K vs Q

Core Concept

K is the equilibrium constant and describes the ratio of product concentrations to reactant concentrations at equilibrium, with each raised to the power of their stoichiometric coefficients.

Q is the reaction quotient and it is a snapshot of the ratio of product concentrations to reactant concentrations at any point during the reaction, not necessarily at equilibrium.

Key Tips

Using the Wrong Expression: Ensure the Q and K expressions match the balanced equation.
Ignoring States of Matter: Exclude solids and liquids from the Q and K expressions.
Confusing Kc and Kp: Use concentrations for Kc and partial pressures for Kp.

Test Yourself

Assorted Multiple Choice

For the reversible reaction $N_2(g) + 3 H_2(g) \rightleftharpoons 2 NH_3(g)$, the equilibrium constant $K_c$ is $5.0 \times 10^2$ at a certain temperature. A chemist measures the concentrations in a reaction vessel and calculates the reaction quotient, $Q_c$, to be $1.2 \times 10^3$. What will happen to the system to achieve equilibrium?

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Reaction Quotient vs Equilibrium Constant (K vs Q)

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Calculating Q and Determining System Status

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Predicting the Direction of Net Change

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Interpreting the "Distance" from Equilibrium

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Assorted Multiple Choice

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K, Equilibrium Constant

K describes the ratio of product concentrations to reactant concentrations at equilibrium, with each raised to the power of their stoichiometric coefficients.

Expression: $K = \frac{[\text{products}]^{\text{coefficients}}}{[\text{reactants}]^{\text{coefficients}}}$
Key Points:
- K is calculated using equilibrium concentrations or pressures.
- K is constant at a given temperature.

Q, the Reaction Quotient

Definition: Q is a snapshot of the ratio of product concentrations to reactant concentrations at any point during the reaction, not necessarily at equilibrium.
Expression: $Q = \frac{[\text{products}]^{\text{coefficients}}}{[\text{reactants}]^{\text{coefficients}}}$
Key Points:
- Q is calculated using the current (non-equilibrium) concentrations or pressures.
- Q changes as the reaction progresses toward equilibrium.

Comparing K and Q

The relationship between Q and K determines the direction in which the reaction will proceed to reach equilibrium:

  
            Relationship
            Interpretation
            Direction of Reaction
        
            Q < K
            Too few products; too many reactants
            Reaction shifts right (toward products)
        
            Q = K
            Reaction is at equilibrium
            No shift; reaction is balanced
        
            Q > K
            Too many products; too few reactants
            Reaction shifts left (toward reactants)

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