All Class 11 Chemistry formulas

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The formulas from 4 chapters, in one place to revise from or print. Open a chapter's own sheet to see where students usually slip up.

1. Some Basic Concepts of Chemistry

Moles from mass: $\displaystyle n = \frac{\text{given mass}}{\text{molar mass}}$
Moles from particles: $\displaystyle n = \frac{N}{N_A}, \quad N_A = 6.022\times10^{23}\ \text{mol}^{-1}$
Moles of a gas at STP: $\displaystyle n = \frac{V}{22.4\ \text{L}}$
Older value. At 1 bar and 273 K, it's 22.7 L.
Mass per cent: $\displaystyle \text{mass \%} = \frac{\text{mass of element}}{\text{molar mass of compound}}\times100$

Molarity: $\displaystyle M = \frac{\text{moles of solute}}{\text{litres of solution}}$
Molality: $\displaystyle m = \frac{\text{moles of solute}}{\text{kg of solvent}}$
Doesn't change with temperature.
Mole fraction: $\displaystyle x_A = \frac{n_A}{n_A + n_B}$
Dilution: $\displaystyle M_1V_1 = M_2V_2$

Energy of the nth orbit: $\displaystyle E_n = -2.18\times10^{-18}\,\frac{Z^2}{n^2}\ \text{J}$
Radius: $\displaystyle r_n = 52.9\,\frac{n^2}{Z}\ \text{pm}$
Spectral lines: $\displaystyle \bar\nu = 109677\left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)\ \text{cm}^{-1}$

de Broglie wavelength: $\displaystyle \lambda = \frac{h}{mv}$
Heisenberg's uncertainty principle: $\displaystyle \Delta x\cdot\Delta p \ge \frac{h}{4\pi}$
Orbital angular momentum: $\displaystyle L = \sqrt{l(l + 1)}\,\frac{h}{2\pi}$
Nodes: $\displaystyle \text{radial} = n - l - 1, \quad \text{angular} = l$
Maximum electrons: $\displaystyle \text{in a shell} = 2n^2, \quad \text{in a subshell} = 2(2l + 1)$

First law: $\displaystyle \Delta U = q + w$
w is the work done on the system.
Expansion work: $\displaystyle w = -p_{\text{ext}}\Delta V$
Reversible isothermal work: $\displaystyle w = -2.303\,nRT\log\frac{V_2}{V_1}$
Enthalpy: $\displaystyle \Delta H = \Delta U + \Delta n_gRT$
Heat capacities: $\displaystyle C_p - C_V = R$

From formation enthalpies: $\displaystyle \Delta_rH^\circ = \sum\Delta_fH^\circ_{\text{products}} - \sum\Delta_fH^\circ_{\text{reactants}}$
From bond enthalpies: $\displaystyle \Delta_rH = \sum\text{BE}_{\text{reactants}} - \sum\text{BE}_{\text{products}}$
Note this order is reversed.

Entropy change: $\displaystyle \Delta S = \frac{q_{\text{rev}}}{T}$
Gibbs energy: $\displaystyle \Delta G = \Delta H - T\Delta S$
A process is spontaneous when ΔG is negative.
Link to equilibrium: $\displaystyle \Delta G^\circ = -2.303\,RT\log K$

Equilibrium constant: $\displaystyle K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$
Kp and Kc: $\displaystyle K_p = K_c(RT)^{\Delta n_g}$
Reaction quotient: $\displaystyle Q < K: \text{forward}, \quad Q > K: \text{backward}, \quad Q = K: \text{at equilibrium}$

pH: $\displaystyle \text{pH} = -\log[\text{H}^+]$
Water: $\displaystyle K_w = [\text{H}^+][\text{OH}^-] = 10^{-14}$
At 298 K, so pH + pOH = 14.
Conjugate pairs: $\displaystyle K_a\times K_b = K_w$
Weak acid: $\displaystyle \alpha = \sqrt{\frac{K_a}{C}}, \quad [\text{H}^+] = \sqrt{K_aC}$
For small α.
Acidic buffer: $\displaystyle \text{pH} = \text{p}K_a + \log\frac{[\text{salt}]}{[\text{acid}]}$
Basic buffer: $\displaystyle \text{pOH} = \text{p}K_b + \log\frac{[\text{salt}]}{[\text{base}]}$

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