Chemistry HSC Reference Sheet — Explained

$n$ — amount of substance (mol); $m$ — mass (g); $M$ — molar mass (g mol⁻¹).

Any time you are given a mass and need moles, or given moles and need to find the mass of a substance.

$n$ — amount of substance (mol); $N$ — number of particles; $N_A = 6.022 \times 10^{23}$ mol⁻¹ (Avogadro's constant).

Converting between a raw count of atoms, molecules or ions and the amount in moles.

$n$ — amount of substance (mol); $c$ — molar concentration (mol L⁻¹); $V$ — volume (L).

Finding moles of solute from a solution's concentration and volume, or working out what volume contains a required number of moles.

$P$ — pressure (Pa); $V$ — volume (m³); $n$ — amount of substance (mol); $R = 8.314$ J mol⁻¹ K⁻¹; $T$ — absolute temperature (K).

Any gas calculation where pressure, volume, moles and temperature are mixed — including finding an unknown when the other three are known, or finding molar mass from gas density data.

$V$ — volume of gas (L); $n$ — amount of substance (mol); $V_m$ — molar volume at the given conditions (22.71 L mol⁻¹ at 0 °C / 100 kPa; 24.79 L mol⁻¹ at 25 °C / 100 kPa).

Quick calculations at standard temperature and pressure — avoids the full ideal gas law when $T$ and $P$ are fixed at the standard conditions.

$c$ — molar concentration (mol L⁻¹); $n$ — amount of solute (mol); $V$ — volume of solution (L).

Making up a standard solution, calculating concentration from titration data, or determining how many moles are in a given volume.

$c_1$ — initial concentration (mol L⁻¹); $V_1$ — initial volume (L); $c_2$ — final concentration (mol L⁻¹); $V_2$ — final volume (L).

Calculating the concentration after diluting a solution, or determining what volume of stock solution to dilute.

$m_{\text{solute}}$ — mass of solute (g); $m_{\text{solution}}$ — mass of solution (g). 1 ppm = 1 mg kg⁻¹ = 1 mg L⁻¹ (for dilute aqueous solutions where $\rho \approx 1$ g mL⁻¹).

Expressing very low concentrations of trace substances (pollutants, minerals in water, blood alcohol).

$q$ — heat transferred (J); $m$ — mass of solution or substance (g); $c$ — specific heat capacity (J g⁻¹ K⁻¹; for water $c = 4.18$ J g⁻¹ K⁻¹); $\Delta T$ — temperature change (K or °C, same magnitude).

Calorimetry experiments: calculating heat released or absorbed by a reaction from the temperature change of the surrounding water/solution.

$\Delta H$ — molar enthalpy change (kJ mol⁻¹); $q$ — heat transferred to surroundings (kJ, from $q = mc\Delta T$); $n$ — moles of the substance specified in the equation.

Converting a calorimetry result ($q$) into a standard molar enthalpy ($\Delta H$) for reporting, or for comparing with literature values.

$\Delta H_{\text{rxn}}$ — enthalpy change for the target reaction; $\sum \Delta H_{\text{products}}$ — sum of standard enthalpies of formation of products; $\sum \Delta H_{\text{reactants}}$ — sum of standard enthalpies of formation of reactants.

Calculating $\Delta H$ of a reaction from standard enthalpies of formation, or combining thermochemical equations (multiply, reverse, add) when a direct measurement is not available.

$[\text{X}]$ — equilibrium molar concentration of species X (mol L⁻¹); $a, b, c, d$ — stoichiometric coefficients from the balanced equation $a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D}$.

Writing an equilibrium expression for any reversible reaction; comparing $Q$ vs $K$ to predict direction of shift; calculating equilibrium concentrations.

Same form as $K_{eq}$ but uses current (non-equilibrium) concentrations. $Q < K$: reaction proceeds forward; $Q > K$: reaction proceeds in reverse; $Q = K$: system is at equilibrium.

Predicting which direction a reaction will shift when conditions change (e.g., adding or removing a reactant).

$[\text{H}^+]$ — molar concentration of hydrogen ions (mol L⁻¹, at 25 °C); equivalently $[\text{H}_3\text{O}^+]$.

Converting between $[\text{H}^+]$ and pH, or calculating pH of a strong acid/base solution.

$\text{pOH} = -\log_{10}[\text{OH}^-]$; the sum equals 14 because $K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}$ at 25 °C, so $-\log K_w = \text{pH} + \text{pOH} = 14$.

Finding pH of a strong base solution (calculate pOH first, then subtract from 14), or finding $[\text{OH}^-]$ from pH.

$K_w$ — ionic product of water (mol² L⁻²); $[\text{H}^+]$ and $[\text{OH}^-]$ in mol L⁻¹ at 25 °C.

Finding $[\text{OH}^-]$ from $[\text{H}^+]$ or vice versa; understanding neutral pH = 7 at 25 °C.

$K_a$ — acid dissociation constant; $[\text{H}^+]$ — hydrogen ion concentration (mol L⁻¹); $[\text{A}^-]$ — conjugate base concentration (mol L⁻¹); $[\text{HA}]$ — equilibrium concentration of weak acid (mol L⁻¹).

Calculating the pH of a weak acid solution, finding degree of dissociation, or comparing acid strengths (larger $K_a$ = stronger weak acid).

$\text{p}K_a$ — negative base-10 logarithm of the acid dissociation constant. Smaller $\text{p}K_a$ = stronger acid.

Comparing relative strengths of weak acids; calculating $K_a$ from a tabulated $\text{p}K_a$ value.

$Q$ — electric charge (coulombs, C); $I$ — current (amperes, A); $t$ — time (seconds, s).

Calculating the total charge passed through an electrolytic cell from current and time, as the first step in all electrolysis yield calculations.

$n_e$ — moles of electrons transferred; $Q$ — charge (C); $F = 96\,485$ C mol⁻¹ (Faraday constant).

Converting total charge into moles of electrons, then using the half-equation stoichiometry to find moles of substance deposited or dissolved at an electrode.

$m$ — mass of substance deposited or dissolved (g); $n$ — moles of substance (mol, from Faraday's law and half-equation stoichiometry); $M$ — molar mass (g mol⁻¹).

Final step in electrolysis calculations — converting moles of deposited/dissolved substance to mass.