How this instrument works
The reaction quotient Q is calculated from the law-of-mass-action expression — products raised to their stoichiometric coefficients, divided by reactants raised to theirs — using whatever concentrations happen to be present at the instant you evaluate it. That is the entire difference between Q and the equilibrium constant K: K uses this identical formula but only describes the system once it has actually reached equilibrium, while Q can be computed at the very start of a reaction, partway through, or at equilibrium itself.
That flexibility is what makes Q genuinely useful rather than just a renamed K. Comparing a freshly computed Q against a known K tells you which direction a reaction still needs to shift to reach equilibrium: if Q is less than K, there aren't yet enough products relative to reactants, so the reaction proceeds forward, consuming more reactants; if Q is greater than K, there are more products than the equilibrium ratio allows, so the reaction runs in reverse; and if Q equals K, the system is already at equilibrium and there is no net shift left to happen.
This site also offers separate equilibrium constant (Kc) and Kp calculators, which share this exact mathematical form but describe genuinely different situations, not just relabeled versions of the same number. Kc and Kp report the fixed ratio a system settles into at equilibrium, computed from concentrations or partial pressures respectively, and that value doesn't change unless temperature changes. Q reports the same style of ratio but at whatever moment you choose to check it — often not at equilibrium at all — which is precisely what makes Q the tool for predicting reaction direction rather than describing equilibrium itself.
- Enter the concentration of your first product into Product 1 concentration (M) and its coefficient from the balanced equation into Product 1 coefficient.
- For a second product, use Product 2 concentration (M, 1 if none) and Product 2 coefficient (0 if none); otherwise leave them at 1 and 0 so that term drops out.
- Enter the concentration of your first reactant into Reactant 1 concentration (M) and its coefficient into Reactant 1 coefficient.
- For a second reactant, use Reactant 2 concentration (M, 1 if none) and Reactant 2 coefficient (0 if none); otherwise leave them at 1 and 0.
- Read Reaction quotient Q beneath the fields, then compare it to your reaction's known K: Q < K means the reaction still runs forward, Q > K means it runs in reverse, and Q = K means it's already at equilibrium.
Worked example — one product squared over one reactant
Enter 0.5 into Product 1 concentration (M) with coefficient 2, leave Product 2 concentration (M, 1 if none) at 1 with coefficient 0 (no second product), enter 0.2 into Reactant 1 concentration (M) with coefficient 1, and leave Reactant 2 concentration (M, 1 if none) at 1 with coefficient 0 (no second reactant). Reaction quotient Q reads 1.25.
That comes from Q = 0.5^2 / 0.2^1 = 0.25 / 0.2 = 1.25. If this reaction's known equilibrium constant Kc happens to also equal 1.25 at this temperature, the snapshot represents a system already at equilibrium; if Kc were instead, say, 3, a Q of only 1.25 would mean the reaction still has further to run forward, building up more product before Q rises to meet K.
Questions
What's the actual difference between Q and K?
The formula is identical — products raised to their coefficients over reactants raised to theirs. The difference is entirely about when you're allowed to evaluate it: K is only meaningful once the system has reached equilibrium, where it becomes a fixed number for that temperature, while Q can be calculated from concentrations at any point in the reaction, before, during, or at equilibrium. Every K is technically a Q, but not every Q equals K.
How do I use Q to predict which way a reaction will shift?
Compare your calculated Q against the reaction's known equilibrium constant K at that temperature. If Q < K, the reaction proceeds forward, converting more reactants into products, until Q rises to meet K. If Q > K, the reaction runs in reverse, converting products back into reactants, until Q falls to meet K. If Q = K already, the system is at equilibrium and there's no further net shift.
How is Q different from Kp?
Kp uses the same law-of-mass-action structure as Q and Kc, but it's built from partial pressures of gaseous species rather than molar concentrations, and like Kc it only describes a system genuinely at equilibrium. Q as calculated here uses concentrations and can be evaluated at any point in a reaction — a gas-phase reaction actually has a Q version built from partial pressures too, evaluated the same way Kp is, just not necessarily at equilibrium.
Why do unused product or reactant slots default to a concentration of 1 and coefficient of 0?
Because raising a concentration to the power of 0 always equals 1, which makes that entire factor vanish from the multiplication without needing a separate 'ignore this term' option. Setting the coefficient to 0 for a species that doesn't appear in your reaction, with any placeholder concentration, cleanly removes it from the calculation.
Does Q need every concentration to be at equilibrium to be valid?
No — that's precisely the point of Q as distinct from K. Q is defined for whatever concentrations are actually present at the moment you check, including right after mixing reactants with no product formed yet, or partway through a reaction. Its usefulness comes specifically from being computable outside of equilibrium, so it can be compared against K to see how far from equilibrium the system currently sits.