Chemical reactions and equations



We observe many chemical changes taking place in our daily life. The milk turns sour if kept for a long time at room temperature, milk changes to curd, rusting of iron, digestion of food in our body are examples of chemical changes. In such changes, the nature and the properties of the substances change and we say a  chemical reaction has taken place. A chemical reaction is represented by a chemical equation which is a convenient way of  describing a chemical reaction with the help of symbols of elements and formulae of chemical compounds. In this chapter, we shall discuss about chemical formulae, chemical equations, balancing  of chemical equations and types of chemical reactions.

  • Physical change:– A change in which the physical properties of the substance changes but the chemical composition does not change. The substance is restored to its original state as soon as the cause of change is withdrawn.
  • Chemical change:– In a chemical change, at least one of the reacting substance changes into a new substances with a different composition. The new substances can not be changed back to the original substance even if the cause of change is withdrawn.

Difference between physical & chemical change

Chemical reaction

The processes, in which a substance or substances undergo a chemical change to produce new substance or substances, with entire new properties, are known as chemical reaction. The nature and identity of products totally changes from the reactants.

Some important characteristics of chemical reactions are:
(i) Change in state: The physical state of the substances normally changes.
e.g (a) Formation of solid MgO from solid Mg and gaseous O2.
(b) Formation of solid Pbl(ppt) from liquid solutions of PbNO3 and Kl.
(c) Formation of H2 gas from the reaction of solid Zn with liquid H2SO4.

(ii) Change in colour: In some of the chemical reactions change in colour can be observed.
(a) Formation of brown rust on black iron nails.
(b) Formation of yellow ppt. of lead iodide from colourless solution of PbNO3 and Kl.

(iii) Evolution on a gas: In some cases, a gas may be evolved.
e.g. (a) Evolution of H2 gas, in the reaction between Zn and dil HCL
(b) Evolution of CO2 gas, during burning of any fuel, which contains carbon.

(iv) Change in temperature: Most of the reactions are accompanied by temperature change. i.e. increase or decrease in temperature.
e.g. (a) In the reaction between Zn and H2SO4, flask was found to be warm. Thus rise in temperature has taken place.
(b) If a reaction between barium hydroxide, Ba(OH)2 and ammonium chloride, NH4Cl is carried out in a test tube, it is observed that bottom of test tube becomes cooler.


Word equation

A chemical equation which represents a chemical reaction briefly in words is called word equation.
Example: For the example the word equation is
Sodium + water  ? Sodium hydroxide + Hydrogen


  • Rules for writing a word equation

    (i) The substances taking part in chemical reaction, reactants are always written on the left hand side of arrow.
    (ii) The substances formed after the chemical reaction, products are always written on the right hand side of arrow.
    (iii) A plus sign (+) is put in between the reactants or between the products. If their number is two or more.
    (iv) An arrow (?) is put between the reactants and products, the arrow shows the direction of the reaction in which the reaction proceeds. The arrow is read as “to yield” or “to form”. 

    In the word equation when symbols and chemical formulae of the reactants and products are used then it is called as chemical equation. 
    Example : Na + H2O ? NaOH + H2


Important Terms and Concepts

I. Formulae of Ions

1. Valency.
The number of electrons shared by an atom is called its valency. It is also called the combining capacity of an atom, e.g., Cl atom can share one valence electron, its valency is 1, Oxygen can share two valence electrons, its valency is 2. Nitrogen can share 3 valence electrons, its valency is 3, Carbon can share 4 valence electrons, therefore its valency is 4 and so on. It means if Carbon combines with Chlorine, Carbon will share four valence electrons with four Chlorine atoms, therefore the molecular formula of the covalent compound will be

Some more examples are :

*The elements show more than one valency. So a Roman numeral shows theft valency in a bracket.

2. Chemical Equations. “A chemical equation is a symbolic notation that uses formulae of compounds and symbols of elements to represent a chemical reaction”, e.g., Copper oxide reacts with Carbon to form Copper and Carbon monoxide. The reaction may be represented as
CuO  +  C  ? Cu + Co

3. Writing of a Chemical Equation.
(i) The symbols of elements and the formulae of reacting substances (reactants) are written on the left hand side and plus (+) sign is written between them.
(ii) The symbols and formulae of the substances formed (products) are written on the right hand side with a plus sign (+) between them.
(iii) An arrow (?) sign in put between the reactants and products, e. g.,
Mg + H2SO4  ? MgSO4 + H2

(iv) The physical states of the reactants and products are also mentioned in a chemical equation. The notations g, l, s, aq. are written in brackets along with symbols/formulae of reactants and products. These symbols stand for gaseous, liquid, solid and aqueous solution respectively, e.g.,
Mg (s) + H2SO4(aq)  MgSO4(aq) + H2 (g)
Zn (s) + H2SO4 (aq)  ZnSO4(aq) + H2 ­?


The symbol (?­) may also be used to represent a gaseous product. The symbol (?) is used to represent the formation of a precipitate (water insoluble) or a sparingly soluble substance formed during the reaction which settles down mostly, e.g.,

NaCl (aq) + AgNO3 (aq)  ?  AgCl (?) + NaNO3 (aq)

(v) Sometimes, the temperature, pressure and catalyst of the reaction are indicated above and or below the arrow in the equation, e.g.,
CO (g) + 2H2 (g)  \mathrel{\mathop{\kern0pt\longrightarrow}\limits_{340atm,heat}^{Zno/C{r_2}{O_3}}}     CH3OH (g)

(vi) A chemical equation represents an actual chemical reaction in which the reactants and products are known, e.g.,
2 KMnO4 (s)    \buildrel {heat} \over\longrightarrow   K2MnO4 (s) + MnO2 (s) + O2 (g)
2 KClO3 (s)  \mathrel{\mathop{\kern0pt\longrightarrow}\limits_{Mn{O_2}}^{heat}}    2 KCl (s) + 3 O2 (g)

4. Balancing of chemical equation. Observe the following two chemical equations :
Zn + H2SO4    ?     ZnSO4 + H2                                                      ……(i)
Na + H2O  NaOH + H2                                                             …….(ii)
In equation (i), the number of atoms of Zn, H, S, and O are equal on both sides, i.e., the equation is balanced.

5. Balanced Equations. The equations in which atoms of various elements on the reactant’s and the product’s side are equal.
Equation (ii) is not balanced because the number of hydrogen atoms is not equal on both sides. It is called a skeleton chemical equation.

6. Reason of Balancing Equations. The number of atoms of elements on both sides of a chemical equation should be equal in accordance with the law of conservation of mass.

7. Balancing. The process of making atoms of various elements equal in an equation on either side is called balancing.

8. Steps in Balancing of Chemical Equations. A number of steps are involved in balancing a chemical equation, e.g.,
Na + H2O    ?    NaOH   + H2


Step 1: Examine the number of atoms of different elements present in unbalanced equation.

Step 2: Pick an element to balance the equation. In the above equation Na and O are  balanced, Hydrogen is not.
Step: To balance Hydrogen on both sides we need to multiply H2O by 2 which makes Hydrogen atoms equal to 4 on the reactants’ side. To make Hydrogen 4 on the products’ side, multiply NaOH by 2. Now oxygen has become 2 on both sides. But Sodium atom has become two on the product’s side. Multiply Na by 2 on the reactant’s side so that they become equal on both sides. The steps are as follows:


(i) Na + 2H2O  →  NaOH + H2
(ii) Na + 2 H2O   →  2 NaOH + H2
(iii) 2 Na + 2 H2O  →  2 NaOH + H2

The equation is now balanced.
Example : Fe + H2O  →  Fe3O4 + H2


Step 1 :

Step 2 :  Pick up the compound which has the maximum number of atoms whether a  reactant or a product, and in that compound select the element which has the highest number of atoms, e.g., we select Fe3O4 in the above equation : To balance oxygen atoms,

Fe (s) + 4 H2O (g)  →  Fe3O4 (s) + H2 (g) (Partly balanced)

Step 3: Pick up the second element to balance this partly balanced equation. Let us try to balance hydrogen atoms. bIn partly balanced equation. Atoms of Hydrogen

To equalise the number of Hydrogen atoms, we use 4 as the coefficient of H2 in the  products.
Fe (s) + 4H2O (g)  → Fe3O4 (s) 4H2


Step 4: Pick up the third element to be balanced. The element which is left to be balanced is Fe.

To equalise iron, we use 3 as coefficient of Fe in reactants.
3 Fe + 4 H2O   →  Fe3O4 + 4H2


Step 5: Check the correctness of the balanced equation.

The equation is balanced because atoms of all the elements are equal on both sides. This method of balancing equation is known as hit and trial method.


9. Balancing of Ionic Equations. In these equations, charge balancing is also done along  with balancing of atoms on both sides of the equation, e.g.,

Initial    Cu2+ (aq) + H2S  →  CuS (s) + H+ (aq).
Balanced     Cu2 + (aq) + H2S   →   CuS (s) + 2 H+ (aq)

We have balanced the charges. It was + 2 on LHS and we have made + 2 on RHS. Number of Hydrogen atoms, Cu and Sulphur atoms are also balanced on both sides.


Types of Chemical Reactions

The chemical reactions are classified into various categories depending upon the types of changes taking place. The different types of reactions are as follows :

(i) Combination Reaction.

The reactions in which two or more substances combine to form a single new substance are called combination reaction. Combination may take place,
(i) Between two or more elements.
(ii) Between two or more compounds.
(iii) Between elements and compounds.
Some more examples of combination reactions:

(a) Between two elements
(i) Burning of Coal C(s)             +            O2(g)         →         CO2(g)
Carbon                    Oxygen                   Carbon dioxide
(ii) Formation of Water
2H2(g)            +             O2(g)          →          2H2O(l)
Hydrogen                   Oxygen                            Water
(iii) Burning of Magnesium in air
2Mg(s)            +                 O2(s)         →        2MgO(s)
Magnesium                        Oxygen                   Magnesium oxide
(iv) Formation of Iron sulphide
Fe(s)                +                    S(s)          Δ→          FeS(s)
Iron                                   Sulphur               Iron sulphide

(b) Between 2 compounds
(i) Formation of Ammonium chloride
NH3(g)      +      HCl(g)    →    NH4Cl(s)
Ammonia       Hydrogen        Ammonium
Chloride             Chloride

(ii) Formation of Calcium Carbonate
CaO(s)                 +                 CO2(g)      →       CaCO3(s)
Calcium oxide                     Carbon                  Calcium
(Quick lime)                         dioxide                  carbonate

(c) Between an element and a compound
(i) Reaction of carbon monoxide with oxygen
2CO(g) + O2(g)  →    2CO2(g)
This is also an exothermic reactions.

So, we can say that respiration is an exothermic reaction.


(ii) Decomposition Reaction.

A reaction in which a single compound breaks down to produce two or more simpler substances. i.e., a compound decomposes into simpler substances.  It is opposite to combination reactions.

There are three ways in which decomposition reactions can be carried out, i.e., energy required in decomposition reaction can be supplied in the following ways:
(i) Heat
(ii) Electricity
(iii) Light

(1) Electrolysis.  When decomposition reaction is carried out with the help of electric current, the process is called electrolysis (‘electro’ means electric, ‘lysis’ means break down), e.g., when electric current is passed through acidified water (water mixed with a few drops of acid so as to make it a good conductor), it decomposes into Hydrogen and Oxygen gases.

2 H2O (l)      \buildrel {ElectricCurrern} \over\longrightarrow   2 H2 (g) + O2 (g)


(2) Thermal Decomposition. When decomposition reaction is carried out by heating, it is called thermal decomposition reaction, e.g.,
CaCO3 (s)   \buildrel {heat} \over\longrightarrow    CaO (s) + CO2 (g)
[Limestone]                             [Quick lime]

FeSO4 (S)    \buildrel {heat} \over\longrightarrow    Fe2O3 + SO2(g) + SO3(g)
[Ferric oxide]

2Pb(NO3)2(S)  \buildrel {heat} \over\longrightarrow 2PbO + 4NO2(g) + O2(g)

Zn CO3 (s)  \buildrel {heat} \over\longrightarrow  ZnO (s) + CO2 (g)

The process of heating ZnCO3 (Calamine), an ore of zinc in absence of air to form Zinc oxide (ZnO) and CO (g) is also called calcination.


(3) Photochemical Decomposition:
Chemical reaction in which a compound decomposes into simpler substances on the absorption of light energy is called photo-decomposition reaction.

2Agcl(s)    \buildrel {Sunlight} \over\longrightarrow   2Ag(s) + Cl2(g)
Silver Chloride                Silver
(White)                              (Grey in colour)

2AgBr(s)     \buildrel {Sunlight} \over\longrightarrow     2Ag(s) + Br2(g)
Silver Bromide              Silver         Bromine

  • The Decomposition of a Compound with light is called “Photolysis.” 
  • All Decomposition reaction requires energy i.e. these reactions are “Endothermic reactions.”
  • These reactions are used in extractions of metals.

These reactions are photochemical reactions which are used in black and white photography. Another important example of decomposition reaction in our body is digestion of food. When we eat rice, wheat or potatoes, the starch gets decomposed to simple sugar and proteins get converted into simple substances called amino acids in our body.
(C6H10O5)n + H2O        →            C12H22O11 
C12H22O11 + H2O      →           2C6H12O
Proteins                              →           Amino acids

We have observed all the decomposition reactions require energy either in form of heat, light or electricity for breaking down of reactants. Therefore, they are endothermic reactions.

Endothermic Reactions: Those reactions in which heat is absorbed are called endothermic reactions.

(iii) Displacement Reactions
. Those reactions in which a more reactive element displaces a less reactive element from a compound are called displacement reactions.

These reactions mostly occur in solution form, e.g.,
Zn (s) + CuSO4 (aq)   →     ZnSO4 (aq) + Cu (s)
Colourless      Reddish brown

Pb + CuSO4   →     (aq)  PbSO4  ↓ + Cu (s)
It is a displacement reaction. Other examples are :
Mg + CuSO4      →      MgSO4(aq) + Zn (s)


It shows magnesium is more reactive than Cu because it can displace Copper from Copper sulphate solution.  Mg(s) + ZnSO4 (aq) →  MgSO4 (aq) + Zn (s)
Mg (s) + FeSO4 (aq)  →  MgSO4 (aq) + Fe (s)

These reactions show that Mg is more reactive than
Zn and Fe. Zn (s) + FeSO4 (aq)   →   ZnSO4 (aq) + Fe (s)]

It shows Zn is more reactive than Fe.
On the basis of the above reactions, we can conclude Mg > Zn >Fe > Pb > Cu > Ag is the order of reactivity.

Zn (s) + H2SO4 (dil.)  ZnSO4 (aq) + H2 (g)
Mg (s) + H2SO4 (dil.) MgSO4 (aq) + H2 (g)

These reactions show that Zn and Mg are more reactive than Hydrogen because they displace Hydrogen from dilute acids. These are also examples of displacement reactions.



(iv) Double Decomposition Reactions (Double Displacement Reactions).

Those reactions in which two different atoms or groups of atoms are displaced by other atoms or groups of atoms, i.e., two compounds exchange their ions and one of the products formed is insoluble, e.g.,
BaCl2 (aq) + Na2SO4 (aq)    →  BaSO4 (s) + 2NaCl (aq)

Here, S2- ions are displacing Cl ions and Cl ions are displacing SO2-4 ions. Since it involves displacement of two species, therefore, is called double displacement reactions.

  • If one of the products formed in the reaction is insoluble, it is also called double decomposition reaction.
  • These reactions usually occur in between ionic compounds when they are dissolved in water i.e., in aqueous solution.
  • These reactions are fast reactions and take place within fraction of a second.


AgNO3 (aq)                 +     NaBr (aq)             →         AgBr ↓                      +        NaNO3 (aq)
AgNO3 (aq)                 +     KI (aq)                    →         AgI ↓                          +        KNO3 (aq)
FeSO4 (aq)                   +      2 NaOH (aq)      →         Fe (OH)2 ↓             +        Na2SO4 (aq)
Cr2(SO4)3                   +     6 NaOH                  →        2 Cr(OH)3 ↓          +        3 Na2SO4 (aq)
FeCl3                               +     3 NaOH                →          Fe(OH)3 ↓             +         3 NaCl (aq)
AlCl                               +    3 NaOH               →           Al(OH)3 ↓              +          3 NaCl (aq)
CuSO                           +     H2S(g)                 →            CuS ↓                       +          H2SO4 (aq)
MnSO4                          +     H2S (g)                →             MnS ↓                     +           H2SO4 (aq.)
NiCl2                               +     H2S (g)               →             NiS ↓                       +           2 HCl (aq)
Pb (NO3)2                    +     2 HCl (aq)         →             PbCl2 (s)              +           2 HNO3 (aq)
CaCO                          +    2 NaCl (aq)      →              CaCl2 (aq)            +           Na2CO3(aq)
Ca(NO3)                    +    (NH4)2CO3    →       CaCO3 ↓                       +         2 NH4NO3
ZnSO4                            +    H2S                    →              ZnS ↓                        +          H2SO4 (aq)


(v) Neutralization Reactions. Those reactions in which acid or acidic oxide reacts with base or basic oxide to form salt and water are called neutralization reactions, e.g.,

NaOH                      +                   HNO3             →                    NaNO3                   +                   H2O
2 NaOH                 +                    H2SO4           →                    Na2SO4                 +               2 H2O
KOH                        +                    HCl                  →                     KCl                            +                  H2O
KOH                       +                     HNO3            →                     KNO3                      +                  H2O
2KOH                    +                     H2SO         →                     K2SO4                     +              2H2O
CH3COOH        +                     NaOH            →                      CH3COONa        +                  H2O
                 +                    2NaOH                                                                          +                        2H2O

When salt of weak acid reacts with strong acid, it is also called neutralization reaction.
Na2CO3             +                  2HCl               →              2NaCl                         +                H2O + CO2
CH3COONa    +                  HCl                 →               CH3COOH             +                 NaCl

When acidic salt reacts with base to form salt and water, it is also called neutralization reaction.
NaHCO3            +                NaOH              →                 Na2CO3               +                    H2O



(vi) Oxidation and Reduction

(1) Oxidation.
(a) It is a process in which Oxygen or an electronegative element is added.
(b) It can also be defined as a process in which Hydrogen or an electropositive element is  removed.
(c) In terms of electronic concept, oxidation is a process in which loss of electrons takes place.
4Na + O2      →       2 Na2O (Addition of oxygen)
2Mg + O2    →    2 MgO (Addition of oxygen)
Fe2+        →            Fe3+ + e- (Loss of electron)
2 Fe + 3 Cl2  →    2 FeCl3  (Addition of electronegative element)
Hg2Cl2      →     Hg + HgCl2  (Removal of electropositive element)
CH3CH2OH       \mathrel{\mathop{\kern0pt\longrightarrow}\limits_{300^\circ C}^{Cu}}       CH3CHO + H2 (Removal of Hydrogn)


(2) Reduction. 

(a) It is a process in which addition of Hydrogen or an electropositive element  takes place.
(b) It is also defined as a process in which Oxygen or an electronegative element is removed.
(c) In electronic concept, reduction process involves gain of electrons.

2 Na                       +                 H2             →          2NaH                      (Addition of Hydrogen)
CuO                       +                H2              →        Cu + H2O                (Removal of Oxygen and  addition of hydrogen)
Fe3+                      +                  e-              →          Fe2+                         (Gain of electron)
Hg                          +               HgCl        →        Hg2Cl                   (Addition of an electropositive element)
Cu                          +              CuCl2          →       Cu2Cl2                      (Addition of an electropositive element)
AuCl3                 →                  AuCl          +                Cl2                       (Removal of an electronegative element)


Redox Reactions : Those reactions in which oxidation and reduction take place simultaneously are called redox reactions, e.g.,


Mn4++ 2e-                Mn2+            (Reduction) [Gain of electrons]
2 Cl–                             Cl2 + 2e     ( Oxidation) [Loss of electron]

(ii) Cu (s) + I2 (s)          →         CuI2 (s)
Cu (s)                             →            Cu2+ (aq) + 2e–                     (Oxidation) [Loss of electron]
I2 (s) + 2e-                    →           2I-                                               (Reduction) [Gain of electron]
Cu (s) + I2 (s)              →           CuI2 + 2e– 
Cu (s) + I2 (s)               →          CuI2 (s)                      is a Redox reaction.


Effects of oxidation in everyday life

Oxidation has damaging effect on metals as well as on food. The damaging effect of oxidation on metals is studied as corrosion and that on food is studied as rancidity.

Thus there are two common effects of oxidation reactions are as
(I) Corrosion of metals
(II) Rancidity of food

(i) Corrosion of metals:– Corrosion is the process of deterioration of metals as a result of its reaction with air, moisture and acids. (Present in environment) surrounding it. The corrosion causes damage to buildings, bridges, ships and many other articles especially made of iron.
Rust: Iron corrode readily when exposed to moisture and gets covered with a brown flaky substance called rust. It is called rusting of iron, Rust is a hydrated Iron (III) oxide. [Fe2O3 · 2H2O]

Rusting of iron takes place under the following conditions –
(a) Presence of air (or oxygen)
(b) Presence of water (or moisture) It has been observed that
(c) Presence of impurities in the metal speed up the rusting process. Pure iron does not rust.
(d) Presence of electrolytes in water also speeds up the process of rusting
(e) The position of the metal in the electrochemical series determines the extent of corrosion. More the reactivity of the metal, there will be more possibility of the metal getting corroded.


Other examples of corrosion are –
(i) Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface and acquires a green coating of basic copper carbonate.
(ii) Silver articles become black after sometime when exposed to air because it reacts with sulphur to form a coating of silver sulphide.
(iii) Lead or stainless steel lose their lusture due to corrosion.
(iv) Unreactive metals such as Gold, Platinum, Palladium, Titanium etc. do not corrode.


  •  Prevention of Rusting.
    1. The iron articles should be painted.
    2. The machine parts should be oiled and greased.
    3. Galvanised iron pipes are used for water supply.
    4. Iron can be coated with chromium to prevent rusting.

(II) Rancidity. The oxidation of oils or fats in food, resulting into a bad taste and bad smell is called rancidity. It is caused due to prolonged exposure of food in air. Oxygen present in air oxidise fats/oil present in food and form volatile substances, which have bad odour.


Prevention of rancidity :–
(i) Rancidity can be prevented by adding antioxidants to foods containing fats and oils. Antioxidants are reducing agents so when they are added to food it do not get oxidised easily and hence do not turn rancid. The two common anti oxidants are –
(a) BHA (Butylated Hydroxy Anisole)
(b) BHT (Butylated Hydroxy Toluene)

(ii) Rancidity can be prevented by packaging fat and oil containing foods in nitrogen gas.
(iii) It can be retarted by keeping food in refrigerator.
(iv) It can also be retarded by storing food in air tight containers.
(v) It can be retarded by storing foods away from light.



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IBPS Clerk 2017 Video Lecturesx