An understanding of the relationship between products and reactants when a chemical reaction reaches equilibrium is provided by the equilibrium constant of a chemical reaction (which is typically denoted by the symbol K). For example, the equilibrium constant of concentration (denoted by Kc) of a chemical reaction at equilibrium can be defined as the ratio of the concentration of products to the concentration of reactants, both raised to their respective stoichiometric coefficients. Remember that there are several different types of equilibrium constants, each of which provides relationships between the products and reactants of equilibrium reactions in terms of a different number of different units.

As the ratio between the amounts of reactant and product in a chemical reaction, the equilibrium constant can be defined as the quantity that determines the chemical behaviour of the reaction.

At equilibrium, Rate of the forward reaction = Rate of the backward reaction

i.e. rf = rb Or, kf × α × [A]a[B]b =  kb × α × [C]c [D]d

During a specific temperature range, the rate constants are always the same. As a general rule, the relationship between the rate constants for forward reaction and backward reaction should be constant, and this is referred to as an equilibrium constant (Kequ).

Equilibrium Constant Formula

Kequ = kf/kb = [C]c [D]d/[A]a [B]b = Kc

where Kc is the equilibrium constant expressed in moles per litre and represents the equilibrium constant.

In the case of gaseous reactions, the following is true: In terms of partial pressure, the equilibrium constant formula will be as follows:

Kequ = kf/kb = [[pC]c [pD]d]/[[pA]a [pB]b] = Kp

When expressed in terms of partial pressures, Kp denotes the equilibrium constant formula in terms of partial pressures.

• Higher Kc/Kp values indicate greater product formation as well as higher percentage conversion rates.

• The lower the Kc/Kp values, the less product is formed and the lower the percentage conversion.

• Kc/Kp values in the middle range indicate optimal product formation.

Units of Equilibrium Constant

The equilibrium constant is defined as the ratio of the concentrations raised to the stoichiometric coefficients in a system. As a result, the unit of the equilibrium constant is equal to [Mole L-1]△n.

where ∆n = the sum of the stoichiometric coefficients of the products – sum of the stoichiometric coefficients of the reactants

Importance of Chemical Equilibrium

It is useful in a variety of industrial processes, including, for example,

• When nitrogen reacts with hydrogen to form ammonia, the yield of ammonia is higher at lower temperatures and higher pressures, and when iron serves as a catalyst, the yield of ammonia is higher still. • Production of ammonia by the Haber’s process
• This process is characterised by the oxidation of sulphur dioxide to sulphur trioxide, which is the fundamental reaction in the formation of sulfuric acid. Chemical equilibrium is required for this.

Industrial application

This are:

• Vegetable oil hydrogenation is a common industrial application of hydrogenation. 
• Vegetable oil has long unsaturated carbon chains that are converted into saturated fatty acids, resulting in the production of vegetable ghee.

Applications of Equilibrium Constants

• The expression for the equilibrium constant is valid only when all of the reactants and products have reached a constant concentration at the equilibrium state.
• Its value is unaffected by the initial concentrations of the reactants and products in the reaction mixture.
• It is affected by the temperature. It demonstrates that a balanced equation has a unique value at a given temperature.
• The equilibrium constant for the reverse reaction is equal to the inverse of the equilibrium constant for the forward reaction in the forward reaction.
• Whenever a reaction has reached equilibrium, the equilibrium constant of that reaction can be calculated from the equation of the original reaction, which is obtained by multiplying or dividing it by a small integer (in this case, 1).
• It is employed in order to predict the extent of a reaction based on the magnitude of the reaction.
• It is employed in order to predict the direction of the response.
• When calculating equilibrium concentrations, it is used as a reference point.

Equilibrium Constant for Predicting the Direction of a Reaction

The equilibrium constant can be used to predict the direction of a reaction by calculating the rate of reaction. A term called the reaction quotient (Qc expressed in terms of concentrations or Qp expressed in terms of partial pressures) is required, which is analogous to the equilibrium constant, with the exception that the conditions are not in equilibrium.

The equation for a balanced reaction is aA + bB = cC + dD.

The reaction quotient (Qc or Qp) is expressed as follows:

Qc = [C]c[D]d/[A]a[B]b

Qp = pcC × pdD / paA × pbB

When compared to Kc, and when considering the direction of Reaction,

• In the case where Q = Kc, the reaction is in equilibrium [where Kc is defined as the equilibrium constant].
• If Q is greater than Kc, Q will tend to decrease until it is equal to K. As a result, the reaction will proceed in the opposite direction of its initial direction.
• If Q is greater than Kc, Q will tend to increase until it is equal to K. As a result, the reaction will proceed in the direction of the forward arrow.

Factors Affecting Equilibrium Constant

The following are some of the factors that influence the equilibrium constant: 

1. Changes in the concentration of any product or reactant
2. Pressure in the system changes as a result of this.
3. Temperature fluctuation in the system is the third factor.
4. The addition of inert gas.
5. Incorporating a catalyst.

What is hydrogenation? What is its industrial application?

When a catalyst is present, hydrogenation occurs, which is an addition reaction between hydrogen and other compounds.

As an illustration, the hydrogenation of ethene involves the addition of two hydrogen atoms across the double bond of ethene, with the resultant formation of saturated ethane. Because the energy of the reactants is greater than the energy of the product, the reaction is more favourable than it would otherwise be. The reaction is an exothermic reaction, which means it generates heat.

CH2 = CH2 → CH3CH3

Conclusion

‘Equilibrium’ refers to the state of a process in which the properties of the system (such as its temperature, pressure, and concentration) do not change with the passing of time.

Whenever a reaction reaches chemical equilibrium, the rates of the forward reaction and the backward reaction will be the same.

It is referred to as an equilibrium mixture when it contains a mixture of reactants and products that has reached equilibrium.