Chemical equilibrium is a fundamental concept in JEE Chemistry covering reversible reactions, dynamic equilibrium, and the quantitative expression of equilibrium constants. Understanding this topic helps solve various complex problems related to reaction spontaneity, ionic balance, and solubility equilibria.
In a reversible reaction, the forward and backward reactions occur simultaneously. When the rate of forward reaction equals the rate of backward reaction, the system is said to be in dynamic equilibrium.
For the general reaction,
aA + bB \rightleftharpoons cC + dD
at equilibrium, concentrations satisfy the law of mass action.
The equilibrium constant quantifies the position of equilibrium:
K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}
Where square brackets denote molar concentrations.
Similarly, for gases, we use partial pressures to define:
K_p = \frac{P_C^c P_D^d}{P_A^a P_B^b}
The relation between \( K_p \) and \( K_c \) is given by:
K_p = K_c (RT)^{\Delta n}
Where \( \Delta n = (c + d) - (a + b) \) is the change in moles of gas, \( R \) is the gas constant, and \( T \) is temperature in Kelvin.
Le Chatelier’s principle states that if a change is imposed on a system at equilibrium, the system adjusts to partially counteract the change and establish a new equilibrium.
Increasing reactant concentration shifts equilibrium towards products, while increasing product concentration shifts it towards reactants.
Changing pressure affects equilibria involving gases. Increasing pressure shifts equilibrium toward the side with fewer moles of gas.
Temperature change shifts equilibrium depending on reaction enthalpy:
Catalyst speeds up both forward and backward reactions equally, so it does not affect equilibrium position but helps attain equilibrium faster.
Ionic equilibrium deals with equilibria involving ions in aqueous solutions — acid-base equilibria, hydrolysis, salt solutions, and buffer systems.
Arrhenius, Bronsted-Lowry, and Lewis definitions describe acids and bases.
For a weak acid \( HA \), dissociation is:
HA \rightleftharpoons H^+ + A^-
Acid dissociation constant \( K_a \):
K_a = \frac{[H^+][A^-]}{[HA]}
Similarly, for weak bases:
B + H_2O \rightleftharpoons BH^+ + OH^-
Base dissociation constant \( K_b \):
K_b = \frac{[BH^+][OH^-]}{[B]}
The ionization constant of water \( K_w \) is \( 1.0 \times 10^{-14} \) at 25°C:
K_w = [H^+][OH^-]
And,
K_a \times K_b = K_w
pH is the negative logarithm of hydrogen ion concentration:
pH = -\log[H^+]
Similarly, pOH:
pOH = -\log[OH^-]
At 25°C:
pH + pOH = 14
A buffer resists pH changes upon addition of small amounts of acid or base. Typically composed of a weak acid and its conjugate base or vice versa.
Henderson-Hasselbalch equation for buffer pH:
pH = pK_a + \log \frac{[A^-]}{[HA]}
Sparingly soluble salts establish equilibrium between solid and ions in solution.
Example:
AgCl (s) \rightleftharpoons Ag^+ (aq) + Cl^- (aq)
Solubility product constant:
K_{sp} = [Ag^+][Cl^-]
For salt \( MX \) dissociating into \( M^+ \) and \( X^- \):
MX \rightleftharpoons M^+ + X^-
Let solubility be \( s \) mol/L, then:
K_{sp} = s \times s = s^2
So, solubility \( s = \sqrt{K_{sp}} \).
Presence of a common ion suppresses solubility due to Le Chatelier’s principle. For example, solubility of AgCl decreases in presence of \( Cl^- \) ions.
Use ICE tables (Initial, Change, Equilibrium) to calculate concentrations of species.
Reaction quotient \( Q \) is calculated using initial concentrations. Comparing \( Q \) with \( K \) determines the direction:
For endothermic reactions, increasing temperature increases \( K \); for exothermic, it decreases \( K \).
| Concept | Formula | Notes |
|---|---|---|
| Equilibrium Constant (Kc) | \( K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b} \) | Molar concentration terms |
| Equilibrium Constant (Kp) | \( K_p = \frac{P_C^c P_D^d}{P_A^a P_B^b} \) | Partial pressure terms |
| Relation Kp and Kc | \( K_p = K_c (RT)^{\Delta n} \) | \( \Delta n \) = change in moles of gas |
| Acid dissociation constant | \( K_a = \frac{[H^+][A^-]}{[HA]} \) | For weak acids |
| Base dissociation constant | \( K_b = \frac{[BH^+][OH^-]}{[B]} \) | For weak bases |
| Water ion product | \( K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \) | At 25°C |
| pH and pOH | pH = -log[H+], pOH = -log[OH-] | pH + pOH = 14 (at 25°C) |
| Solubility product | \( K_{sp} = [M^+]^m [X^-]^x \) | For sparingly soluble salts |
| Henderson-Hasselbalch equation | \( pH = pK_a + \log \frac{[A^-]}{[HA]} \) | Buffer solutions |
Equilibrium is a challenging but rewarding topic in JEE Chemistry. With consistent practice and conceptual clarity, you can excel in this important area.