Grade 12 Chemistry formula sheet — printable

Moles from Mass (n = m/M)

n = \frac{m}{M}

Particles from Moles (Avogadro's Number)

N = n\,N_A

Sensible Heat (Q = mcΔT)

Q = m c \Delta T

Latent Heat

Q = m L

Heat of Reaction

q = n \Delta H

Molarity (C = n/V)

C = \frac{n}{V}

Hess's Law (Three-Step Sum)

\Delta H_{\text{rxn}} = \Delta H_1 + \Delta H_2 + \Delta H_3

Standard Enthalpy of Reaction from Formation Enthalpies

\Delta H^{\circ}_{\text{rxn}} = \sum \Delta H^{\circ}_{f,\text{prod}} - \sum \Delta H^{\circ}_{f,\text{react}}

Power-Law Reaction Rate

r = k\,C_A^{\,n}

Arrhenius Equation

k = A\,e^{-E_a/RT}

Arrhenius Two-Temperature Form

\ln\frac{k_2}{k_1} = \frac{E_a}{R}\left(\frac{1}{T_1} - \frac{1}{T_2}\right)

Activation Energy from an Arrhenius Plot

E_a = -R \times \text{slope}

Zero-Order Integrated Rate Law

[\mathrm{A}] = [\mathrm{A}]_0 - kt

First-Order Integrated Rate Law

[\mathrm{A}] = [\mathrm{A}]_0\,e^{-kt}

Half-Life Decay

N = N_0 \left(\frac{1}{2}\right)^{t/t_{1/2}}

Half-Life and Decay Constant

t_{1/2} = \frac{\ln 2}{\lambda}

Second-Order Integrated Rate Law

\frac{1}{[\mathrm{A}]} = \frac{1}{[\mathrm{A}]_0} + kt

Half-Life of a Second-Order Reaction

t_{1/2} = \frac{1}{k\,[\mathrm{A}]_0}

Equilibrium Constant Kc (A + B ⇌ C + D)

K_c = \frac{[\mathrm{C}][\mathrm{D}]}{[\mathrm{A}][\mathrm{B}]}

Reaction Quotient Q (aA + bB ⇌ cC)

Q = \frac{[\mathrm{C}]^{c}}{[\mathrm{A}]^{a}\,[\mathrm{B}]^{b}}

Kp from Kc (Kp = Kc(RT)^Δn)

K_p = K_c (RT)^{\Delta n}

Ideal Gas Law

P V = n R T

Solubility Product of a 1:1 Salt

K_{sp} = s^{2}

Solubility Product of an AB₂ Salt

K_{sp} = 4s^{3}

Gibbs Free Energy Change (ΔG = ΔH − TΔS)

\Delta G = \Delta H - T\,\Delta S

Gibbs Free Energy and the Equilibrium Constant

\Delta G^{\circ} = -RT\ln K

pH from Hydrogen Ion Concentration

\mathrm{pH} = -\log_{10}\,[\mathrm{H^+}]

pH and pOH Relation

\mathrm{pH} + \mathrm{pOH} = 14

Ka and Kb Relation through Kw

K_a \, K_b = K_w

pKa from Acid Dissociation Constant

\mathrm{p}K_a = -\log_{10} K_a

pKb from Base Dissociation Constant

\mathrm{p}K_b = -\log_{10} K_b

pH of a Weak Acid from Ka

\mathrm{pH} = -\log_{10}\sqrt{K_a\,C}

Percent Ionization of a Weak Acid

\%\,\text{ion} = \frac{[\mathrm{H^+}]}{C} \times 100\%

Henderson–Hasselbalch Equation (Weak Acid Buffer)

\mathrm{pH} = \mathrm{p}K_a + \log_{10}\!\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}

Henderson–Hasselbalch Equation (Weak Base Buffer)

\mathrm{pOH} = \mathrm{p}K_b + \log_{10}\!\frac{[\mathrm{BH^+}]}{[\mathrm{B}]}

Dilution Equation (C1V1 = C2V2)

C_1 V_1 = C_2 V_2

Titration: Concentration of an Unknown

C_a = \frac{n\,C_b V_b}{V_a}

Standard Cell Potential from Half-Cells

E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}

Nernst Equation

E = E^{\circ} - \frac{RT}{nF}\ln Q

Electric Charge (Q = It)

Q = I t

Faraday's Law of Electrolysis (m = QM/nF)

m = \frac{Q M}{n F}

Boyle's Law

P_1 V_1 = P_2 V_2

Charles's Law

\frac{V_1}{T_1} = \frac{V_2}{T_2}

Gay-Lussac's Law

\frac{P_1}{T_1} = \frac{P_2}{T_2}

Combined Gas Law

\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}

Gas Volume at STP

V = n\,V_m

Gas Density from Molar Mass

\rho = \frac{PM}{RT}

Partial Pressure from Mole Fraction

P_i = x_i \, P_{\text{total}}

Mole Fraction

x_1 = \frac{n_1}{n_1 + n_2}

Graham's Law of Effusion

\frac{r_1}{r_2} = \sqrt{\frac{M_2}{M_1}}

Molality (b = n/m)

b = \frac{n}{m_{\text{solvent}}}

Mass Percent of a Solution

c = \frac{m_{\text{solute}}}{m_{\text{solution}}} \times 100\%

Boiling-Point Elevation

\Delta T_b = K_b \, b

Freezing-Point Depression

\Delta T_f = K_f \, b

Osmotic Pressure (Π = MRT)

\Pi = M R T

Raoult's Law

P = x \, P^{0}

Henry's Law (Gas Solubility)

C = H\,P

Clausius–Clapeyron Equation (Two-Point Form)

\ln\!\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_{vap}}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)