Strong and Weak Acids and Bases

Overview

Acid or base strength describes the position of an ionization equilibrium. A strong species ionizes essentially completely in water at ordinary introductory-chemistry concentrations. A weak species ionizes only partly and remains in equilibrium with its conjugate partner. Strength is not the same property as concentration: a dilute solution of HCl can be strongly ionized, while a concentrated solution of acetic acid can still contain mostly un-ionized CH₃COOH.

Learning objectives

By the end of this lesson, you should be able to compare strength and concentration, identify the dominant species in strong and weak solutions, use equilibrium constants to compare strength, and solve representative pH questions without applying an unchecked approximation.

Prerequisites

• Review the pH and pOH scale, including logarithms and the relationship between pH and hydrogen-ion concentration.
• An Arrhenius acid increases H⁺ (more precisely H₃O⁺) in water; a Brønsted–Lowry acid donates a proton and a Brønsted–Lowry base accepts one.
• Equilibrium means the forward and reverse processes continue at equal rates; see the chemical equilibrium topic.

Strength is not concentration

Property	What it tells you	Example question
Strength	Fraction of dissolved molecules that ionize and the equilibrium position	Does HA mostly become H3O+ + A−?
Concentration	Amount of solute per volume of solution	How many moles of HA are present per litre?

Strong does not mean concentrated, and weak does not mean dilute. For example, 0.0010 M HCl is dilute but is treated as fully dissociated in an introductory calculation. A 1.0 M CH₃COOH solution is concentrated relative to a 0.0010 M solution, but acetic acid remains a weak acid and most molecules are still CH₃COOH. The pH also depends on both properties: changing concentration changes the amount of H⁺, while changing strength changes the fraction that is produced.

Definitions and ionization equations

• A strong acid ionizes nearly completely in water. HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq) is written with a one-way arrow for the introductory treatment.
• A weak acid ionizes partially and establishes equilibrium: CH₃COOH(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CH₃COO⁻(aq).
• A strong base dissociates essentially completely when it dissolves, and its soluble portion supplies OH⁻. NaOH(aq) → Na⁺(aq) + OH⁻(aq).
• A weak base reacts only partly with water: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq).

“Complete” and “near-complete” are useful calculation models, not a claim that a real solution has literally zero molecules of the reverse species. The solvent, temperature, activity effects, and course convention can affect how a strength list is defined.

Comparing strong and weak species

Feature	Strong acid/base	Weak acid/base
Degree of ionization	Nearly complete for the dissolved strong species	Partial; a measurable fraction remains molecular
Equilibrium position	Far toward ions in water	Noticeably toward reactants for a weak acid/base
Electrolyte behavior	Strong electrolyte; high ion concentration at a given formal concentration	Weak electrolyte; fewer ions at the same formal concentration
pH at equal formal concentration	More extreme acidic or basic pH	Closer to neutral because fewer ions are formed
Representative examples	HCl, HBr, HI, HNO3, HClO4; the first ionization of H2SO4; soluble Group 1 hydroxides	CH3COOH, HF, H2CO3, NH3, and amines

The common strong-acid list depends on solvent and course convention. Strong bases include soluble Group 1 hydroxides and the dissolved portion of heavier Group 2 hydroxides; “soluble” matters because a sparingly dissolving solid does not create a high OH⁻ concentration simply by being labelled strong.

Ka, Kb, pKa, and pKb

For a weak acid HA in water,

$K_a = \frac{[\mathrm{H_3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$

For a weak base B,

$K_b = \frac{[\mathrm{BH^+}][\mathrm{OH^-}]}{[\mathrm{B}]}$

Larger K_a means a stronger acid; larger K_b means a stronger base. Logarithmic forms make comparisons easier:

$pK_a = -\log K_a \qquad pK_b = -\log K_b$

Thus, a smaller pK_a means a stronger acid, and a smaller pK_b means a stronger base. For a conjugate acid–base pair at a stated temperature,

$K_aK_b = K_w$

At 25 °C, pK_w ≈ 14.00, so pK_a + pK_b ≈ 14.00. Do not use 14.00 automatically at another temperature; K_w and pK_w change with temperature.

Worked examples

1. pH of a monoprotic strong acid

Find the pH of 0.0100 M HNO₃. Treat HNO₃ as fully ionized in this introductory model:

1. Stoichiometry gives [H³O⁺] ≈ 0.0100 M.
2. pH = −log(0.0100) = 2.00.

The concentration is dilute, but the acid is strong. Those statements are independent.

2. pH of a strong base

Find the pH of 0.0250 M Ba(OH)₂, assuming the dissolved portion dissociates completely:

1. Ba(OH)₂ → Ba²⁺ + 2OH⁻, so [OH⁻] = 2(0.0250) = 0.0500 M.
2. pOH = −log(0.0500) = 1.301.
3. At 25 °C, pH = 14.000 − 1.301 = 12.699 ≈ 12.70.

The factor of 2 is stoichiometry before the logarithm.

3. Weak monoprotic acid and the small-x check

Find the approximate pH of 0.100 M CH₃COOH when K_a = 1.8 × 10⁻⁵.

For CH₃COOH ⇌ H⁺ + CH₃COO⁻, let x be the amount ionized:

	CH3COOH	H+	CH3COO−
Initial	0.100	0	0
Change	−x	+x	+x
Equilibrium	0.100 − x	x	x

K_a = x²/(0.100 − x). The small-x approximation gives x ≈ √(K_aC) = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³ M. Therefore pH ≈ −log(1.34 × 10⁻³) = 2.87.

Check the approximation: (x/C) × 100% = (0.00134/0.100) × 100% = 1.34%, below the common 5% limit, so the approximation is valid. If the percentage is too large, solve the quadratic or use an exact equilibrium method.

4. Comparing strength from Ka or pKa

Acid A has K_a = 1.0 × 10⁻³ (pK_a = 3.00); acid B has K_a = 1.0 × 10⁻⁵ (pK_a = 5.00). Acid A is stronger because its K_a is larger. The two-unit pK_a difference corresponds to a 100-fold difference in K_a.

Common mistakes

• Treating weak-acid ionization as complete and setting [H⁺] equal to the formal acid concentration.
• Confusing the amount or concentration of solute with its strength.
• Ignoring stoichiometry before equilibrium when mixing solutions or using a polyprotic/ionic formula.
• Applying the small-x approximation without checking whether x is less than about 5% of the initial concentration.
• Using pH + pOH = 14 without stating the usual 25 °C context.
• Comparing acids by pH without controlling for concentration; K_a or pK_a is the strength comparison.

Decision guide

1. Identify whether the acid/base is strong or weak using the course convention.
2. Perform any stoichiometric reaction first: dissociation, neutralization, dilution, or mixing.
3. For a strong species, use the resulting ion concentration directly when the approximation is appropriate.
4. For a weak species, write the equilibrium expression and use an ICE table.
5. Check the approximation, units, significant figures, and whether temperature changes pK_w.

Practice

Try each question before opening the answer.

1. Is 0.0010 M HCl strong, concentrated, both, or neither?
2. Write the equilibrium equation for HF in water and identify the conjugate base.
3. Calculate the pH of 0.0020 M NaOH.
4. For a 0.050 M weak acid with K_a = 1.0 × 10⁻⁶, estimate x and check the 5% rule.
5. Which is stronger: an acid with pK_a = 2.4 or one with pK_a = 4.8? By what factor do their K_a values differ?
6. A solution is made by mixing strong acid with a weak acid. What must be done before using the weak-acid equilibrium expression?

Show final answers

1. Strong and dilute; concentration and strength are different properties.
2. HF + H₂O ⇌ H₃O⁺ + F⁻; F⁻ is the conjugate base.
3. pOH = 2.699, so pH = 11.301 ≈ 11.30 at 25 °C.
4. x ≈ √(K_aC) = √(5.0 × 10⁻⁸) = 2.24 × 10⁻⁴ M; 0.448% ionized, so the approximation passes.
5. The pK_a = 2.4 acid is stronger; ΔpK_a = 2.4 means K_a is about 10^2.4 ≈ 250 times larger.
6. Complete the stoichiometric reaction first, update concentrations and volume, then set up the weak-acid equilibrium.

Related resources

Use the pH and pOH lesson, equilibrium constant formula, pH calculator, and weak acid/base calculator for follow-up practice.