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.
Understanding Chemistry: The Equilibrium Constant
The equilibrium constant is most useful when it is compared with the reaction quotient, often called Q. Q is calculated in the same pattern as the equilibrium constant, but it uses the amounts present at one particular moment. If Q is smaller than K, the mixture has too few products for its equilibrium state.
The forward reaction then has a net effect. If Q is larger than K, there are too many products, so the reverse reaction has a net effect.
If Q equals K, there is no net change in the overall amounts. This comparison lets chemists predict the direction of change before the reaction has settled.
The coefficients in a balanced equation matter greatly. They become powers in the equilibrium calculation, so a small concentration change can have a large effect. For example, if a substance has a coefficient of two, its concentration is squared.
Students often make mistakes by using the subscripts inside a chemical formula instead of the coefficients in front of it. A subscript tells how atoms are arranged in one particle.
A coefficient tells the reaction ratio between substances. Write and check the balanced equation before building any equilibrium expression.
Equilibrium constants follow useful rules when equations are changed. Reversing a reaction makes the new constant the reciprocal of the old one. Multiplying every coefficient in an equation by a number raises the constant to that same power.
Adding two reactions means multiplying their constants. These rules help when a target reaction is built from known reactions. They are especially important in thermochemistry and electrochemistry, where equations are regularly rearranged.
The size of K describes the final balance, but it does not tell how quickly equilibrium is reached. A reaction can have a very large K yet proceed extremely slowly.
Temperature has a special role because changing it changes the value of K itself. For a reaction that releases heat, raising the temperature usually lowers K. For a reaction that absorbs heat, raising the temperature usually raises K.
Changes in concentration, pressure, or volume do not alter K at a fixed temperature. They only disturb the mixture and change Q, causing a shift until Q matches K again. A catalyst does not change K either.
It speeds up both directions, so equilibrium is reached sooner without changing the final composition. In real systems such as the manufacture of ammonia, chemists choose pressure and temperature carefully because they must balance product yield, reaction speed, energy cost, and safe operation.
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.