Equilibrium Constant

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definition

Also known as: Keq, K, equilibrium expression

Grade 9-12

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The ratio of product concentrations to reactant concentrations at equilibrium, each raised to its stoichiometric coefficient. The equilibrium constant predicts reaction yield, determines solubility of ionic compounds, and governs acid-base behavior.

Definition

The ratio of product concentrations to reactant concentrations at equilibrium, each raised to its stoichiometric coefficient.

💡 Intuition

K is the scoreboard at the end of the game — it tells you which side won. Large K means products dominate; small K means reactants dominate.

🎯 Core Idea

Changing concentrations does NOT change K — it changes Q. Only temperature changes K. This is the most common point of confusion.

Example

For the reaction N₂O₄ ⇌ 2NO₂, K_c = \frac{[NO_2]^2}{[N_2O_4]}. If K = 0.2 at 25°C, the colorless N₂O₄ dominates — the mixture stays mostly clear.

Formula

K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b} \quad \text{for } aA + bB \rightleftharpoons cC + dD

Notation

K_c uses molar concentrations (mol/L) in square brackets. K_p uses partial pressures (atm) for gaseous equilibria. The subscript tells you which units the expression uses.

🌟 Why It Matters

The equilibrium constant predicts reaction yield, determines solubility of ionic compounds, and governs acid-base behavior. In industry, knowing K helps chemists decide whether a reaction is feasible and what conditions will maximize product formation.

💭 Hint When Stuck

When writing or using an equilibrium expression, follow these steps carefully. First write the balanced equation and identify all species and their states. Then construct K_{eq} with products in the numerator and reactants in the denominator, each raised to its coefficient power. Finally, exclude pure solids and pure liquids — only include aqueous and gaseous species.

Formal View

For gaseous systems, K_p = \frac{(P_C)^c(P_D)^d}{(P_A)^a(P_B)^b}. The two forms are related by K_p = K_c(RT)^{\Delta n} where \Delta n = (c+d)-(a+b) is the change in moles of gas.

🚧 Common Stuck Point

Pure solids and liquids are excluded from the equilibrium expression.

⚠️ Common Mistakes

  • Including pure solids or pure liquids in the equilibrium expression — only aqueous (aq) and gaseous (g) species appear in K
  • Forgetting to raise concentrations to the power of their coefficients — for 2A \rightleftharpoons B, K = \frac{[B]}{[A]^2}, not \frac{[B]}{[A]}
  • Thinking changing concentration changes K — adding or removing a substance changes the reaction quotient Q, but K remains constant unless temperature changes

Frequently Asked Questions

What is Equilibrium Constant in Chemistry?

The ratio of product concentrations to reactant concentrations at equilibrium, each raised to its stoichiometric coefficient.

What is the Equilibrium Constant formula?

K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b} \quad \text{for } aA + bB \rightleftharpoons cC + dD

When do you use Equilibrium Constant?

When writing or using an equilibrium expression, follow these steps carefully. First write the balanced equation and identify all species and their states. Then construct K_{eq} with products in the numerator and reactants in the denominator, each raised to its coefficient power. Finally, exclude pure solids and pure liquids — only include aqueous and gaseous species.

How Equilibrium Constant Connects to Other Ideas

To understand equilibrium constant, you should first be comfortable with equilibrium and concentration. Once you have a solid grasp of equilibrium constant, you can move on to le chateliers principle.

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