What Is the Equilibrium Constant?
The equilibrium constant, denoted K, measures the relative amounts of reactants and products present when a chemical system reaches equilibrium. For a general reaction aA + bB ⇌ cC + dD, the constant expresses how far the reaction proceeds toward products or remains biased toward reactants.
Two common forms exist:
- Kc — based on molar concentrations (mol/L) at equilibrium
- Kp — based on partial pressures of gases, typically in atmospheres
A large K value indicates products are strongly favoured; a small K suggests reactants predominate. K = 1 means reactants and products are equally balanced.
The Equilibrium Constant Expression
For any reversible reaction, raise each concentration or pressure to the power of its stoichiometric coefficient, multiply products together, divide by reactants multiplied together:
K = ([C]c × [D]d) ÷ ([A]a × [B]b)
K— Equilibrium constant (dimensionless or unit-dependent on reaction)[A], [B]— Molar concentrations of reactants at equilibrium (mol/L)[C], [D]— Molar concentrations of products at equilibrium (mol/L)a, b, c, d— Stoichiometric coefficients from the balanced equation
Practical Applications in Chemistry and Biochemistry
The equilibrium constant is indispensable across multiple fields. In industrial chemistry, K guides optimisation of synthesis conditions: sulphuric acid production relies on shifting the SO₂ + O₂ equilibrium toward SO₃. In biochemistry, haemoglobin's oxygen binding and the pH buffering systems in blood depend on precise equilibrium constants for function.
Medical diagnostics also use K concepts: transferrin saturation, an indicator of iron metabolism and potential anemia, reflects equilibrium between bound and free iron in plasma. Environmental chemists apply K to predict pollutant speciation in water systems, while pharmaceutical researchers use it to understand drug–protein binding affinity.
How to Measure Concentrations for K Calculation
Determining K requires accurate experimental measurement of all reactant and product concentrations once equilibrium has been established. Common analytical techniques include:
- Potentiometry — voltage measurement using electrodes, particularly useful for acid–base equilibria
- Spectrophotometry — measuring light absorption by coloured species at specific wavelengths
- NMR spectroscopy — identifying and quantifying molecules via nuclear magnetic resonance
- Gas chromatography — separating and measuring volatile reactants and products
- Calorimetry — inferring concentrations from heat changes during the reaction
Once concentrations are known, substitute them into the K expression with their stoichiometric exponents.
Common Pitfalls When Calculating K
Avoid these mistakes to ensure accurate equilibrium constant determination:
- Forgetting stoichiometric exponents — The coefficients in the balanced equation must become exponents in the K expression. For 2NO + O₂ ⇌ 2NO₂, K = [NO₂]² ÷ ([NO]² × [O₂]), not [NO₂] ÷ ([NO] × [O₂]). Omitting exponents leads to wildly incorrect values.
- Measuring concentrations before equilibrium is reached — K only applies at true equilibrium. If you measure too early, concentrations are still changing and your calculated K will be inconsistent with later measurements. Always allow sufficient time and confirm equilibrium by checking that concentrations remain stable.
- Confusing K with reaction quotient Q — Q is calculated the same way as K but uses non-equilibrium concentrations. When Q < K, the reaction shifts forward; when Q > K, it shifts backward. Only Q = K indicates equilibrium.
- Ignoring temperature effects — K changes significantly with temperature. A K value is only valid at the specific temperature where it was measured. Heating or cooling shifts equilibrium, requiring recalculation of K or use of the van't Hoff equation.