How this instrument works
The equilibrium constant, Kc, is the ratio of product concentrations to reactant concentrations at chemical equilibrium, with each concentration raised to the power of its coefficient in the balanced equation. This relationship — the law of mass action, first proposed by Guldberg and Waage in 1864 — holds that for any reversible reaction at a given temperature, this particular ratio always settles to the same value once the forward and reverse reaction rates become equal, no matter what concentrations the reaction started from.
A large Kc (much greater than 1) means the equilibrium favors products — at equilibrium, most of the reactants have converted forward. A small Kc (much less than 1) means the equilibrium favors reactants — the reaction barely proceeds forward before settling. Kc is specific to a given temperature; changing the temperature generally changes Kc itself, sometimes substantially, which is why equilibrium constants are always reported alongside the temperature they were measured or calculated at.
It's worth being precise about what makes this Kc specifically, rather than a related but different quantity that shares the same formula shape. Kc uses concentrations (in molarity) and, critically, requires the system to actually be at equilibrium — plug in concentrations from a reaction still in progress and you get the reaction quotient, Qc, not Kc, even though the arithmetic looks identical. A separate constant, Kp, uses partial pressures of gases instead of concentrations and is related to Kc but numerically different in general. This instrument computes Kc specifically: concentration-based, and assumed to be evaluated at true equilibrium.
- Enter Reactant 1 concentration (M) and Reactant 1 coefficient — the equilibrium molar concentration of your first reactant and its coefficient in the balanced equation.
- If there's a second reactant, enter Reactant 2 concentration (M, 1 if none) and Reactant 2 coefficient (0 if none); otherwise leave them at 1 and 0 so it doesn't affect the ratio.
- Enter Product 1 concentration (M) and Product 1 coefficient the same way for your first product.
- If there's a second product, enter Product 2 concentration (M, 1 if none) and Product 2 coefficient (0 if none), following the same 'leave at 1 and 0' pattern if there isn't one.
- Read Equilibrium constant Kc directly — this assumes all concentrations entered are equilibrium values for the same reaction at the same temperature.
Worked example — a squared product over a single reactant
A reaction at equilibrium has one product at a concentration of 0.5 M with a coefficient of 2 (so it appears in the equation as 2P), and one reactant at a concentration of 0.2 M with a coefficient of 1. Enter 0.5 into Product 1 concentration (M), 2 into Product 1 coefficient, 1 into Product 2 concentration (M, 1 if none), 0 into Product 2 coefficient (0 if none), 0.2 into Reactant 1 concentration (M), 1 into Reactant 1 coefficient, and 1 and 0 into the second reactant's fields. Equilibrium constant Kc reads 1.25.
That comes from squaring the product concentration first — 0.5² = 0.25 — then dividing by the reactant concentration raised to its own coefficient — 0.2¹ = 0.2 — giving 0.25 / 0.2 = 1.25. A Kc slightly above 1 like this indicates the equilibrium sits only modestly toward the product side, not overwhelmingly in either direction.
Questions
What's the difference between Kc and Qc, the reaction quotient?
They use the exact same formula — products over reactants, each raised to its coefficient — but Kc specifically requires the concentrations to be measured at true equilibrium, while Qc can be calculated from concentrations at any point in the reaction, equilibrium or not. Comparing Qc to Kc at any moment tells you which direction the reaction still needs to shift: if Qc is less than Kc, the reaction proceeds forward; if Qc is greater than Kc, it proceeds backward. This instrument computes Kc, so it assumes the concentrations entered are already at equilibrium.
What's the difference between Kc and Kp?
Kc is built from molar concentrations of species in solution or gas phase, while Kp is built the same way but from partial pressures of gaseous species instead. For a reaction involving gases, the two are related through the ideal gas law, but they are numerically different constants in general — you cannot substitute a Kp value where a Kc value is needed, or vice versa, without converting between them first.
What does a Kc much greater than 1 tell me about a reaction?
It means the equilibrium lies far toward the products — at equilibrium, the reaction has converted the large majority of reactants into products, leaving relatively little reactant concentration behind. Conversely, a Kc much less than 1 means the reverse: the equilibrium favors reactants, and only a small fraction has converted forward, no matter how the reaction was initially set up.
Why do I need to enter '1' for a concentration field I'm not using, rather than 0?
Because a concentration of 1 raised to a coefficient of 0 (also entered as 0 in that case) evaluates to 1, which multiplies into the ratio without changing it — exactly the behavior you want for a species that isn't part of your reaction. Entering 0 for an unused concentration would instead force the whole numerator or denominator to zero, which would break the calculation rather than neutrally ignore the unused term.
Does Kc change if I change the concentrations I start the reaction with?
No — Kc is a constant for a given reaction at a given temperature, and the whole point of the law of mass action is that whatever starting concentrations you begin with, the system settles at equilibrium into concentrations that satisfy this same ratio. What does change with starting concentrations is how far the reaction has to shift, and in which direction, to reach that fixed Kc — not the value of Kc itself.
Why does temperature matter for Kc if it's supposedly 'constant'?
Kc is constant at a fixed temperature, but changing the temperature generally changes Kc's value, sometimes substantially, because temperature affects the relative rates of the forward and reverse reactions differently. That's why any reported equilibrium constant is only meaningful alongside the temperature it applies to — a Kc measured at 25°C doesn't necessarily hold at 100°C for the same reaction.