pH & Ionic Equilibrium — Chemistry Mnemonic
Target: mnemonic for pH pOH buffer solution ionic equilibrium
Why is this hard to memorize?
Ionic equilibrium covers acids, bases, pH, buffers, solubility product, and common ion effect. NEET and JEE test: pH = −log[H⁺], pOH = −log[OH⁻], pH + pOH = 14 (at 25°C), Henderson-Hasselbalch equation for buffers (pH = pKa + log[salt/acid]), Ksp calculations, and the common ion effect on solubility. The pH scale runs from 0 (strongly acidic) to 14 (strongly basic), with 7 being neutral. Most exam errors come from log calculation mistakes and forgetting that Kw = 10⁻¹⁴ only at 25°C.
Classic mnemonics you should know
"pH = −log[H⁺]. pOH = −log[OH⁻]. pH + pOH = 14. Kw = [H⁺][OH⁻] = 10⁻¹⁴ (at 25°C)"
These four equations solve 80% of pH problems. pH < 7 = acidic, pH = 7 = neutral, pH > 7 = basic. Remember: lower pH = more acidic (more H⁺ ions). Each pH unit = 10× difference in [H⁺].
"Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]) — "pH = pKa + log(salt/acid)""
For acidic buffer (weak acid + salt): pH = pKa + log([conjugate base]/[weak acid]). When [A⁻] = [HA], pH = pKa (maximum buffer capacity). For basic buffer: pOH = pKb + log([salt]/[base]).
"Strong acids (HCl, H₂SO₄, HNO₃): pH = −log(C). Weak acids: pH = ½(pKa − logC) — "use Ka for weak!""
Strong acids dissociate completely: [H⁺] = concentration. Weak acids partially dissociate: use Ka and ICE table, or the shortcut pH = ½(pKa − logC) when C >> Ka.
The complete list
- pH = −log[H⁺]
- pOH = −log[OH⁻]
- pH + pOH = 14 (at 25°C)
- Kw = [H⁺][OH⁻] = 10⁻¹⁴
- Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA])
- Strong acid pH = −log(C)
- Weak acid pH ≈ ½(pKa − log C)
- Ksp = product of ion concentrations
Frequently asked questions
How do I calculate pH of a weak acid?
For weak acid HA with concentration C and dissociation constant Ka: Set up ICE table. At equilibrium: [H⁺] = [A⁻] = x, [HA] = C−x. Ka = x²/(C−x). If Ka << C (common case), approximate C−x ≈ C: x = √(Ka×C). Then pH = −log(x). Shortcut: pH = ½(pKa − log C).
What is a buffer solution and how does it work?
A buffer resists pH changes when small amounts of acid or base are added. Acidic buffer = weak acid + its conjugate base (e.g., CH₃COOH + CH₃COONa). Adding H⁺: the conjugate base neutralizes it (A⁻ + H⁺ → HA). Adding OH⁻: the weak acid neutralizes it (HA + OH⁻ → A⁻ + H₂O). pH stays nearly constant.
How do I use Ksp to predict precipitation?
Calculate the ionic product (Q) from actual ion concentrations. Compare Q with Ksp: If Q > Ksp → precipitation occurs (supersaturated). If Q < Ksp → no precipitation (unsaturated). If Q = Ksp → saturated (at equilibrium). Example: Will PbCl₂ precipitate if [Pb²⁺] = 0.01M and [Cl⁻] = 0.1M? Q = (0.01)(0.1)² = 10⁻⁴. If Ksp = 1.7×10⁻⁵, Q > Ksp → yes, precipitation occurs.
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