Qualitative predictions for chemical equilibrium shifts
The rate and extent of chemical change • Reversible reactions and equilibrium
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Key concepts
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Le Chatelier’s principle - definition
Le Chatelier’s principle states that a system at equilibrium responds to an external change by shifting in the direction that tends to reduce the effect of that change. The principle applies to changes in concentration, pressure, temperature, and the presence of catalysts, provided the system is closed and reversible. Use the balanced equation to determine which side (reactants or products) the shift favours after a disturbance.
Concentration changes - cause and effect
Cause: Increase in concentration of a reactant or product. Effect: Equilibrium shifts toward the opposite side to consume the added substance. Cause: Decrease in concentration. Effect: Equilibrium shifts toward the side with the removed substance to replace it. Limiting factors: Solids and pure liquids do not appear in equilibrium expressions and their amounts do not directly affect Kc, but changing the concentration of species in solution or gas phase does.
Pressure changes for reactions involving gases
Cause: Increase in total pressure (by reducing volume) for a gaseous equilibrium. Effect: Equilibrium shifts toward the side with fewer moles of gas to reduce pressure. Cause: Decrease in pressure (by increasing volume). Effect: Equilibrium shifts toward the side with more moles of gas. Limiting factors: Adding an inert gas at constant volume does not change partial pressures relevant to equilibrium and therefore does not shift the position of equilibrium.
Temperature changes and the equilibrium constant
Cause: Increase in temperature. Effect: Equilibrium shifts in the endothermic direction (absorbs heat) to oppose the temperature rise. Cause: Decrease in temperature. Effect: Equilibrium shifts in the exothermic direction (releases heat). Important distinction: Temperature changes alter the value of the equilibrium constant (K). For exothermic reactions, increasing temperature decreases K. For endothermic reactions, increasing temperature increases K.
Catalysts - rate effect without shifting equilibrium
Cause: Addition of a catalyst. Effect: Both forward and reverse reaction rates increase equally, so the system reaches equilibrium faster but the position of equilibrium and the equilibrium constant remain unchanged. Limiting factors: Catalysts do not affect thermodynamic quantities such as K or the concentrations at equilibrium; they only affect the time needed to reach equilibrium.
Reaction quotient Q and direction of spontaneous shift
Definition: The reaction quotient Q uses the same expression as K but with current concentrations or partial pressures. Comparison rule: If Q < K, the forward reaction is favored and the system shifts toward products. If Q > K, the reverse reaction is favored and the system shifts toward reactants. If Q = K, the system is at equilibrium and no net shift occurs. Use Q vs K to predict spontaneous direction before and after a change.
Heterogeneous equilibria and excluded species
Definition: Heterogeneous equilibrium involves reactants and products in different phases. Limiting factor: Pure solids and pure liquids do not appear in K expressions because their concentrations remain effectively constant. Qualitative predictions focus on changes in gaseous and aqueous species concentrations or temperature and pressure when those species are involved.
Key notes
Important points to keep in mind