Summary of "CHEMICAL EQUILIBRIUM in 1 Shot || All Concepts & PYQs Covered || Prachand NEET"

Core ideas and definitions

Chemical equilibrium (dynamic equilibrium): a stage in a reversible reaction (usually in a closed system) where measurable macroscopic properties (concentration, pressure, etc.) stop changing with time even though forward and backward reactions continue. Net rate of reaction = 0 because rate(forward) = rate(backward).

• “Saturation” is an informal term sometimes used to describe reaching equilibrium; do not confuse this with reactions stopping — at equilibrium reactions continue microscopically (dynamic).
• Reaction quotient Q vs equilibrium constant K:
  • Q < K → reaction shifts forward (more product formed).
  • Q = K → system at equilibrium.
  • Q > K → reaction shifts backward (more reactant formed).
• Size of K:
  • K >> 1 (e.g., > 10^3) → products favored at equilibrium.
  • K << 1 → reactants favored.
  • K ~ 1 → significant amounts of both reactants and products.

Types / classifications of equilibrium

• By change type:
  • Physical equilibrium: phase changes (solid ↔ liquid, liquid ↔ vapor, sublimation, solubility equilibria). No chemical change at molecular level.
  • Chemical equilibrium: chemical conversion (A ⇌ B); includes ionic equilibria (acid–base, dissociation).
• By phases present:
  • Homogeneous equilibrium: reactants and products all in the same phase (e.g., all gases).
  • Heterogeneous equilibrium: more than one phase (e.g., CaCO3(s) ⇌ CaO(s) + CO2(g)).

Active mass / activities (practical rules)

• For gases: active mass ≈ partial pressure (use pressure in Kp expressions).
• For solutions: active mass ≈ concentration (use molarity in Kc expressions).
• For pure solids and pure liquids: activity ≈ 1 (omit these from equilibrium expressions).
• Practical exam convention: if volume is not given in mole problems, you can assume 1 L so moles ≈ molarity to simplify numerics.

Law of Mass Action and equilibrium-constant expressions

• Rate laws (mass action): rate ∝ product of active masses of reactants raised to their stoichiometric coefficients.
• For a general reaction aA + bB ⇌ cC + dD:
  • Kc = [C]^c [D]^d / ([A]^a [B]^b) (concentrations)
  • Kp = (PC)^c (PD)^d / (PA)^a (PB)^b (partial pressures)
  • Kx or K (mole-fraction based) can be written using mole fractions.
• Relating Kp and Kc:
  • Kp = Kc (RT)^{Δn}, where Δn = (moles gas products) − (moles gas reactants). Use R in consistent units and T in K.
  • Units of Kp or Kc depend on Δn—derive units carefully when requested.
• Thermodynamic relation:
  • ΔG° = −RT ln K (connects standard free energy and equilibrium constant).
  • van ’t Hoff (temperature dependence): ln(K2/K1) = −ΔH°/R (1/T2 − 1/T1) (assumes ΔH° constant).
    • If ΔH° > 0 (endothermic): K increases with temperature.
    • If ΔH° < 0 (exothermic): K decreases with temperature.
• Manipulating equilibrium constants:
  • Reverse a reaction → new K = 1 / K(original).
  • Multiply reaction by n → new K = [K(original)]^n.
  • Add reactions → overall K = product of individual Ks.

Le Chatelier’s principle (qualitative predictor of shifts)

If a stress is applied to a system at equilibrium, the system shifts to partially oppose the stress and establish a new equilibrium.

Common stresses and effects:

• Concentration changes:
  • Increasing concentration of a reactant → equilibrium shifts to consume that reactant (toward products).
  • Increasing concentration of a product → shifts toward reactants.
• Temperature changes:
  • Treat temperature as “heat”: for a forward endothermic reaction (heat on reactant side), increasing T favors the forward direction; for exothermic forward reaction, increasing T favors reverse.
  • Temperature changes alter K (not just the position); increasing T favors the side that absorbs heat.
• Pressure/volume changes (gaseous systems only):
  • Increasing pressure (decreasing volume) shifts toward the side with fewer moles of gas.
  • Decreasing pressure (increasing volume) shifts toward the side with more moles of gas.
  • If Δn(gases) = 0, pressure/volume changes have no effect on equilibrium position.
• Addition of an inert (non‑reactive) gas:
  • At constant volume: no effect on equilibrium position (partial pressures of reactants/products unchanged).
  • At constant pressure: adding inert gas increases volume → can shift equilibrium (similar to volume increase).
• Catalysts:
  • Catalysts increase rates of both forward and backward reactions equally; they speed attainment of equilibrium but do not change K or the equilibrium composition.

Practical tip: to predict direction, consider which side of the equation will reduce the imposed change.

Reaction quotient Q and direction prediction

• Compute Q the same way as K but using current (instantaneous) pressures/concentrations.
• Compare Q to K to predict spontaneous shift:
  • Q < K → forward
  • Q = K → at equilibrium
  • Q > K → backward

Graphical illustrations (qualitative)

• Concentration vs time: reactant concentration falls, product concentration rises, both become constant at equilibrium.
• Rate vs time: forward and reverse rates approach each other and become equal at equilibrium.

Worked‑problem methodology / ICE approach

General ICE (Initial — Change — Equilibrium) steps:

1. I (initial): write initial concentrations or pressures.
2. C (change): express changes using stoichiometric multiples (e.g., −2x, +3x).
3. E (equilibrium): write equilibrium expressions (e.g., [A] = A0 − 2x).
4. Substitute equilibrium expressions into the K expression and solve for x.

Common formats and notes:

• Initial concentrations given + K given → use ICE with concentrations and Kc.
• Initial pressures given + Kp given → same ICE method using partial pressures (or convert from mole fractions & Ptotal).
• Initial moles given + Ptotal/Kp given → convert moles → mole fractions xi = ni / n_total, then Pi = xi · Ptotal and use Kp.
• If volume not provided, it’s common to assume V = 1 L to convert moles → molarity for Kc problems.
• If solids/pure liquids are present, omit their activity (set = 1) in the K expression.
• Use Dalton’s law when needed: Pi = (ni / n_total) · P_total.

Conversions and care:

• Stoichiometric coefficients are the exponents in K expressions.
• When reversing/multiplying/adding reactions, adjust K appropriately (reciprocal, power, product).
• For Kp, be consistent with units of pressure (atm commonly used). Useful conversions: 1 atm = 760 mmHg = 1.013×10^5 Pa; 1 bar = 10^5 Pa.

Thermodynamics and K

• ΔG° = −RT ln K relates standard free-energy change and K.
  • If ΔG° = 0 → K = 1.
• Use van ’t Hoff equation for how K changes with temperature:
  • ln(K2/K1) = −ΔH°/R (1/T2 − 1/T1) (use for numerical estimates when ΔH° is approximately constant).

Practical exam tips

• Memorize concise NCERT definitions/lines for exam answers.
• In numericals:
  • Remember common unit conversions (atm, mmHg, Pa).
  • If initial moles are given without volume, assume 1 L unless instructed otherwise.
  • Identify the unknown x and set up equations with correct stoichiometric factors; omit pure solids/liquids from K.
• For multi‑step reaction combinations: overall K = product of individual Ks.
• Practice using calculators during study (note exam rules may differ).

Common example systems (for intuition)

• N2 + 3H2 ⇌ 2NH3 (illustrates Δn and Kp/Kc relation)
• PCl5 ⇌ PCl3 + Cl2 (illustrates gaseous equilibria and pressure/volume effects)
• Solid ⇌ liquid at the melting point (rate of melting = rate of crystallization)
• Liquid ⇌ vapor at the boiling point (vaporization rate = condensation rate)
• Dissolution/crystallization (saturation: rate of dissolution = rate of crystallization; solubility varies with temperature and, for gases, with pressure)

Problems and practice focus areas

• Writing correct K expressions (products raised to stoichiometric powers over reactants); omit solids/pure liquids.
• Predicting direction using Q vs K.
• ICE table practice for:
  • Initial concentrations
  • Initial pressures
  • Initial mole numbers & total pressure
  • K given in different forms (Kc, Kp, etc.)
• Algebraic manipulation of K for reversed/multiplied/combined reactions.
• Numerical use of van ’t Hoff relation to relate K at different temperatures.
• Units of K: derive from Δn (difference in gaseous moles).

Key formulas (concise)

• Kc = [C]^c [D]^d / ([A]^a [B]^b)
• Kp = (PC)^c (PD)^d / (PA)^a (PB)^b)
• Kp = Kc (RT)^{Δn}, Δn = (gas moles products) − (gas moles reactants)
• ΔG° = −RT ln K
• van ’t Hoff: ln(K2/K1) = −ΔH°/R (1/T2 − 1/T1)
• Q defined the same as K but using instantaneous concentrations/pressures

Speakers / sources referenced

• Main lecturer/instructor: Prachand (NEET / classroom teacher — video presenter).
• NCERT (text referenced for definitions and standard statements).
• Law of Mass Action: Guldberg & Waage.
• Le Chatelier: Le Chatelier’s principle.
• Dalton: Dalton’s law of partial pressures.
• Henry: Henry’s law for gas solubility in liquids.
• Thermodynamics relations: standard textbook thermodynamics.
• (Various cultural analogies used by the lecturer — e.g., Dhoni, movie references — not scientific sources.)

Optional supplementary materials you might prepare

• One‑page ICE table template and a worked example (step‑by‑step) for each common problem type (initial concentrations, initial pressures, initial moles).
• A short set of representative NCERT/NEET practice problems with full solutions.

Category

Educational

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