Chemical Equilibrium Called Dynamic Equilibrium

Equilibrium is a state of rest due to the equal action of opposing forces. It is dynamic, that is, both forward and reverse reactions continue to take place even when the reaction has stopped. Whenever a chemical reaction takes place, the reaction can go in both directions i.e., either in a forward or backward direction, or it can go only in one order. The reactions that go in two directions are called reversible reactions. For example, consider the chemical reaction below:

H2O(l) ⇌ H+(aq) + OH–(aq)

Dynamic equilibrium is a state of chemical reaction where there is no net change in the number of products and reactions. However, the reaction continuously keeps going on.

Overview of Dynamic Equilibrium 

Dynamic equilibrium can be defined as a system state where reversible reactions occurring in it stop changing the ratio of reactants and products. The reaction continues to occur i.e., there is a movement of substances from reactants to products at an equal rate, and there is no net change in the ratio of reactants and products. It exists only in reversible reactions. The above equation is dynamic because the rate of reactants and products is equal and stable. However, the reaction is continuously occurring. 

Why is Chemical Equilibrium Called Dynamic Equilibrium?

Now, suppose we consider a reaction, say: A + B ⇌ C + D. In this reaction, as you can see, if we combine both the substances, A and B, in a closed container, then the forward reaction occurs, and it produces C and D. In this process, the concentration of A and B constantly decreases, whereas the concentration of C and D simultaneously increases. Hence, the forward reaction rate decreases when the pace of the reverse reaction increases, and then a point comes when both the reactors equalise, achieving the state of equilibrium.

When A + B ⇌ C + D achieve equilibrium, the concentration of all the four components remains constant over the period. Although it may look like the balance is lost at that point, in actuality, it has reached the state of dynamic equilibrium. Therefore, the reverse and forward processes reach equilibrium, however the concentration remains unaffected.       

If we see the above kinetic molecular model, we can easily understand that when A and B collide with each other, then C and D are formed. On the other hand, the same process continues as C and D return to the state of A and B. Although in the process, the equilibrium remains the same.   

Law of Chemical Equilibrium

Consider a reversible homogeneous reaction at equilibrium at a particular temperature, such as:

A+B ⇌ C+D

Let us assume the active masses of A, B, C, and D are [A], [B], [C], and [D], respectively.

As we know, according to the law of mass action, the rate of the forward reaction shall be ∝ [A] [B]. Thus, the rate of the forward reaction = kf [A] [B] where kf is the velocity for the forward reaction.

Similarly, the rate of the backward reaction shall be ∝ [C] [D]. Thus, the rate of the backward reaction = kb [C] [D] where kb is the velocity for the backward reaction.

Now, at equilibrium constant, K= Rate of the forward reaction = Rate of the backward reaction, i.e. kf [A] [B]= kb [C] [D]

Kf / kb= [C] [D] / [A] [B]

Let’s suppose kf  / kb as a new constant, i.e. K.

Kf / kb= K = [C] [D] / [A] [B]

This K is the equilibrium constant. 

Equilibrium constant for molar concentration Kc

Consider a reaction             aA+bB ⇌ cC+dD

Where a, b , c and d are moles of A, B,C and D respectively.

Kc = [ C ]c·[ D ]d / [ A ]a·[ B ]b

Equilibrium constant for gaseous Kp 

Kp= (PCcPDd)/(PAaPBb)

Relation between Kc and Kp

Deriving the ideal gas equation, 

PV = nRT

And, P = (n/V)RT 

P = CRT, where C is the number of moles per litre (molar concentration).

By inserting PA=CART, PB=CBRT, PC=CCRT, and PD=CDRT in the equation of equilibrium constant for gaseous Kp, we get,

Kp = Kc (RT)Δn   

WhereΔn=(c+d)-(a+b),the difference in the sums of the coefficients for the gaseous products and reactants.   

Gibbs free energy and chemical equilibrium 

ΔG = ΔH -TΔS,

Thus, Gibbs free energy is the enthalpy minus the product of absolute temperature with its entropy. The relation of Gibbs free energy with the equilibrium constant is as follows:

G° = -RT ln Keq

Where, T= temperature

R= universal gas constant, and 

Keq= equilibrium constant.

Types of Chemical Equilibrium

The law of chemical equilibrium is of two types: homogeneous equilibrium and heterogeneous equilibrium.

Homogeneous Equilibrium

In this type of equilibrium, all the reacting components are stated in one phase of matter, such as solid, gas, or liquid. These types of reactions are classified in three different ways:

1. The reaction when no mole number changes the net of the system (Δn = 0).

2. Mole number will increase due to reaction time (Δn = +ve).

3. Mole numbers will decrease due to reaction time. (Δn = -ve).

Example

H2 (g) + I2 (g) ⇌ 2HI (g)  , (Δn = 0)

PCl5  ⇌ PCl3 + Cl2  , (Δn = +ve)

N2 + 3H2 ⇌ 2NH3  , (Δn = -ve)

Heterogeneous Equilibrium

In this type of equilibrium, the reacting components do not stay in the same matter phase. For example, calcium carbonates decompose to calcium oxide and dioxide.

 CaCO3 (s) ⇌ CaO (s) + CO2

The equation includes the three different phases of chemical equilibrium. 

Conclusion

The equilibrium state is one in which there is no net change in the centralisations of reactants and items. Regardless of the way that there is no evident change in harmony, this does not imply that all compound responses have stopped. The law of chemical equilibrium may be used to describe how solutions behave in an equilibrium that is dynamic.