Class 11 Chemistry Notes on Chapter 7 – Equilibrium

Class 11 Chemistry Notes on [Chapter 7 – Equilibrium]

Equilibrium is one of the most captivating chapters in class 11 chemistry. It bridges our understanding of how reversible reactions occur and settle into balance, whether in physical systems or chemical processes. Let’s break down the essentials of class 11 chemistry notes on [Chapter 7 – Equilibrium], exploring key topics step by step to make it easier for you to ace your exams.

Introduction

Imagine a seesaw perfectly balanced with two equal weights. That’s what equilibrium in chemistry looks like—a state where opposing forces or reactions balance out. In class 11 chemistry notes on [Chapter 7 – Equilibrium], we delve into how reversible reactions reach this balance and the factors that can tip the scales. This knowledge isn’t just theoretical—it’s the foundation for understanding everything from industrial synthesis to biological processes.

Types of Equilibrium

Dynamic and Static Equilibrium

Dynamic Equilibrium occurs when the forward and reverse reactions happen at the same rate, maintaining constant concentrations of reactants and products.
Example: Water evaporating and condensing in a sealed container.
Static Equilibrium, in contrast, implies no movement or reaction—essentially a state of rest.

Physical and Chemical Equilibrium

Physical Equilibrium involves changes in physical states, like the equilibrium between ice and water.
Chemical Equilibrium occurs in reversible chemical reactions, such as the Haber process for ammonia synthesis.

Key Difference

Dynamic equilibrium involves activity, whereas static equilibrium represents inactivity.

Equilibrium Involving Dissolution of Solids

When a solid dissolves in a solvent, it establishes a dynamic equilibrium between the dissolved ions and the undissolved solid:

\text{Solid} \leftrightarrow \text{Dissolved ions}

Saturated Solutions: These contain maximum solute that can dissolve.
Supersaturation: When more solute dissolves than the equilibrium allows, leading to precipitation.

For example, common salt ( $NaCl$ ) dissolves in water until it reaches a saturation point, balancing between dissolution and precipitation.

General Characteristics of Physical Equilibrium

Dynamic Nature: Even at equilibrium, molecules keep changing their phase or state.
Condition Specificity: It depends on external conditions like temperature and pressure.
No Net Change: While individual molecules shift, their overall concentration remains constant.

Example

At boiling point, water’s rate of evaporation equals the rate of condensation, establishing equilibrium.

Equilibrium Involving Chemical Processes

Chemical equilibrium occurs in reversible reactions where forward and backward reaction rates are equal:

aA + bB \leftrightarrow cC + dD

Key Features

Concentration Stability: Reactant and product concentrations remain constant at equilibrium.
Dependence on Conditions: Temperature, pressure, and catalysts influence equilibrium.

Example: In the synthesis of ammonia ( $N_2 + 3H_2 \leftrightarrow 2NH_3$ ), increasing pressure shifts equilibrium toward ammonia production.

Law of Mass Action

Statement

The rate of a reaction is proportional to the product of the concentrations of reactants, each raised to a power equal to its stoichiometric coefficient.

For the reaction:

aA + bB \leftrightarrow cC + dD

The equilibrium constant ( $K$ ) is:

K = \frac{[C]^c [D]^d}{[A]^a [B]^b}

Implications

This law predicts how changes in reactant or product concentrations impact equilibrium.

Relationship Between $K_p$ and $K_c$

The relationship between the equilibrium constants $K_p$ (for gases) and $K_c$ (for concentrations) is:

K_p = K_c (RT)^{\Delta n}

Where:

$R$ = Universal gas constant
$T$ = Temperature in Kelvin
$\Delta n$ = Change in moles of gas ( $\text{moles of products - moles of reactants}$ )

This equation highlights how pressure and concentration interconnect in gas-phase reactions.

Characteristics of Equilibrium Constant ( $K$ )

Temperature Dependence: $K$ changes with temperature, indicating the reaction's favorability under varying conditions.
Direction Indicator: A high $K$ favors products, while a low $K$ favors reactants.
Dimensionless Value: Though derived from concentrations or pressures, $K$ itself is a ratio.

Characteristics of Chemical Equilibrium

Reversible Nature

Chemical equilibrium is established in reversible reactions where the forward and backward processes occur simultaneously. For instance, the reaction between hydrogen and iodine gas to form hydrogen iodide ( $H_2 + I_2 \leftrightarrow 2HI$ ) demonstrates equilibrium.

Constant Concentrations

At equilibrium, the concentrations of reactants and products remain constant over time, though molecules are dynamically interacting.

Condition Sensitivity

Changes in pressure, temperature, or concentration can shift equilibrium, as explained by Le Chatelier's principle.

Effect of Temperature on $K$

Endothermic Reactions

For reactions absorbing heat ( $\Delta H > 0$ ), increasing the temperature increases $K$ , favoring the forward reaction.

Exothermic Reactions

For reactions releasing heat ( $\Delta H < 0$ ), increasing the temperature decreases $K$ , shifting equilibrium toward the reactants.

Practical Implication

Industrial processes, like the Haber process, are carefully controlled by adjusting temperature to optimize yield while considering $K$ .

Equilibrium Constant and Chemical Equation

Stoichiometric Influence

The value of the equilibrium constant depends on how the reaction equation is written. Doubling the coefficients squares $K$ , while halving them takes the square root:

\text{Original Reaction: } K = \frac{[C]^c [D]^d}{[A]^a [B]^b}

\text{New Reaction: } K_{\text{new}} = (K)^{n} \text{ (where \(n\) is the factor of change).}

Direction Reversal

Reversing the reaction inverts $K$ :

\text{Forward: } K = \frac{[Products]}{[Reactants]}

\text{Reverse: } K' = \frac{[Reactants]}{[Products]} = \frac{1}{K}

Application

Understanding this relationship is crucial for predicting outcomes in multi-step reactions.

Application of Equilibrium Constant

Predicting Reaction Direction

The reaction quotient ( $Q$ ) compared with $K$ reveals the system's state:

$Q < K$ : Reaction proceeds forward.
$Q > K$ : Reaction proceeds backward.
$Q = K$ : System is at equilibrium.

Calculating Extent of Reaction

Using $K$ , chemists determine how far a reaction will go before reaching equilibrium.

Real-Life Applications

Industrial synthesis of chemicals like sulfuric acid.
Predicting solubility and precipitation in solutions.

Relationship Between $K$ , Reaction Quotient, and Gibbs Free Energy

Connection with Gibbs Free Energy

The spontaneity of a reaction is tied to the equilibrium constant through Gibbs free energy ( $\Delta G$ ):

\Delta G = \Delta G^0 + RT \ln Q

At equilibrium ( $Q = K$ ):

\Delta G = 0 \implies \Delta G^0 = -RT \ln K

Practical Insights

$K > 1$ : $\Delta G^0 < 0$ , reaction is product-favored.
$K < 1$ : $\Delta G^0 > 0$ , reaction is reactant-favored.

Factors Affecting State of Equilibrium: Le Chatelier's Principle

Statement

If a system at equilibrium is disturbed, it will adjust to counteract the disturbance and restore equilibrium.

Applications

Change in Concentration: Adding reactants shifts equilibrium to products, and vice versa.
Change in Pressure: Increasing pressure favors the side with fewer moles of gas.
Change in Temperature: Depends on whether the reaction is exothermic or endothermic.

Industrial Example

In the Haber process, increasing pressure shifts equilibrium to favor ammonia ( $N_2 + 3H_2 \leftrightarrow 2NH_3$ ).

Ionic Equilibrium

Definition

Ionic equilibrium involves the balance of ions in weak electrolytes that partially dissociate in solution.

Key Examples

Dissociation of acetic acid: $CH_3COOH \leftrightarrow H^+ + CH_3COO^-$
Dissociation of ammonia in water: $NH_3 + H_2O \leftrightarrow NH_4^+ + OH^-$

Significance

Understanding ionic equilibrium is critical for acid-base chemistry and buffer systems.

Acids, Bases, and Salts

Arrhenius Concept

Acids release $H^+$ ions, and bases release $OH^-$ ions in water. Example:

$HCl \rightarrow H^+ + Cl^-$
$NaOH \rightarrow Na^+ + OH^-$

Bronsted-Lowry Concept

Acids are proton donors, and bases are proton acceptors. Example:

$NH_3 + H^+ \leftrightarrow NH_4^+$

Lewis Concept

Acids accept electron pairs, while bases donate them. Example:

$BF_3$ (Lewis acid) reacts with $NH_3$ (Lewis base).

Class 11 Chemistry Notes - Chapter 6: Thermodynamics

Introduction

Types of Equilibrium

Dynamic and Static Equilibrium

Dynamic Equilibrium occurs when the forward and reverse reactions happen at the same rate, maintaining constant concentrations of reactants and products.
Example: Water evaporating and condensing in a sealed container.
Static Equilibrium, in contrast, implies no movement or reaction—essentially a state of rest.

Physical and Chemical Equilibrium

Physical Equilibrium involves changes in physical states, like the equilibrium between ice and water.
Chemical Equilibrium occurs in reversible chemical reactions, such as the Haber process for ammonia synthesis.

Key Difference

Dynamic equilibrium involves activity, whereas static equilibrium represents inactivity.

Equilibrium Involving Dissolution of Solids

When a solid dissolves in a solvent, it establishes a dynamic equilibrium between the dissolved ions and the undissolved solid:

\text{Solid} \leftrightarrow \text{Dissolved ions}

Saturated Solutions: These contain maximum solute that can dissolve.
Supersaturation: When more solute dissolves than the equilibrium allows, leading to precipitation.

For example, common salt ( $NaCl$ ) dissolves in water until it reaches a saturation point, balancing between dissolution and precipitation.

General Characteristics of Physical Equilibrium

Dynamic Nature: Even at equilibrium, molecules keep changing their phase or state.
Condition Specificity: It depends on external conditions like temperature and pressure.
No Net Change: While individual molecules shift, their overall concentration remains constant.

Example

At boiling point, water’s rate of evaporation equals the rate of condensation, establishing equilibrium.

Equilibrium Involving Chemical Processes

Chemical equilibrium occurs in reversible reactions where forward and backward reaction rates are equal:

aA + bB \leftrightarrow cC + dD

Key Features

Concentration Stability: Reactant and product concentrations remain constant at equilibrium.
Dependence on Conditions: Temperature, pressure, and catalysts influence equilibrium.

Example: In the synthesis of ammonia ( $N_2 + 3H_2 \leftrightarrow 2NH_3$ ), increasing pressure shifts equilibrium toward ammonia production.

Law of Mass Action

Statement

The rate of a reaction is proportional to the product of the concentrations of reactants, each raised to a power equal to its stoichiometric coefficient.

For the reaction:

aA + bB \leftrightarrow cC + dD

The equilibrium constant ( $K$ ) is:

K = \frac{[C]^c [D]^d}{[A]^a [B]^b}

Implications

This law predicts how changes in reactant or product concentrations impact equilibrium.

Relationship Between Kp and Kc

The relationship between the equilibrium constants $K_p$ (for gases) and $K_c$ (for concentrations) is:

K_p = K_c (RT)^{\Delta n}

Where:

$R$ = Universal gas constant
$T$ = Temperature in Kelvin
$\Delta n$ = Change in moles of gas ( $\text{moles of products - moles of reactants}$ )

This equation highlights how pressure and concentration interconnect in gas-phase reactions.

Characteristics of Equilibrium Constant ( $K$ )

Temperature Dependence: $K$ changes with temperature, indicating the reaction's favorability under varying conditions.
Direction Indicator: A high $K$ favors products, while a low $K$ favors reactants.
Dimensionless Value: Though derived from concentrations or pressures, $K$ itself is a ratio.

Characteristics of Chemical Equilibrium

Reversible Nature

Constant Concentrations

At equilibrium, the concentrations of reactants and products remain constant over time, though molecules are dynamically interacting.

Condition Sensitivity

Changes in pressure, temperature, or concentration can shift equilibrium, as explained by Le Chatelier's principle.

Effect of Temperature on $K$

Endothermic Reactions

For reactions absorbing heat ( $\Delta H > 0$ ), increasing the temperature increases $K$ , favoring the forward reaction.

Exothermic Reactions

For reactions releasing heat ( $\Delta H < 0$ ), increasing the temperature decreases $K$ , shifting equilibrium toward the reactants.

Practical Implication

Industrial processes, like the Haber process, are carefully controlled by adjusting temperature to optimize yield while considering $K$ .

Equilibrium Constant and Chemical Equation

Stoichiometric Influence

The value of the equilibrium constant depends on how the reaction equation is written. Doubling the coefficients squares $K$ , while halving them takes the square root:

\text{Original Reaction: } K = \frac{[C]^c [D]^d}{[A]^a [B]^b}

\text{New Reaction: } K_{\text{new}} = (K)^{n} \text{ (where \(n\) is the factor of change).}

Direction Reversal

Reversing the reaction inverts $\text{Forward: } K = \frac{[Products]}{[Reactants]}$

Application

Understanding this relationship is crucial for predicting outcomes in multi-step reactions.

Application of Equilibrium Constant

Predicting Reaction Direction

The reaction quotient ( $Q$ ) compared with $K$ reveals the system's state:

$Q < K$ : Reaction proceeds forward.
$Q > K$ : Reaction proceeds backward.
$Q = K$ : System is at equilibrium.

Calculating Extent of Reaction

Using $K$ , chemists determine how far a reaction will go before reaching equilibrium.

Real-Life Applications

Industrial synthesis of chemicals like sulfuric acid.
Predicting solubility and precipitation in solutions.

Relationship Between $K$ , Reaction Quotient, and Gibbs Free Energy

Connection with Gibbs Free Energy

The spontaneity of a reaction is tied to the equilibrium constant through Gibbs free energy ( $\Delta G$ ):

\Delta G = \Delta G^0 + RT \ln Q

At equilibrium ( $Q = K$ ):

\Delta G = 0 \implies \Delta G^0 = -RT \ln K

Practical Insights

$K > 1$ : $\Delta G^0 < 0$ , reaction is product-favored.
$K < 1$ : $\Delta G^0 > 0$ , reaction is reactant-favored.

Factors Affecting State of Equilibrium: Le Chatelier's Principle

Statement

If a system at equilibrium is disturbed, it will adjust to counteract the disturbance and restore equilibrium.

Applications

Change in Concentration: Adding reactants shifts equilibrium to products, and vice versa.
Change in Pressure: Increasing pressure favors the side with fewer moles of gas.
Change in Temperature: Depends on whether the reaction is exothermic or endothermic.

Industrial Example

In the Haber process, increasing pressure shifts equilibrium to favor ammonia ( $N_2 + 3H_2 \leftrightarrow 2NH_3$ ).

Ionic Equilibrium

Definition

Ionic equilibrium involves the balance of ions in weak electrolytes that partially dissociate in solution.

Key Examples

Dissociation of acetic acid: $CH_3COOH \leftrightarrow H^+ + CH_3COO^-$
Dissociation of ammonia in water: $NH_3 + H_2O \leftrightarrow NH_4^+ + OH^-$

Significance

Understanding ionic equilibrium is critical for acid-base chemistry and buffer systems.

Acids, Bases, and Salts

Arrhenius Concept

Acids release $H^+$ ions, and bases release $OH^-$ ions in water. Example:

$HCl \rightarrow H^+ + Cl^-$
$NaOH \rightarrow Na^+ + OH^-$

Bronsted-Lowry Concept

Acids are proton donors, and bases are proton acceptors. Example:

$NH_3 + H^+ \leftrightarrow NH_4^+$

Lewis Concept

Acids accept electron pairs, while bases donate them. Example:

$BF_3$ (Lewis acid) reacts with $NH_3$ (Lewis base).

Ionic Product of Water

Definition

Water self-ionizes to produce hydrogen ( $H^+$ ) and hydroxide ( $OH^-$ ) ions:

H_2O \leftrightarrow H^+ + OH^-

The equilibrium constant for this ionization is known as the ionic product of water ( $K_w$ ):

K_w = [H^+][OH^-]

At $25^\circ C$ , $K_w = 1 \times 10^{-14}$ .

Significance

Helps in calculating $pH$ and $pOH$ .
Forms the basis for understanding acid-base equilibria.

Example Calculation

If $K_w = 1 \times 10^{-14}$ , and in neutral water, $[H^+] = [OH^-]$ , then:

[H^+] = [OH^-] = \sqrt{K_w} = \sqrt{1 \times 10^{-14}} = 1 \times 10^{-7} \, \text{mol/L}.

Expressing Hydrogen Ion Concentration: The pH Scale

What is pH?

The pH scale measures the acidity or basicity of a solution:

pH = -\log [H^+]

Neutral solution: $pH = 7$ .
Acidic solution: $pH < 7$ .
Basic solution: $pH > 7$ .

pOH Relationship

The pH and pOH of a solution are related:

pH + pOH = 14

Practical Application

Monitoring pH in chemical reactions.
Ensuring correct pH levels in industries like food and pharmaceuticals.

Acid-Base Equilibrium: Ionization of Acids and Bases

Ionization of Acids

Weak acids partially ionize in water. For example:

CH_3COOH \leftrightarrow H^+ + CH_3COO^-

The equilibrium constant ( $K_a$ ) is given by:

K_a = \frac{[H^+][CH_3COO^-]}{[CH_3COOH]}

Ionization of Bases

Weak bases like ammonia ionize as follows:

NH_3 + H_2O \leftrightarrow NH_4^+ + OH^-

The equilibrium constant ( $K_b$ ) is:

K_b = \frac{[NH_4^+][OH^-]}{[NH_3]}

Significance

The strength of acids and bases is determined by their $K_a$ and $K_b$ values.

Relation Between $K_a$ , $K_w$ , and $K_b$

The dissociation constants of acids and bases are related through the ionic product of water ( $K_w$ ):

K_w = K_a \times K_b

Where:

$K_a$ : Acid dissociation constant.
$K_b$ : Base dissociation constant.

Application

For weak acids and bases, knowing $K_a$ or $K_b$ helps calculate the other, aiding in equilibrium calculations.

Protic and Non-Protic Acids

Protic Acids

Protic acids can donate hydrogen ions ( $H^+$ ):

Monoprotic: Donates one proton (e.g., $HCl$ ).
Diprotic: Donates two protons (e.g., $H_2SO_4$ ).
Triprotic: Donates three protons (e.g., $H_3PO_4$ ).

Non-Protic Acids

These do not release protons but may accept electron pairs, like $BF_3$ (Lewis acid).

Importance

The classification helps in understanding the strength and reactivity of acids in different reactions.

Common Ion Effect

Definition

The common ion effect occurs when a common ion is added to a solution at equilibrium, reducing the solubility of the solute.

Example

Adding $NaCl$ to a saturated $AgCl$ solution shifts equilibrium:

AgCl \leftrightarrow Ag^+ + Cl^-

Extra $Cl^-$ suppresses $AgCl$ dissociation.

Applications

Used in qualitative analysis.
Controls pH in buffer solutions.

Hydrolysis of Salt

What is Salt Hydrolysis?

Salts react with water to form acidic or basic solutions. For example:

$NH_4Cl + H_2O \rightarrow NH_4OH + HCl$ (acidic solution).
$CH_3COONa + H_2O \rightarrow CH_3COOH + NaOH$ (basic solution).

Types of Salts

Neutral salts: Derived from strong acids and bases.
Acidic salts: Derived from strong acids and weak bases.
Basic salts: Derived from weak acids and strong bases.

Buffer Solution

Definition

Buffers resist changes in pH upon adding small amounts of acid or base. They typically consist of:

A weak acid and its salt.
A weak base and its salt.

Example

A buffer of acetic acid and sodium acetate:

CH_3COOH \leftrightarrow H^+ + CH_3COO^-

Importance

Buffers maintain pH stability in biological and industrial processes.

Solubility Product

What is Solubility Product?

The solubility product constant ( $K_{sp}$ ) applies to sparingly soluble salts:

AB \leftrightarrow A^+ + B^-

K_{sp} = [A^+][B^-]

Application

Predicting precipitation and calculating solubility in solutions.

Conclusion

Equilibrium in class 11 chemistry notes on [Chapter 7 – Equilibrium] is a fundamental concept connecting physical and chemical processes. Whether it’s understanding acid-base interactions or the stability of ionic systems, equilibrium principles empower chemists to control reactions and predict outcomes. By mastering these topics, you’ll unlock insights into countless scientific phenomena.

FAQs

1. What is the difference between $K_c$ and $K_p$ ?

$K_c$ is the equilibrium constant based on concentration, while $K_p$ is based on partial pressure. They are related by $K_p = K_c (RT)^{\Delta n}$ .

2. How does temperature affect equilibrium?

For endothermic reactions, increasing temperature shifts equilibrium toward products. For exothermic reactions, it shifts toward reactants.

3. What is Le Chatelier’s principle?

It states that if a system at equilibrium is disturbed, it adjusts to counteract the disturbance and reestablish equilibrium.

4. What is a buffer solution?

A buffer is a solution that resists pH changes when small amounts of acid or base are added.

5. How does the common ion effect influence solubility?

The common ion effect reduces solubility by shifting equilibrium toward the undissolved form.

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