Notes of chapter # 06

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 Chapter #06 

Chemical Equilibrium 

 Prepared by:

 Lecturer  S.Fayyaz Hussain

Chemical Equilibrium

All chemical reactions do not proceed to same extent. Some reactions proceed to completion after sometime, but there are many reactions, which are never completed. On the basis of extent reactions, chemical reactions are classified into:

Ø  Irreversible Reactions

Ø  Reversible Reactions

Irreversible Reactions:

“Those reactions that proceed to complete in a definite direction are called Irreversible Reactions”.

OR

Those reactions in which the reactants react to form product which do not change back are known as “Irreversible Reactions”.

Reversible Reaction:

There are some reactions in which the products again combine to form the reactants. This reaction therefore precedes in two directions i.e. forward & backward directions.

“Such reactions which proceed to both backward & forward directions and are never completed are called Reversible Reactions”.

e.g.

(a)        2HI                              H₂ + I₂

(b)        N₂ + 3H₂                                            2NH₃­

(c)        CH₃COOH + C₂H₅OH                          CH₃COOC₂H₅ + H₂O

The double arrow indicates that the reaction is reversible and both the forward and backward reactions can occur simultaneously.

Equilibrium State:

“A reversible reaction is said to be in a state of Equilibrium, when the rate of its forward reaction equals to the rate of its backward reaction and the concentration of various constituents remains unaltered”.

Explanation:

In a reversible reaction the changes, forward and reverse occur simultaneously. Under these circumstances, a reaction might come to some kind of “Balance” in which the forward and reverse reactions occur at the same rate. For example consider

A + B                          C + D

            In the beginning, forward reaction predominates, but as soon as C & D are formed the reverse reaction builds up until equilibrium position is reached, where the forward as well as the reverse change proceeds with the same rate i.e. At Equilibrium state.                                                                                                       Rate of forward Reaction            =          Rate of backward Reaction

Law of Mass Action:

In 1864, Goldberg and Wage studied the effect of concentration on reversible reactions in equilibrium and put forward a law, which is know as “Law of Mass Action”                              

According to this law,

“The rate of a chemical reactions is directly proportional to the product of active masses (or molar concentrations) of the reactants”.

            The number of moles of a substance in 1dm³ is called its “Molar Concentration” or “Active Mass” and is denoted by square brackets.

Expression of   K_c:

We consider the following reaction,

                     _Forward

mA + nB                     xC + yD

        ^Backward     

According to the law of mass action

or --------- (1)

Similarly,

               --------- (2)

or

At equilibrium state,

Rate of forward = Rate of backward reaction

Where ‘K_c’ is called the “Chemical Equilibrium Constant” and the equation is called the chemical equilibrium constant expression, for the general reversible reaction, where x,y,m & n represents the moles of species and are called co-efficient of chemical equations.

            In the case of gaseous equilibrium, a partial pressure is sued instead of concentration because at a given temperature, partial pressure of a gas is proportional to its concentration. In this case the equilibrium constant is expressed as K_P instead of K_C. e.g. for the following gaseous equilibrium;

A_(g) + B_(g)                    C_(g) + D_(g)

Where P_A, P_B, P_C & P_D are the partial pressures of gases A,B,C, & D respectively.

Application Of Equilibrium Constant:

            The knowledge of chemical equilibrium constant of a chemical reaction is useful for a chemist working in a laboratory or in an industry to predict:

• Direction of the reaction
• Extent of the reaction

I)         Direction Of The Reaction:

            “The value of equilibrium constant K_C is a valuable aid in prediction the direction in which a reaction will shift in order to achieve the equilibrium, provided the initial concentration of the reagent’s involved is known.

            Consider the general reversible reaction

A + B                          C  + D

For which,

Now we come across with three possibilities

1)         If  < K_C

Then, the reaction will proceed to the right, i.e. to the forward direction until equilibrium position is attained.

2)         If  < K_C

Then the reaction will proceed to the left i.e., the backward direction until equilibrium state is obtained.

3)         If  = K_C

Then, the reaction is already in a equilibrium state, and the concentration of reactants and products will remain constant.

ii)         Extent Of A Reaction:

            From the magnitude of equilibrium constant, we can also predict about the extent to which the reaction will take place. In this case, also there may be three possibilities.

1)         When K_C is very large:

e.g. the equilibrium constant, K_C for the following reaction is very large

2O₃                 3O₂

From this large value of K_C, we can conclude that the forward reaction is almost completed. In other words Ozone is very unstable gas and it decomposes spontaneously into oxygen.

2)         When K_C is very small:

            e.g. for the reaction

                        2HF                               H₂ + F₂

            ðK_C = 10¹³

From this low value of K_C, we can conclude that there will be very little tendency for the reaction to occur in the forward direction. In other words, HF is very stable compound.

3)         When K_­C is very moderate:

            e.g. N₂ + 3H₂                           2NH₃

            ðK_C = 10

It indicates that the equilibrium mixture contains both the reactants and products in appreciable quantities and neither of them are reactive nor un reactive.

Le -Chatelier’s Principle:

It is a general principle that gives qualitatively the influence of change in temperature, pressure or concentration on system in equilibrium.

This principle was first enunciated by a French Chemist                      Henn Le-Chatelier in 1884. According to this principle,

“If a system in equilibrium is subjected to a stress, the equilibrium shifts in a direction to minimize or undo the effect of the stress”.

Where “Stress” means change in concentration, temperature or pressure. If one of the factors involved in a chemical equilibrium is altered, the equilibrium shifts towards right or left in order to restore the balance of equilibrium.

1)                  Effect of Concentration Change:

Change in Concentration	Effect on Equilibrium Position A + B                              C + D
Increase in Conc. Of A & B	Equilibrium shifts to right and more C & D is formed
Increase in Conc. of C or D	Equilibrium shifts left & more A & B is forward

2)                  Effect of Temperature Change:

Nature of Reaction	Change in Temperature	Effect on Equilibrium
Exothermic e.g: 2NO+O2            2NO2	Increase	Equilibrium shifts to left i.e. more NO & O2 are formed
	Decrease	Equilibrium shifts towards right & more products are formed
Endothermic e.g: N2 + O2             2NO	Increase	Equilibrium shifts towards right & more products are formed (i.e., yield of NO increases)
	Decrease	Equilibrium shifts to left & more N2 & O2 are present

3.         Effect of Pressure Change:

Volume Involved mA + nB                 xC + yD	Change in pressure	Effect on Equilibrium
If x + y > m + n i.e. Volume of products is less than reactants e.g.: 2SO2 + O2       2SO3	Increase	Equilibrium position moves towards right i.e. more SO3 is formed.
	Decrease	Equilibrium position moves towards left i.e. more SO2 &  O2 is formed
If x + y < m + n i.e. Volume of products is grater than reactants e.g.:  N2O4                   2NO2	Increase	Equilibrium position moves towards left i.e.; Yield of NO2 increases
	Decrease	Equilibrium shifts towards right i.e. more N2O4 is produce.
If x + y = m + n	Increase or decrease	No effect

4.         Effect of Catalyst:

            A catalyst has not effect on the equilibrium position, but it enables equilibrium to the reached more quickly by decreasing the “Energy of activation”. In fact a catalyst affects forward and reverse states equally.

Applications of Le-Chatelier’s Principle:

1.         The Haber’s Process:

            Ammonia can be prepared by the Haber’s process as;

            N₂ + 3HH₂Hhhhaa

H₂                    2NH_­3   (DH = -46.2)

i.e. The reaction between nitrogen and hydrogen to produce ammonia is accompanied by decrease in volume and it is exothermic.

§  Effect of Concentration:

According to Le-Chatelier’s Principle, the addition of more N₂ or H₂ or both will move the reaction to the right, thus more NH₃ gas will be produced till the equilibrium state is reached again.

Similarly addition of NH₃ at the equilibrium state will move the reaction to the left. Thus, a part of NH₃ gas will decompose into N₂ & H₂ gases in order to reach the equilibrium state again.

§  Effect of Temperature:

It is an exothermic reaction i.e. heat is evolved during the reaction. According to Le-Chatelier’s principle, the law temperature shifts the equilibrium to the right ion i.e.

Common Ion Effect:

(i)     CH₃COONa                     CH₃COO^- + Na⁺

         (ii)     CH₃COOH_(aq)                   CH₃COO^- + H⁺

If sodium acetate is added to an acetic acid solution, it ionizes into acetate and sodium ions as

CH₃COONa                         CH₃COO^- + Na⁺

Therefore acetate ion concentration is increased and equilibrium will shift. Since there are more CH₃COO^- ions, so according to Le-Chatelier’s principle the rate of reaction will increase toward acetic acid. Some of the excess acetate ions unite with H⁺ ions to form molecular acetic acid, and hence the degree of ionization of acetic acid will reduce.

The acetate ion is common to both acetic and sodium acetate. The effect of the acetate ion on the acetic acid and solution is called “Common Ion Effect”.

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