Salt Chemistry Calculator - pH and Classification

When a salt dissolves in water, its ions may react with water (hydrolyze) to shift the pH away from 7. Whether a salt is acidic, basic, or neutral depends on the strength of the parent acid and base.

Classification rules:

• Strong acid + strong base: neutral salt. No hydrolysis. pH = 7 (e.g. NaCl, KNO3).
• Strong acid + weak base: acidic salt. The cation hydrolyzes: BH+ + H2O -> B + H3O+. pH < 7 (e.g. NH4Cl, pH about 5).
• Weak acid + strong base: basic salt. The anion hydrolyzes: A- + H2O -> HA + OH-. pH > 7 (e.g. NaCH3COO, sodium acetate, pH about 9).
• Weak acid + weak base: depends on relative Ka and Kb. If Ka > Kb, slightly acidic; if Kb > Ka, slightly basic.

For a basic salt (weak acid + strong base), pH estimate: Ka of the weak acid is known. Kb of conjugate base = Kw / Ka (where Kw = 10^-14). For dilute solution: [OH-] ≈ sqrt(Kb * C), where C is concentration. pH = 14 + log([OH-]) = 14 + 0.5*(log(Kb) + log(C))

For an acidic salt (strong acid + weak base): Ka of conjugate acid = Kw / Kb (base) [H+] ≈ sqrt(Ka * C) pH = 0.5*(pKa - log C) = -0.5*(log(Ka) + log(C))

Those square-root formulas assume the hydrolysis barely touches the concentration, which is only true when Ka or Kb is far below C. When it is not, the approximation overshoots, sometimes badly. Take sodium cyanide at 0.001 mol/L: Ka for HCN is 6.2 x 10^-10, so Kb of the conjugate base is 1.6 x 10^-5, and the square root gives [OH-] = 1.27 x 10^-4. That is more than a tenth of everything you dissolved, so treating C as undisturbed is not defensible.

So this calculator solves the quadratic instead:

x^2 + Kx - KC = 0, giving x = (-K + sqrt(K^2 + 4KC)) / 2

That is exact for the hydrolysis equilibrium at any concentration, and it collapses to the square-root answer whenever the approximation was valid anyway. The result panel reports both numbers along with the gap between them, so you can see when your textbook shortcut would have been fine and when it would not.

One limit worth knowing: below about 10^-6 mol/L, water’s own ionization starts contributing more hydroxide than the salt does, and neither formula here accounts for that. Very dilute solutions of any salt trend toward pH 7 for that reason.