Cue Column | Notes Column
| Question / Keyword | Detail / Answer |
|---|---|
| What is dynamic equilibrium? | Rate of forward reaction = rate of backward reaction; concentrations are constant but NOT equal and NOT zero |
| Write Kc for aA + bB ⇌ cC + dD | — use equilibrium concentrations only |
| How is Kp related to Kc? | where = (moles gaseous products) − (moles gaseous reactants) |
| What is Q and how to use it? | Q = same form as Kc but with non-equilibrium concentrations; Q < K → forward; Q > K → backward; Q = K → at equilibrium |
| How does temperature affect K? | Only temperature changes K; increase T → K increases for endothermic, decreases for exothermic |
| Does a catalyst change K? | No — catalyst speeds up both forward and backward reactions equally; no shift, no change in K |
| Inert gas at constant V vs constant P? | Constant V → no effect; Constant P → shifts toward more moles of gas (volume expands, concentrations drop) |
| Henderson-Hasselbalch equation | ; pOH = pKb + log([salt]/[base]) |
| Common ion effect | Adding a common ion suppresses ionization of weak electrolyte; reduces solubility of sparingly soluble salt |
| When does precipitation occur? | When ionic product (Qsp) > Ksp; if Qsp = Ksp → saturated; Qsp < Ksp → unsaturated |
Summary
Equilibrium (chemical and ionic) is one of NEET's highest-yield Physical Chemistry topics. The equilibrium constant Kc is defined by concentration ratios at equilibrium; Kp = Kc(RT)^ for gaseous reactions. Le Chatelier's principle predicts shifts in response to stress — only temperature changes the value of K itself. Ionic equilibrium covers acid-base theories (Arrhenius, Bronsted-Lowry, Lewis), pH calculations via Ka (weak acids), Henderson-Hasselbalch (buffers), salt hydrolysis, and Ksp (solubility product with common ion effect). The most frequently tested traps: catalyst does not shift equilibrium; inert gas at constant V has no effect; ultra-dilute acid pH is not simply −log[acid].