Chemical Equilibrium and Le Châtelier’s Principle
Students predict how changes in concentration, pressure, and temperature shift a reversible reaction and apply these ideas to optimizing industrial ammonia production.
Illustrations are auto-generated and may be placeholders. They can be refreshed to match the narration.
Dynamic Equilibrium in Reversible Reactions
A reversible reaction can proceed in both directions: reactants form products, and products can react to form reactants. In a closed system, the reaction may reach dynamic equilibrium. At equilibrium, particles continue reacting, but the forward and reverse reaction rates are equal. As a result, the concentrations of reactants and products remain constant, although they are not necessarily equal. For example, colorless dinitrogen tetroxide gas reversibly forms brown nitrogen dioxide: N2O4 ⇌ 2NO2. At equilibrium, N2O4 molecules still separate into NO2 molecules while NO2 molecules continue combining into N2O4. The visible shade of brown stays constant because both processes occur at the same rate. Dynamic equilibrium requires a closed system so that matter cannot enter or escape.
Reading Concentration-versus-Time Graphs
A concentration-versus-time graph shows how reactant and product concentrations change as a reaction approaches equilibrium. Time is plotted on the horizontal axis, and concentration is plotted on the vertical axis. For N2O4 ⇌ 2NO2, a sample beginning mostly with N2O4 shows a decreasing N2O4 curve and an increasing NO2 curve. Both curves eventually become horizontal, indicating that their concentrations are no longer changing. This flat region marks equilibrium, not the end of all reactions. The plateau heights do not need to match because equilibrium requires equal reaction rates rather than equal concentrations. A steeper curve indicates a faster concentration change. If the system is disturbed, a sudden jump may show that a substance was added or removed, followed by gradual changes as the system moves toward a new equilibrium.
Le Châtelier’s Principle
Le Châtelier’s principle predicts how an equilibrium system responds to a disturbance. If concentration, pressure, or temperature changes, the system shifts in the direction that reduces the effect of that change and then establishes a new equilibrium. A shift to the right increases product formation, while a shift to the left increases reactant formation. Consider Fe3+ + SCN− ⇌ FeSCN2+, in which FeSCN2+ produces a deep red solution. Adding Fe3+ causes the system to consume some of the added ion, so equilibrium shifts right and the red color becomes darker. Removing SCN− shifts the system left and makes the color lighter. The response does not completely cancel the disturbance. At constant temperature, the equilibrium constant remains unchanged even though equilibrium concentrations change.
Effects of Concentration, Pressure, and Temperature
Different disturbances affect equilibrium in predictable ways. Adding a reactant or removing a product generally shifts a reaction toward products; removing a reactant or adding a product shifts it toward reactants. For gas reactions, decreasing volume raises pressure and favors the side with fewer moles of gas. Increasing volume favors the side with more gas particles. Temperature must be treated as energy: heating favors the endothermic direction, while cooling favors the exothermic direction. In N2 + 3H2 ⇌ 2NH3 + heat, removing NH3 shifts equilibrium right. Compressing the mixture also shifts it right because four moles of reactant gas form two moles of product gas. Raising temperature shifts it left because the forward reaction releases heat. A catalyst speeds both directions equally and reaches equilibrium sooner, but it does not shift equilibrium.
Optimizing the Haber Process
The Haber process produces ammonia for fertilizers using N2 + 3H2 ⇌ 2NH3 + heat. High pressure increases equilibrium ammonia yield because the product side has fewer gas moles. Low temperature also favors ammonia because the forward reaction is exothermic, but very low temperature makes production too slow. Industry therefore uses a compromise of about 400–500°C, pressures commonly near 150–250 atmospheres, and an iron-based catalyst. The catalyst increases production rate without changing equilibrium yield. Cooling the product mixture removes liquid ammonia, while unreacted nitrogen and hydrogen are recycled, shifting continued production toward ammonia. Economic decisions compare marginal benefits and costs. A small pressure increase may produce more ammonia, but its added benefit must be weighed against extra energy use, stronger equipment, maintenance, and safety risks. The best operating conditions maximize useful output without making the next improvement cost more than it adds in value.
