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Schooling · Class 11 · Chemistry · Equilibrium · Chapter on a page

CBSE Class 11 · Chemistry · NCERT Chemistry Part-I

Equilibrium: chapter on a page

A cut of Shishya's notes for this chapter: the first paragraph of the introduction and the key ideas, formulas and facts, common mistakes and quick revision lists, word for word. Nothing is added or reworded; the worked examples are left out (they are on the chapter page).

About the notes: Written by Shishya's AI on 27 Sept 2026 from the chapter's title and class level, in Shishya's own words — not a copy or summary of the textbook. Read the official chapter for the book's own text, activities and exercises.

Fact-checked on 21 Oct 2026: an AI-assisted check of this note against cited sources; the errors it found were corrected that day. Some errors may remain.

Mind map

Drawn from the same notes by a fixed rule (no AI): each label is the item's own words — its bold term, its label before a colon, a short formula as written, or its first words.

Mind map of Equilibrium: Key ideas, Formulas and facts, Worked examples, Watch out forKey ideasDynamic equilibriumEquilibrium constant (Kc and Kp)Relation Kp = Kc (RT)^ΔnReaction quotient QLe Chatelier's principleIonic product of waterpH and pOHFormulas and factsKc = [products]^coefficients /…Kp = Kc (RT)^ΔnpH = −log₁₀[H⁺]Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °CFor a weak acid HA with dissociation…Henderson–Hasselbalch equation for a…Ksp for AgCl ⇌ Ag⁺ + Cl⁻ is Ksp =…Worked examplesCalculating KcFinding pH of a weak acidPredicting precipitationWatch out forForgetting to raise concentrations…Using moles instead of molar…Thinking equilibrium means…Ignoring the sign of Δn when…Assuming pH = 7 means neutral…Equilibrium
The mind map as a list
  • Key ideas
    • Dynamic equilibrium
    • Equilibrium constant (Kc and Kp)
    • Relation Kp = Kc (RT)^Δn
    • Reaction quotient Q
    • Le Chatelier's principle
    • Ionic product of water
    • pH and pOH
  • Formulas and facts
    • Kc = [products]^coefficients /…
    • Kp = Kc (RT)^Δn
    • pH = −log₁₀[H⁺]
    • Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C
    • For a weak acid HA with dissociation…
    • Henderson–Hasselbalch equation for a…
    • Ksp for AgCl ⇌ Ag⁺ + Cl⁻ is Ksp =…
  • Worked examples
    • Calculating Kc
    • Finding pH of a weak acid
    • Predicting precipitation
  • Watch out for
    • Forgetting to raise concentrations…
    • Using moles instead of molar…
    • Thinking equilibrium means…
    • Ignoring the sign of Δn when…
    • Assuming pH = 7 means neutral…

What this chapter is about

This chapter introduces the idea that many chemical reactions do not go to completion. Instead, they reach a state called equilibrium, where the forward and reverse reactions occur at equal rates and the concentrations of reactants and products remain constant over time. You learn how to describe this balance quantitatively using the equilibrium constant, and how factors such as concentration, pressure and temperature shift the position of equilibrium according to Le Chatelier's principle.

Key ideas

  • Dynamic equilibrium: At equilibrium the forward and reverse reactions continue, but their rates are equal, so macroscopic properties (concentration, colour, pressure) stay constant.
  • Equilibrium constant (Kc and Kp): For aA + bB ⇌ cC + dD, Kc = [C]^c [D]^d / [A]^a [B]^b. A large K means products dominate; a small K means reactants dominate.
  • Relation Kp = Kc (RT)^Δn: Here Δn = (moles of gaseous products) − (moles of gaseous reactants), R = 0.0831 bar L mol⁻¹ K⁻¹ (pressures in bar, concentrations in mol L⁻¹), T in kelvin.
  • Reaction quotient Q: Calculated the same way as K but with current (non-equilibrium) concentrations. If Q < K the reaction proceeds forward; if Q > K it proceeds backward.
  • Le Chatelier's principle: When a system at equilibrium is disturbed, it shifts to counteract the change. Adding reactant pushes equilibrium toward products; raising temperature favours the endothermic direction.
  • Ionic product of water: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. Pure water has [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol L⁻¹.
  • pH and pOH: pH = −log₁₀[H⁺]; pOH = −log₁₀[OH⁻]; pH + pOH = 14 at 25 °C.
  • Solubility product (Ksp): For a sparingly soluble salt MX₂ ⇌ M²⁺ + 2 X⁻, Ksp = [M²⁺][X⁻]². If the ionic product exceeds Ksp, precipitation occurs.

Formulas and facts to remember

  • Kc = [products]^coefficients / [reactants]^coefficients (each concentration in mol L⁻¹).
  • Kp = Kc (RT)^Δn, with Δn = sum of gaseous product coefficients minus sum of gaseous reactant coefficients.
  • pH = −log₁₀[H⁺]; pOH = −log₁₀[OH⁻]; pH + pOH = 14 (at 25 °C).
  • Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.
  • For a weak acid HA with dissociation constant Ka, degree of dissociation α ≈ √(Ka / C) when α is small.
  • Henderson–Hasselbalch equation for a buffer: pH = pKa + log₁₀([salt]/[acid]).
  • Ksp for AgCl ⇌ Ag⁺ + Cl⁻ is Ksp = [Ag⁺][Cl⁻]; precipitation begins when ionic product > Ksp.

Common mistakes

  • Forgetting to raise concentrations to the power of their coefficients → always use exponents from the balanced equation.
  • Using moles instead of molar concentrations in the Kc expression → divide by volume in litres first.
  • Thinking equilibrium means concentrations are equal → equilibrium means they are constant, not necessarily equal.
  • Ignoring the sign of Δn when relating Kp and Kc → count only gaseous species and keep track of whether Δn is positive or negative.
  • Assuming pH = 7 means neutral regardless of temperature → pH of pure water changes with temperature because Kw changes.

Quick revision

  • Equilibrium is dynamic: forward rate = reverse rate; concentrations stay constant.
  • K large → products favoured; K small → reactants favoured.
  • Le Chatelier: the system opposes any external change.
  • pH = −log₁₀[H⁺]; neutral water at 25 °C has pH 7.
  • Precipitation occurs when ionic product exceeds Ksp.
  • Buffers resist pH change; use Henderson–Hasselbalch to calculate pH.