the concept of dynamic equilibrium for reversible reactions, the equilibrium law expression and equilibrium constant Kc (including the meaning of Q vs Kc and the units of Kc), and the qualitative application of Le Chatelier's principle to predict the effect on equilibrium of changes in concentration, gas pressure (volume), temperature and the addition of a catalyst

What this dot point is asking

VCAA wants the dynamic equilibrium concept (reversible reactions with equal forward and reverse rates), the equilibrium law expression with the equilibrium constant Kc (including units and the meaning of Q vs Kc), and the qualitative application of Le Chatelier's principle to predict shifts when concentration, pressure (volume), or temperature changes, or a catalyst is added.

The answer

Dynamic equilibrium

A reaction is at dynamic equilibrium when:

1. It is occurring in a closed system (no matter enters or leaves).
2. The forward and reverse reactions are still happening (the equilibrium is "dynamic", not static).
3. The rates of the forward and reverse reactions are equal, so the concentrations of reactants and products stay constant.

A reaction at equilibrium has not stopped; it is occurring in both directions at the same rate. There is no net change in the amount of any species, but molecules are still being interconverted continually.

The equilibrium expression and Kc

For the general equilibrium:

aA + bB ⇌ cC + dD

the equilibrium law expression is:

Kc = [C]^c × [D]^d / ([A]^a × [B]^b)

with square-bracketed concentrations measured at equilibrium in mol L^-1. The product of product concentrations (raised to their coefficients) is on the top, and the product of reactant concentrations (raised to their coefficients) is on the bottom.

Rules:

• Only (g) and (aq) species appear in the expression. Pure (l) and (s) species are omitted because their effective concentrations are constant.
• The value of Kc depends only on temperature for a given reaction.
• Units of Kc depend on the stoichiometry. For aA + bB ⇌ cC + dD, units are (mol L^-1)^((c+d)-(a+b)). Many marking schemes accept "no units" if the convention is applied consistently.

Reading the value of Kc

• Kc >> 1: products dominate at equilibrium (the position lies far to the right).
• Kc << 1: reactants dominate at equilibrium (position lies far to the left).
• Kc ≈ 1: comparable amounts of reactants and products.

Q vs Kc

The reaction quotient Q has the same algebraic form as Kc but uses current concentrations (which may or may not be at equilibrium).

Q = [products] / [reactants] (using current concentrations)

Comparison with Kc tells you which direction the reaction will move:

When Q is computed before equilibrium is reached, the system shifts in whichever direction will move Q towards Kc.

Le Chatelier's principle

Le Chatelier's principle: if a system at equilibrium is disturbed, it shifts to partially counteract the disturbance and restore equilibrium.

Four disturbances are examined in VCE:

1. Change in concentration

• Adding more of a reactant: shift forward (consume the added reactant).
• Removing a product: shift forward (replace the removed product).
• Adding a product: shift reverse.
• Removing a reactant: shift reverse.

Kc is unchanged because temperature is unchanged.

2. Change in pressure (by changing volume)

Only matters for reactions with gases. Compare the moles of gas on each side.

• Increasing pressure (decreasing volume): shift to the side with fewer moles of gas.
• Decreasing pressure (increasing volume): shift to the side with more moles of gas.
• If both sides have the same number of moles of gas, no shift occurs.

Kc is unchanged.

Adding an inert gas at constant volume does not shift the equilibrium because partial pressures (and concentrations) of reactants and products are unchanged.

3. Change in temperature

The only disturbance that changes Kc.

• Increasing temperature: shift in the endothermic direction (the direction that absorbs heat). For an exothermic forward reaction, this means shifting to the left, decreasing Kc.
• Decreasing temperature: shift in the exothermic direction. For an exothermic forward reaction, this means shifting right, increasing Kc.

Treat heat as a "reactant" or "product":

• Exothermic forward: A + B ⇌ C + D + heat. Adding heat shifts left.
• Endothermic forward: A + B + heat ⇌ C + D. Adding heat shifts right.

4. Adding a catalyst

A catalyst lowers the activation energy of both the forward and reverse reactions equally, so equilibrium is reached faster but the position and Kc are unchanged.

Visualising on a concentration vs time graph

A common SAC/exam graph plots concentration on the y-axis and time on the x-axis. Reactant and product concentrations level off when equilibrium is reached. A disturbance shows up as a vertical jump or drop in the concentration of the disturbed species, then a re-equilibration as the system shifts.

For example, adding more H2 to the N2/H2/NH3 equilibrium shows a vertical jump in [H2], then a fall in [H2] and [N2] (consumed) and a rise in [NH3] (produced) until new constant levels are reached.

Worked example

For the Haber process:
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ mol^-1

Predict and explain the effect on (a) the equilibrium yield of NH3 and (b) the value of Kc, for each change.

Industrial Haber process compromise: high pressure (favours NH3) and moderate temperature (about 400 to 500°C) plus an iron catalyst. The temperature compromise sacrifices yield for an acceptable rate; the catalyst speeds the approach to equilibrium without affecting yield.

Common traps

Saying equilibrium means "the reaction has stopped". It has not. Forward and reverse rates are equal but non-zero.

Including solids or pure liquids in Kc. Only (g) and (aq) species appear in the expression.

Saying a catalyst increases yield. A catalyst affects rate only, not the equilibrium position or Kc.

Confusing "shift left" with "decrease Kc" for concentration changes. Only temperature changes Kc. A concentration disturbance shifts the position but the system returns to the same Kc.

Forgetting to compare moles of gas for pressure changes. The shift direction depends on which side has fewer moles of gas, not on which side has more "atoms" or "mass".

Treating "adding an inert gas at constant volume" as a pressure change for the reaction. The partial pressures of the reactants and products do not change, so no shift occurs.

Applying Le Chatelier without checking ΔH. For temperature changes you must know the sign of ΔH for the forward reaction. Without it, you cannot predict the direction.

In one sentence

At dynamic equilibrium in a closed system the forward and reverse rates are equal so concentrations are constant; Kc is the temperature-dependent ratio of product to reactant concentrations (each raised to its coefficient), Q compared with Kc tells you which way the system will shift, and Le Chatelier's principle predicts qualitative shifts when concentration, pressure (volume), or temperature change (catalysts change only rate, not position).