Heterogeneous equilibrium comes into play for reactants and products in different phases.
What is Heterogeneous Equilibrium?
When reactants and products are present in multiple phases, a chemical system experiences a heterogeneous equilibrium. Solids, liquids, and gasses could coexist in these phases until the system achieves equilibrium. For example, a reaction might occur when a solid and a gas combine to form a liquid.
Here’s a simple example to understand it better:
CaCO3(s) ↔ CaO(s) + CO2(g)
In this reaction:
Calcium carbonate (CaCO3) is a solid.
Calcium oxide (CaO) is also a solid.
Carbon dioxide (CO2) is a gas.
Rate at which calcium carbonate breaks down equals the rate at which calcium oxide and carbon dioxide are formed, the system is said to be in balance.
How is Heterogeneous Equilibrium Different?
All of the reactants and products, such as all liquids or all gases, are in the same phase when there is a homogeneous equilibrium. However, there are various phases are in a heterogeneous equilibrium. Because it affects how we calculate the equilibrium constant (K), this difference is essential.
The concentrations of pure solids and pure liquids are regarded as constant for heterogeneous equilibria and are not included in the equilibrium equation. Only the gas and aqueous solution concentrations.
The Equilibrium Constant for Heterogeneous Equilibrium
To calculate the equilibrium constant, we only consider the gaseous and aqueous components. For the reaction:
CaCO3(s)↔CaO(s)+CO2(g)
The equilibrium constant (Kp) is expressed as: Kp = PCO2
Here, PCO2 represents the partial pressure of carbon dioxide. The concentrations of solids (CaCO3 and CaO) are omitted because their activity is constant.
Characteristics of Heterogeneous Equilibrium
1.Multiple Phases: The coexistence of at least two distinct phases is the unique characteristic.
2. Constant Concentrations of Solids and Liquids: The equilibrium constant is unaffected by pure solids or pure liquids.
3. Dependency on Gaseous or Aqueous Species: Only species present in the gas phase or solution have an impact on the equilibrium constant.
Examples of Heterogeneous Equilibria
1.Dissociation of Ammonium Chloride:
NH4Cl(s) ↔ NH3(g) + HCl(g)
Here, the solid ammonium chloride dissociates into gaseous ammonia and hydrogen chloride.
2. Formation of Carbon Dioxide: C(s)+O2(g) ↔ CO2(g)
In this reaction, solid carbon reacts with gaseous oxygen to produce carbon dioxide.
Factors Affecting Heterogeneous Equilibrium
1. Temperature: Le Chatelier’s Principle states that variations in temperature have the potential to move the equilibrium position, just like it can with any other equilibria.
2. Pressure: Variations in pressure may have an impact on the equilibrium if gases are there.
3. Surface Area: Increasing the surface area of solid-phase reactions can speed up the process without changing the equilibrium state.
Applications of Heterogeneous Equilibria
1. Industrial Processes: Heterogeneous equilibrium principles are essential to many industrial operations, including the creation of cement and ammonia.
2. Environmental Chemistry: Heterogeneous equilibria are concerned in processes such as the dissolution of minerals in water and the interaction of air gases with the earth’s surface.
3. Biological Systems: Heterogeneous equilibria are found in a variety of biological reactions, ie.g; gas exchange in the lungs.
Units o equilibrium constant: Concentration of a substance is measured in terms of moles / litre, therefore, unit of Kc is (mol L-1) Δng, similarly, partial pressure is measured in terms of atmosphere, hence, unit of Kp is (atm)ng .
If Δng = 0 both Kc and Kp have no units.
If Δng > 0 , unit of Kc = (mol L-1) Δng, unit of
Kp = (atm) Δng
If Δng < 0, unit of Kc (L mol-1) Δng, unit of Kp = (atm-1) Δng
Equilibrium constants can also be expressed as dimensionless quantities if the standard states of reactants and products are specified.
Value of equilibrium constant explains the thermodynamic stability of products. Greater the value of K, greater will be the stability of the products and instability of reactants.
Note :-
How chemical reactions operate in practical situations, we must have a idea of heterogeneous equilibrium. It demonstrates how equilibrium can be reached even in complicated systems and interactions between various phases. On this idea, we will enable to evaluate and forecast the behavior of multi-phase reactions.
Heterogeneous equilibrium occurs in a chemical system where the reactants and products are present in different physical phases (e.g., solids, liquids, and gases). For example, in the reaction:
CaCO3(s) ↔ CaO(s) + CO2(g) the equilibrium are solids and a gas.
All reactants and products (such as all gases or liquids) are in the same phase when there is homogeneous equilibrium. When substances are in heterogeneous equilibrium, they can be solids, liquids, or gasses.
Since the densities of pure solids and liquids do not change much during the reaction, their concentrations remain constant. As a result, the equilibrium constant statement only takes into account gases and aqueous solutions, leaving them out.
For a heterogeneous equilibrium, the equilibrium constant includes only the concentrations or partial pressures of gases and aqueous species. For example: CaCO3(s) ↔ CaO(s)+CO2(g)
The equilibrium constant (Kp) is written as: Kp = PCO2
Yes, variations in pressure can cause the equilibrium position to change when gases are present. The side of the reaction with fewer moles of gas is favored by increasing pressure, whereas the side with more moles of gas is favored by reducing pressure.
The equilibrium position is unaffected by the surface area of solids, but the rate of reaction is affected. Although it does not alter the equilibrium constant, increasing surface area enables more contacts at the solid’s surface, accelerating the reaction.
Examples include:
The decomposition of limestone in the cement industry:
CaCO3(s) ↔ CaO(s) + CO2(g)
The interaction of carbon with oxygen in combustion:
C(s)+O2(g)↔CO2(g)
Dissolution of salts in water where a solid salt forms an equilibrium with its ions in solution.
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