Homogeneous Equilibrium

Introduction

Chemical reactions are performed to get a particular product by some chemical process happening in the reactants. In some cases, the product formed is again converted back to the reactant. That is a reverse chemical reaction is taking place. But not all chemical reactions will not undergo such a reverse reaction.

The state in which both forward and reverse reaction takes place simultaneously is called equilibrium. The term equilibrium was introduced in the year 1803 by the discovery of scientist Berthollet in which some reactions are moving in the reverse direction. The rates of such chemical reactions are also the same. Based on the phase of chemical reactions there are two types of equilibrium, homogeneous and heterogeneous equilibrium.

What is Homogeneous Equilibrium?

When all the chemical species that are the product and the reactant involved in the typical chemical reaction that has achieved chemical equilibrium are all in the same phase is a homogeneous equilibrium. This means that if the reactant is in the gas phase the product will also be in the gas phase itself.

• For example, gaseous or homogeneous equilibrium is achieved in the below reaction which involves the conversion of Sulfur dioxide to Sulfur trioxide. The reaction is,

$$\mathrm{2SO_{2}(g)\:+\:O_{2}(g)\:\leftrightarrows\:2SO_{3}(g)}$$

This is a reaction taking place in the contact process where Sulfuric acid is prepared.

Homogeneous Equilibrium- Example.

Some of the examples where homogeneous equilibrium is achieved pointed below.

• An example of homogeneous equilibrium is the reaction between Nitrogen and Hydrogen results in the formation of Ammonia during the Ammonia manufacture in the Haber process. The reaction is,

$$\mathrm{N_{2}(g)\:+\:3H_{2}(g)\:\leftrightarrows\:2NH_{2}(g)}$$

• Liquid or homogeneous equilibrium is achieved in the below reaction where acetic acid is reacting with alcohol. And the process is esterification. The reaction is

$$\mathrm{CH_{3}COOH(l)\:+\:CH_{3}CH_{2}OH(l)\leftrightarrows\:CH_{3}COOCH_{2}CH_{3}(l)\:+\:H_{2}O(l)}$$

• Another example is the formation of Nitrogen monoxide by the reaction of Nitrogen and Oxygen. The reaction is,

$$\mathrm{N_{2}(g)\:+\:O_{2}(g)\:\rightarrow\:2NO(g)}$$

Homogeneous Reaction- Equilibrium Constant

A homogeneous equilibrium constant can be derived by considering a chemical reaction to be in a homogeneous equilibrium. For example,

$$\mathrm{aA\:+\:bB\leftrightarrows\:cC\:+\:dD}$$

The equilibrium constant can be derived by dividing the product of products raised with their exponents by the product of reactants raised with their exponents. The equation is,

$$\mathrm{Kc\:=\:[C]^{c}[D]^{d}/[A]^{a}[B]^{b}}$$

This is the formula for the equilibrium constant when a chemical reaction has achieved homogeneous equilibrium. And is also similar to the equilibrium constant of all the reactions. So, the equilibrium constant does not depend on the phase or state of the chemical species involved in chemical reactions.

Calculate Equilibrium Constant

The equilibrium constant for a reaction can be calculated with the help of the corresponding equation. For example, the below reaction where Sulfur trioxide is formed is an example of homogeneous equilibrium. The reaction is

$$\mathrm{2SO_{2}(g)\:+\:O_{2}(g)\:\leftrightarrows\:2SO_{3}(g)}$$

The equation used for the calculation of the equilibrium constant is,

$$\mathrm{Kc\:=\:[C]^{c}[D]^{d}/[A]^{a}[B]^{b}}$$

That is,

$$\mathrm{Kc\:=\:[SO_{3}]^{2}/[SO_{2}]^{2}[O_{2}]}$$

By substituting corresponding concentrations of chemical species, the value of the equilibrium constant can be easily obtained.

Differences between KC and KP

Kc and Kp are both equilibrium constants for reactions that will not be completely converted into corresponding products. Both the reactant and product are in a reaction where interconversion is possible. The main difference between Kc is an equilibrium associated with the concentration of chemical species while Kp is the equilibrium constant associated with pressure. Some of the differences between these two terms are tabulated below.

Kc	Kp
It is associated with the concentration of chemical species and can be calculated if the concentration is known.	It is associated with the pressure of chemical species involved.
The value can be calculated for both liquids and gases.	The value can be calculated only for chemicals in the gas phase.
It is expressed by the units of concentration.	It is expressed by the units of pressure.
The equation used for calculating Kc is, $\mathrm{Kc\:=\:[C]^{c}[D]^{d}/[A]^{a}[B]^{b}}$	The equation used for calculating the value of Kp is, $\mathrm{kp\:=\:pR^{r}.pS^{s}/pP^{p}.pQ^{q}}$ Partial pressure is p. .

There is an equation that relates both these terms. And is,

$$\mathrm{kp\:=\:kc(RT)^{\Delta\:n}}$$

$$\mathrm{\Delta\:n\:=\:no\:of\:moles.}$$

$$\mathrm{R\:=\:0.082062\:L.atom.K^{-1}mol^{-1}}$$

This relation is obtained from the ideal gas equation,

$$\mathrm{PV\:=\:nRT}$$

Conclusion

Homogeneous equilibrium is achieved by chemical reactions where the reaction and the production involved are all in the same phase. The state of reactants and products is always in the gas phase or the solution phase. There are many chemical reactions of this type. The equilibrium constant of this type of reaction is calculated by using the equation $\mathrm{Kc\:=\:[C]^{c}[D]^{d}/[A]^{a}[B]^{b}}$ . And this value depends greatly on the concentration of corresponding reactants and products. Kp is also another equilibrium constant associated with chemical reactions. The value of Kp is associated with the partial pressure of chemical species while Kc is associated with concentration.

FAQs

1. What factor does the equilibrium constant depend on?

Three factors will affect the equilibrium of a chemical reaction. They are the concentration of the chemical species involved, temperature, and pressure. Will get destroyed in a chemical reaction if one of the factors has been changed.

2. Does catalyst affect equilibrium?

The addition of a catalyst to a reaction where the forward and backward reaction is taking place simultaneously will not has any effect. This is because of the reason that catalyst will alter the rate of forwarding and reverse reaction to the same extent. So, there will not be any change in the equilibrium of that reaction.

3. Give an example of heterogeneous equilibrium?

The chemical reactions where the reactants and products are at different phases are heterogeneous in equilibrium. In most cases, the solid phase is involved. An example is the reaction of steam with carbon results in the formation of Hydrogen gas,

$$\mathrm{H_{2}O(g)\:+\:C(s)\:\longleftrightarrow\:H_{2}(g)\:+\:CO(g)}$$

4. Give the Le chatelier's principle?

The change in any of the factors supplied to a chemical reaction where it has achieved a chemical equilibrium, which will result in the nullification of that change by the system. Thereby reducing the effect of change that has been achieved by the system.

5. Are reaction rate and equilibrium related?

The rate of chemical reactions and equilibrium don't have any relation. Instead, there is a relation in which the rate of forwarding and reverse reactions is taking place at the same rate.

6. What are some examples of heterogeneous equilibrium?

The disintegration of solid calcium carbonate to create 𝐶𝑎𝑂 solid and $\mathrm{CO_{2}}$ gas in this phase with different reactants and products are instances of a heterogeneous equilibrium. Similar to how solid ferrous oxide forms solid iron and gaseous carbon dioxide when combined with gaseous carbon monoxide.