AP Chemistry Strong Acids and Bases
The lists to memorize, plus pH calculations for acids, bases, and strong acid-base mixtures.
Strong acids and bases are the one place in acid-base chemistry where you skip the equilibrium work: they react completely, so the concentration you start with tells you the pH. That only helps if you can tell strong from weak on sight, so this page puts the lists first and then runs four calculations, one of which ends in a mixture of an acid and a base. For the surrounding Unit 8 material, see the Unit 8 review.
The Strong Acid and Strong Base Lists to Memorize
These are not given on the exam, so they have to come from memory.
| Strong acids (6) | Strong bases | Common weak ones mistaken for strong |
|---|---|---|
| HCl, HBr, HI HNO3 HClO4 H2SO4 |
Group 1 hydroxides: LiOH, NaOH, KOH, RbOH, CsOH Group 2: Ca(OH)2, Sr(OH)2, Ba(OH)2 |
HF (unlike HCl, HBr, HI) CH3COOH H3PO4 NH3 (a weak base) |
Sulfuric acid needs one caution: only its first ionization is complete, and the second is partial. Check how your course treats it in pH problems; the examples here avoid the question. Anything not on the strong list, including every carboxylic acid and every nitrogen base such as ammonia, is weak and needs an ICE table instead of a shortcut.
How to Find the pH of a Strong Acid or Base
- Find the ion concentration from the formula. Complete dissociation means [H3O+] or [OH−] equals the starting concentration times the number of H+ or OH− in the formula.
- Take the negative logarithm: pH = −log[H3O+] or pOH = −log[OH−].
- Convert if needed with pH + pOH = 14.00 at 25°C.
For pH, the number of digits after the decimal point equals the number of significant figures in the concentration, a rule explained in the significant figures guide.
Worked Example 1: pH of a Strong Acid (HNO3)
What is the pH of 3.0 × 10−3 M HNO3?
HNO3 is on the strong list and gives one H+ per formula unit, so
[H3O+] = 3.0 × 10−3 M.
pH = −log(3.0 × 10−3) = 2.52.
As a check, pOH = 14.00 − 2.52 = 11.48, so [OH−] = 3.3 × 10−12 M, tiny, as
expected for an acid.
Worked Example 2: pH of a Strong Base That Releases Two OH− (Ba(OH)2)
What is the pH of 0.015 M Ba(OH)2?
Ba(OH)2 → Ba2+ + 2 OH−, so
[OH−] = 2 × 0.015 = 0.030 M.
pOH = −log(0.030) = 1.52.
pH = 14.00 − 1.52 = 12.48.
The mistake to avoid is using 0.015 M as [OH−], which would give 12.18.
Worked Example 3: Finding a Concentration From a Given pH
A NaOH solution has pH 12.40. What is its concentration?
pOH = 14.00 − 12.40 = 1.60.
[OH−] = 10−1.60 = 0.025 M.
NaOH releases one OH− per formula unit, so [NaOH] = 0.025 M.
The same steps run in reverse for an acid: [H3O+] = 10−pH. If the base had been Ba(OH)2, you would then divide by 2. To start from a concentrated stock instead, use the dilution calculations guide first.
Worked Example 4: pH After Mixing a Strong Acid With a Smaller Amount of Strong Base
50.0 mL of 0.100 M HCl is mixed with 30.0 mL of 0.100 M NaOH. What is the pH of the mixture?
Moles of each: H+ = (0.0500 L)(0.100 M) = 5.00 × 10−3 mol;
OH− = (0.0300 L)(0.100 M) = 3.00 × 10−3 mol.
React them 1:1: H+ + OH− → H2O. The acid is in excess by
5.00 − 3.00 = 2.00 × 10−3 mol H+.
New volume: 50.0 + 30.0 = 80.0 mL = 0.0800 L, so
[H+] = 2.00 × 10−3 / 0.0800 = 0.0250 M.
pH = −log(0.0250) = 1.60.
Be careful to divide by the total volume, not by either starting volume. This is the strong-acid, strong-base case of the neutralization that the net ionic equation H+ + OH− → H2O describes, and the same bookkeeping drives the strong-acid portions of a titration curve.
Common Strong Acid and Base Mistakes
- Calling HF strong because it is a hydrogen halide. HCl, HBr, and HI are strong; HF is weak.
- Forgetting the second hydroxide in Ba(OH)2, Ca(OH)2, or Sr(OH)2. Multiply the concentration by 2 before taking the log.
- Using pH = −log[OH−]. That gives pOH. Subtract from 14.00 to reach pH.
- Dividing by the wrong volume after mixing. Use the total volume of the combined solution.
- Setting up an ICE table for a strong acid. It adds work and no accuracy; complete ionization makes [H3O+] equal the starting concentration.
Whether a species is a strong acid also decides what its conjugate base does, which is covered in Bronsted-Lowry acids, bases, and conjugate pairs. When your practice problems are done, you can check where the rest of your preparation stands using the AP Chem Score Calculator.
Related Resources
- AP Chem Score Calculator
- Unit 8 Review: Acids and Bases
- AP Chemistry Bronsted-Lowry Acids and Bases
- AP Chemistry Titration Curve pH Calculations
- AP Chemistry Henderson-Hasselbalch Equation
- How to Write Net Ionic Equations
- AP Chemistry Dilution Calculations
- AP Chemistry Molarity
- AP Chemistry Study Guide
Frequently Asked Questions
What are the strong acids in AP Chemistry?
The six to memorize are hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI), nitric acid (HNO3), perchloric acid (HClO4), and sulfuric acid (H2SO4). Each ionizes essentially completely in water.
What are the strong bases?
The hydroxides of group 1 (LiOH, NaOH, KOH, RbOH, CsOH) and the group 2 hydroxides Ca(OH)2, Sr(OH)2, and Ba(OH)2 dissociate completely when they dissolve, releasing OH- ions.
Is HF a strong acid?
No. Although the other hydrogen halides (HCl, HBr, HI) are strong, hydrofluoric acid (HF) is a weak acid: it ionizes only slightly in water, so you need its Ka and an equilibrium calculation to find its pH.
Why does Ba(OH)2 give twice the hydroxide concentration?
Each formula unit of Ba(OH)2 contains two hydroxide ions, and it dissociates completely, so [OH-] is twice the concentration of the dissolved base. A 0.015 M solution has [OH-] = 0.030 M.
Do you use an ICE table for a strong acid?
No. A strong acid ionizes completely, so [H3O+] equals the acid's starting concentration and you can take the negative log directly. ICE tables are for weak acids and bases, where only some of the molecules react.
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