Equilibrium Notes for NEET and JEE Main 2027: Key Points and PDF

Equilibrium is Unit 6 of NCERT Class 11 Chemistry Part I and Unit 6 of both the NEET UG 2026 and JEE Main 2026 chemistry syllabi. These notes cover physical and chemical equilibrium, Kc and Kp, reaction quotient, Gibbs energy, Le Chatelier’s principle, acids and bases, pH, Ka and Kb, salt hydrolysis, buffers and Ksp, with a free PDF.

Equilibrium Notes for NEET and JEE Main

Section-wise key points from NCERT Unit 6, formula tables, high-yield facts and common mistakes for NEET and JEE Main 2027 revision.

Equilibrium at a Glance

The topic list follows the NEET UG 2026 syllabus released by the National Medical Commission (NMC) on 22 December 2025 and the JEE Main 2026 syllabus published by NTA. These stay the reference until the 2027 syllabi are out. Full lists: NEET syllabus and JEE Main syllabus.

What is dynamic equilibrium?

A system is in equilibrium when the rates of the forward and reverse processes become equal, so the concentrations (or other measurable properties) stop changing. The equilibrium is dynamic: both processes keep going at the same rate.

• Equilibrium can be set up for physical processes (phase changes, dissolution) and for chemical reactions, and can be reached from either direction.
• The mixture of reactants and products at equilibrium is called the equilibrium mixture.
• NCERT sorts reactions into three groups: those that go nearly to completion, those that form only small amounts of product, and those where reactant and product concentrations are comparable at equilibrium.
• Equilibrium is possible only in a closed system at a given temperature. An open watch glass of water never reaches equilibrium because the vapour disperses.
• Normal melting point: temperature at which solid and liquid coexist at 1 atm (1.013 bar). Normal boiling point: temperature at which liquid and vapour are in equilibrium at 1 atm. Boiling point falls at high altitude.
• Evidence of dynamic nature: radioactive sugar dropped into a saturated sugar solution soon shows up in both the solution and the solid; ammonia synthesis with D₂ gives NH₃, NH₂D, NHD₂ and ND₃ when mixtures are combined (scrambling of H and D).

Henry’s law: the mass of a gas dissolved in a given mass of solvent at any temperature is proportional to the pressure of the gas above the solvent; this amount decreases as temperature rises. That is why an opened soda bottle goes flat.

Law of Chemical Equilibrium and the Equilibrium Constant Kc

For aA + bB ⇌ cC + dD, Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ, using equilibrium molar concentrations. Guldberg and Waage proposed this in 1864; it is also called the law of mass action (concentration was once called ‘active mass’).

• Kc has a fixed value for a reaction at a given temperature; it does not depend on the initial concentrations.
• Because Kc depends on how the equation is written, always quote the balanced equation with the value.
• H₂(g) + I₂(g) ⇌ 2HI(g): NCERT’s six experiments at 731 K show that only [HI]²/[H₂][I₂] stays constant.

Kp and Its Relation to Kc

For gaseous reactions the equilibrium constant can be written with partial pressures (Kp). Using p = cRT, the two are related by Kp = Kc(RT)^Δn.

Homogeneous and Heterogeneous Equilibria

• Pure solids and pure liquids have constant molar concentration, so they are left out of the K expression.
• For CaCO₃(s) ⇌ CaO(s) + CO₂(g), Kp = pCO₂. At 1100 K, pCO₂ = 2.0 × 10⁵ Pa, so Kp = 2.00.
• The pure solid or liquid must still be present (however small the amount) for the heterogeneous equilibrium to exist.

How is K used to predict extent and direction?

The size of K tells how far a reaction goes; comparing the reaction quotient Q with K tells which way it will move. K says nothing about how fast equilibrium is reached.

• Qc has the same form as Kc but uses concentrations at any moment, not necessarily equilibrium values.
• To find equilibrium concentrations: write the balanced equation, tabulate initial, change (in terms of x) and equilibrium values, substitute into K, solve for x and choose the root that makes chemical sense, then check.

Equilibrium Constant and Gibbs Energy

K is linked to thermodynamics, not to the rate: ΔG = ΔG° + RT ln Q. At equilibrium ΔG = 0 and Q = K, so ΔG° = −RT ln K, or K = e^(−ΔG°/RT).

Le Chatelier’s Principle: Factors Affecting Equilibrium

Le Chatelier’s principle: a change in any factor that determines the equilibrium conditions makes the system shift so as to reduce or counteract the effect of the change. It applies to all physical and chemical equilibria.

• Removing NH₃ (liquefied) in the Haber process and CO₂ from a lime kiln keeps Qc below Kc, so the reaction keeps moving forward.
• Fe³⁺ + SCN⁻ ⇌ [Fe(SCN)]²⁺ (deep red): oxalic acid removes Fe³⁺ and HgCl₂ removes SCN⁻, so the red colour fades; adding KSCN deepens it.
• C(s) + CO₂(g) ⇌ 2CO(g): higher pressure pushes the reaction backward because gas moles increase in the forward direction.
• 2NO₂ (brown) ⇌ N₂O₄ (colourless), ΔH = −57.2 kJ mol⁻¹: cooling favours N₂O₄ and lightens the colour; heating favours NO₂ and darkens it.
• Haber process: N₂ + 3H₂ ⇌ 2NH₃, ΔH = −92.38 kJ mol⁻¹. Low temperature favours yield but is slow, so an iron catalyst is used; optimum conditions are around 500 °C and 200 atm.
• A catalyst lowers the activation energy of the forward and reverse reactions by exactly the same amount and does not appear in the K expression. Contact process: Pt or V₂O₅ catalyses 2SO₂ + O₂ ⇌ 2SO₃ (Kc = 1.7 × 10²⁶ but slow). A catalyst is of little help if K is exceedingly small.

Ionic Equilibrium: Electrolytes and Acid-base Concepts

• Faraday classified substances as electrolytes (conduct in aqueous solution) and non-electrolytes, and electrolytes as strong (almost completely ionized) and weak (partially ionized).
• NaCl solution is almost 100% ionized; acetic acid is less than 5% ionized. In a weak electrolyte, ions are in equilibrium with unionized molecules: this is ionic equilibrium.
• Water’s high dielectric constant (80) cuts the attraction between Na⁺ and Cl⁻ by a factor of 80, and hydration keeps the ions apart.
• A bare proton cannot exist in water; it forms the trigonal pyramidal hydronium ion H₃O⁺, further hydrated to H₅O₂⁺, H₇O₃⁺ and H₉O₄⁺.

• Water, HCO₃⁻, HSO₄⁻ and NH₃ can act as both Brönsted acids and bases (conjugate acids: H₃O⁺, H₂CO₃, H₂SO₄, NH₄⁺; conjugate bases: OH⁻, CO₃²⁻, SO₄²⁻, NH₂⁻).
• Strong acids in NCERT: HClO₄, HCl, HBr, HI, HNO₃, H₂SO₄. Strong bases: LiOH, NaOH, KOH, CsOH, Ba(OH)₂.

Ionization of Water and the pH Scale

• pH paper (four strips) measures pH 1 to 14 with about 0.5 accuracy; a pH meter measures to 0.001.
• For very dilute strong acid (e.g. 10⁻⁸ M HCl), H⁺ from water must also be counted; the pH works out just below 7 (6.98 in NCERT), never above 7.

Weak Acids and Bases: Ka, Kb and Acid Strength

• Larger Ka means a stronger acid; larger Kb means a stronger base. Percent dissociation is another measure of weak-acid strength.
• Down a group, H-A bond strength decides acidity: HF < HCl < HBr < HI, and H₂S is a stronger acid than H₂O.
• Across a period, bond polarity (electronegativity of A) decides: CH₄ < NH₃ < H₂O < HF.
• Common ion effect: adding acetate ions to acetic acid (or H⁺ from outside) suppresses its ionization. It is an application of Le Chatelier’s principle. 0.05 M acetic acid with 0.05 M acetate gives pH 4.74.

Hydrolysis of Salts and the pH of Their Solutions

Hydrolysis is the reaction of a salt’s cation, anion or both with water. Ions of strong acids and strong bases (Na⁺, K⁺, Ca²⁺, Ba²⁺, Cl⁻, Br⁻, NO₃⁻, ClO₄⁻) are only hydrated and do not hydrolyse.

For a salt of a weak acid and a weak base, the degree of hydrolysis does not depend on the concentration of the solution. With pKa 4.76 and pKb 4.75, ammonium acetate has pH 7.005.

Buffer Solutions

A buffer resists a change in pH on dilution or on adding small amounts of acid or alkali. Body fluids such as blood and urine have definite pH values, and many medical and cosmetic formulations need a fixed pH.

• When [salt] = [acid], pH = pKa, so choose an acid whose pKa is close to the required pH.
• Diluting a buffer does not change its pH, because the ratio inside the log term stays the same.

Solubility Equilibria and Solubility Product (Ksp)

• A salt dissolves when its solvation enthalpy can overcome its lattice enthalpy; non-polar solvents give too little solvation, so ionic salts do not dissolve in them.
• BaSO₄(s) ⇌ Ba²⁺ + SO₄²⁻: Ksp = [Ba²⁺][SO₄²⁻] = 1.1 × 10⁻¹⁰ at 298 K, so molar solubility S = 1.05 × 10⁻⁵ mol L⁻¹.
• General salt MₓXᵧ: Ksp = (xS)ˣ(yS)ʸ = xˣ·yʸ·S^(x+y). For Zr₃(PO₄)₄, Ksp = (3S)³(4S)⁴ = 6912 S⁷; for A₂X₃, Ksp = 108 S⁵.
• Compare solubilities through S, not Ksp, when salts have different formula types: NCERT shows Ni(OH)₂ (Ksp 2.0 × 10⁻¹⁵) is more soluble than AgCN (Ksp 6 × 10⁻¹⁷).
• Qsp (ionic product at any moment) decides the process: precipitation when Qsp > Ksp, dissolution when Qsp < Ksp.
• Common ion effect lowers solubility. Passing HCl gas into saturated NaCl solution precipitates very pure NaCl.

Most-asked Facts for NEET and JEE Main: Equilibrium

Revise these one-line facts from the NCERT unit just before the exam.

• At equilibrium the forward and reverse rates are equal; equilibrium is dynamic.
• Kp = Kc(RT)^Δn; Δn counts gaseous species only.
• Reverse reaction: 1/K. Equation × n: Kⁿ. Reactions added: K₁ × K₂.
• Pure solids and liquids do not appear in K; for CaCO₃ ⇌ CaO + CO₂, Kp = pCO₂.
• Qc < Kc: forward; Qc > Kc: backward.
• ΔG° = −RT ln K; K > 1 when ΔG° is negative.
• Only temperature changes the value of K.
• Inert gas at constant volume: no change in equilibrium.
• Catalyst: equilibrium reached faster, composition unchanged.
• Haber process: iron catalyst, about 500 °C and 200 atm; ΔH = −92.38 kJ mol⁻¹.
• Kw = 1.0 × 10⁻¹⁴ at 298 K; pH + pOH = 14.
• Ka × Kb = Kw for a conjugate pair; pKa + pKb = 14.
• Lewis acids without a proton: BF₃, AlCl₃, metal ions.
• Acid strength: HF < HCl < HBr < HI; CH₄ < NH₃ < H₂O < HF.
• CH₃COONa solution is basic; NH₄Cl solution is acidic.
• Weak acid-weak base salt: pH = 7 + ½(pKa − pKb).
• Acetate buffer about pH 4.75; ammonium buffer about pH 9.25.
• Ksp of BaSO₄ = 1.1 × 10⁻¹⁰; Ksp of AxBy = xˣyʸS^(x+y).

Common Mistakes to Avoid

Watch for these mix-ups when you answer questions on this unit.

• Calling equilibrium static. Both reactions continue at equal rates.
• Thinking a catalyst increases yield. It only helps reach equilibrium faster.
• Saying pressure or concentration changes alter K. They shift the position; only temperature changes K.
• Counting solids, liquids or water (as solvent) in Δn or in the K expression.
• Using Kp = Kc(RT)^Δn with pressure in atm and R = 0.0831. That value of R needs pressure in bar.
• Assuming adding an inert gas always shifts equilibrium. At constant volume it has no effect.
• Forgetting that K for an exothermic reaction falls as temperature rises.
• Writing pH = 8 for 10⁻⁸ M HCl. An acid solution cannot be basic; include H⁺ from water (pH about 6.98).
• Ranking salt solubility by Ksp alone when the salts have different formula types. Work out S first.
• Confusing the two Henderson-Hasselbalch forms: for an acidic buffer it is log [salt]/[acid], not the reverse.

Equilibrium Notes PDF and Related Pages

Download the AglaSem notes for Equilibrium (key notes and important questions with answers) free: View PDF

• All chapters: Chemistry notes for NEET and JEE Main
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• Practice: NEET question papers and JEE Main question papers
• Source text: NCERT Unit 6 (NCERT, Class 11 Chemistry Part I)

Chemistry Notes for NEET and JEE Main: All Chapters

• Some Basic Concepts of Chemistry
• Structure of Atom
• Chemical Kinetics
• Equilibrium
• Organic Chemistry: Some Basic Principles and Techniques
• Electrochemistry
• Amines
• Solutions
• Thermodynamics
• Coordination Compounds
• Haloalkanes and Haloarenes
• p-Block Elements
• Alcohols, Phenols and Ethers
• Hydrocarbons
• Aldehydes, Ketones and Carboxylic Acids

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