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Chapter: 11th 12th std standard Class Organic Inorganic Physical Chemistry Higher secondary school College Notes

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Balancing chemical equation in its molecular form

A chemical equation is called balanced equation only when the numbers and kinds of molecules present on both sides are equal. The several steps involved in balancing chemical equation are discussed below.

Balancing chemical equation in its molecular form

 

A chemical equation is called balanced equation only when the numbers and kinds of molecules present on both sides are equal. The several steps involved in balancing chemical equation are discussed below.

Example 1

Hydrogen combines with bromine giving HBr

H2 + Br2   -> HBr

 

This is the skeletal equation. The number of atoms of hydrogen on the left side is two but on the right side it is one. So the number of molecules of HBr is to be multiplied by two. Then the equation becomes

 

H2 + Br2    -> 2HBr

This is the balanced (or) stoichiometric equation.

Example 2

 

Potassium permanganate reacts with HCl to give KCl and other products. The skeletal equation is

KMnO4 + HCl   -> KCl + MnCl2 + H2O + Cl2

 

If an element is present only one substance in the left hand side of the equation and if the same element is present only one of the substances in the right side, it may be taken up first while balancing the equation.

 

According to the above rule, the balancing of the equation may be started with respect to K, Mn, O (or) H but not with Cl.

 

There are   

L.H.S.,  R.H.S

K = 1, 1

Mn = 1,       1

O = 4, 1

So the equation becomes

 

KMnO4 + HCl   ->  KCl + MnCl2 + 4H2O + Cl2

 

Now there are eight hydrogen atoms on the right side of the equation, we must write 8 HCl.

 

KMnO4 + 8HCl    -> KCl + MnCl2 + 4H2O + Cl2

 

Of the eight chlorine atoms on the left, one is disposed of in KCl and two in MnCl2 leaving five free chlorine atoms. Therefore, the above equation becomes

KMnO4+8HCl    -> KCl+MnCl2+4H2O+5/2 Cl2

Equations  are  written  with  whole  number  coefficient  and  so  the  equation is multiplied through out by 2 to become

2KMnO4+16 HCl ->  2KCl+2 MnCl2+8H2O+5Cl2

Redox reactions [Reduction - oxidation]

 

In our daily life we come across process like fading of the colour of the clothes, burning of the combustible substances such as cooking gas, wood, coal, rusting of iron articles, etc. All such processes fall in the category of specific type of chemical reactions called reduction - oxidation (or) redox reactions. A large number of industrial processes like, electroplating, extraction of metals like aluminium and sodium, manufactures of caustic soda, etc., are also based upon the redox reactions. Redox reactions also form the basis of electrochemical and electrolytic cells. According to the classical concept, oxidation and reduction may be defined as,

 

Oxidation is a process of addition of oxygen (or) removal of hydrogen

 

Reduction is a process of removal of oxygen (or) addition of hydrogen.

 

Example

Reaction of Cl2 and H2S

H2S  +  Cl2  - >  2HCl  +  S

Oxidation : H2S   -> S

Reduction : Cl2 -> 2HCl 

In the above reaction, hydrogen is being removed from hydrogen sulphide (H2S) and is being added to chlorine (Cl2). Thus, H2S is oxidised and Cl2 is reduced.

Electronic concept of oxidation and Reduction

 

According to electronic concept, oxidation is a process in which an atom taking part in chemical reaction loses one or more electrons. The loss of electrons results in the increase of positive charge (or) decrease of negative of the species. For example.

Fe2+   -> Fe3+ + e- [Increase of positive charge]

Cu  ->  Cu2+ + 2e- [Increse of positive charge]

The species which undergo the loss of electrons during the reactions are called reducing agents or reductants. Fe2+ and Cu are reducing agents in the above example.

 

Reduction

 

Reduction is a process in which an atom (or) a group of atoms taking part in chemical reaction gains one (or) more electrons. The gain of electrons result in the decrease of positive charge (or) increase of negative charge of the species. For example,

Fe3+ + e-  -> Fe2+ [Decrease of positive charges]

Zn2+ + 2e-   ->  Zn [Decrease of positive charges]

 

The species which undergo gain of electrons during the reactions are called oxidising agents (or) oxidants. In the above reaction, Fe3+ and Zn2+ are oxidising agents.

 

Oxidation Number (or) Oxidation State

 

Oxidation number of the element is defined as the residual charge which its atom has (or) appears to have when all other atoms from the molecule are removed as ions.

 

Atoms can have positive, zero or negative values of oxidation numbers depending upon their state of combination.

General Rules for assigning Oxidation Number to an atom

 

The following rules are employed for determining oxidation number of the atoms.

 

1.    The oxidation number of the element in the free (or) elementary state is always Zero.

Oxidation number of Helium in  He = 0

Oxidation number of chlorine in Cl2 = 0

 

2.    The oxidation number of the element in monoatomic ion is equal to the charge on the ion.

 

3.   The oxidaton number of fluorine is always - 1 in all its compounds.

 

4.    Hydrogen is assigned oxidation number +1 in all its compounds except in metal hydrides. In metal hydrides like NaH, MgH2, CaH2, LiH, etc., the oxidation number of hydrogen is -1.

 

5.    Oxygen is assigned oxidation number -2 in most of its compounds, however in peroxides like H2O2, BaO2, Na2O2, etc its oxidation number is -

 

6.     Similarly the exception also occurs in compounds of Fluorine and oxygen like OF2 and O2F2 in which the oxidation number of oxygen is +2 and +1 respectively.

 

7.    The oxidation numbers of all the atoms in neutral molecule is Zero. In case of polyatomic ion the sum of oxidation numbers of all its atoms is equal to the charge on the ion.

 

8.    In binary compounds of metal and non-metal the metal atom has positive oxidation number while the non-metal atom has negative oxidation number. Example. Oxidation number of K in KI is +1 but oxidation number of I is - I.

9.    In binary compounds of non-metals, the more electronegative atom has negative oxidation number, but less electronegative atom has positive oxidation number. Example : Oxidation number of Cl in ClF3 is positive (+3) while that in ICl is negative (-1).

Problem

 

Calculate the oxidation number of underlined elements in the following species.

CO2, Cr2O72-, Pb3O4, PO43-

Solution

 

1. C in CO2. Let oxidation number of C be x. Oxidation number of each O atom = -2. Sum of oxidation number of all atoms = x+2 (-2) x - 4.

As it is neutral molecule, the sum must be equal to zero.

 

\ x-4 = 0 (or) x = + 4

 

2. Cr in Cr2O72-. Let oxidation number of Cr = x. Oxidation number of each oxygen atom =-2. Sum of oxidation number of all atoms

2x + 7(-2)  =  2x - 14

Sum of oxidation number must be equal to the charge on the ion.

Thus, 2x - 14  =  -2

2x  =  +12

x  =  12/2

x  =  6


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