Definition of dynamic equilibrium reaction

Definition of dynamic equilibrium in chemistry

The definition of equilibrium in chemistry is a dynamic point where no further action apart in the chemical reactions. But the product of the chemical reaction can not be cited if the favorable external condition not maintained. Hence the reaction proceeds to some extent than stop by converting some portion of reactant to product.

If hydrogen and iodine vapor kept at a constant temperature in a closed vessel. The reaction proceeds to some extent and then stop, by converting some portion of hydrogen and iodine to hydrogen iodide.

Definition of dynamic equilibrium point
Equilibrium point

Because after this dynamic point, the reactant and product attained equilibrium.

H2 + I2 ⇆ 2HI

But if you take some hydrogen iodide in the closed vessel at the same temperature. A faction of hydrogen iodide converted into hydrogen and iodine and the rest of the hydrogen iodide remains unchanged.

2HI ⇆ H2 + I2


But in both the experiment, the amount of hydrogen, iodine, and hydrogen iodide remains unchanged. This dynamic point of the chemical reactions is called equilibrium.

So the equilibrium of the chemical reaction maintains the following criteria

  1. Approachability from both ends
  2. Permanency of the equilibrium
  3. The incompleteness of the reaction
  4. Dynamic equilibrium

Reversible reaction and equilibrium point

When water added to bismuth chloride, a milky solution appears due to the formation of bismuth oxychloride.

BiCl3 + H2O ⇆ BiOCl + 2HCl

But when hydrochloric acid solution added to this milky solution, the milkiness of the solution disappears. This shows the reversible nature of this chemical reaction.

Why do chemical reactions occur?

This leading question can not be answered in a simple sentence. Because details study reveals that the two primary factors responsible for the feasibility of a chemical reaction.

  1. The potential energy of the reacting system must be lowered by the reaction.
  2. Reactants posses potentially to react because they must find a suitable path to react at a participle rate under specific conditions.

The fast concern with thermodynamics because it studies under chemical equilibrium but the second one constitutes the study of kinetics.

Thermodynamics of chemical equilibrium

At the equilibrium, all reversible reactions, the total thermodynamics entropy change is zero.

ΔStotal = ΔSsystem + ΔSsurr

But for all irreversible processes or natural chemical reactions the value of total entropy increases and greater than zero.

Thus for an irreversible reaction, the total entropy goes to increases and when it reached an equilibrium point, the entropy attains maximum value.

Because at equilibrium no further reaction takes place or no further possible entropy change occurs. Thus at equilibrium

ΔStotal = 0

Endothermic vs exothermic reactions

All the natural reaction follow a general trend because they take place in a direction which results in an ultimate decrease in the chemical energy of the universe. Thus the exothermic vs endothermic reaction differ by definition as

  1. Exothermic chemical reactions are accompanied by the emission of energy.
  2. But endothermic chemical reactions are accompanied by the absorption of energies.

Many exothermic reactions are reversible thus the reverse process must be endothermic in nature.

For example, the synthesis of ammonia is an endothermic reversible reaction but the reverse process going vs exothermic.

Forward reaction

2NH3 → N2 + 3H2 (ΔH = +92.22 KJ mol-1)

Backword reaction

N2 + 3H2 → 2NH3 (ΔH = -92.22 KJ mol-1)

Gibbs energy change and equilibrium

For the spontaneous process, two factors control the ultimate energy change of the chemical reaction.

  1. Enthalpy change (ΔH).
  2. Entropy change (ΔS).

But these two changes collectively determined by another fundamental property of the system which known Gibbs free energy (ΔG). Van’t Hoff equation proposed the relation between free energy and the equilibrium constant of a reaction.

 He used the Gibbs – Helmholtz equation to derive this energy equation for equilibrium.

Why energy, whether a catalyst used or not?

The heat of a reaction or enthalpy is a state function because ΔH does not change if the initial state and final state are the same.

Thus a catalyst cannot change the initial and final state of a reaction, hence ΔH remains the same whether a catalyst used or not.

Definition of the equilibrium constant

Norwegian Physicists, Guldberg and Waage in 1867 developed the quantitative relation between the amount of the product and reactant at the equilibrium point and this relation is known as mass action law.

When temperature kept constant, the rate of a reaction at equilibrium, proportional to the active masses of the reacting system.

  1. If the solution dilutes, molar concentration or moles/liter used.
  2. Thus when the reactants and product in the gaseous state, palatial pressure used.
  3. For pure solid and pure liquid, active mass assumed to be unity.