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Generated on: July 20, 2026 at 09:33 PM

Reaction Quotient Calculator

Inputs

Concentration of A0.1 M
Stoichiometric Coefficient a1
Concentration of B0.2 M
Stoichiometric Coefficient b1
Concentration of C0.5 M
Stoichiometric Coefficient c1
Concentration of D1 M
Stoichiometric Coefficient d0
Equilibrium Constant K10

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Chemistry

Reaction Quotient Calculator

Calculate the reaction quotient Q from current concentrations and compare it to the equilibrium constant K to predict which way the reaction aA + bB ⇌ cC + dD will shift.

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\begin{aligned} Q &= \dfrac{[\text{C}]^{c}\,[\text{D}]^{d}}{[\text{A}]^{a}\,[\text{B}]^{b}} \\ &= \dfrac{ 0.5^{1}\,1^{0} }{ 0.1^{1}\,0.2^{1} } \\ &= ? \end{aligned}

Since Q < K, the reaction shifts forward (to the right), consuming reactants and forming more products until Q rises to equal K.

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What the reaction quotient tells you

The reaction quotient Q measures the relative amounts of products and reactants present at any moment during a reversible reaction. It is calculated with the same expression as the equilibrium constant K, but from the concentrations that exist right now rather than the concentrations at equilibrium. Comparing Q to K reveals whether a reaction still has room to move forward, has overshot toward products, or has already settled at equilibrium.

The reaction quotient expression

For the general reaction:

aA+bB⇌cC+dDa\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D}

the reaction quotient in terms of molar concentrations is:

Qc=[C]c[D]d[A]a[B]bQ_c = \frac{[\text{C}]^c [\text{D}]^d}{[\text{A}]^a [\text{B}]^b}

Square brackets denote current molar concentrations (mol/L), and the exponents are the stoichiometric coefficients from the balanced equation. As with K, pure solids and pure liquids are omitted because their concentrations are effectively constant.

Comparing Q to K

Once Q is known, a single comparison with K predicts the direction of the shift:

ComparisonDirection of shiftMeaning
Q < KForward (to the right)Too few products; the reaction forms more product
Q = KNo net shiftThe system is already at equilibrium
Q > KReverse (to the left)Too many products; the reaction re-forms reactants

This is a direct expression of Le Chatelier's principle: the system always moves in the direction that brings Q closer to K.

Worked example

Consider the reaction A + B ⇌ C with an equilibrium constant K = 10. Suppose the current concentrations are:

[A]=0.10 M,[B]=0.20 M,[C]=0.50 M[\text{A}] = 0.10\ \text{M}, \quad [\text{B}] = 0.20\ \text{M}, \quad [\text{C}] = 0.50\ \text{M}

All coefficients equal 1, so:

Q=[C][A][B]=0.50(0.10)(0.20)=0.500.020=25\begin{aligned} Q &= \frac{[\text{C}]}{[\text{A}][\text{B}]} \\[6pt] &= \frac{0.50}{(0.10)(0.20)} \\[6pt] &= \frac{0.50}{0.020} \\[6pt] &= 25 \end{aligned}

Because Q = 25 is greater than K = 10, there is more product than equilibrium allows. The reaction shifts in reverse, consuming C and re-forming A and B until Q falls to 10.

Excluding a species: setting a coefficient to zero

Any species with a stoichiometric coefficient of 0 contributes a factor of [X]⁰ = 1, which drops out of the expression. This makes it easy to handle reactions with fewer than two reactants or products — set the unused coefficient to 0, and its concentration no longer matters.

Q, K, and Gibbs free energy

The comparison between Q and K also connects to the thermodynamics of the reaction. The relationship

ΔG=ΔG∘+RTln⁡Q\Delta G = \Delta G^\circ + RT \ln Q

shows that when Q < K the free-energy change ΔG is negative and the forward reaction is spontaneous; when Q > K, ΔG is positive and the reverse reaction is favored; and when Q = K, ΔG = 0 and the system is at equilibrium. The reaction quotient is therefore the bridge between measured concentrations and the driving force of a reaction.

Common pitfalls

Q must be evaluated with concentrations expressed in the same units used to define K — typically mol/L for Qc, or partial pressures for Qp. Mixing units, forgetting to raise each concentration to its stoichiometric power, or including pure solids and liquids are the most frequent mistakes. Because Q depends on the current state and K depends only on temperature, a reaction with a fixed K can have Q above, below, or equal to K depending entirely on the concentrations present at that instant.

Frequently Asked Questions (FAQ)

What is the formula for the reaction quotient Q?

For the general reaction aA + bB ⇌ cC + dD, the reaction quotient is Q = ([C]^c · [D]^d) / ([A]^a · [B]^b), where the square brackets denote the current molar concentrations and the exponents are the stoichiometric coefficients from the balanced equation. It has exactly the same form as the equilibrium constant expression, but the concentrations are measured at the moment of interest rather than at equilibrium.

What is the difference between Q and K?

Q and K share the same algebraic expression, but K uses equilibrium concentrations while Q uses concentrations at any arbitrary point in time. K is fixed for a given reaction at a given temperature; Q changes continuously as the reaction proceeds and equals K only when the system reaches equilibrium.

How does comparing Q to K predict the direction of a reaction?

If Q < K, there are relatively too few products, so the reaction proceeds forward (to the right) to form more products. If Q > K, there are relatively too many products, so the reaction proceeds in reverse (to the left) to re-form reactants. If Q = K, the forward and reverse rates are already balanced and the system is at equilibrium.

What does it mean when Q equals K?

When Q = K, the system is at equilibrium: the ratio of product to reactant concentrations already matches the equilibrium constant, so there is no net change in either direction. The forward and reverse reactions continue at equal rates, keeping the concentrations constant.

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