Equilibrium: Comprehensive NEET Chemistry Notes

1. Introduction to Equilibrium

Equilibrium is a fundamental concept in chemistry, referring to the state where the concentrations of reactants and products remain constant over time. This balance occurs when the forward and reverse reactions happen at the same rate.

1.1 Dynamic Nature of Equilibrium

Equilibrium in chemical reactions is dynamic, meaning that reactions continue to occur, but there is no net change in the concentrations of reactants and products.

Did You Know?

Equilibrium involving oxygen molecules and hemoglobin is crucial for oxygen transport in the body.

Real-life Application:

Understanding vapor pressure equilibrium helps in designing efficient cooling systems.

2. Types of Equilibria

Equilibria can be established in both physical and chemical processes. It varies with the nature of the processes and the experimental conditions.

2.1 Physical Equilibria

Examples include solid-liquid equilibrium, liquid-vapour equilibrium, and solid-vapour equilibrium.

Example:

Ice and water coexist at 0°C under atmospheric pressure, showing solid-liquid equilibrium.

2.2 Chemical Equilibria

Chemical equilibria involve reactions where reactants form products, and products reform reactants.

Example:

The equilibrium between nitrogen dioxide (NO2) and dinitrogen tetroxide (N2O4).

3. The Law of Chemical Equilibrium

The law of chemical equilibrium states that at a given temperature, the ratio of the product of the concentrations of the products to the product of the concentrations of the reactants, each raised to the power of their respective coefficients in the balanced equation, is constant.

3.1 Equilibrium Constant (Kc)

For a reaction: $aA+bB\leftrightarrow cC+dD$ The equilibrium constant is given by: $K_c=\frac{[C{]}^c[D{]}^d}{[A{]}^a[B{]}^b}$

Mnemonic:

"Products over reactants, each raised to their coefficients."

4. Factors Affecting Equilibrium

Several factors can affect the state of equilibrium in a chemical reaction.

4.1 Concentration Changes

Changing the concentration of reactants or products shifts the equilibrium to counteract the change.

Example:

Adding more hydrogen to the system$H_2(g)+I_2(g)\leftrightarrow 2HI(g)$ shifts the equilibrium to the right, producing more HI.

4.2 Temperature Changes

The effect of temperature on equilibrium depends on whether the reaction is exothermic or endothermic.

NEET Tip:

For exothermic reactions, increasing temperature shifts equilibrium to the left, while for endothermic reactions, it shifts to the right.

4.3 Pressure Changes

Changes in pressure affect gaseous equilibria. Increasing pressure shifts equilibrium toward the side with fewer gas molecules.

Example:

For the reaction $N_2(g)+3H_2(g)\leftrightarrow 2N{H}_3(g)$, increasing pressure favors the formation of ammonia.

5. Homogeneous and Heterogeneous Equilibria

Equilibria can be homogeneous or heterogeneous, depending on the phases of the reactants and products.

5.1 Homogeneous Equilibria

All reactants and products are in the same phase.

Example:

The equilibrium between gaseous nitrogen dioxide and dinitrogen tetroxide.

5.2 Heterogeneous Equilibria

Reactants and products are in different phases.

Example:

The equilibrium between calcium carbonate, calcium oxide, and carbon dioxide.

6. Le Chatelier’s Principle

Le Chatelier's principle predicts the effect of a change in conditions on the position of equilibrium.

6.1 Application of Le Chatelier’s Principle

• Concentration: Adding or removing substances shifts the equilibrium to counteract the change.
• Temperature: Changing temperature shifts equilibrium in the direction that absorbs or releases heat.
• Pressure: Increasing pressure shifts equilibrium toward the side with fewer gas molecules.

Common Misconception:

Le Chatelier’s principle only applies to changes in concentration and pressure, not to catalysts.

Quick Recap

• Equilibrium is a dynamic state with no net change in reactant and product concentrations.
• Physical equilibria include solid-liquid, liquid-vapour, and solid-vapour equilibria.
• Chemical equilibrium is governed by the equilibrium constant, which remains constant at a given temperature.
• Factors affecting equilibrium include concentration, temperature, and pressure changes.
• Le Chatelier’s principle helps predict the direction of the shift in equilibrium in response to changes.

Practice Questions

1. Question: Define dynamic equilibrium in chemical reactions. Solution: Dynamic equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant concentrations of reactants and products.
2. Question: What is the effect of increasing the concentration of a reactant in an equilibrium reaction? Solution: Increasing the concentration of a reactant shifts the equilibrium to the right, favoring the formation of products.
3. Question: How does temperature affect the equilibrium of an exothermic reaction? Solution: Increasing temperature shifts the equilibrium to the left, favoring the reactants in an exothermic reaction.
4. Question: Write the expression for the equilibrium constant for the reaction: $2S{O}_2(g)+O_2(g)\leftrightarrow 2S{O}_3(g)$. Solution: $K_c=\frac{[S{O}_3{]}^2}{[S{O}_2{]}^2[O_2]}$
5. Question: Explain the effect of pressure change on the equilibrium of the reaction: $N_2(g)+3H_2(g)\leftrightarrow 2N{H}_3(g)$. Solution: Increasing pressure shifts the equilibrium to the right, favoring the formation of ammonia.
6. Question: What is the significance of Le Chatelier’s principle in industrial processes? Solution: Le Chatelier’s principle helps optimize conditions to maximize product yield, as seen in the Haber process for ammonia synthesis.
7. Question: Differentiate between homogeneous and heterogeneous equilibria. Solution: Homogeneous equilibria involve reactants and products in the same phase, while heterogeneous equilibria involve different phases.
8. Question: Calculate the equilibrium constant for the reaction: $H_2(g)+I_2(g)\leftrightarrow 2HI(g)$ if at equilibrium, the concentrations are [H2] = 0.2 M, [I2] = 0.1 M, and [HI] = 0.4 M. Solution: $K_c=\frac{[HI{]}^2}{[H_2][I_2]}=\frac{(0.4{)}^2}{(0.2)(0.1)}=8.0$
9. Question: Describe the effect of removing a product from an equilibrium reaction. Solution: Removing a product shifts the equilibrium to the right, favoring the formation of more products.
10. Question: How does the presence of a catalyst affect the equilibrium position? Solution: A catalyst speeds up the rate of both the forward and reverse reactions equally, without affecting the position of equilibrium.

Quick Reference Guide and Glossary

• Equilibrium: A state where the concentrations of reactants and products remain constant over time.
• Dynamic Equilibrium: A state of balance where the rates of the forward and reverse reactions are equal.
• Le Chatelier’s Principle: States that a system at equilibrium will adjust to counteract any imposed change in concentration, temperature, or pressure.
• Equilibrium Constant (Kc): The ratio of the concentration of products to reactants at equilibrium, each raised to the power of their coefficients.
• Homogeneous Equilibria: Equilibria involving reactants and products in the same phase.
• Heterogeneous Equilibria: Equilibria involving reactants and products in different phases.