How to Use an ICE Table: Step-by-Step Guide With Examples

An ICE table is a three-row setup for solving equilibrium problems. The rows are INITIAL concentrations, the CHANGE as the reaction proceeds, and EQUILIBRIUM concentrations. You write the change in terms of x, using the coefficients from the balanced equation. Then you plug the equilibrium row into the K expression and solve for x. Almost every equilibrium problem in AP Chemistry and General Chemistry, including every weak acid pH problem, uses this same setup.

Watch 2 free Unit 15 lessons →

If equilibrium problems feel like a different puzzle every time, the ICE table is the fix. It turns every problem into the same four moves.

What does ICE stand for?

• I: Initial. The concentrations before the reaction shifts. These come from the problem.
• C: Change. How much each concentration goes up or down on the way to equilibrium. You write these in terms of x.
• E: Equilibrium. Initial plus change. These are the only numbers that go into the K expression.

The ICE table is just bookkeeping. The chemistry you need is already in the balanced equation. The table keeps you from losing track of it.

What are the steps for any ICE table problem?

1. Write the balanced equation and the K expression. Leave out solids and pure liquids.
2. Fill in the I row with starting molarities (or partial pressures, if you’re working with K_p).
3. Fill in the C row using the coefficients as multipliers on x. Reactants lose (−), products gain (+). If you’re not sure which direction the reaction goes, calculate Q first. More on that below.
4. Fill in the E row by adding I and C.
5. Substitute the E row into the K expression and solve for x.
6. Answer the actual question. This is critical. The most common trap is solving for x without knowing what x represents. Calculate the concentrations or pH the problem asked for, not just x.

Worked example 1: Finding equilibrium concentrations

At a certain temperature, K_c = 64 for the reaction below. If 1.00 M H₂ and 1.00 M I₂ are placed in a container, what are the equilibrium concentrations?

H₂(g) + I₂(g) ⇌ 2HI(g)

K_c = [HI]² / ([H₂][I₂])

	H₂	I₂	HI
I	1.00	1.00	0
C	−x	−x	+2x
E	1.00 − x	1.00 − x	2x

Notice the +2x under HI. That coefficient of 2 in the balanced equation becomes the multiplier. This is the most common setup mistake in the whole topic.

Substitute:

64 = (2x)² / (1.00 − x)²

Both sides are perfect squares, so take the square root:

8 = 2x / (1.00 − x)
8 − 8x = 2x
x = 0.80

Now answer the question:

• [H₂] = [I₂] = 1.00 − 0.80 = 0.20 M
• [HI] = 2(0.80) = 1.60 M

Check: (1.60)² / (0.20 × 0.20) = 2.56 / 0.040 = 64 ✓

Worked example 2: Finding K from equilibrium values

0.500 M N₂O₄ is placed in a flask. At equilibrium, [NO₂] = 0.200 M. Find K_c.

N₂O₄(g) ⇌ 2NO₂(g)

	N₂O₄	NO₂
I	0.500	0
C	−x	+2x
E	0.500 − x	2x = 0.200

Here the problem hands you one equilibrium value, so you solve for x directly. Because the coefficient on NO₂ is 2, the equation is 2x = 0.200, so x = 0.100.

Then [N₂O₄] = 0.500 − 0.100 = 0.400 M, and:

K_c = (0.200)² / 0.400 = 0.100

This problem type trips students up because they try to plug 0.200 into the [Initial] row. Read carefully: does the number describe the start or the end?

Worked example 3: Finding the pH of a weak acid

Find the pH of 0.10 M acetic acid. K_a = 1.8 × 10⁻⁵.

CH₃COOH(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CH₃COO⁻(aq)

Water is a pure liquid, so it stays out of the expression.

	CH₃COOH	H₃O⁺	CH₃COO⁻
I	0.10	≈0	0
C	−x	+x	+x
E	0.10 − x	x	x

1.8 × 10⁻⁵ = x² / (0.10 − x)

Solving this exactly would take the quadratic formula. But K_a is tiny, which means very little acid ionizes, so x is small compared to 0.10. Since this reaction barely proceeds, we can treat x as negligible next to 0.10 and drop it from the denominator:

1.8 × 10⁻⁵ ≈ x² / 0.10
x² = 1.8 × 10⁻⁶
x = 1.34 × 10⁻³ M = [H₃O⁺]

pH = −log(1.34 × 10⁻³) = 2.87

Check the approximation with the 5% rule: (1.34 × 10⁻³ / 0.10) × 100 = 1.3%. That’s under 5%, so the shortcut is valid. If it came out over 5%, you’d go back and use the quadratic formula.

This is why the ICE table matters beyond the equilibrium unit. Weak acids, weak bases, buffers, and K_sp problems all use this same setup.

What if you don’t know which direction the reaction goes?

If the problem gives you starting amounts of both reactants and products, you can’t assume the reaction runs forward. Calculate Q, the reaction quotient, using the initial concentrations in the K expression:

• Q < K: too few products. The reaction shifts forward. Reactants get −x, products +x.
• Q > K: too many products. The reaction shifts in reverse. Flip the signs.
• Q = K: already at equilibrium. No table needed.

What are the most common ICE table mistakes?

1. Ignoring coefficients in the C row. If the equation says 2HI, the change is 2x, and that term gets squared in the K expression.
2. Including solids or pure liquids. They never appear in the table’s math or in K.
3. Using moles instead of molarity. If the problem gives moles and a volume, divide first.
4. Plugging initial values into K. Only the E row goes into the K expression.
5. Stopping at x. The question almost never asks for x. It asks for a concentration, a pH, or a percent.

Frequently asked questions

What is an ICE table used for?

An ICE table organizes the Initial, Change, and Equilibrium concentrations of every species in a reaction so you can solve for an unknown equilibrium concentration, an equilibrium constant, or a pH.

Do solids and liquids go in an ICE table?

No. Pure solids and pure liquids, including water as a solvent, are left out of the K expression, so they are left out of the ICE table math too.

Do you use moles or molarity in an ICE table?

Use molarity for K_c problems (or partial pressures for K_p). If a problem gives you moles and a volume, divide to get molarity before you fill in the table.

When can you ignore x in an ICE table?

When K is very small, x is usually tiny compared to the starting concentration, so you can drop it from the denominator. Check with the 5% rule: if x divided by the initial concentration is under 5%, the shortcut is valid. If not, use the quadratic formula.

How do you know which direction the reaction shifts?

Calculate the reaction quotient Q from the starting concentrations. If Q is less than K, the reaction shifts forward. If Q is greater than K, it shifts in reverse. If Q equals K, the system is already at equilibrium.

Want to see it worked out step by step?

The ICE table is built step by step in Unit 15: Equilibrium, with worked problems for K, Q, and Le Chatelier’s principle. Weak acid ICE tables continue in Unit 16: Acid/Base Equilibrium and Buffers. The first two lessons of every unit are free. For the “which way does it shift” side of equilibrium, read Le Chatelier’s Principle Explained with Real Examples, and if you are still shaky on the mole ratios behind the C row, start with What Is Stoichiometry?