Question #5b9c3

An ICE Table (you'll sometimes see this referred to as a RICE Table) is a notation system used to help keep track of how the concentrations of the species involved in an equilibrium rection change during the course of a reaction.

The ICE initials come from

• I - Initial
• C - Change
• E - Equilibrium

This notation system uses the initial concentrations of the species, the way these concentrations change in accordance to their stoichiometric coefficients, and the concentrations you can expect once equilibrium is established.

So, for a generic equilibrium reaction

#A + color(red)(2)B rightleftharpoons C + color(blue)(3)D#

you start with reactants, #[A]_0# and #[B]_0#, and want to know what the equilibrium concentrations of all the species will be.

Since the initial concentrations of the products will be zero, the first line of the ICE table will look like this

#" "A" " + " "color(red)(2)B " "rightleftharpoons" " C + " "color(blue)(3)D#
I....#[A]_0#............#[B]_0#.................0...............0

Now look for how thse concentrations will change once the reaction proceeds. Notice that 1 mole of #A# needs #color(red)(2)# moles of #B# in order for the reaction to take place.

Moreover, 1 mole of #A# will form 1 mole of #C# and #color(blue)(3)# moles of #D#.

This means that, if the concentration of #A# will drop by an unknown value #x#, the rest of the concentrations will change in accordance to those mole ratios.

The second line of the table will be

#" "A" " + " "color(red)(2)B " "rightleftharpoons" " C + " "color(blue)(3)D#
I....#[A]_0#............#[B]_0#.................0...............0
C...(-x)...............(-#color(red)(2)#x)................(+x)..........(+#color(blue)(3)#x)

The final line of the table is simply the sum of the above two lines, i.e. what remains after the change takes place.

#" "A" " + " "color(red)(2)B " "rightleftharpoons" " C + " "color(blue)(3)D#
I....#[A]_0#............#[B]_0#.................0...............0
C...(-x)...............(-#color(red)(2)#x)................(+x)..........(+#color(blue)(3)#x)
E..#[A]_0#-x........#[B]_0#-#color(red)(2)#x...............x.............#color(blue)(3)#x

You'd then go on to use the respective reaction's equilibrium constant, #K_(eq)#, and solve for the equilibrium concentrations of the species.