Reaction Quotient
Definition and meaning of Reaction Quotient in chemistry.
The reaction quotient (Q) is the ratio of products to reactants at any moment during a reaction. It is calculated using the same form as the equilibrium constant K but with current concentrations instead of equilibrium values.
In more detail
The reaction quotient predicts how a reaction will shift to reach equilibrium. When Q is less than K, the reaction proceeds forward to form more products. When Q is greater than K, the reaction shifts backward toward reactants. When Q equals K, the system is at equilibrium and no net change occurs.
Key facts
| Symbol | Q |
|---|---|
| General form | For aA + bB ⇌ cC + dD, Q = [C]^c[D]^d / [A]^a[B]^b |
| Key principle | When Q does not equal K, the reaction shifts to reach equilibrium |
| Field | Physical Chemistry |
For N2(g) + 3H2(g) ⇌ 2NH3(g), with [N2] = 0.1 M, [H2] = 0.3 M, and [NH3] = 0.5 M, Q = (0.5)² / [(0.1)(0.3)³] ≈ 93. If K = 500 at this temperature, then Q less than K means the reaction shifts forward toward products.
Frequently asked questions
What is the difference between Q and K?
K is the equilibrium constant (fixed for a reaction at a given temperature), while Q is the reaction quotient (varies during the reaction). Both use the same mathematical form, but K uses equilibrium concentrations and Q uses current concentrations. At equilibrium, Q equals K.
How does Q tell you which direction a reaction will shift?
Comparing Q to K reveals the reaction's direction: if Q is less than K, the reaction proceeds forward (toward products); if Q is greater than K, it proceeds backward (toward reactants); if Q equals K, the system is at equilibrium with no net shift.