Concentration changes and chemical equilibrium
The rate and extent of chemical change • Reversible reactions and equilibrium
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Dynamic equilibrium and concentration
Dynamic equilibrium occurs in a closed system when forward and reverse reaction rates are equal, while concentrations of reactants and products remain constant. The constant concentrations at equilibrium result from continuous microscopic reactions occurring in both directions. A change in concentration disturbs the balance of rates. An increase in a reactant increases the forward rate relative to the reverse, causing a net shift that produces more products until the rates balance again.
Le Chatelier’s principle for concentration changes
Le Chatelier’s principle states that a system at equilibrium responds to a disturbance by shifting in the direction that counteracts the disturbance. Adding a reactant causes the equilibrium to shift toward products; removing a reactant causes a shift toward reactants. Cause → effect explanations follow directly: increase reactant concentration → forward reaction is favoured → product concentrations rise. Decrease product concentration → reverse reaction is favoured → product concentration replenishes.
Reaction quotient Q and equilibrium constant Kc
The reaction quotient Q uses the same expression as Kc but with current concentrations rather than equilibrium values. Comparison of Q and Kc predicts the direction of shift: Q < Kc causes a forward shift; Q > Kc causes a reverse shift; Q = Kc indicates equilibrium. Cause → effect interpretation uses inequality: Q < Kc → forward reaction has greater capacity to form products than required for equilibrium → reaction moves to produce more products until Q equals Kc.
Interpreting data to predict shifts (HT)
Given concentration or mole data, calculation of Q and comparison with Kc gives a definite prediction of shift direction. The calculation requires substituting the given concentrations into the expression for Q and simplifying to a numeric value. Cause → effect reasoning then follows: if Q is less than the quoted Kc, the forward reaction increases product concentration; if Q is greater, the reverse reaction increases reactant concentration. Temperature must match the Kc reference value for the comparison to hold.
Limiting factors and conditions
Temperature changes alter Kc and therefore change the equilibrium position in a way that cannot be predicted by concentration alone. Pressure changes affect equilibria containing gases according to changes in total gas moles, while catalysts alter the rates but not the position of equilibrium. Cause → effect limitations: concentration changes shift equilibrium only within the constraints of constant temperature and a closed system. Open systems, large volume changes, or temperature variation invalidate direct Q versus Kc comparison.
Stepwise method for HT predictions from data
Step 1: Write the balanced equation and the Kc expression. Step 2: Substitute the given concentrations into the Q expression and calculate Q. Step 3: Compare Q with Kc and state the direction of shift. Step 4: Describe the expected changes in concentration qualitatively or calculate new equilibrium concentrations using an ICE table if required. Cause → effect emphasis: calculation of Q (cause) leads to a clear predicted shift (effect). Accurate interpretation requires attention to units, stoichiometric coefficients, and whether concentrations or partial pressures are used in the equilibrium expression.
Key notes
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