NaI).

            Therefore, Cl (chlorine) is more reactive than I (Iodine).

Displacement of Hydrogen From Acids By Active Metals

            When zinc reacts with sulphuric acid, hydrogen gas is liberated and zinc sulphate is formed.

Zn (s)         +     H₂SO₄ (aq)      →       Zn SO₄ (aq)  +     H₂ (g)    

Zinc              Sulphuric acid                  Zinc sulphate     Hydrogen

            Conclusion

            Zn (Zinc) is more reactive than H (hydrogen).

Double Displacement Reactions

            The reactions given under displacement reactions involve displacement of an atom or group of atoms in a molecule. But there are some reactions in which two different atoms or groups of atoms are displaced by other atoms or groups of atoms. Such reactions in which two compounds react by an exchange or displacement of ions to form new compounds are called double displacement reactions. These reactions generally occur in solutions and in some cases, one of the products being insoluble, precipitates out as a solid and settles down. Let us perform an activity to have a clear understanding of double displacement reactions.

Activity

            Take 2ml each of solution of sodium sulphate and barium chloride in two test tube. Mix the two solutions and shake (figure). You will observe the formation of a white substance which is insoluble in water. This insoluble substance is known as a precipitate and is due to the formation of barium sulphate.

            When barium chloride solution is added to a solution of sodium sulphate, a white precipitate of barium sulphate is formed alongwith the formation of sodium chloride solution.

Na₂SO₄ (aq)    +     BaCl₂ (aq)  →     2NaCl (aq)       +            BaSO₄ (s)

Sodium sulphate   Barium chloride     Sodium chloride       Barium sulphate

                    (White ppt.)

            In this reaction, SO₄^2– ions displace Cl^– ions and Cl^– ions displace SO₄^2– ions. Since the reaction involves the displacement of two chemical species, therefore, it is known as double displacement reaction.

            These reactions usually occur in ionic compounds. It may be noted that in the reaction discussed above, barium sulphate is formed as an insoluble white solid, known as white precipitate.

            A few more examples of double decomposition reactions are given below :

            Example 1.

            If we add silver nitrate solution to sodium chloride solution, a white precipitate of silver chloride is formed along with the formation of sodium nitrate solution.

AgNO₃ (aq)     +      NaCl (aq)  →        AgCl (s)        +            NaNO₃(aq)

Silver nitrate       Sodium chloride        Silver chloride      Sodium

              (White ppt.)          Nitrate

            In this reaction, NO₃^– ions displace Cl^– ions and Cl^– ions displace NO₃^– ions. It may be noted that in this reaction, silver chloride is formed as an insoluble white solid known as white precipitate.

            Example 2.

            If ammonium hydroxide solution is added to aluminium chloride solution, a white precipitate of aluminium hydroxide along with ammonium chloride solution is obtained.

 AlCI₃(aq)       +    3NH₄OH(aq) →      AI(OH)₃ (s)      +           3NH₄CI(aq)

Aluminium            Ammonium                 Aluminium           Ammonium

chloride               hydroxide         hydroxide (White ppt.)      chloride

            In this double displacement reaction, Cl^– ions and OH^– ions have displaced each other to form insoluble aluminium hydroxide and ammonium chloride solution.

            The acid-base reactions are also double displacement  reactions as discussed below :

            If we add sodium hydroxide to hydrochloric acid, we obtain sodium chloride and water.

HCl(aq)        +     NaOH(aq) →        NaCI(aq)        +              H₂O(l)

Hydrochloric acid    Sodium hydroxide    Sodium chloride       Water

            In this reaction, CI^– ions displace OH^– ions and OH^– ions displace Cl^– ions. It may be noted that in this case, no precipitate of sodium chloride is formed because sodium chloride is soluble in water. 

            The double decomposition reactions can be further classified in two types :

            (a) Precipitation reactions                                   (b) Neutralisation reactions

            Let us discuss these double decomposition reactions individually. 

(a)        Precipitation reactions

            Those reactions in which two clear and transparent solutions on mixing result in the formation of an insoluble product are known as precipitation reactions, and  the insoluble product is known as precipitate.

            Some examples of precipitation reactions are given below :

            (Combination of aqueous solution of sodium sulphide and lead acetate). On mixing these two solutions, a black precipitate of lead sulphide (PbS) is formed. The precipitation reaction may be represented as:

Na₂S(aq)        +           (CH₃COO)₂Pb(aq)    →     PbS (s)        +         2CH₃COONa (aq)

Sodium sulphide          Lead acetate             Lead sulphide             Sodium acetate

                                                                 (Black precipitate)

            Now, can you explain what will happen when an aqueous solution of silver nitrate is added to an aqueous solution of sodium iodide?

AgNO₃(aq)         +       NaI(aq)    →       AgI (s)         +           NaNO₃ (aq)

Silver nitrate           Sodium iodide    Silver iodide         Sodium nitrate

                                (Yellow precipitate)

            On mixing the aqueous solutions of silver nitrate and sodium iodide, a yellow coloured precipitate of silver iodide will be obtained as shown above. In  this reaction, iodide ions (I^–) have replaced  nitrate ions (NO₃^–) and vice versa.

(b)        Neutralization Reactions

            When an aqueous solution of hydrochloric acid is mixed with an aqueous solution of sodium hydroxide in equal  amounts, a reaction takes place to form sodium chloride and water

HCl (aq)       +     NaOH(aq)  →      NaCl(aq)        +      H₂O(l)

Hydrochloric acid    Sodium hydroxide     Sodium chloride      Water

            Such a reaction is termed as a neutralization reaction. The hydrogen (H⁺) ions which were responsible for the acidic properties of HCl have reacted with hydroxyl (OH^–) ions which were responsible for the basic properties of  NaOH, there occurs a chemical change and appear in the form of crystalline sodium chloride on evaporation.

            Such reactions in which an acid and a base react with each other to produce salt and water are known as neutralization reactions.

            Some other examples of neutralization reactions are :

(i)              NaOH(aq)      +      HNO₃(aq)  →    NaNO3(aq)      +     H₂O(l)

            Sodium hydroxide      Nitric acid      Sodium nitrate             Water

(ii)              KOH(aq)       +       HCl(aq)  →        KCl(aq)         +     H₂O(l)

  Potassium hydroxide      Hydrochloric acid    Potassium chloride  Water

Exothermic and Endothermic Reactions

            On the basis of energy changes chemical reaction are classified under two categories that are exothermic and endothermic reactions.

(a)        Exothermic Reaction

            The  chemical reactions in which formation of products is accompanied  by evolution of heat are known as exothermic reaction.

            In such cases, the sign “ + Heat”  is written alongwith the products i.e.

            Reaction → Products  +  Heat

            Examples :

            (1)  C(s) + O₂ (g)  → CO₂ (g) + Heat

            (2)  CH₄(g) + O₂(g)  → CO₂ (g) + 2H₂O (l) + Heat

            (3)  2Mg(s) + O₂ (g)  → 2MgO(s) + Heat

            (4)  2H₂S(g) + 3O₂(g)  → 2H₂O(l) + 2SO₂ (g) + Heat

            (5)  CaO(s)  + H₂O  →Ca(OH)₂ (aq) + Heat

            (6)  C₆H₁₂O₆ (aq) + 6O₂  → 6O₂ (aq) + 6H₂O (l) + Energy

(b)        Endothermic Reaction

            The chemical reactions in which formation of products is accompanied by the absorption of heat are known as endothermic reactions.

            In such cases, the sign “+ Heat”  is written alongwith the reactants or ‘Heat’  with the products

                                                Reactants → Products – Heat

            Example :        

(1)    N₂ (g) + O₂ (g) + Heat → 2NO (g)

    Nitrogen     Oxygen        Nitric Oxide   

                  This is the only combustion reaction which is endothermic in nature.

(2)  Ba(OH)₂ + 2NH₄Cl  → BaCl₂  +  2NH₄OH  – Heat

(3)  C(s) + H₂O(v) → CO(g) + H₂(g) – Heat 

Hot coke    Steam     Water gas

(4)   NH₄CI + H₂O → NH₄OH + HCI – Heat

(5) H₂(g)   +   I₂ (g)    2HI (g)  – Heat

(6)  C(s) + 2S(g) →  CS₂ (l) – Heat

Oxidation and Reduction reactions  

            Redox Reaction

            In our daily life we come across  processes like rusting of objects made of iron, fading of the colour of the clothes, burning of the combustible substances such as cooking gas, wood, coal, etc. All such processes fall in the category of specific type of chemical reactions called oxidation – reduction reactions or redox reactions.

            Most of the elements are reactive and react with oxygen and hydrogen. Initially, on the basis of addition of oxygen and hydrogen, the chemical reactions were considered as oxidation and reduction reactions but afterwards, the definition was expanded, on the basis of addition or displacement of other elements except O₂ and H₂, which are as follows–

Oxidation

            Definitions :  The oxidation of a substance taken place when :

(a)     There is addition of oxygen to a substance.

(i)     The chemical reactions in which oxygen combines with some other substance.

(1)   C (s)       +      O₂ (g)    →     CO₂  (g)

(2)   2Mg (s)   +      O₂ (g)    →     2MgO (g)

(3)   P₄(s)       +      5O₂    →      2P₂O₅ (s)               

(b)        There is removal of hydrogen from a substance

(ii)    The chemical reactions in which hydrogen is lost from a substance.

(1)   H₂S    →     H₂ +  S

(2)   CH₃CH₂OH         CH₃CHO           +     H₂

Ethyl alcohol             Acetaldehyde

In ethyl alcohol, 6 hydrogen atoms are present and 4 hydrogen atoms are present in acetaldehyde, so formed.

     Hence, there is a loss of two hydrogen atoms so this reaction is an oxidation reaction.

(3)   H₂S (aq)  +  Br₂  (aq)  →  2HBr  (aq)   + S

              

(4)     

                         

(5)        2H₂S + O₂ → 2H₂O +2S     

                            

 (c)       There is addition of an electronegative element to a substance

            Examples

(1)                 

(2)      SnCI₂ + CI₂ → SnCI₄

            In  these reactions, numbers of atoms of electronegative element chlorine is increased in Fe and SnCl₂ and forms FeCl₃ and SnCl₄ respectively

(d)        There is removal of electropositive element from a substance

            Examples :                               

(1)       2KI +CI₂ → 2KCI + I₂

            The electropositive element K (Potassium) is lost and in this way oxidation of KI taken place.

(2)      2NaBr + CI₂ (g) → 2NaCI + Br₂ (g)

            In this reaction NaBr chages to Br₂  . This involves removal of electropositive atom i.e. oxidation thus NaBr is oxidised to Br₂ 

            So,  in short oxidation is a chemical reaction in which substance combine with oxygen or an electronegative elment or loses hydrogen or an electropositive element.

Reduction

            The reduction of a substance takes place when :

            There is addition of hydrogen to a substance.

            Examples

(a)        (1) H₂ + S  →  H₂S

            (2) Cl₂ + H₂S → 2HCl + S

            Hydrogen sulphide (H₂S) when reacted with chlorine (Cl₂) gets oxidised to sulphur where as chlorine gets reduced to HCl

            (3)    H₂ + Cl₂  → 2HCl

            (4)    2Na + H₂  → 2NaH

            (5)    2Na + H₂ → 2NaH

(b)        These  is removal of oxygen from a substance

            Examples :

(1)    H₂O  +   C   → CO + H₂ 

       Steam   coke          water gas

In this case, water  has been reduced to hydrogen by the removal of oxygen

(2)    ZnO  +  H₂   →  Zn  +  H₂O

Here Zinc oxide has been reduced to zinc by the removal of oxygen

(3)    ZnO + C → Zn + CO

     

(4)  CuO + H₂ → Cu + H₂O

                    

(5)  2MgO → 2Mg + O₂

   

(c)        There is addition of electropositive element to a substance.

            Examples

(1)   CI₂ + Mg → MgCI₂

(2)   I₂ + Hg → HgI₂

            In above reaction, Cl₂  is reduced to MgCl₂ and I₂ is reduced to HgCl₂

(d)        There is removal of an electronegative element from a substance.

            Examples :

(1)   2FeCI₃ → 2FeCI₂ + CI₂

(2)   FeS →  Fe + S

            In  the above reactions change of  FeCl₃ and FeS to FeCl₂ and Fe respectively are examples of reduction.

            So, in short, reduction is a chemical reaction in which substance combine with hydrogen or an electropositive element or loses oxygen or an electro negative element.

            A substance  that brings about oxidation that is addition of oxygen or electronegative element and removal of hydrogen or electropositive element is called oxidising agent.

            On the other hand, a substance that brings about reduction, that is removal of oxygen or electronegative element and addition of hydrogen or electropositive element is called reducing agent.

            Consider the reaction 

 CuO + H₂ → Cu + H₂O      

            In this reaction, hydrogen removes oxygen from copper oxide. Thus, CuO is reduced and hydrogen behaves as reducing agent.

            Copper oxide  gives oxygen to hydrogen and hydrogen is oxidised to water by CuO. Therefore, CuO is acting as oxidising agent

 

 CuO + H₂ → Cu + H₂O       

                Oxidising    Reducing       

                                                                       agent     agent

               

            CuO  makes oxidation to occur  → Oxidising agent

            H₂  makes reduction to occur  →  Reducing agent

Note :   (i)     The  substance to which oxygen is  added or substance from which hydrogen is removed is said to be oxidised

            (ii)    The substance from which oxygen is removed substance to which hydrogen is added is said to be reduced  

            →     The substance  which gets oxidised acts as on reducing agent

            →     The substances which gets reduced acts as oxidising agent

Modern concept of oxidation and  reduction

            Besides earlier concepts of oxidation and reduction, two modern concepts of oxidation and reduction are more common.

(i)         Electronic  concept

(ii)         Oxidation number concept

            Here we will study only electronic concept.

Oxidation

            According to electronic concept, those chemical reactions in which an atom, ion or molecule loses electron are known as Oxidation reaction. This is also known as de-electronation.

(i)         Oxidation of atom

                      Na        →      Na⁺          +        e^–

                      Ca        →      Ca²⁺       +       2e^–

                      Zn        →     Zn²⁺       +       2e^–

(ii)        Oxidation of ion (Cation)

                     Fe²⁺      →     Fe³⁺       +        e^–

                     Sn²⁺     →      Sn⁴⁺       +       2e^–

                      Cu⁺      →     Cu²⁺       +        e^–

(iii)       Oxidation of ion (Anion)

                     2Cl^–      →       Cl²        +       2e^–

                   MnO₄^2–    →     MnO^–₄      +        e^–

                     2O^2–      →       O²        +       4e^–

(iv)       Oxidation of molecule

                     H₂S      →        H₂         +         S

                     H₂S      →        2H⁺        +       S^2–

            In oxidation reaction, electrons are lost due to which, positive charge on substance is increased and negative charge on substance is decreased.

Reduction

            Reduction is the chemical reaction in which an atom, ion or molecule gains one or more electron. This is also known as electronation.

(i)         Reduction of atom

                       Cl            +           e^–      →       Cl^–

                       S            +          2e^–    →       S^2–

(ii)        Reduction of ion

                     Mg²⁺          +          2e^–    →        Mg

                     Pb⁴⁺          +          2e^–    →       Pb²⁺

                   Cr₂O₇^2–    →      2Cr³⁺

                    NO ₃^–    →       NO ₂^–

(iii)       Reduction of molecule

                    MnO₂    →     Mn²⁺

            In reduction reaction, gain of electrons takes place due to which there is increase in negative charge or decrease  in positive charge.

Redox Reactions

            Generally, oxidation and reduction reaction takes place simultaneously because in a chemical reaction, one atom, ion or molecule of a substance loses electron while another atom, ion or molecule gains electron. In this way oxidation of one substance while reduction of another substance takes place. Thus such reaction are known as Redox reactions.

            Example

(1)              

         2Mg +O₂ → 2MgO (Mg²⁺ O^-2)                

                         

           

(2)    Fe  +  S    FeS  (Fe²⁺  S^–2)

          

Oxidizing agent according to electronic concept

            Those substances which gains electrons or get reduced are known as oxidizing agent.

Reducing agent according to electronic concept

            Those substances which loses electrons or get oxidised are known as reducing agent,

                   CuSO₄        +           Fe        →        FeSO₄         +         Cu

                                                              Oxidation  

      

                     Cu²⁺          +          Fe⁰      →         Fe²⁺           +         Cu⁰

                                                                      

                                                            Reduction

            In this reaction, Fe loses electrons and get oxidised so it is reducing agent, while CuSO₄ gains electrons and get reduced, so clearly it is an oxidizing agent.

Effects of Oxidation Reaction In Everyday Life  

            There are a number of oxidation reactions taking place around us which affect our everyday life. Two of these are briefly described below :

Corrosion    

            The process of slowly eating up of the metals due to attack of atmospheric gases such as oxygen, carbon dioxide, hydrogen sulphide, water vapour etc. on the surface of the metals so as to convert the metal  into oxide, carbonate, sulphide etc. is known as corrosion.

            The most common example of corrosion is rusting i.e. corrosion of iron. When an iron article remains exposed to moist air for a long time, its surface is covered with a brown, flaky (non-sticky) substance called rust. Rust is mainly hydrated ferric oxide (Fe₂O₃ .xH₂O). It is formed due to attack of oxygen gas and water vapour present in the air on the surface of iron.

2Fe(s) + 3/2O₂(g) + xH₂O(I) → Fe₂O₃.xH₂O(s) 

Iron       (From air)    Moisture      Hydrated ferric oxide (Rust)                                           

            Similarly , copper objects lose their lusture or shine after some time.  The surface of these objects acquire a green coating of basic carbonate, CuCO₃.Cu(OH)₂ when exposed to air. This is due to attack of  O₂, CO₂ and water vapour present in the air on the surface of copper.

2Cu(s) + CO₂(g) + O₂(g) + H₂O(I) → CuCO₃.Cu(OH)₂                       

Copper       Basic copper carbonate     

                                                                (Green)

            Likewise, the surface of silver metal gets tarnished (i.e. loses lustre and becomes dull) on exposure to air. This is due to the formation of a coating of black silver sulphide (Ag₂S)  on its surface by the action of H₂S present in the air.

2Ag(s) + H₂S(g) → Ag₂S(s) + H₂(g) 

Silver  Hydrogen sulphide   Silver sulphide   

                 (From air)                   (Black)                                        

            Rusting is a serious problem because it weakens the structure of bridges, iron railings, automobile parts etc. Every year a large amount of money is spent to replace rusted iron and steel structure. The reason is that the reddish brown crust of rust does not stick to the surface. It peels off (or falls down) exposing fresh surface for rusting. Thus, corrosion of iron is a continuous process which ultimately eats up the whole iron object.

            Methods  to prevent rusting. Rusting can be prevented if iron object are not allowed to come in contact with the damp air. Some common methods generally used are given below.

        (i)   By painting the iron articles such as window grills, iron gates, steel furniture, railway coaches bodies of cars, buses etc.

        (ii)   By greasing and oiling the iron articles such as mechanical tools, machine parts etc.

        (iii)  By galvanisation, i.e. coating the surface of iron objects with a thin layer of zinc.

Rancidity

            Oxidation also has damaging effect on foods containing fats and oils. When the food materials prepared in fats and oils are kept for a long time, they start giving unpleasant smell and taste. The fat and oil containing food materials which give unpleasant smell and taste are said to have become rancid (sour or stale). This happens as follows :

            When the fats and oils present in food materials get oxidised by the oxygen (of air), their oxidation products have unpleasant smell and taste. Due to this, the smell and taste of food materials containing fats and oils change and become very unpleasant (or obnoxious). The condition produced by aerial oxidation of fats and oils in foods marked by unpleasant smell and taste is called rancidity. Rancidity spoils the food materials prepared in fats and oils which have been kept for a considerable time and makes them unfit for eating. The characteristics of a rancid food are that it gives out unpleasant smell and also has an unpleasant taste. Rancidity is called ‘vikritgandhita’ in Hindi.

            The development of rancidity of food can be prevented or retarded (slowed down) in the following ways :

1. Rancidity can be prevented by adding anti-oxidants to foods containing fats and oils.

            Anti-oxidant is a substance (or chemical) which prevents oxidation. Anti-oxidants are actually reducing agents. When anti-oxidants are added to foods, then the fats and oils present in them do not get oxidised easily and hence do not turn rancid. So the foods remain good to eat for a much longer time. The two common anti-oxidants used in foods to prevent the development of rancidity are BHA (Butylated Hydroxy-Anisole) and BHT (Butylated Hydroxy-Toluene).

2. Rancidity can be prevented by packaging fat and oil containing foods in nitrogen gas.

            When the packed food is surrounded by an unreactive gas like nitrogen, there is no oxygen (of air) to cause its oxidation and make it rancid. The manufacturers of potato chips (and other similar food products) fill the plastic bags containing chips with nitrogen gas to prevent the chips from being oxidised and turn rancid.

3. Rancidity can be retarded by keeping food in a refrigerator.

            The refrigerator has a low temperature inside it. When the food is kept in a refrigerator, the oxidation of fats and oils in it is slowed down due to low temperature. So, the development of rancidity due to oxidation is retarded.

4. Rancidity can be retarded by storing food in air-tight containers.

            When food is stored in air-tight containers, then there is little exposure to oxygen of air. Due to reduced exposure to oxygen, the oxidation of fats and oils present in food is slowed down and hence the development of rancidity is retarded.

5. Rancidity can be retarded by storing foods away from light.

            In the absence of light, the oxidation of fats and oils present in food is slowed down and hence the development of rancidity is retarded.