PC-06 — 30 Essential NEET Facts — Summary

At dynamic equilibrium, rates of forward and backward reactions are equal; concentrations are constant, not equal.
Kc uses equilibrium molar concentrations; pure solids and liquids are excluded from K expressions.
Kp = Kc(RT)^ $\Delta n$ ; $\Delta n$ counts only gaseous moles (products minus reactants).
Q < K → forward reaction; Q > K → backward reaction; Q = K → at equilibrium.
$\Delta G$ ° = −RT ln K; K > 1 means $\Delta G$ ° < 0 (products favoured thermodynamically).
Only temperature changes K — concentration, pressure, and catalyst changes do not affect K.
Catalyst: no shift in equilibrium position, no change in K; only increases rate of attaining equilibrium.
Inert gas at constant V → no effect on equilibrium; at constant P → shifts toward more gas moles.
Increasing pressure → shifts toward side with fewer gas moles (if $\Delta n$ ≠ 0).
Adding reactant → shifts forward; adding product → shifts backward.
Arrhenius: $H^{+}$ / $OH^{-}$ in water. Bronsted-Lowry: proton donor/acceptor. Lewis: electron-pair acceptor/donor.
Kw = [ $H^{+}$ ][ $OH^{-}$ ] = 10^{-14} at 25°C; pH + pOH = 14 (at 25°C only).
pH = 7 is neutral ONLY at 25°C; at higher T, Kw > 10^{-14} and neutral pH < 7.
Weak acid: [ $H^{+}$ ] = √(Ka · C) when degree of ionisation α ≪ 1.
Ka × Kb = Kw = 10^{-14} for any conjugate acid-base pair at 25°C.
Henderson-Hasselbalch: pH = pKa + log([salt]/[acid]) for acidic buffer.
When [salt] = [acid], pH = pKa — most effective buffer point.
Salt hydrolysis: SA + SB → pH 7; SA + WB → acidic; WA + SB → basic.
Common ion effect: suppresses ionisation of weak electrolyte; reduces solubility of sparingly soluble salt.
Ksp = [cation]^m [anion]^n; precipitation occurs when ionic product Qsp > Ksp.
AgCl: Ksp = $s^{2}$ → s = √Ksp; $Ag_{2}CrO_{4}$ : Ksp = 4 $s^{3}$ → s = (Ksp/4)^(1/3).
AgCl in 0.1 M NaCl: s = Ksp/0.1 = $1.8 \times 10^{-9}$ M (vs $1.34 \times 10^{-5}$ M in pure water).
pH of 10^{-8} M HCl ≈ 6.98 (not 8!) — water's autoionisation contribution cannot be ignored.
The salt with the lower Ksp precipitates first during selective precipitation.
pKa of acetic acid = 4.74; Ka = $1.8 \times 10^{-5}$ .
pH of 0.1 M acetic acid = 2.87 (using [ $H^{+}$ ] = √( $1.8 \times 10^{-6}$ )).
Haber process: $N_{2}$ + $3H_{2}$ ⇌ $2NH_{3}$ is exothermic; high pressure favours $NH_{3}$ ( $\Delta n$ = −2); catalyst = Fe.
Contact process: $2SO_{2}$ + $O_{2}$ ⇌ $2SO_{3}$ ; exothermic; compromise T = ~450°C with $V_{2}O_{5}$ catalyst.
Blood buffer: C $O_{2}$ / $HCO_{3}^{-}$ ; pH = 6.1 + log([ $HCO_{3}^{-}$ ]/[C $O_{2}$ ]) ≈ 7.4.
Kp has units of (atm)^ $\Delta n$ ; Kc has units of (mol/L)^ $\Delta n$ ; both are dimensionless in standard thermodynamic treatment.