Acid-Base Chemistry Calculators

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Acid-base chemistry is the study of proton transfer reactions between chemical species. In the Brønsted-Lowry definition, an acid is a proton donor and a base is a proton acceptor. The strength of acids and bases is quantified by Ka, Kb, and pKa values, and the balance of acids and bases in solution determines pH. Acid-base reactions are fundamental in biochemistry (enzyme catalysis, blood buffering), analytical chemistry (titrations, indicators), environmental science (acid rain, ocean acidification), and industrial chemistry.

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Definitions of Acids and Bases

  • Arrhenius: Acid produces H⁺ in water; base produces OH⁻
  • Brønsted-Lowry: Acid = proton donor; base = proton acceptor. Conjugate pairs differ by one H⁺.
  • Lewis: Acid = electron pair acceptor; base = electron pair donor. Most general definition.

pH, pOH, and Kw

pH = −log₁₀[H⁺]; pOH = −log₁₀[OH⁻]
At 25°C: pH + pOH = 14 (Kw = 10⁻¹⁴)
Acidic: pH < 7; neutral: pH 7; basic: pH > 7

Strong vs. Weak Acids/Bases

  • Strong acids (HCl, HNO₃, H₂SO₄, HBr, HI, HClO₄): fully dissociate. [H⁺] = C.
  • Weak acids: Partial dissociation. [H⁺] ≈ √(Ka × C). pH = ½(pKa − log C)
  • Strong bases (NaOH, KOH): fully dissociate
  • Weak bases (NH₃): Kb governs extent of ionization

Buffer Systems

Buffer = weak acid + conjugate base. Henderson-Hasselbalch:
pH = pKa + log([A⁻]/[HA])
Max buffer capacity at pH = pKa. Effective range: pKa ± 1 pH unit.

Biological Buffers

  • Bicarbonate (pKa 6.1): blood CO₂/HCO₃⁻ system — maintains blood pH 7.4
  • Phosphate (pKa 7.2): intracellular buffer
  • Histidine imidazole (pKa ~6): enzyme active site buffer

Acid-Base Titrations

Equivalence point: moles H⁺ = moles OH⁻. pH at equivalence: strong-strong → 7; weak acid-strong base → > 7; strong acid-weak base → < 7.

Glossary

Brønsted-Lowry Acid
A substance that donates a proton (H⁺) in a chemical reaction. Its conjugate base differs by loss of one H⁺. Contrasts with Lewis acid (electron pair acceptor).
Buffer
A solution containing a weak acid and its conjugate base that resists pH changes. Effective within ±1 pH unit of pKa. Described by Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]).
Equivalence Point
The point in a titration where moles of acid equal moles of base. pH at equivalence: 7 (strong-strong), >7 (weak acid-strong base), <7 (strong acid-weak base). Detected by indicator or pH meter.

Frequently Asked Questions

Brønsted-Lowry: acid = proton (H⁺) donor; base = proton acceptor. Applies to reactions involving H⁺ transfer. Lewis: acid = electron pair acceptor; base = electron pair donor. More general — includes reactions without proton transfer (e.g., BF₃ + NH₃). All Brønsted-Lowry acids/bases are also Lewis acids/bases, but not vice versa. Brønsted-Lowry is most used in biochemistry; Lewis concepts are important in coordination chemistry and enzyme metal-center reactions.

For strong acids (HCl, HNO₃, HBr, H₂SO₄, HI, HClO₄, HClO₃), assume complete dissociation: [H⁺] = acid concentration. pH = −log₁₀[H⁺]. Example: 0.01 M HCl → [H⁺] = 0.01 M → pH = −log(0.01) = 2. At concentrations below ~10⁻⁷ M, the contribution from water's self-ionization becomes significant and must be included.

A buffer resists pH changes when acid or base is added. It contains a weak acid (HA) and its conjugate base (A⁻). When H⁺ added: A⁻ + H⁺ → HA (absorbs proton). When OH⁻ added: HA + OH⁻ → A⁻ + H₂O (donates proton). pH = pKa + log([A⁻]/[HA]). Max capacity at pH = pKa; effective range pKa ± 1 unit. Buffer capacity increases with higher concentration of the weak acid/conjugate base pair.

At the equivalence point, moles of acid exactly equal moles of base. pH depends on what is formed: strong acid + strong base → pH 7 (water and neutral salt); weak acid + strong base → pH > 7 (conjugate base A⁻ hydrolyzes: A⁻ + H₂O → HA + OH⁻); strong acid + weak base → pH < 7 (conjugate acid BH⁺ hydrolyzes: BH⁺ → B + H⁺). The indicator must change color near the equivalence point pH for accurate endpoint detection.