Chemical Equilibrium: The Balance of Reactions ⚖️

Chemical reactions can go both ways! When the forward and reverse reactions happen at the same speed, we reach dynamic equilibrium 🌀. At this point:

  • Concentrations of reactants and products stay constant over time ⏱️.
  • It doesn’t matter if you start with reactants or products – you’ll reach the same equilibrium mixture! 🔄
  • Isotope experiments (like using deuterium in ammonia synthesis) prove atoms keep swapping due to ongoing reactions – equilibrium isn’t “frozen”! ❄️➡️🔥

The Equilibrium Constant (Kc) 🔑

For any reversible reaction:

\[ aA + bB \rightleftharpoons cC + dD \]

The equilibrium constant Kc is given by:

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

Where [A], [B], [C], [D] are equilibrium concentrations (usually in mol/L).

Example 🧪: For \( H_2(g) + I_2(g) \rightleftharpoons 2HI(g) \):

\[ K_c = \frac{[HI]^2}{[H_2][I_2]} \]

Data shows this ratio is constant (~46.4–47.6) at 731K, no matter the starting concentrations! 📊

Key Rules for Kc 📜

  • Reverse reaction: \( K_{c,\text{reverse}} = \frac{1}{K_{c,\text{forward}}} \) 🔄
  • Multiply reaction by n: \( K_{c,\text{new}} = (K_c)^n \) ✖️
  • Ignore solids/liquids in Kc expressions! 🚫🧊

Practice Problems 💡

Problem 1: For \( N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \) at 500K:
[N2] = 1.5 × 10−2 M, [H2] = 3.0 × 10−2 M, [NH3] = 1.2 × 10−2 M.
Solution:

\[ K_c = \frac{[NH_3]^2}{[N_2][H_2]^3} = \frac{(1.2 \times 10^{-2})^2}{(1.5 \times 10^{-2})(3.0 \times 10^{-2})^3} = 3.55 \times 10^2 \]

Problem 2: For \( N_2(g) + O_2(g) \rightleftharpoons 2NO(g) \) at 800K:
[N2] = 3.0 × 10−3 M, [O2] = 4.2 × 10−3 M, [NO] = 2.8 × 10−3 M.
Solution:

\[ K_c = \frac{[NO]^2}{[N_2][O_2]} = \frac{(2.8 \times 10^{-3})^2}{(3.0 \times 10^{-3})(4.2 \times 10^{-3})} = 0.622 \]

Homogeneous Equilibria ⚗️

All reactants/products are in the same phase:

  • Gases: \( N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \) 🌬️
  • Aqueous solutions: \( Fe^{3+}(aq) + SCN^-(aq) \rightleftharpoons Fe(SCN)^{2+}(aq) \) 💧

Gaseous Systems & Partial Pressures 🌡️

For gas-phase reactions, K can also be written using partial pressures (Kp). The ideal gas law relates concentration and pressure:

\[ p = \frac{n}{V} RT \]

where p = pressure, n/V = concentration, R = gas constant, T = temperature.

NEET Super-Important Concepts! 🚨

  1. Dynamic Equilibrium: Rates of forward/reverse reactions are equal ⚖️.
  2. Kc Expression: Products over reactants, raised to their coefficients 📐.
  3. Kc vs. Reverse Reaction: \( K_{c,\text{reverse}} = \frac{1}{K_{c,\text{forward}}} \) 🔄
  4. Calculating Kc: Plug equilibrium concentrations into the expression 🧮.
  5. Homogeneous Equilibria: All species in same phase (gas/solution) 🌪️💧.

Keep practicing – equilibrium is all about balance! 🧘‍♂️✨