CBSE · Class 11 · Chemistry

Chemistry Formula Sheet

67 formulas across 9 chapters — with variables explained and exam tips where needed.

Ch 1Basic Concepts of Chemistry(8 formulas)

Number of moles

mol

n = m/M

m = mass (g), M = molar mass (g/mol)

Number of particles

N = n × Nₐ

Nₐ = 6.022×10²³ mol⁻¹ (Avogadro's number)

Molarity

mol/L

M = n_solute / V_solution(L)

Mole fraction

χ_A = n_A / (n_A + n_B + ...)

Parts per million

ppm = (mass of solute / mass of solution) × 10⁶

Percentage yield

% yield = (actual yield / theoretical yield) × 100

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Theoretical yield is based on limiting reagent

Limiting reagent identification

Divide moles of each reactant by its stoichiometric coefficient — smallest ratio → limiting reagent

Empirical formula steps

% → g (assume 100g) → mol (÷M) → simple ratio → empirical formula

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Multiply ratios to get whole numbers if needed

Ch 2Structure of Atom(8 formulas)

Bohr's radius (hydrogen)

rₙ = n² × 0.529 Å

n = principal quantum number

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For hydrogen-like ions: rₙ = (n²/Z) × 0.529 Å

Energy of nth Bohr orbit

Eₙ = −13.6 / n² eV

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Negative sign indicates bound state; energy increases (less negative) as n increases

Rydberg formula

1/λ = R_H (1/n₁² − 1/n₂²)

R_H = 1.097 × 10⁷ m⁻¹, n₁ < n₂

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n₁=1: Lyman (UV), n₁=2: Balmer (visible), n₁=3: Paschen (IR)

de Broglie wavelength

λ = h / mv = h / p

h = 6.626 × 10⁻³⁴ J·s, m = mass, v = velocity

Heisenberg uncertainty principle

Δx · Δp ≥ h / 4π

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Δx = uncertainty in position, Δp = uncertainty in momentum; cannot simultaneously know both precisely

Max electrons in a shell

2n²

n = principal quantum number (K=2, L=8, M=18, N=32)

Max electrons in a subshell

2(2l + 1)

l = azimuthal quantum number (s:2, p:6, d:10, f:14)

Number of orbitals in subshell

(2l + 1)

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s: 1 orbital, p: 3, d: 5, f: 7

Ch 3Classification of Elements and Periodicity(7 formulas)

Effective nuclear charge

Z_eff = Z − σ

Z = atomic number, σ = shielding constant (from Slater's rules)

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Electrons farther from nucleus experience lower Z_eff due to shielding

Electronegativity (Mulliken scale)

EN = (IE₁ + EA) / 2

IE₁ = first ionisation enthalpy, EA = electron affinity

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Pauling scale is more common; F has highest EN = 4.0

Atomic radius trend (period)

Decreases left → right (↑ Z_eff pulls electrons closer)

Atomic radius trend (group)

Increases top → bottom (new shells added)

Ionisation enthalpy trend (period)

Increases left → right (with exceptions at Group 2→3 and 5→6)

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IE₁ < IE₂ < IE₃... always; large jump in IE indicates valence shell is exhausted

Electron gain enthalpy trend

Most negative for Group 17 (halogens); Cl more negative than F due to small size of F

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More negative = more energy released = higher tendency to gain electron

Metallic character trend

Decreases across period, increases down group

Ch 4Chemical Bonding and Molecular Structure(7 formulas)

Bond order (MO theory)

BO = (Nᵦ − Nₐ) / 2

Nᵦ = bonding electrons, Nₐ = antibonding electrons

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Higher bond order → shorter bond length, higher bond enthalpy, more stable molecule

Formal charge

FC = V − N − B/2

V = valence electrons, N = non-bonding electrons, B = shared (bonding) electrons

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Structures with FC closest to zero on each atom are most stable

Dipole moment

μ = q × d

q = charge (in esu), d = bond length (cm); 1 Debye = 3.336 × 10⁻³⁰ C·m

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Net dipole = vector sum of all bond dipoles; symmetrical molecules (BF₃, CCl₄) have μ = 0

VSEPR notation

AXₙEₘ

A = central atom, X = bonded atoms (n), E = lone pairs (m)

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AX₄E₀ = tetrahedral; AX₃E₁ = pyramidal; AX₂E₂ = bent (V-shape)

Hybridisation and geometry

sp: linear | sp²: trigonal planar | sp³: tetrahedral | sp³d: trigonal bipyramidal | sp³d²: octahedral

Bond length order

Triple bond < Double bond < Single bond

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More electrons shared → atoms pulled closer → shorter bond

Resonance: stability

More resonance structures → greater delocalisation → more stable molecule

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Resonance hybrid has lower energy than any single contributing structure

Ch 5Thermodynamics(7 formulas)

Relation between ΔH and ΔU

ΔH = ΔU + Δn_g RT

Δn_g = moles of gaseous products − moles of gaseous reactants

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For reactions with no gas: ΔH ≈ ΔU

Hess's law

ΔH_reaction = Σ ΔH_products − Σ ΔH_reactants (using standard enthalpies)

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Enthalpy is a state function — path doesn't matter

Bond enthalpy method

ΔH_rxn = Σ BE(bonds broken) − Σ BE(bonds formed)

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Energy is absorbed to break bonds and released to form bonds

Gibbs free energy

ΔG = ΔH − TΔS

T = temperature in K, ΔS = entropy change

Spontaneity condition

ΔG < 0 → spontaneous | ΔG = 0 → equilibrium | ΔG > 0 → non-spontaneous

ΔG and equilibrium constant

ΔG° = −RT lnK

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K > 1 → ΔG° < 0 → products favoured

First law of thermodynamics

ΔU = Q + W (IUPAC sign convention)

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Note: W = −PΔV; work done ON system is positive

Ch 6Equilibrium(10 formulas)

Equilibrium constant Kc

Kc = [products]^stoich / [reactants]^stoich (at equilibrium)

Equilibrium constant Kp

Kp = Kc(RT)^Δn_g

R = 0.0821 L·atm/mol·K, T in K

Ionic product of water

Kw = [H⁺][OH⁻] = 1×10⁻¹⁴ at 25°C

Acid dissociation constant

Ka = [H⁺][A⁻] / [HA]

Base dissociation constant

Kb = [BH⁺][OH⁻] / [B]

Relation Ka × Kb

Ka × Kb = Kw (conjugate acid-base pair)

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Stronger acid has weaker conjugate base

pH definition

pH = −log[H⁺] | pOH = −log[OH⁻] | pH + pOH = 14

Henderson-Hasselbalch equation

pH = pKa + log([A⁻]/[HA])

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Used to find pH of buffer solutions

Degree of dissociation (weak acid)

α = √(Ka/C)

C = initial concentration of weak acid

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Valid when α << 1

pH of weak acid

pH = ½(pKa − log C)

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Easier form of Henderson equation for pure weak acid

Ch 7Redox Reactions(8 formulas)

Oxidation number: free element

ON = 0 (e.g. Fe, Cl₂, S₈)

Oxidation number: monoatomic ion

ON = charge of ion (e.g. Na⁺ → +1, Cl⁻ → −1)

Oxidation number: oxygen

ON(O) = −2 usually; −1 in peroxides (H₂O₂); +2 in OF₂

Oxidation number: hydrogen

ON(H) = +1 usually; −1 in metal hydrides (NaH)

n-factor (acid-base)

n-factor = basicity of acid or acidity of base

n-factor (redox)

n-factor = change in oxidation number per formula unit

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In ionic equations: n-factor = total electrons transferred per formula unit

Equivalents (for calculations)

Equivalents = moles × n-factor | N = M × n-factor

Half-reaction method (acidic)

Add H₂O to balance O, then H⁺ to balance H, then e⁻ to balance charge

Ch 8Organic Chemistry — Basic Principles and Techniques(6 formulas)

Index of Hydrogen Deficiency (IHD)

IHD = (2C + 2 + N − H − X) / 2

C = carbons, N = nitrogens, H = hydrogens, X = halogens; O and S do not affect IHD

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IHD = 1: one ring or one double bond; IHD = 2: one triple bond or two degrees of unsaturation

Inductive effect (+I order)

(CH₃)₃C− > (CH₃)₂CH− > CH₃CH₂− > CH₃− > H−

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+I groups donate electrons; −I groups (halogens, NO₂, CN) withdraw electrons

Carbocation stability

3° > 2° > 1° > methyl

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Hyperconjugation and inductive effect of alkyl groups stabilise higher-order carbocations

Carbanion stability

methyl > 1° > 2° > 3° (opposite of carbocations)

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Electron-withdrawing groups stabilise carbanions; alkyl groups destabilise them

Free radical stability

3° > 2° > 1° > methyl (same order as carbocations)

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Hyperconjugation stabilises free radicals; used in halogenation of alkanes

IUPAC priority for functional groups

COOH > CHO > C=O > OH > NH₂ > C=C > C≡C > halogen

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Highest-priority group gives suffix; others become prefixes

Ch 9Hydrocarbons(6 formulas)

Markovnikov's rule

In addition to unsymmetrical alkene: H adds to C with more H atoms; X adds to C with fewer H atoms

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Explains regioselectivity; anti-Markovnikov occurs with HBr in presence of peroxides (radical mechanism)

Hückel's rule (aromaticity)

Aromatic if: planar, fully conjugated ring, (4n + 2) π electrons (n = 0, 1, 2...)

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Benzene: n=1, 6π electrons; naphthalene: 10π; cyclopentadienyl anion: 6π

Free radical halogenation selectivity

Reactivity: F₂ > Cl₂ > Br₂ > I₂ | Selectivity: I₂ > Br₂ > Cl₂ > F₂

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Br₂ is most useful: high selectivity for 3° > 2° > 1° H atoms

Baeyer's reagent test

Alkene/Alkyne + cold dil. KMnO₄ (alk.) → colourless (purple decolourises) = unsaturation confirmed

Ozonolysis

C=C + O₃ → [ozonide] → with Zn/H₂O: two carbonyl fragments | with H₂O₂: carboxylic acids

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Used to locate double bond position by identifying carbonyl fragments

EAS — directing effects

ortho/para directors (activating): −OH, −NH₂, −CH₃, halogens | meta directors (deactivating): −NO₂, −CHO, −COOH, −CN

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Halogens are deactivating yet o/p directors — remember as an exception

More for CBSE Class 11

All Formula Sheets →Important Derivations →Named Reactions →Important Diagrams →Previous Year Papers →

Chemistry Formula Sheet — Class 11