Chemical equilibrium occurs when a reversible reaction continues in both directions but the concentrations of reactants and products stop changing. The equilibrium constant gives a numerical snapshot of that balance at a specific temperature. It helps chemists predict whether products or reactants are favored and compare different reactions.
This idea matters in acid-base chemistry, gas reactions, industrial synthesis, and many biological processes.
For a general reaction aA + bB ⇌ cC + dD, the equilibrium expression uses product concentrations over reactant concentrations, each raised to its coefficient. Kc is written using molar concentrations, while Kp is written using partial pressures for gases. Pure solids and pure liquids are not included because their effective concentrations stay essentially constant.
The relationship Kp = Kc(RT)^Δn connects the two forms when gases are involved.
Key Facts
- For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / ([A]^a[B]^b).
- For gas reactions, Kp = (P_C)^c(P_D)^d / ((P_A)^a(P_B)^b).
- Large K means products are favored at equilibrium, usually K >> 1.
- Small K means reactants are favored at equilibrium, usually K << 1.
- Kp = Kc(RT)^Δn, where Δn = moles of gaseous products minus moles of gaseous reactants.
- Pure solids and pure liquids are omitted from equilibrium expressions.
Vocabulary
- Chemical equilibrium
- A state in which the forward and reverse reaction rates are equal and concentrations remain constant.
- Equilibrium constant
- A number that describes the ratio of products to reactants at equilibrium for a reaction at a given temperature.
- Kc
- The equilibrium constant written using molar concentrations of dissolved or gaseous species.
- Kp
- The equilibrium constant written using partial pressures of gaseous species.
- Partial pressure
- The pressure a single gas in a mixture would exert if it occupied the container by itself.
Common Mistakes to Avoid
- Including pure solids or pure liquids in K. This is wrong because their activities are treated as constant and are already built into the value of K.
- Forgetting to raise each concentration or pressure to its coefficient. This gives the wrong product-to-reactant ratio because stoichiometric coefficients become exponents in the expression.
- Using initial concentrations instead of equilibrium concentrations. K must be calculated only from amounts present after the system has reached equilibrium.
- Assuming K changes when concentrations or pressures are changed. K changes only with temperature, while the reaction quotient Q changes as the system shifts.
Practice Questions
- 1 For N2(g) + 3H2(g) ⇌ 2NH3(g), write the Kc expression and calculate Kc if [N2] = 0.20 M, [H2] = 0.10 M, and [NH3] = 0.40 M at equilibrium.
- 2 For CO(g) + 2H2(g) ⇌ CH3OH(g), Kc = 14.5 at 500 K. Calculate Kp using R = 0.0821 L·atm/(mol·K).
- 3 A reaction has K = 2.0 x 10^-5 at a certain temperature. Explain whether reactants or products are favored at equilibrium and what that means for the position of the equilibrium scale.