Deep-Dive Guides · Ionic Equilibria
Ionic Equilibria & Buffers — the complete deep dive
One guide dense enough to take you from "pH = −log[H⁺]" to CSIR-NET Part-C fluency. The organising idea: aqueous ionic equilibria are just equilibrium constants applied to protons and ions in water, held together by three conservation laws — mass balance, charge balance, and the water autoionisation Kw.
Free extract — the opening section of the guide, in full.
0. How to use this guide — the master map ▲ top
Every pH or solubility problem is solved by imposing the same three conservation laws on the relevant equilibrium constants. Fix the map first:
| Situation | Governing relation | Fast result |
|---|---|---|
| Strong acid/base | full dissociation | pH = −log cacid (or 14 + log cbase) |
| Weak acid HA | Ka = [H⁺][A⁻]/[HA] | [H⁺] ≈ √(Kac) |
| Weak base B | Kb = [BH⁺][OH⁻]/[B] | [OH⁻] ≈ √(Kbc) |
| Buffer (HA + A⁻) | Henderson–Hasselbalch | pH = pKa + log([A⁻]/[HA]) |
| Salt of weak acid | hydrolysis Kh = Kw/Ka | basic; pH > 7 |
| Sparingly soluble salt | Ksp = ∏[ion](coefficient) | s from Ksp; common-ion suppresses |
| Always true (the relation; not the value) | Kw = [H⁺][OH⁻], which is 1.0×10⁻¹⁴ at 25 °C only | pH + pOH = pKw = 14.00 at 25 °C |
Read the rest of this guide
The remaining 9 sections — Water, pH and weak acids/bases, Buffers, Acid–base titrations & indicators, Solubility equilibria & Ksp and more — are part of ChemVidya Full Access, along with the other guides, the books, the study notes and 6,000+ practice questions with worked solutions.