The Reaction Quotient, Q HL Only
Quick Notes
- The reaction quotient (Q) has the same form as the equilibrium constant (K) but uses non-equilibrium concentrations.
- Formula:
![IB Chemistry reaction quotient expression showing Q = [C]^c [D]^d / [A]^a [B]^b.](images/qexpression.png)
- Comparison with K:
- Q < K → reaction shifts right (towards products)
- Q > K → reaction shifts left (towards reactants)
- Q = K → system at equilibrium (no shift)
Full Notes
What Is the Reaction Quotient, Q?
The reaction quotient (Q) is a snapshot of a reaction’s progress. It is calculated just like the equilibrium constant (K), but it uses concentration values at a specific point in time, which might not be equilibrium values.
General formula for a reaction:
The expression for Q:
aA + bB ⇌ cC + dD
Q = [C]c [D]d / [A]a [B]b
Comparing Q and K
If left for long enough in a closed system, a reversible reaction will reach equilibrium and the ratio of products to reactants in the system will match the equilibrium constant, K.
By comparing Q to K, we can predict in what direction the reaction will proceed in order to reach equilibrium.
| Comparison | Shift | Direction |
|---|---|---|
| Q < K | Shifts right | Favour products |
| Q > K | Shifts left | Favour reactants |
| Q = K | No shift | System at equilibrium |
Worked Example
Using the following information, calculate the reaction quotient Q and determine the direction in which the reaction will shift to reach equilibrium.
For the reaction:
N2(g) + 3H2(g) ⇌ 2NH3(g)
The equilibrium constant K = 0.50 at a given temperature.
At one point during the reaction, the concentrations are:
[N2] = 0.20 mol dm⁻³
[H2] = 0.60 mol dm⁻³
[NH3] = 0.10 mol dm⁻³
- Write the Q expression:
Q = [NH3]2 / ([N2][H2]3) - Substitute values:
Q = (0.10)2 / (0.20 × 0.603) - Calculate:
Q = 0.01 / (0.20 × 0.216) = 0.01 / 0.0432 ≈ 0.231 - Compare Q to K:
Since Q (0.231) < K (0.50), the reaction will proceed to the right, favouring the formation of more NH3.
Importance of Q
- Q values can help predict the reaction direction before equilibrium is reached.
- Useful in industrial and laboratory settings to determine whether a reaction has progressed far enough.
- Q does not require the system to be at equilibrium — that is the key difference from K.
Summary
- Q is calculated like K but uses non-equilibrium concentrations.
- Comparing Q with K predicts the direction the system will shift.
- Q < K shifts right, Q > K shifts left, Q = K means equilibrium.
- Q is useful for monitoring reaction progress in real time.