Acids, Bases, and Buffers

Acids and bases are central to chemistry and physiology. Buffers maintain stable pH, which is critical for biological function.

Acid and Base Definitions

Acids and bases can be defined in several ways. The two most useful introductory definitions are the Arrhenius definition and the Brønsted–Lowry definition.

An Arrhenius acid increases the concentration of hydrogen ions, H⁺, in an aqueous solution. Because free H⁺ ions do not normally remain alone in water, they associate with water molecules to form hydronium ions, H₃O⁺. HCl is an Arrhenius acid because it increases the hydronium-ion concentration.
An Arrhenius base increases the concentration of hydroxide ions, OH⁻, in an aqueous solution. NaOH is an Arrhenius base because it increases the hydroxide-ion concentration.
The Arrhenius example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l). HCl increases hydrogen-ion concentration and NaOH increases hydroxide-ion concentration in water, forming water and salt. The Arrhenius definition applies mainly to aqueous solutions, requires acids to increase H⁺ or H₃O⁺, requires bases to increase OH⁻, and does not describe every acid–base reaction.
A Brønsted–Lowry acid is a proton donor. An acid donates H⁺.
A Brønsted–Lowry base is a proton acceptor. Base accepts H⁺.
Brønsted–Lowry example: Hydrochloric acid and water. HCl donates H⁺ and acts as the acid. H₂O accepts H⁺ and acts as the base. Cl⁻ is the conjugate base of HCl. H₃O⁺ is the conjugate acid of H₂O.

Hydronium Ions Although acid reactions are often written using H⁺, hydrogen ions do not normally exist independently in water. Instead, H⁺ associates with a water molecule: H⁺ + H₂O → H₃O⁺. Therefore, H⁺ and H₃O⁺ are often used interchangeably in introductory acid–base chemistry.

Conjugate Acid–Base Pairs When an acid donates H⁺, it becomes its conjugate base. When a base accepts H⁺, it becomes its conjugate acid. Members of a conjugate acid–base pair differ by one proton (H⁺).

Example: HCl + H₂O → H₃O⁺ + Cl⁻

HCl → acid
Cl⁻ → conjugate base
H₂O → base
H₃O⁺ → conjugate acid

Conjugate pairs: HCl/Cl⁻ and H₂O/H₃O⁺

Amphiprotic Substances An amphiprotic substance can act as either an acid or a base. Water is the most common example.

Water acting as a base: HCl + H₂O → H₃O⁺ + Cl⁻ (water accepts H⁺).
Water acting as an acid: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (water donates H⁺).

Strong and Weak Acids Acid strength describes how readily an acid donates a proton.

Strong Acids: Strong acids ionize almost completely in water. Examples include: HCl, HBr, HI, HNO₃, HClO₄, and H₂SO₄. They produce relatively high concentrations of H₃O⁺ and their equilibrium lies strongly toward the products.
Weak Acids: Weak acids ionize only partially in water. Examples include: HF, acetic acid (CH₃COOH), and carbonic acid (H₂CO₃). Weak-acid reactions establish a dynamic equilibrium containing both ionized and nonionized forms.

Strong and Weak Bases

Strong Bases: Strong bases dissociate almost completely in water. Examples include: NaOH, KOH, and Ca(OH)₂.
Weak Bases: Weak bases accept protons only partially. Examples include ammonia (NH₃), methylamine (CH₃NH₂), and many nitrogen-containing organic bases.

Acid Ionization Constant: Ka Ka measures the extent to which an acid ionizes in water.

Large Ka = Greater ionization, products favored, stronger acid.
Small Ka = Less ionization, reactants favored, weaker acid.

Base Ionization Constant: Kb Kb measures the extent to which a base accepts a proton from water.

Larger Kb means a stronger base.
Smaller Kb means a weaker base.

pKa and pKb

pKa = −log(Ka)

pKb = −log(Kb)

Lower pKa → Stronger acid; the proton is donated more readily.
Higher pKa → Weaker acid.
Lower pKb → Stronger base.
Higher pKb → Weaker base.

For a conjugate acid–base pair at 25°C: Ka × Kb = Kw. Since Kw = 1.0 × 10⁻¹⁴, the logarithmic relationship is: pKa + pKb = 14. These equations allow the strength of a conjugate acid to be calculated from the strength of its conjugate base, or vice versa.

pH pH is a logarithmic measure of hydrogen-ion or hydronium-ion concentration. pH = −log[H₃O⁺]

Because the pH scale is logarithmic, a change of 1 pH unit represents a 10-fold change in H₃O⁺ concentration.

Example: pH 5 → pH 6 = 10 times lower [H₃O⁺]

At approximately 25°C:

pH below 7: acidic
pH equal to 7: neutral
pH above 7: basic or alkaline

Pure water has a hydronium concentration of approximately 1.0 × 10⁻⁷ M and a pH of 7 at 25°C.

pOH pOH measures hydroxide-ion concentration. pOH = −log[OH⁻] At 25°C: pH + pOH = 14.

Buffers A buffer is a solution that resists large changes in pH when small amounts of acid or base are added. Buffers do not prevent pH changes completely; they minimize changes in pH when relatively small amounts of acid or base are added.

A typical buffer contains:

A weak acid and its conjugate base, or
A weak base and its conjugate acid.

Henderson–Hasselbalch Equation The Henderson–Hasselbalch equation estimates the pH of a buffer: pH = pKa + log([A⁻]/[HA]) Where [A⁻] is the conjugate-base concentration and [HA] is the weak-acid concentration. When [A⁻] = [HA], the logarithmic term equals zero, so pH = pKa.

Blood pH Normal arterial blood pH is maintained within a narrow range: 7.35–7.45. pH below 7.35 is associated with acidemia. pH above 7.45 is associated with alkalemia. The bicarbonate system, lungs, and kidneys work together to regulate blood pH. The lungs regulate carbon dioxide, while the kidneys regulate bicarbonate and hydrogen-ion handling.

General Chemistry

Acid and Base Definitions

Strong and Weak Bases

pKa and pKb

Key Points