First-year university chemistry: what the modules cover and how to practise

What first-year chemistry modules cover

In the US this is usually called general chemistry, with organic chemistry as a separate course; other universities may split the same material into separate physical, inorganic and organic modules. Check your module handbook for the exact syllabus and exam format.

The core topics

• Atomic structure: quantum numbers, atomic orbitals and electron configurations.
• Bonding: Lewis structures, VSEPR, hybridisation and molecular orbital theory.
• Thermodynamics: enthalpy, entropy and Gibbs free energy.
• Kinetics: rate laws, integrated rate equations, half-lives, activation energy and catalysis.
• Equilibria: equilibrium constants, acids and bases, buffers and solubility.
• Electrochemistry: standard potentials, cell potentials and the link to ΔG.
• Inorganic chemistry: coordination compounds and transition metals.
• Organic chemistry: stereochemistry, substitution, elimination and addition mechanisms, and spectroscopy.

Worked examples

Gibbs free energy

For N₂ + 3H₂ → 2NH₃, ΔH = −92 kJ mol⁻¹ and ΔS = −199 J K⁻¹ mol⁻¹. At 298 K:

The reaction is spontaneous at 298 K. It stops being spontaneous above about 92 ÷ 0.199 ≈ 462 K, where ΔG becomes positive (assuming ΔH and ΔS stay constant).

First-order kinetics

A first-order reaction has k = 0.0231 min⁻¹, so its half-life is t½ = ln 2 ÷ k ≈ 30 minutes, whatever the starting concentration.

Electrochemistry

In a zinc-copper cell, E°cell = E°(Cu²⁺/Cu) − E°(Zn²⁺/Zn) = 0.34 − (−0.76) = 1.10 V. With n = 2, ΔG° = −nFE° = −2 × 96,485 × 1.10 ≈ −212 kJ mol⁻¹.

Molecular orbital theory

In O₂, the valence electrons fill 8 bonding and 4 antibonding orbital places, so the bond order is (8 − 4) ÷ 2 = 2. Two of the antibonding electrons are unpaired, which is why O₂ is paramagnetic, something a simple Lewis structure cannot explain.

Use a rate equation in our Reaction Kinetics course, then press Mark to check your answer in the course:

Use the daily paper as practice

Uni-Sine publishes a free daily University Chemistry: Year 1 paper. It is a short timed practice paper: 8 questions worth 25 to 35 marks, with a suggested 40 to 55 minutes, marked when you finish. It is not your university's exam, so use it alongside your module's problem sheets and past papers.

• What it covers well: atomic orbitals and quantum numbers, molecular orbital theory, thermodynamics and Gibbs free energy, kinetics and mechanisms, electrochemistry, coordination chemistry, stereochemistry, organic reaction mechanisms and spectroscopy.
• What it does not cover: lab skills, and gives less space to acids and bases, solubility, gases and main group chemistry than many courses do.

Common mistakes in first-year chemistry

• Mixing kJ and J in ΔG = ΔH − TΔS. Convert ΔS to kJ K⁻¹ mol⁻¹ first.
• Assuming rate laws come from the equation. Orders are found by experiment.
• Multiplying E° by the coefficients. Standard potentials don't scale with the equation.
• Curly arrows going the wrong way. Arrows show electrons moving from a lone pair or bond.

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