Chemistry 30 Diploma cheat sheet

ΔH is negative for exothermic reactions (energy released to the surroundings) and positive for endothermic reactions (energy absorbed).

Calorimetry: q = mcΔT, where c is the specific heat capacity in J/(g·°C); the reaction's heat is q_reaction = −q_surroundings.

Molar enthalpy: ΔH = −q/n, where n is the moles of the substance the enthalpy value is quoted per mole of.

Hess's law: ΔH_total equals the sum of the step enthalpies; reversing a step changes the sign of its ΔH, and multiplying an equation multiplies its ΔH.

Standard enthalpy of formation ΔH°f is the energy needed to form 1 mol of a compound from its elements in standard states; ΔH°f = 0 for an element in its standard state.

ΔH°rxn = Σ n·ΔH°f(products) − Σ n·ΔH°f(reactants), with the coefficients n taken from the balanced equation.

Average bond energies estimate ΔH ≈ Σ bonds broken − Σ bonds formed, applied to gaseous species.

On a potential energy diagram, ΔH = H_products − H_reactants and the activation energy is the peak height measured above the reactants.

During a phase change the temperature remains constant; q = n·ΔH_fus for melting/freezing and q = n·ΔH_vap for boiling/condensing.

Oxidation is loss of electrons (oxidation number increases); reduction is gain of electrons (oxidation number decreases).

Oxidation always occurs at the anode and reduction at the cathode; electrons travel through the external wire from anode to cathode.

Polarity: in a galvanic cell the anode is negative and the cathode is positive; in an electrolytic cell the anode is positive and the cathode is negative.

E°cell = E°cathode − E°anode using standard reduction potentials from the table; the reaction as written is spontaneous when E°cell > 0.

A higher standard reduction potential means a stronger oxidizing agent; the species with the lowest E°(reduction) is the strongest reducing agent.

Cell notation: anode | anode ion || cathode ion | cathode, with a single bar for a phase boundary and a double bar for the salt bridge.

ΔG° = −nFE° with F = 9.65 × 10^4 C/mol; a positive E°cell corresponds to a negative ΔG°.

In the salt bridge (or porous cup), anions migrate toward the anode compartment and cations toward the cathode compartment to maintain electrical neutrality.

Electrolysis: moles of electrons = It/F; use the half-reaction stoichiometry to convert moles of electrons into moles of product.

General formulas: alkane CnH2n+2, alkene CnH2n, alkyne CnH2n−2, cycloalkane CnH2n.

IUPAC naming: the longest chain gives the parent name; number to give the lowest locants to the principal functional group, then to multiple bonds, then to substituents, and list substituents alphabetically.

Boiling point rises with chain length (stronger London dispersion forces) and falls with increased branching (less surface contact between molecules).

Alcohols boil higher than alkanes or ethers of similar molar mass because of hydrogen bonding; water solubility decreases as the hydrocarbon portion grows.

Alkenes undergo addition: H2 (Ni or Pt catalyst), Br2 (decolourises bromine water), HX, and H2O (acid catalyst) all convert the C=C bond to a single bond.

Markovnikov's rule: in HX addition, the hydrogen goes to the carbon that already has more hydrogens and the halogen to the more substituted carbon.

Alkanes react by substitution with halogens under UV light, whereas alkenes and alkynes react by addition at the multiple bond.

Oxidation: primary alcohol → aldehyde → carboxylic acid; secondary alcohol → ketone; tertiary alcohol resists oxidation; acidified dichromate changes from orange to green.

Esterification: carboxylic acid + alcohol ⇌ ester + water with a strong acid catalyst; heating an alcohol with concentrated H2SO4 eliminates water to give an alkene.

For aA + bB ⇌ cC + dD, K = [C]^c[D]^d / ([A]^a[B]^b); pure solids and pure liquids are omitted from the expression.

K depends only on temperature; a catalyst changes how fast equilibrium is reached but not the equilibrium position or the value of K.

Le Châtelier: adding a reactant shifts the system toward products; decreasing volume (raising pressure) shifts the system toward the side with fewer moles of gas.

Raising temperature increases K for an endothermic reaction and decreases K for an exothermic reaction.

Kw = [H3O+][OH−] = 1.0 × 10^−14 at 25 °C; pH = −log[H3O+], pOH = −log[OH−], and pH + pOH = 14 at 25 °C.

Ka·Kb = Kw for a conjugate pair, so the stronger the acid, the weaker its conjugate base.

Weak acid calculation: Ka = x^2/(c − x) with x = [H3O+] from an ICE table; if c/Ka > 1000, assume c − x ≈ c.

Buffer pH = pKa + log([conjugate base]/[weak acid]); a buffer works best when the required pH lies within about one unit of the pKa.

Equivalence point pH: 7 for strong acid–strong base, greater than 7 for weak acid–strong base, less than 7 for strong acid–weak base; pick an indicator whose pKa is near the equivalence pH.