One app for all your study abroad needs
NCERT Notes Class 11 Chemistry (Part-1) Chapter 6: Equilibrium introduces the concept of equilibrium in both physical and chemical processes in a more comprehensive way. You will learn how reversible reactions reach a dynamic balance and how to express this using equilibrium constants. The chapter also covers factors affecting equilibrium, such as temperature and pressure, along with Le Chatelier’s Principle, all of which are essential for understanding reactions in chemistry and real-life applications.
Contents
- 1 Importance of Chemical Equilibrium
- 2 Equilibrium in Physical Processes
- 3 Classification of Reactions Based on the Extent of Equilibrium
- 4 Factors Affecting the Extent of Equilibrium
- 5 Equilibrium in Physical Processes
- 6 General Characteristics of Equilibria Involving Physical Processes
- 7 Reversible and Irreversible Reactions
- 8 Law of Chemical Equilibrium and Equilibrium Constant
- 9 Applications of Equilibrium Constants
- 10 Gibbs Free Energy and Chemical Reactions
- 11 Factors Affecting Equilibria
- 12 Le Chatelier’s Principle
- 13 Effect of Concentration Change
- 14 Effect of Pressure Change
- 15 Effect of Inert Gas Addition
- 16 Effect of Temperature Change
- 17 Ionic Equilibrium in Solutions
- 18 Acids, Bases, and Salts
- 19 Ionisation of Acids and Bases
- 20 Factors Affecting Acid Strength
- 21 Common Ion Effect in the Ionisation of Acids and Bases
- 22 Hydrolysis of Salts and the pH of their Solutions
- 23 Buffer Solutions
- 24 Solubility Equilibria of Sparingly Soluble Salts
- 25 Important Formulas in NCERT Notes Class 11 Chemistry (Part-I) Chapter 6: Equilibrium
- 26 FAQs
Explore Notes of Class 11 Chemistry
Importance of Chemical Equilibrium
Chemical equilibria are significant in various biological and environmental processes. A key example is the equilibrium between O₂ molecules and hemoglobin, which plays a vital role in the transport and delivery of oxygen from the lungs to the muscles. Similarly, an equilibrium involving carbon monoxide (CO) and hemoglobin is responsible for CO toxicity, as CO interferes with oxygen transport in the body.
Equilibrium in Physical Processes
When a liquid evaporates in a closed container, molecules with higher kinetic energy escape from the liquid surface into the vapour phase. At the same time, vapour molecules collide with the surface and return to the liquid phase. Eventually, a stage is reached where the rate of molecules leaving the liquid equals the rate of return from the vapour; this is known as the equilibrium state. At equilibrium: Rate of evaporation = Rate of condensation
Classification of Reactions Based on the Extent of Equilibrium
Reactions can be grouped into three categories based on how far they proceed before equilibrium is established:
- Reactions that proceed nearly to completion
- Only negligible amounts of reactants remain.
- Sometimes, the remaining reactants are undetectable experimentally.
- Reactions that form very small amounts of products
- Most reactants remain unchanged even at equilibrium.
- Reactions where concentrations of reactants and products are comparable
- Both are present in significant amounts at equilibrium.
Factors Affecting the Extent of Equilibrium
- The extent to which a reaction proceeds depends on various experimental conditions, such as:
- Concentration of reactants
- Temperature
- Other system-specific factors
- In industrial and laboratory settings, optimising these operational conditions is essential to:
- Shift the equilibrium toward the desired product
- Maximise efficiency and yield
Equilibrium in Physical Processes
Physical equilibria occur in phase transformations where two states of matter coexist in a closed system under specific conditions of temperature and pressure. These include:
- Solid ⇌ Liquid
- Liquid ⇌ Gas
- Solid ⇌ Gas
Solid–Liquid Equilibrium
- Example: Ice and water at 273 K in a perfectly insulated thermos flask (no heat exchange with surroundings) are in equilibrium.
- Observations:
- The mass of ice and water does not change with time.
- Temperature remains constant.
- However, the equilibrium is dynamic, not static:
- Liquid water molecules collide with ice and adhere to it.
- Simultaneously, ice molecules escape into the liquid phase.
- Since the rate of melting = rate of freezing, there is no net change in mass.
- This equilibrium occurs only at a specific temperature and pressure.
- The temperature at which solid and liquid coexist at atmospheric pressure is called the normal melting point or freezing point.
Inferences:
- Both opposing processes (melting and freezing) occur simultaneously.
- Both occur at the same rate, keeping the mass of each phase constant.
Liquid–Vapour Equilibrium
Experimental setup: A transparent box with a U-tube manometer and a drying agent like anhydrous CaCl₂ is used. After placing a watch glass containing water inside the box and removing the drying agent:
- The mercury level in the manometer’s right limb increases and becomes constant, indicating a rise in pressure.
- Water volume decreases as it evaporates.
- Initially, the box contains very little or no water vapour. As water evaporates:
- Pressure rises due to the formation of water vapour.
- With time, condensation begins, balancing evaporation.
- Finally, the rate of evaporation = the rate of condensation:
H₂O (l) ⇌ H₂O (vap)
- The constant pressure at equilibrium is called the equilibrium vapour pressure.
- Vapour pressure is temperature-dependent and is higher for more volatile liquids.
- Comparison:
- Different liquids like acetone, methyl alcohol, and ether have different vapor pressures at the same temperature.
- A higher vapor pressure implies: Greater volatility and a Lower boiling point.
- Observation in open systems:
- When liquids are exposed to the atmosphere, evaporation occurs, but condensation is negligible. Equilibrium is not attained due to continuous dispersion of vapour.
- Boiling point:
- For water at 1.013 bar (atmospheric pressure), the boiling point = 100°C.
- The normal boiling point is defined as the temperature at which liquid and vapour are at equilibrium at 1.013 bar.
- Boiling point varies with altitude due to atmospheric pressure differences.
Solid–Vapour Equilibrium
Example: When solid iodine is placed in a closed container:
- Violet vapours fill the vessel.
- Intensity increases over time, then becomes constant, indicating equilibrium.
- The equilibrium can be represented as:
I₂ (solid) ⇌ I₂ (vapour) - Other examples:
- Camphor (solid) ⇌ Camphor (vapour)
- NH₄Cl (solid) ⇌ NH₄Cl (vapour)
Equilibrium Involving Dissolution of Solids or Gases in Liquids
Solids in Liquids: Dissolving sugar in water.
- Only a limited amount of sugar dissolves at room temperature.
- On cooling a hot saturated solution, sugar crystallises.
- A saturated solution is one where no more solute dissolves at a given temperature.
- At this stage, a dynamic equilibrium exists between the solid sugar and sugar in solution.
- Sugar (solution) ⇌ Sugar (solid)
- Rate of dissolution = Rate of crystallisation
- Confirmed using radioactive sugar: When added to a saturated solution of non-radioactive sugar.
Gases in Liquids: Soda water bottle.
- When opened, CO₂ escapes due to lower pressure outside.
- There is an equilibrium between dissolved CO₂ and gaseous CO₂:
CO₂ (gas) ⇌ CO₂ (in solution) - This is governed by Henry’s Law:
- The mass of gas dissolved is proportional to the pressure of gas above the liquid.
- Solubility decreases with temperature.
- Soda bottles are sealed under pressure to maintain high CO₂ solubility.
- On opening:
- Pressure drops
- CO₂ escapes
- A new equilibrium is established
- Eventually, soda goes flat as gas escapes continuously into the open air.
Also Read: NCERT Notes Class 11 English Snapshots Chapter 1: The Summer of the Beautiful White Horse (Free PDF)
General Characteristics of Equilibria Involving Physical Processes
All physical equilibria discussed above share these common features:
- Equilibrium is possible only in a closed system at a fixed temperature.
- Opposing processes occur at equal rates, leading to a dynamic but stable state.
- Measurable properties remain constant during equilibrium.
- Equilibrium is characterised by constant values of measurable parameters like:
- Vapour pressure
- Boiling/melting point
- Solubility, etc.
- The magnitude of these parameters reflects the extent to which the physical process has occurred before reaching equilibrium.
Reversible and Irreversible Reactions
A chemical reaction is said to be reversible if the products formed can react again to give back the original reactants. In a closed system, reversible reactions reach chemical equilibrium, where: Rate of forward reaction = Rate of backward reaction. At this point, concentrations of reactants and products remain constant over time. Representation of a general reversible reaction:
A + B ⇌ C + D
Law of Chemical Equilibrium and Equilibrium Constant
For the general reversible reaction: aA + bB ⇌ cC + dD
The equilibrium constant in terms of concentration (Kc) is given by: Kc = [C]ᶜ [D]ᵈ / [A]ᵃ [B]ᵇ. Square brackets [ ] represent the molar concentration of the substances at equilibrium. Kc remains constant at a fixed temperature for a particular reaction.
Homogeneous Equilibria
In homogeneous equilibria, all the reactants and products are in the same physical phase, typically gaseous or aqueous.
Examples:
- N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Equilibrium constant:
Kc = [NH₃]² / [N₂][H₂]³ - H₂(g) + I₂(g) ⇌ 2HI(g)
Equilibrium constant:
Kc = [HI]² / [H₂][I₂]
Heterogeneous Equilibria
In heterogeneous equilibria, reactants and products exist in different phases. The concentrations of pure solids and pure liquids are not included in the expression for the equilibrium constant.
Example: CaCO₃(s) ⇌ CaO(s) + CO₂(g)
Equilibrium constant:
Kc = [CO₂]
Applications of Equilibrium Constants
Let Qc be the reaction quotient (same form as Kc, but using current concentrations):
- If Qc < Kc, the reaction proceeds in the forward direction.
- If Qc > Kc, the reaction proceeds in the reverse direction.
- If Qc = Kc, the system is at equilibrium.
Gibbs Free Energy and Chemical Reactions
The spontaneity of a chemical reaction is determined by the Gibbs free energy change (ΔG).
- If ΔG < 0, the reaction is spontaneous.
- If ΔG > 0, the reaction is non-spontaneous.
- If ΔG = 0, the system is at equilibrium.
Gibbs Energy and Equilibrium Constant
For a general reaction:
A + B ⇌ C + D, the relation between Gibbs energy and equilibrium constant is:
ΔG = ΔG° + RT ln Q
Where:
- ΔG = Gibbs energy change at non-equilibrium conditions
- ΔG° = Standard Gibbs energy change
- R = Gas constant = 8.314 J mol⁻¹ K⁻¹
- T = Temperature in Kelvin
- Q = Reaction quotient
At equilibrium, ΔG = 0 and Q = K, so:
ΔG° = -RT ln K
Thus,
- If K > 1, then ln K is positive, so ΔG° < 0 → reaction is spontaneous.
- If K < 1, then ln K is negative, so ΔG° > 0 → reaction is non-spontaneous.
- If K = 1, then ΔG° = 0 → system is at equilibrium.
This relation connects thermodynamics (ΔG°) and equilibrium (K).
Factors Affecting Equilibria
Once equilibrium is achieved in a chemical system, changes in conditions (such as concentration, pressure, or temperature) can disturb the equilibrium. The system then adjusts itself to re-establish a new equilibrium state. These changes and the system’s response are governed by Le Chatelier’s Principle.
Le Chatelier’s Principle
If a system at equilibrium is disturbed by a change in concentration, pressure, or temperature, the system will shift its equilibrium position to counteract the disturbance. The principle helps in:
- Predicting the direction of shift in equilibrium.
- Optimizing reaction conditions in industrial and laboratory processes.
Effect of Concentration Change
The following are the effects of a concentration change.
- Adding a reactant increases its concentration:
- The equilibrium shifts in the forward direction to reduce the added concentration.
- Removing a reactant or adding a product:
- Equilibrium shifts in the reverse direction.
- Example:
- In the equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Effect of Pressure Change
Pressure changes affect gaseous reactions involving a change in the number of moles.
Increase in Pressure
If pressure is increased by decreasing volume, the equilibrium shifts in the direction with fewer gas molecules. Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (4 moles ⇌ 2 moles). Increased pressure shifts the equilibrium toward NH₃.
Decrease in Pressure
- If pressure is decreased, the equilibrium shifts in the direction with more gas molecules.
- If the number of moles on both sides is equal, the pressure change has no effect.
Effect of Inert Gas Addition
Addition of inert gas at constant volume: No effect on equilibrium as partial pressures remain unchanged. Addition of inert gas at constant pressure: Volume increases, partial pressures decrease. Equilibrium shifts toward the side with more gas molecules.
Effect of Temperature Change
Endothermic reaction (heat is absorbed): A + B + heat ⇌ C + D
- An increase in temperature shifts the equilibrium forward.
- A decrease in temperature shifts it backward.
Exothermic reaction (heat is released): A + B ⇌ C + D + heat
- An increase in temperature shifts the equilibrium backward.
- A decrease in temperature shifts it forward.
Ionic Equilibrium in Solutions
Chemical reactions in aqueous solutions often involve the formation of ions and reach equilibrium. This equilibrium is referred to as ionic equilibrium. It deals with the behavior of acids, bases, salts, and the extent of their ionisation in water.
Acids, Bases, and Salts
The concepts of acids, bases, and salts are given below.
- Arrhenius Concept:
- Acids: substances that increase the H⁺ concentration in water.
- Bases: substances that increase OH⁻ concentration in water.
- Brønsted-Lowry Concept:
- Acids: proton (H⁺) donors.
- Bases: proton (H⁺) acceptors.
- Lewis Concept:
- Acids: electron-pair acceptors.
- Bases: electron-pair donors.
- Salts are products formed from acid-base neutralization.
Ionisation of Acids and Bases
The following are the effects of the ionization of acids and bases.
- Strong acids and bases completely ionise in aqueous solutions.
- Weak acids and bases only partially ionise.
- Ionisation equilibrium is established between unionised molecules and ions.
- Example for a weak acid:
CH₃COOH ⇌ CH₃COO⁻ + H⁺
Ionisation Constants of Weak Acids
The acid dissociation constant (Ka) quantifies the ionisation of a weak acid:
Ka = [H⁺][A⁻] / [HA]. Larger the Ka, the stronger the acid. The degree of ionization increases with:
- Decrease in the concentration of acid.
- Increase in dielectric constant of the solvent.
Ionization of Weak Bases
The base dissociation constant (Kb) is used for weak bases:
Kb = [BH⁺][OH⁻] / [B]. A higher Kb value indicates a stronger base.
Relation between Ka and Kb
For a conjugate acid-base pair: Ka × Kb = Kw
- Where Kw is the ionic product of water:
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² L⁻² at 298 K - Also: pKa + pKb = 14
Factors Affecting Acid Strength
The following are the key factors that affect the acid strength
- More electronegative atoms bonded to H make the acid stronger.
- Larger size of the conjugate base leads to a weaker H–X bond → stronger acid.
- Inductive effect, resonance, and hydration also influence strength.
Common Ion Effect in the Ionisation of Acids and Bases
Suppression of ionisation of a weak electrolyte by adding a strong electrolyte having a common ion. Example: Addition of HCl suppresses the ionization of CH₃COOH due to the common ion H⁺.
Hydrolysis of Salts and the pH of their Solutions
Reaction of salt with water to produce acidic or basic solutions. Depends on the nature of the acid and base that form the salt:
- Salt of strong acid + strong base → neutral solution.
- Salt of weak acid + strong base → basic solution.
- Salt of strong acid + weak base → acidic solution.
- Salt of weak acid + weak base → pH depends on Ka and Kb.
Buffer Solutions
Buffer solution: Resists change in pH on addition of small amounts of acid or base. Types:
- Acidic buffer: Weak acid + its salt (e.g., CH₃COOH + CH₃COONa)
- Basic buffer: Weak base + its salt (e.g., NH₄OH + NH₄Cl)
- Henderson-Hasselbalch equation (for acidic buffer): pH = pKa + log([salt]/[acid])
Solubility Equilibria of Sparingly Soluble Salts
Sparingly soluble salts establish an equilibrium between solid and dissolved ions. Example: BaSO₄(s) ⇌ Ba²⁺(aq) + SO₄²⁻(aq). Solubility product constant (Ksp): Ksp = [Ba²⁺][SO₄²⁻]. Used to predict precipitation and calculate solubility.
Important Formulas in NCERT Notes Class 11 Chemistry (Part-I) Chapter 6: Equilibrium
Here are the important formulas and equations from NCERT Class 11 Chemistry Chapter 6: Equilibrium:
- Equilibrium Constant (Concentration):
Kc = [C]^c [D]^d / [A]^a [B]^b - Equilibrium Constant (Partial Pressure):
Kp = (PC)^c (PD)^d / (PA)^a (PB)^b - Relation Between Kp and Kc:
Kp = Kc × (RT)^Δn
where Δn = moles of gaseous products − moles of gaseous reactants - Gibbs Free Energy and Equilibrium Constant:
ΔG° = −RT ln K - Reaction Quotient (Qc):
Qc = [C]^c [D]^d / [A]^a [B]^b - Acid Ionisation Constant:
Ka = [H⁺][A⁻] / [HA] - Base Ionisation Constant:
Kb = [BH⁺][OH⁻] / [B] - Ionic Product of Water:
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² L⁻² at 298 K - Henderson-Hasselbalch Equation:
pH = pKa + log([Salt] / [Acid]) - Solubility Product Constant (for AB ⇌ A⁺ + B⁻):
Ksp = [A⁺][B⁻]
Explore Notes of Class 11 Chemistry
Download the Solutions of Other Chapters of Class 11 Chemistry
Related Reads
Source- Munil Sir
Explore notes on other subjects in the NCERT Class 11
FAQs
Answer: Physical equilibrium involves physical changes such as phase transitions (e.g., liquid ⇌ vapour, solid ⇌ liquid), where the substance remains the same but changes state. Chemical equilibrium involves reversible chemical reactions where reactants form products and vice versa, with no net change in concentrations once equilibrium is established.
Answer: Le Chatelier’s Principle helps predict how a change in conditions (concentration, pressure, or temperature) will shift the equilibrium position.
In industries, it is used to:
Maximise product formation (e.g., ammonia in the Haber process).
Minimise waste.
Optimize temperature and pressure conditions for better yield and efficiency.
Answer: The spontaneity of a chemical reaction is determined by the Gibbs free energy change (ΔG), which is given by the equation:
ΔG = ΔH – TΔS, where ΔH is the enthalpy change, T is the temperature, and ΔS is the entropy change.
If ΔG < 0, the reaction is spontaneous.
If ΔG > 0, the reaction is non-spontaneous.
If ΔG = 0, the system is at equilibrium.
For more topics, follow LeverageEdu NCERT Study Material today!
