What is Chemical Equilibrium

Chemical equilibrium in a reversible reaction can be defined as the state at which two opposing reactions occur at the same speed. The reaction doesn’t stop.

Let us consider an example of chemical equilibrium:

2NO2 (g)→ N2O4 (g)

NO2 is reddish-brown coloured, while N2O4 is colourless. NO2 contained in a sealed, evacuated glass vessel at 25℃ starts converting to colourless N2O4. However, gradually, it is seen that the brown colour intensity remains constant, signifying that the concentration of NO2 is no longer changing. At this stage, a state of equilibrium is reached.

If the reaction is of type: jA + kB ⇌ iC +mD.

The law of mass action can be represented by the equilibrium expression given below:

 K = [C]i [D]m∕ [A]j [B]k.

Here, K is the equilibrium constant. 

i,j,m and k are stoichiometric coefficients of A, B, C and D respectively.

Types and examples of chemical equilibrium

There are two types of chemical equilibria: 

(i) Homogeneous

(ii) Heterogeneous.

Homogeneous equilibrium:

The reversible reactions in which all the reactants and products are present in the same state, i.e., only one phase is present, are homogeneous reactions. Equilibrium taking place in such reactions is known as homogeneous equilibrium.

Homogeneous reactions can be further classified into three types:

First type: The number of molecules does not change in a reaction. Examples include:

(a)   H2 (g) + I2 (g) ⇌ 2HI (g).

(b)   CH3COOH(l) + CH3CH2OH (l) ⇌ CH3COOCH2CH3 (l) + H2O (l).

Second type: There is an increase in the number of molecules in a reaction. Examples include:

(a)   PCl5 (g)⇌ PCl3 (g) + Cl2 (g).

(b)    2NH3 (g) ⇌ N2(g) + 3H2(g).

Third type: There is a decrease in the number of molecules in a reaction. Examples include:

(a)    N2(g) + 3H2(g) ⇌ 2NH3 (g).

(b)   2SO2 (g) + O2 (g) ⇌ 2SO3 (g).

The value of the equilibrium constant is also dependent on certain factors in these reactions:

• The reaction’s mode of representation: Concentrations of products are in the numerator, whereas the concentrations of reactants are in the denominator.

A + B ⇌ C+ D.

Here, equilibrium constant K can be written as,

K= [C][D]/[A][B].

However, if we consider the reverse reaction,

C + D⇌ A+ B.

Now, K’ = [A][B]/[C][D] = 1/ K.

• Stoichiometric representation: When a reversible reaction can be written in 2 or more stoichiometric equations, the value of K will differ in each case. If we multiply a balanced equation with any value n, the new equilibrium constant will now be equal to Kn.

A+ B⇌ C+ D,  K= [C][D]/[A][B].

nA +nB ⇌ nC + nD,  K’ = [C]n[D]n/ [A]n[B]n = Kn.

• Use of partial pressures in place of concentrations: When the reaction occurs in the gaseous phase, partial pressure can be used in place of concentration as the partial pressure of a substance is proportional to its concentration in the gas phase.

From the ideal gas equation, PV= nRT,

Or P= (n/V) RT = CRT.

Where, C is the number of moles of gas per unit volume.

For the general reaction, 

jA + kB⇌ lC + mD.

Let the partial pressures be PA, PB, PC and PD, respectively, at equilibrium.

So, Kp = (Pcl)(PDl) / ( PAj)( PBk).

           = ( CC× RT)l ( CD RT)m / (CA RT)j ( CBRT)k.

                 = [C]l[D]m/[A]j[B]k ×(RT)(l+m)-(j+k) = K (RT)∆n.

Where, ∆n = difference in the sums of coefficients for gaseous reactants and products.

Units of the equilibrium constant:

K has no unit for a reaction where the reactants and products are equimolar.

In general, the unit of K = [M]∆n.

Where, M= mol litre-1 and ∆n= number of moles of gaseous products-number of moles of gaseous reactants.

Similarly, the unit of Kp = [atm]∆n.

 Heterogeneous equilibrium: 

The reversible reactions in which more than one phase is present are known as heterogeneous reactions. The equilibrium taking place in heterogeneous reactions is called heterogeneous equilibrium.

Examples include:

(a)    CaCO3 (s) ⇌ CaO (s) + CO2 (g).

(b)   H2O (g) + C (s) ⇌ H2 (g) + CO (g).

If we consider the example of decomposition of calcium carbonate and try to derive its equilibrium expression, we get K’ = [CO2][ CaO]/ [CaCO3].

However, experimentally, it has been determined that the concentrations of pure solid and pure liquid phases don’t appear in the expression of an equilibrium constant for a heterogeneous reaction. Here, CaCO3 is a pure solid; its activity is taken as unity. Similarly, the activity of CaO is also unity as the reference states are pure CaCO3 and pure CaO, respectively.

So, the equilibrium expression can be written as,

K= [CO2](1)/ (1) = [CO2].

And Kp = PCO2 (1) / (1) = PCO2

To generalise, we can say that in the case of a pure solid or a pure liquid, activity is always 1.

Conclusion

To conclude, we can say that there are two chemical reaction types, i.e., reversible and irreversible reactions. Chemical equilibrium is a characteristic property of a reversible reaction. The equilibrium state is the state where the rate of a forward reaction and a backward or reverse reaction is equal.

However, the reaction doesn’t stop, so it is dynamic in nature. 

There are 2 types of chemical equilibrium – heterogeneous and homogeneous.

The equilibrium taking place in a reaction where every component is in the phase is called homogeneous equilibrium, whereas, if two or more phases are present, the equilibrium is heterogeneous.