QiLing Tools

☯️ICE Table Calculator - Equilibrium Concentrations, Kc to Kp, Le Chatelier

The standard way to find equilibrium concentrations is an ICE table: list the Initial amounts, the Change in terms of the extent x, and the Equilibrium amounts, then substitute into the equilibrium constant expression and solve for x. For H₂(g) + I₂(g) ⇌ 2HI(g) with Kc = 50.5 at 448 °C and 1.0 M of each reactant, the equilibrium amounts are [H₂] = [I₂] = 1.0 − x and [HI] = 2x, so (2x)²/(1 − x)² = 50.5 gives x ≈ 0.780 and [HI] ≈ 1.561 M — about 78% conversion. This calculator balances the reaction, writes the Kc or Kp expression (leaving out pure solids and liquids), computes the reaction quotient Q from the initial amounts to decide the direction of the shift, and then solves for x exactly by bisection on the interval where every amount stays non-negative. It does not rely on the small-x approximation, so it stays accurate even for very large or very small K, but it also reports the approximate x and whether the 5% rule would have held. Kc and Kp are converted with Kp = Kc(RT)^Δn in atm or bar. The Le Chatelier tab uses each species' side, the change in gas moles Δn and the reaction enthalpy ΔH to predict the shift for adding or removing species, compressing or expanding, adding inert gas, heating, cooling or adding a catalyst, and the van 't Hoff equation gives K at another temperature. Textbook presets include the Haber process, HI formation, N₂O₄ dissociation, the water-gas shift, PCl₅ decomposition and SO₂ oxidation.

Free · No sign-up · Runs in your browserLast updated:

Loading tool...

How to use Chemical equilibrium (ICE)

  1. 1Type a reversible reaction (state symbols allowed) or pick a preset, choose Kc or Kp and enter K and the temperature
  2. 2Enter initial concentrations or pressures in the ICE table and read Q, the direction, x, the equilibrium amounts, percent conversion and the 5% rule check
  3. 3Open the Le Chatelier tab to see how each change shifts the equilibrium, or use the van 't Hoff equation to find K at another temperature

FAQ

When can I use the small-x approximation?
When K is small and the reaction barely proceeds, 0.10 − x can be approximated as 0.10. The usual test is the 5% rule: x divided by the initial concentration should be no more than 5%. For H₂ + I₂ ⇌ 2HI with Kc = 50.5 the approximate x is far larger than the initial concentration, so the approximation fails completely. The tool always gives the exact answer and shows the approximate one alongside for comparison.
How do I convert between Kc and Kp?
For gas-phase reactions Kp = Kc(RT)^Δn, where Δn is the moles of gaseous products minus the moles of gaseous reactants, R is 0.08206 L·atm/(mol·K) for pressures in atm or 0.08314 L·bar/(mol·K) for bar, and T is in kelvin. When Δn = 0, as in H₂ + I₂ ⇌ 2HI, Kp equals Kc.
Why are pure solids and liquids left out of the equilibrium expression?
Their activity is taken as 1, so the amount present does not affect the position of equilibrium; they do not appear in K and adding or removing them causes no shift. Setting a species' state to (s) or (l) in the tool excludes it automatically. Water used as the solvent is usually written H₂O(l) for the same reason.
Why does temperature change K while concentration and pressure do not?
The equilibrium constant depends only on temperature. By the van 't Hoff equation, heating an endothermic reaction (ΔH > 0) increases K and shifts it to the right, while heating an exothermic reaction decreases K and shifts it to the left. Changing concentration or pressure only changes Q, so the system shifts until Q equals the unchanged K again; a catalyst only helps it get there faster.