Chemical Equilibrium class 11 Notes

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Chemical Equilibrium: Dynamic Balance in Reactions

Chemical Equilibrium: Dynamic Balance in Reactions

Reaction Dynamics

Understanding the state where forward and reverse reactions occur at equal rates, maintaining constant concentrations

What is Chemical Equilibrium?

Chemical equilibrium is the dynamic state where the rates of the forward and reverse reactions are equal, resulting in no net change in the concentrations of reactants and products. It represents a balance, not a cessation, of chemical activity.

Dynamic Nature

Reactions continue occurring but at equal rates in both directions

$$ \text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g) $$

Equilibrium Constant

Quantitative measure of the equilibrium position

$$ K_c = \frac{[\text{products}]}{[\text{reactants}]} $$

Le Chatelier’s Principle

Systems at equilibrium respond to minimize disturbances

The Equilibrium Constant (K)

Writing Equilibrium Expressions

  1. Products in numerator, reactants in denominator
  2. Concentrations raised to their stoichiometric coefficients
  3. Pure solids/liquids are omitted (activity = 1)
  4. K depends only on temperature

Example: Haber Process

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

$$ K_c = \frac{[\text{NH}_3]^2}{[\text{N}_2][\text{H}_2]^3} $$

Interpreting K Values

K Value Interpretation
K > 10³ Products favored (reaction goes nearly to completion)
10⁻³ < K < 10³ Significant amounts of both reactants and products
K < 10⁻³ Reactants favored (very little product formed)

Relationship Between Kₚ and K꜀

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

Where Δn = moles gaseous products – moles gaseous reactants

Reaction Quotient (Q) vs Equilibrium Constant (K)

Predicting Reaction Direction

  • Q < K: Forward reaction favored to reach equilibrium
  • Q = K: System is at equilibrium
  • Q > K: Reverse reaction favored to reach equilibrium

Interactive Q Calculator

Le Chatelier’s Principle

Stress Applied
System Response
↑ [Reactant]
Shifts toward products to consume added reactant
↓ [Product]
Shifts toward products to replace removed product
↑ Pressure
Shifts toward side with fewer moles of gas
↑ Temperature
Shifts to absorb heat (endothermic direction)
Catalyst Added
No shift (speeds both forward and reverse equally)

Industrial Application: Haber Process

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol

  • High pressure (200 atm): Favors product side (fewer gas moles)
  • Moderate temperature (450°C): Compromise between kinetics and equilibrium
  • Iron catalyst: Increases rate without affecting equilibrium

Equilibrium Calculations

ICE Tables

Systematic approach to solve equilibrium problems:

  1. Initial concentrations
  2. Change in concentrations
  3. Equilibrium concentrations

Example Problem

For the reaction H₂(g) + I₂(g) ⇌ 2HI(g), K꜀ = 50. If initial [H₂] = [I₂] = 0.5M, find equilibrium concentrations.

Approximation Method

When K is very small (< 10⁻³), the change (x) is negligible compared to initial concentrations:

$$ \text{If } K \text{ is small: } [\text{Reactant}]_{\text{eq}} \approx [\text{Reactant}]_{\text{initial}} $$

Example

For N₂(g) + O₂(g) ⇌ 2NO(g), K꜀ = 4.7×10⁻³¹ at 25°C. If [N₂]₀ = 0.8M, [O₂]₀ = 0.2M:

x ≈ √(K꜀ × [N₂] × [O₂]) = √(4.7×10⁻³¹ × 0.8 × 0.2) ≈ 2.7×10⁻¹⁶ M

Solubility Equilibria

Solubility Product Constant (Kₛₚ)

$$ \text{For } \text{A}_x\text{B}_y(s) \rightleftharpoons x\text{A}^{y+}(aq) + y\text{B}^{x-}(aq) $$ $$ K_{sp} = [\text{A}^{y+}]^x[\text{B}^{x-}]^y $$
  • Only for sparingly soluble ionic compounds
  • Higher Kₛₚ = more soluble
  • Used to predict precipitation

Common Ion Effect

Adding a common ion decreases solubility:

$$ \text{PbCl}_2(s) \rightleftharpoons \text{Pb}^{2+}(aq) + 2\text{Cl}^-(aq) $$

Adding NaCl (source of Cl⁻) shifts equilibrium left, reducing Pb²⁺ concentration

Practical Applications

Pharmaceutical Formulation

  • Optimizing drug solubility
  • Controlled release formulations
  • Buffer systems for pH stability

Water Treatment

  • Controlling calcium carbonate scaling
  • Heavy metal precipitation
  • Disinfection byproduct management

Industrial Synthesis

  • Ammonia production (Haber process)
  • Sulfuric acid (Contact process)
  • Methanol synthesis

Biological Systems

  • Oxygen binding to hemoglobin
  • Enzyme-substrate interactions
  • Acid-base balance in blood

Conclusion

Chemical equilibrium represents a fundamental concept in chemistry with wide-ranging applications from industrial synthesis to biological systems. Understanding how to manipulate equilibrium conditions through Le Chatelier’s Principle allows chemists to optimize reaction yields, while quantitative analysis using equilibrium constants provides predictive power. The dynamic nature of equilibrium reveals the continuous molecular activity underlying apparently static chemical systems.

Previous: Chemical Kinetics Next: Chemical Thermodynamics

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