Concentration at equilibrium

An ICE Table can be used to find the concentrations of all aqueous and gaseous reactants and products when a chemical reaction achieves equilibrium. It is a method of organizing stoichiometric calculations, and its letters stand for the following:

All three are measured in mol/L, and they are related by I `+` C `=` E.

Example

In the reaction H2(g) + I2(g) ⇌ 2 HI(g), 2.00 mol of H2(g) and 3.00 mol of I2(g) are placed in a 1.00 L container. Calculate the other two equilibrium concentrations if I2(g) has an equilibrium concentration of 1.30 mol/L.

Here is the I calculation for H2(g) (you don’t need to show all of them):

`c = n/V = (2.00\ "mol")/(1.00\ "L") = 2.00\ "mol/L"`.

Let `x` represent the absolute value of the change in concentration of H2(g). This can also be written more concisely like this: let `x = |Delta ["H"_2]|`.

When writing the let statement for `x`, always choose a reactant or product that has a coefficient of 1. This way, you can simply fill in all the C values with `-ax` for reactants and `+ax` for products, where `a` is the coefficient for that reactant or product.

H2(g) I2(g) 2 HI(g)
I `2.00` `3.00` `0`
C `-x` `-x` `+2x`
E `2.00 − x` `3.00 − x` `2x`

The E value for I2(g) is known to be 1.30 mol/L, but our table tells us that it is also `x` subtracted from 3.00 mol/L, therefore we can set them equal:

`1.30\ "mol/L" = 3.00\ "mol/L" - x`,

`x = 1.70\ "mol/L"`.

By substituting 1.70 mol/L for `x` into the E expressions for H2(g) and HI(g), we can easily find their concentrations at equilibrium as well (0.30 mol/L and 3.40 mol/L respectively).