The Concept of 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 (that is, either forward or backward direction), or it can go in only 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.

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 (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. 

Relation of rate constant with dynamic equilibrium 

Whenever the reaction is in a state of dynamic equilibrium, it will always have a specific rate constant, i.e., equilibrium constant (Keq). It is a coefficient that shows the number of products and reactants in a chemical reaction at a given point in time.

Let us take an example, where reactants are [A] and Products are [B], and the reaction is at dynamic equilibrium, then the equilibrium constant will be given by,

For the reaction 

[A] ⇌ [B]

The equilibrium constant K is:

K = [B]eq/ [A]eq

The value of the Equilibrium constant tells the amount of product and reactant at equilibrium at a given point of time.

If Keq > 1000 at equilibrium, then there will be primarily products.

If Keq is between .001 and 1000, there will be many reactants and products.

If Keq <.001, then mostly there will be reactants.

Henry’s law : 

It was given by William Henry, English physician and chemist, in 1803. It states that the weight of gas dissolved by a liquid is directly proportional to the pressure of the gas on the liquid. Also, the partial pressure helps detect/predict the tendency to dissolve because the gases with high partial pressure have more molecules and bounce in the solution than those with low partial pressure.

Not only for gaseous substances, but Henry’s law is also applicable to other substances that are not gaseous, for example, applied in the equilibrium of organic pollutants in water based on the concentration of pollutants present in media in which it is suspended in.

The Mathematical expression of Henry’s law is as follows:

p ∝ C (or) p = KH.C

where,

p= is the partial pressure of solute in the gas above the solution;

C = concentration of solute;

KH = Henry’s constant depends on the solute, solvent, and temperature.

Solubility predicts the tendency of a substance to go towards equilibrium in solution, which explains why gases with the same partial pressure may/may not have different tendencies to dissolve. 

Other applications of Henry’s law 

The other utmost important application of Henry’s law is in respiration. It is used to predict how the gases dissolve in alveoli in the bloodstream during gas exchange. The amount of dissolved oxygen in the bloodstream is directly proportional to the partial pressure of oxygen in alveolar air. Oxygen has higher partial pressure in alveolar air than in deoxygenated blood, so oxygen has a high tendency to dissolve into deoxygenated blood. 

Vice-versa, Carbon dioxide has more extensive partial pressure in deoxygenated blood than in alveolar air; it will diffuse out of solution and go back to gaseous form again. Since partial pressures in the bloodstream and alveoli are very low for carbon dioxide compared to oxygen. Carbon dioxide will have a much higher solubility in plasma than blood; oxygen, having large partial pressure gradients to diffuse quickly in the bloodstream, will not have any effect due to its lower solubility in blood during the gaseous exchange. 

Therefore, partial pressure and solubility of carbon dioxide and oxygen help determine how they will behave during a gaseous exchange in respiration.

Examples of dynamic equilibrium 

• Whenever we take a new aerated drinks bottle, there is a fixed concentration of carbon dioxide in the bottle, but when once opened, half of the drink is poured out. Again the bottle is sealed. Over time, the carbon dioxide will leave the liquid phase and will keep on decreasing over and over. At the same time, the partial pressure of carbon dioxide in the gas phase keeps on increasing until they reach the equilibrium state. Once the equilibrium is reached, at that time, the rate of transfer of carbon dioxide will be the same from gas to liquid and liquid to gas.
• The industrial synthesis of ammonia by Haber’s process is also an application.

             N2(g) + 3H2(g) ⇌ 2NH3

Conclusion 

A chemical reaction can occur forward, backward, or maybe in both directions. The reactions that occur in both directions are called reversible reactions. Dynamic equilibrium occurs only in the case of reversible reactions. At every reaction in dynamic equilibrium, there is a specific rate of constant, known as the equilibrium constant represented by Keq. It tells the number of reactants and products at a given point in time. Henry’s law is also an example of dynamic equilibrium, which states that the amount of gas dissolved in a liquid is directly proportional to the partial pressure of the gas above the liquid at a constant temperature.