What is Chemical Equilibrium?-Factors, Types, And Applications

Chemical Equilibrium Definition

Chemical equilibrium refers to the state in which the concentrations of reactants and products remain constant over time in a chemical reaction. While the forward and reverse reactions continue to occur, they do so at rates that are equal, resulting in no net change to the amount of reactants and products present.

This is an example of a dynamic equilibrium, where there is balance despite ongoing change at the molecular level.

For a general reaction:

aA + bB ⇌ cC + dD

The system is at equilibrium when the rate of the forward reaction (aA + bB → cC + dD) equals the rate of the reverse reaction (cC + dD → aA + bB). The equilibrium concentrations of A, B, C, and D remain constant, though reactions are still occurring in both directions.

Factors Affecting the Chemical Equilibrium

There are several ways to disturb this state of equilibrium and cause a net shift towards products or reactants.

Concentration

When the concentration of any of the reactants or products in an equilibrium reaction is changed, the equilibrium mixture’s composition changes in order to minimize the effect of the concentration change. This is known as Le Chatelier’s principle.

For example, consider the following equilibrium reaction:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

If the concentration of N₂ is increased, the equilibrium will shift to the right to produce more NH₃. This is because the system is trying to reduce the concentration of N₂. Conversely, if the concentration of NH₃ is increased, the equilibrium will shift to the left to consume some of the NH₃.

Temperature

Temperature also affects the equilibrium position of a chemical reaction. According to Le Chatelier’s principle, when the temperature of a system is increased, the equilibrium will shift to the side of the reaction that is endothermic (absorbs heat). Conversely, when the temperature of a system is decreased, the equilibrium will shift to the side of the reaction that is exothermic (releases heat).

For example, the following equilibrium reaction is endothermic:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Therefore, if the temperature of the system is increased, the equilibrium will shift to the right to produce more NH₃. Conversely, if the temperature of the system is decreased, the equilibrium will shift to the left to consume some of the NH₃.

Pressure

Pressure only affects the equilibrium position of a chemical reaction when there is a difference in the number of moles of gas between the reactants and products. According to Le Chatelier’s principle, when the pressure of a system is increased, the equilibrium will shift to the side of the reaction with fewer moles of gas.

Conversely, when the pressure of a system is decreased, the equilibrium will shift to the side of the reaction with more moles of gas.

For example, the following equilibrium reaction has fewer moles of gas on the right side:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Therefore, if the pressure of the system is increased, the equilibrium will shift to the right to produce more NH₃. Conversely, if the pressure of the system is decreased, the equilibrium will shift to the left to consume some of the NH₃.

Catalysts

A catalyst is a substance that increases the rate of a chemical reaction without being consumed by the reaction itself. Catalysts do not affect the equilibrium position of a chemical reaction, but they can help the system reach equilibrium more quickly.

For example, the following equilibrium reaction is catalyzed by iron:

3H₂(g) + N₂(g) ⇌ 2NH₃(g)

If iron is added to the system, the rate of both the forward and reverse reactions will increase. However, the equilibrium position of the reaction will not change.

Examples of Chemical Equilibrium

An everyday example of chemical equilibrium is the dissociation of carbonic acid (H₂CO₃) into hydrogen ions (H⁺) and bicarbonate ions (HCO^3-) in water:

H₂CO₃ ⇌ H⁺ + HCO^3-

It is a reversible reaction. Initially, a small amount of H₂CO₃ splits and dissociates into ions. But as [H⁺] and [HCO^3-] increase, the reverse reaction is favored, until equilibrium is established. If carbon dioxide is bubbled into the water, more H₂CO₃ forms, disturbing the equilibrium. The reaction then shifts forward again to produce more ions until reaching the new equilibrium state.

Understanding chemical equilibria allows chemists to predict the extent and yields of chemical reactions based on experimental conditions like concentrations and temperature. It is a foundational concept in chemical kinetics and thermodynamics.

Types of Chemical Equilibrium

There are two types of chemical equilibrium.

• Homogeneous Equilibrium
• Heterogeneous Equilibrium

Homogeneous Equilibrium

In a homogeneous equilibrium, the reactants and products exist in the same phase, usually gaseous or aqueous.

For Example

• Example: N₂(g) + O₂(g) ⇌ 2NO(g)

This allows molecules to interact uniformly, with reaction occurring throughout the bulk phase rather than at an interface. Kinetics are faster as molecular collisions are unhindered by phase boundaries. Equilibrium establishes rapidly. Pressure and temperature changes can shift the equilibrium position per Le Chatelier’s principle.

Heterogeneous Equilibrium

A heterogeneous equilibrium has reactants and products in different phases, typically gas-solid, liquid-solid or gas-liquid.

For example

• Example: Fe(s) + H₂O(g) ⇌ FeO(s) + H₂(g)

Reaction occurs at the interface between phases, with only interfacial molecules able to collide and react. Kinetics are slower, as the reaction rate depends on the contact area between phases. Equilibrium takes longer to establish, with temperature mainly affecting vapor pressure or solubility. Pressure has little effect. Recognizing the equilibrium type is key, as it impacts the rate, thermodynamics, and appropriate models used to describe the chemical system.

Importance of chemical equilibrium

Chemical equilibrium is important for a number of reasons, including:

• It allows chemical reactions to take place at a constant rate. Without chemical equilibrium, the rate of reaction would be constantly changing, making it difficult to produce products consistently.
• It is essential for many industrial processes. For example, the Haber process for producing ammonia and the contact process for producing sulfuric acid both rely on chemical equilibrium.
• It plays a role in many biological processes. For example, the equilibrium between carbon dioxide and oxygen in the blood is essential for respiration.

Here are some specific examples of the importance of chemical equilibrium:

• Production of fertilizers: Ammonia is a key ingredient in fertilizers, and it is produced using the Haber process, which is an equilibrium reaction. The Haber process is highly efficient and produces large quantities of ammonia at a relatively low cost.
• Production of sulfuric acid: Sulfuric acid is one of the most important industrial chemicals, and it is produced using the contact process, which is also an equilibrium reaction. Sulfuric acid is used in a wide variety of applications, including the production of fertilizers, batteries, and plastics.
• Respiration: When we breathe in, oxygen from the air diffuses into our blood and binds to hemoglobin. This reaction is reversible, and the equilibrium between oxygen and hemoglobin ensures that our blood is always carrying enough oxygen to our tissues.
• Digestion: Many digestive enzymes work by catalyzing reversible equilibrium reactions. For example, the enzyme lactase catalyzes the breakdown of lactose (milk sugar) into glucose and galactose. The equilibrium between lactose and glucose + galactose ensures that our bodies can absorb all of the nutrients from milk.