A Clear Explanation on the Types of Chemical Equilibrium

Chemical equilibrium is defined as the state in which both products and reactants are present in the present concentrations that have no further tendency to change with time, such that there is no observable change in the properties of the system that has been established. 

As a result, the forward reaction proceeds at the same pace as the reverse reaction, resulting in the state described above. When comparing backward and forward reflexes, it is important to note that they are not always comparable. 

As a result, there are no significant changes in the concentrations of reactants and products. A state like this is referred to as a dynamic equilibrium state.

Chemical Equilibrium Can Be Classified into Several Types

However, although the macroscopic equilibrium concentrations remain constant in time, there are reactions taking place at the molecular level.

If the acetic acid is dissolved in water, it generates the acetate and hydronium ions 

CH₃CO₂H + H₂O→CH₃CO₂H + H₃O⁺          

which is acetate and hydronium.

Chemical Equilibrium can be classified into several types.

It is possible to classify chemical equilibrium into two categories:

• Equilibrium in a homogeneous environment
• Equilibrium in a Heterogeneous Environment

Chemical Equilibrium in a Homogeneous Environment

Homogeneous Equilibrium is characterised by the fact that the reactants and products of chemical equilibrium are all in the same phase. A further division can be made into two categories:

a) Reactions in which the number of molecules of the products is equal to the number of molecules of the reactants are considered to be equilibrium.

As an illustration,

H₂ (g) plus I₂ (g) equals 2HI (g)

2NO = N₂(g) + O₂(g) = N₂O (g)

b) Reactions in which the number of molecules produced by the products is less than the total number of molecules produced by the reactants. 

Chemical Equilibrium in a Heterogeneous Environment

Heterogeneous Equilibrium is characterised by the presence of reactants and products of chemical equilibrium in various phases at the same time.

Examples of heterogeneous equilibrium are provided in the next section.

2CO = CO₂ (g) + C (s) = CO₂ (g)

When CaCO₃ (s) is combined with CO₂ , the result is CO₂ (s) (g)

The Most Important Factors in Chemical Equilibrium

Temperature, pressure, the presence or absence of a catalyst, and the concentration of the solution are all factors that influence equilibrium. 

Some of the most important factors that influence chemical equilibrium are addressed in greater detail below:-

Change in Pressure

The change in pressure occurs as a result of the change in volume.

It is possible to influence the gaseous reaction.

If there is a change in pressure, this is due to the fact that the total quantity of gaseous reactants and products has changed. 

Le Chatelier’s principle states that the change in pressure in both liquids and solids can be ignored in heterogeneous chemical equilibrium because the volume is independent of pressure in heterogeneous chemical equilibrium.

Change in Temperature

According to Le-Principle, Chatelier’s the influence of temperature on chemical equilibrium is dependent on the sign of the H of the reaction.

A drop in the equilibrium constant of an exothermic reaction occurs as the temperature of the reaction increases.

Temperature increases an endothermic reaction’s equilibrium constant, which causes the equilibrium constant to grow.

The change in temperature has an effect on both the rate of reaction and the equilibrium constant, as well as the equilibrium constant itself.

According to Le Chatelier’s principle, as the temperature of an exothermic reaction rises, the equilibrium changes towards the reactant side.

However, when the temperature of an endothermic reaction rises, the equilibrium shifts towards the product side, indicating that the process is endothermic.

Concentration

A change in the concentration of one of the substances in an equilibrium system is normally caused by the addition or removal of one of the reactants or products from the equilibrium system. Consider the following description of the Haber-Bosch process for the industrial manufacture of ammonia from nitrogen and hydrogen gases:

N₂(g)+3H₂(g)⇌2NH₃(g)

Increasing the concentration of a single substance in a system causes the system to respond by favouring the reaction that consumes the substance at the higher concentration.

Catalysts are used to accelerate the rate of a reaction, and it would be reasonable to expect that this would have an impact on the equilibrium state. 

The forward and reverse rates, on the other hand, are both affected equally by catalysts, therefore for a system that is already at equilibrium, these two rates remain equal as well. 

Even though a catalyst speeds up the process of reaching equilibrium, the equilibrium position itself is unaffected by the presence of the catalyst.

Conclusion

Chemical equilibrium is the thermodynamic equilibrium in a system in which both direct and reverse chemical reactions are conceivable, according to the definition given above.

When chemical equilibrium is achieved in a system, the rates of all reactions proceeding in two opposite directions are equal to one another.

As a result, at a given temperature, the macroscopic properties of the system stay constant, and the connection between the concentrations of reacting chemicals remains constant.

It is useful in a variety of industrial operations, such as the creation of ammonia with the aid of Haber’s method, which makes use of it. 

Ammonia is formed when nitrogen gas reacts with hydrogen gas in this location. The production of ammonia is enhanced at low temperatures and high pressures, as well as in the presence of iron, which acts as a catalyst. In the manufacture of sulphuric acid through the contact process, it is employed as follows:

The primary reaction in this process is the oxidation of sulphur dioxide to form sulphur trioxide.