Class 11 Chemistry: Essential Formulas & Concepts (PDF) — All Important Formulas Free
Free class 11 chemistry: essential formulas & concepts — all 64 key formulas on one page, downloadable as PDF for fast revision before CBSE board exams. No signup.
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TL;DR: Free class 11 chemistry: essential formulas & concepts — all 64 key formulas on one page, downloadable as PDF for fast revision before CBSE board exam…
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Master the foundational concepts of chemistry: atomic structure, mole theory, gas laws, thermodynamics, equilibrium, and redox reactions. All formulas in one reference sheet for quick problem-solving.
Mole Concept & Stoichiometry
| Formula | Expression |
|---|---|
| Number of moles from mass | n = given mass (g) / molar mass (g/mol) — n is always in moles; divide grams by atomic/molecular weight |
| Number of moles from particles | n = N / Nₐ — N = number of particles, Nₐ = Avogadro's number = 6.022 × 10²³ |
| Molar mass from density | M = dRT / P — d = density (g/L), R = 0.0821 L·atm/(mol·K), T in Kelvin, P in atm |
| Percentage composition | % element = (atomic mass × no. of atoms / molar mass) × 100 — Sum of all percentages = 100% |
| Empirical formula mass | n = molar mass / empirical formula mass — Molecular formula = (empirical formula)ₙ |
| Stoichiometric ratio | mol A / mol B = coefficient A / coefficient B — From balanced chemical equation |
| Limiting reagent | Compare moles of reactant / stoichiometric coefficient — Smallest ratio is limiting; controls theoretical yield |
| Percent yield | % yield = (actual yield / theoretical yield) × 100 — Actual from experiment; theoretical from stoichiometry |
Concentration Terms
| Formula | Expression |
|---|---|
| Molarity | M = n / V(L) = moles of solute / litres of solution — Temperature-dependent; liters of final solution, not solvent |
| Molality | m = n / W(kg) = moles of solute / kg of solvent — Temperature-independent; uses mass of solvent only |
| Normality | N = n × equivalents / V(L) — Equivalents = n × valency for acids/bases/salts |
| Mole fraction | χₐ = nₐ / (nₐ + nᵦ + ...) — Sum of all mole fractions = 1; dimensionless |
| Mass percent (w/w) | w/w % = (mass of solute / mass of solution) × 100 — Solution = solute + solvent |
| Volume percent (v/v) | v/v % = (volume of solute / volume of solution) × 100 — For liquids; may not be additive |
| Relation: Molarity & Molality | m = M / (1 - M × M / 1000) — M = molar mass of solute; holds for dilute solutions |
| Dilution formula | M₁V₁ = M₂V₂ — Moles solute conserved; works for any molarity scale |
Gas Laws & Ideal Gas Equation
| Formula | Expression |
|---|---|
| Ideal gas equation | PV = nRT — P (Pa or atm), V (m³ or L), n (mol), R = 8.314 J/(mol·K) or 0.0821 L·atm/(mol·K), T (K) |
| Boyle's Law | P₁V₁ = P₂V₂ — At constant T and n; inverse relationship |
| Charles' Law | V₁/T₁ = V₂/T₂ — At constant P and n; direct relationship |
| Gay-Lussac's Law | P₁/T₁ = P₂/T₂ — At constant V and n; direct relationship |
| Combined gas law | P₁V₁/T₁ = P₂V₂/T₂ — No mole change needed; compare two states |
| Dalton's law of partial pressures | Pₜₒₜₐₗ = P₁ + P₂ + P₃ + ... — Each gas contributes its own pressure independently |
| Henry's Law | Pgas = Kₕ × χsolute (in solution) — Kₕ = Henry's law constant; Pgas = partial pressure |
| Graham's law of diffusion | r₁/r₂ = √(M₂/M₁) — r = rate of diffusion; M = molar mass; lighter gas diffuses faster |
| Density of gas at STP | d = PM / RT — d (g/L); relates density to molar mass |
Atomic Structure
| Formula | Expression |
|---|---|
| Wavelength-frequency relation | c = λν — c = 3 × 10⁸ m/s, λ = wavelength (m), ν = frequency (Hz) |
| Photon energy | E = hν = hc/λ — h = 6.626 × 10⁻³⁴ J·s; higher frequency = higher energy |
| Rydberg equation (hydrogen) | 1/λ = R∞(1/n₁² − 1/n₂²) — R∞ = 1.097 × 10⁷ m⁻¹; n₁ < n₂ for emission; n₁ = 1 is Lyman, 2 is Balmer |
| Energy levels (Bohr model) | Eₙ = −13.6/n² eV — n = 1, 2, 3...; negative energy means bound state |
| de Broglie wavelength | λ = h / p = h / mv — All particles have wavelike behavior; h = Planck constant |
| Uncertainty principle (Heisenberg) | Δx × Δp ≥ h / 4π — Cannot know position and momentum simultaneously with certainty |
| Ionization energy (hydrogen-like) | IE = 13.6 × Z² / n² eV — Z = atomic number, n = principal quantum number |
| Electron in orbital probability | ψ² = probability density at a point in space — ψ = wave function; ψ² gives electron density |
Thermodynamics
| Formula | Expression |
|---|---|
| Heat absorbed/released | q = mcΔT — m = mass (g), c = specific heat capacity (J/g·°C), ΔT = temperature change (K or °C) |
| Enthalpy change | ΔH = q_p (at constant pressure) — ΔH > 0 endothermic, ΔH < 0 exothermic |
| Hess's Law | ΔH_reaction = Σ ΔH_products − Σ ΔH_reactants — Enthalpy is a state function; independent of reaction path |
| Standard enthalpy of formation | ΔH°f = enthalpy change for 1 mol pure compound from elements — ΔH°f(element) = 0; sum of ΔH°f gives ΔH°rxn |
| Entropy change | ΔS = q_rev / T — S > 0 disorder increases; state function |
| Gibbs free energy | ΔG = ΔH − TΔS — ΔG < 0 spontaneous, ΔG > 0 non-spontaneous at that T |
| Gibbs energy and equilibrium | ΔG° = −RT ln Kₑq — R = 8.314 J/(mol·K); Kₑq = equilibrium constant |
| Calorimetry (coffee cup) | q_solution + q_calorimeter + q_reaction = 0 — Heat lost = heat gained (closed system) |
Chemical Equilibrium
| Formula | Expression |
|---|---|
| Equilibrium constant (Kc) | Kc = [C]^c[D]^d / [A]^a[B]^b — For aA + bB ⇌ cC + dD; square brackets = molar concentration at equilibrium |
| Equilibrium constant (Kp) | Kp = P_C^c × P_D^d / P_A^a × P_B^b — P = partial pressure (atm or Pa) |
| Relation between Kp and Kc | Kp = Kc(RT)^Δn — Δn = (c + d) − (a + b) = change in moles of gas |
| Ion product of water | Kw = [H⁺][OH⁻] = 10⁻¹⁴ at 25°C — Always 10⁻¹⁴ at 25°C regardless of acid/base |
| pH definition | pH = −log[H⁺]; pOH = −log[OH⁻] — pH + pOH = 14 at 25°C; pH < 7 acidic, pH > 7 basic |
| Acid dissociation constant | Ka = [H⁺][A⁻] / [HA] — Larger Ka = stronger acid; pKa = −log Ka |
| Base dissociation constant | Kb = [BH⁺][OH⁻] / [B] — Ka × Kb = Kw for conjugate acid-base pair |
| Henderson-Hasselbalch equation | pH = pKa + log([A⁻]/[HA]) — For buffer solutions; [A⁻] = conjugate base, [HA] = weak acid |
| Buffer capacity | Buffer resists pH change when small amounts of acid/base are added — Best when pH ≈ pKa; ratio [A⁻]/[HA] ≈ 1 |
| Common ion effect | Adding common ion shifts equilibrium left, decreases solubility — E.g. adding NaCl to NaCl(aq) ⇌ Na⁺ + Cl⁻ |
Redox Reactions & Electrochemistry
| Formula | Expression |
|---|---|
| Oxidation number rules | Element = 0; monatomic ion = charge; O usually −2; H usually +1; F always −1 — Sum of oxidation numbers = 0 (neutral) or charge (ion) |
| Oxidation and reduction | Oxidation = loss of e⁻ (O.N. increases); Reduction = gain of e⁻ (O.N. decreases) — OIL RIG; redox pair involved in every redox reaction |
| Balancing redox (half-reaction method) | Balance atoms except O & H → balance O with H₂O → balance H with H⁺ or OH⁻ → balance charge with e⁻ — Multiply half-reactions to equalize electrons lost/gained |
| Equivalent mass | Eq.M = molar mass / valency change per atom — For redox, valency change = no. of e⁻ transferred per formula unit |
| Normality for redox | N = n / (Eq.M × V) × 1000; n_e⁻ = N × V — Electrons gained = electrons lost in balanced equation |
| Electrochemical cell notation | Anode (−) | salt bridge | Cathode (+); read −→ + for conventional current — Oxidation at anode (left), reduction at cathode (right) |
| Standard cell potential | E°cell = E°cathode − E°anode — E° > 0 spontaneous, E° < 0 non-spontaneous; look up in tables |
| Gibbs energy and cell potential | ΔG° = −nFE°cell — n = moles of e⁻, F = Faraday constant = 96,500 C/mol |
Quick Key Concepts Review
| Formula | Expression |
|---|---|
| Avogadro's number | Nₐ = 6.022 × 10²³ particles/mol — Defines the mole; used to count atoms, molecules, ions |
| Molar volume at STP | Vm = 22.4 L/mol (at 0°C, 1 atm) — Useful for gas stoichiometry; nearly 24 L/mol at ~25°C, 1 atm |
| Gas constant R | R = 8.314 J/(mol·K) = 0.0821 L·atm/(mol·K) = 2 cal/(mol·K) — Use first form for SI; second for gas law calculations |
| Faraday's constant | F = 96,500 C/mol e⁻ — Charge of 1 mole of electrons; coulomb = amp × second |
| Atomic mass unit | 1 u = 1.66 × 10⁻²⁷ kg — Mass of one nucleon (proton or neutron); ¹²C standard = 12.000 u |
FAQs
How do I find the empirical formula from percentage composition?
Assume 100 g sample. Convert % to grams, divide by atomic masses to get moles, divide by smallest to get ratio, multiply by a whole number if needed. Example: C 80%, H 20% → 100 g sample → 80/12 = 6.67 mol C, 20/1 = 20 mol H → divide by 6.67 → 1:3 ratio → empirical formula CH₃.
When should I use Kc vs Kp?
Use Kc when concentrations (mol/L) are given; Kp when partial pressures (atm/Pa) are given. For gases, they relate by Kp = Kc(RT)^Δn where Δn is the change in moles of gas. Both describe the same equilibrium.
What is the difference between an exothermic and endothermic reaction?
Exothermic releases heat (ΔH < 0, q > 0); reactants have more energy than products. Endothermic absorbs heat (ΔH > 0, q < 0); products have more energy than reactants. Combustion is always exothermic; melting ice is endothermic.