CHEM10007 Fundamentals of Chemistry
The Complete Exam Bible · S1 2026CHEM10007 Overview
- University of Melbourne
- S1 2026
- 11 chapters
- Chemistry
CHEM10007 Fundamentals of Chemistry at the University of Melbourne takes you from zero chemistry knowledge through atoms, bonding, the mole and stoichiometry, gases and energy, equilibrium, acids/bases and redox, to organic structure across 11 weeks.
- Assessed by Online Feedback Quizzes (FBQ-1 to FBQ-5) 10% · Sustainability Independent Learning Task (ILT) 5% + 3 more — Practical work (laboratory) is a hurdle
- Key terms Mole (n), Limiting reagent, Enthalpy change (ΔH), Standard cell potential (E°cell)
- How to prepare Treat the appendices booklet as a working tool, not a crutch: print the formula sheet early and practise choosing the right relationship for…
- Most asked Is the CHEM10007 exam open or closed book?
How CHEM10007 is assessed
| Component | Weight | Format |
|---|---|---|
| Online Feedback Quizzes (FBQ-1 to FBQ-5) | 10% | 5 quizzes (2% each), 30-min online MCQ, open across the relevant weeks |
| Sustainability Independent Learning Task (ILT) | 5% | Online interactive quiz on sustainability (content not covered in lectures), due Week 12 |
| Practical work (laboratory) · hurdle | 20% | 6 experiments F1–F6; 6 lab worksheets/reports; report cannot be submitted without doing the experiment |
| Pre-Laboratory Quizzes | 5% | 5 online quizzes, each completed before its practical class (no quiz = no lab entry) |
| Exam | 60% | On-campus, closed-book, 2 h writing + 15 min reading; Section A MCQ (60 marks) + Section B extended written (60 marks); appendices booklet provided |
What CHEM10007 covers
The whole subject → one exam-ready map. Each topic links to its free chapter guide.
Atoms, the Periodic Table & Electronic Structure
Matter, isotopes, atomic mass, the Bohr & quantum-mechanical models, electron configuration and periodic trends (Wks 1-2)The opening fortnight builds chemistry's vocabulary from nothing: what matter is, how atoms are structured, and why the Periodic Table is arranged the way it is.
- Matter: elements, compounds and mixtures; empirical vs molecular formulae
- The Periodic Table: groups (alkali metals, alkaline earths, halogens, noble gases), periods, metals/non-metals/metalloids
- Atomic structure: atomic number Z, mass number A = p + n, isotopes
- Relative atomic mass as an abundance-weighted average: A<sub>r</sub> = Σ(isotope mass × fractional abundance)
- Bohr model: fixed energy levels, line spectra, ΔE = hν = hc/λ, and its limitations
- Quantum-mechanical model: shells (n) → sub-shells (s, p, d, f) → orbitals; 2 electrons per orbital
Chemical Bonding, Lewis Structures & Shape
Ionic, covalent and metallic bonding; common ions; naming; drawing Lewis structures and molecular shape (Wk 2)This chapter explains why atoms join and what shape the result takes. You distinguish ionic, covalent and metallic bonding, learn the charges of common monatomic and polyatomic ions, and name ionic and covalent compounds.
- Ionic bonding: metal + non-metal, electron transfer, ionic lattices
- Covalent bonding: non-metal + non-metal, electron sharing; molecules vs covalent lattices
- Metallic bonding: a 'sea' of delocalised electrons
- Common ions: monatomic charges from group position; polyatomic ions (NO<sub>3</sub><sup>−</sup>, SO<sub>4</sub><sup>2−</sup>, CO<sub>3</sub><sup>2−</sup>, PO<sub>4</sub><sup>3−</sup>, NH<sub>4</sub><sup>+</sup>, OH<sup>−</sup>)
- Naming: ionic compounds (Roman numerals for variable-charge metals) and covalent compounds (Greek prefixes)
- Lewis structures: count valence electrons, octet rule, lone vs bonding pairs, multiple bonds
Intermolecular Forces, Chemical Reactions & the Mole
Polarity & IMF, balancing equations, precipitation/net-ionic reactions, the mole concept, molar mass and empirical formula (Wks 3-4)Two big ideas meet in this chapter. First, intermolecular forces — dispersion, dipole–dipole and hydrogen bonding — explain bulk properties such as boiling point once you can judge a molecule's polarity.
- Polarity: polar bonds (electronegativity) vs polar molecules (geometry); why CO<sub>2</sub> and CH<sub>4</sub> are non-polar
- Intermolecular forces ranked weakest → strongest: dispersion < dipole–dipole < hydrogen bonding (H to N, O or F)
- IMF strength and bulk properties (boiling point, viscosity)
- Chemical vs physical change; writing and balancing equations with state symbols
- Precipitation reactions: molecular → complete-ionic → net-ionic; identifying spectator ions
- The mole: n = m/M, molar mass M, Avogadro's number to count particles
Stoichiometry, Solutions & Gases
Mass-mass stoichiometry, limiting reagents, percentage yield, significant figures, molarity, the gas laws and ideal-gas calculations (Wks 4-5)This is the most calculation-heavy and most-tested chapter.
- Mass-mass roadmap: mass A → n(A) ÷ M → n(B) × mole ratio → mass B × M
- Limiting reagent: divide each reactant's moles by its coefficient; the smallest is limiting
- Percentage yield = (actual yield / theoretical yield) × 100
- Significant figures: × ÷ take the fewest sig figs; + − take the fewest decimal places
- Solubility (g per 100 g water); saturated, unsaturated and supersaturated solutions
- Molarity c = n/V (mol L<sup>−1</sup>); dilution c<sub>1</sub>V<sub>1</sub> = c<sub>2</sub>V<sub>2</sub>; solution/titration stoichiometry
Thermochemistry, Calorimetry & Hess's Law
Energy, heat and work; enthalpy; calorimetry (q = mcDeltaT) and Hess's Law (Wks 5-6)Energy bookkeeping for chemical reactions. You define heat, work and enthalpy, classify reactions as exothermic (ΔH < 0) or endothermic (ΔH > 0), and use the first law ΔU = q + w.
- Kinetic vs potential energy; heat (q) vs temperature; work (w); system vs surroundings
- Exothermic (ΔH < 0, heat released) vs endothermic (ΔH > 0, heat absorbed)
- First law of thermodynamics: ΔU = q + w
- Enthalpy change ΔH of reaction
- Calorimetry: q = mcΔT for a solution and q = C<sub>cal</sub>ΔT for a calorimeter
- Specific heat of water = 4.18 J g<sup>−1</sup> °C<sup>−1</sup>
Redox & Electrochemistry
Oxidation numbers, balancing redox half-equations, galvanic cells and EMF, and electrolysis (Wks 6-7)Reactions that move electrons. You assign oxidation numbers, identify oxidant and reductant, and balance redox half-equations in acidic solution.
- Oxidation Is Loss, Reduction Is Gain of electrons (OIL RIG); oxidant is reduced, reductant is oxidised
- Assigning oxidation numbers; a change in oxidation number signals a redox reaction
- Balancing acidic redox: balance atoms, then O with H<sub>2</sub>O, H with H<sup>+</sup>, charge with e<sup>−</sup>; combine half-equations
- Galvanic (voltaic) cell: anode = oxidation, cathode = reduction; salt bridge keeps electroneutrality
- Electron flow through the external wire from anode → cathode; cation migration to the cathode
- E°<sub>cell</sub> = E°<sub>reduction</sub> + E°<sub>oxidation</sub>; E°<sub>cell</sub> > 0 means spontaneous
Reaction Kinetics
Reaction rate, factors affecting rate, activation energy and catalysts (Wk 7)How fast reactions go and what changes that speed.
- Reaction rate = change in concentration per unit time, rate = −Δ[reactant]/Δt = +Δ[product]/Δt
- Stoichiometric scaling of rates between reactants and products
- Factors affecting rate: concentration, temperature, surface area, catalyst
- Collision theory: effective collisions need correct orientation and energy ≥ E<sub>a</sub>
- Activation energy E<sub>a</sub> as the energy barrier to reaction
- Temperature effect: higher T means more molecules exceed E<sub>a</sub>, so faster reaction
Chemical Equilibrium
Dynamic equilibrium, Kc, the reaction quotient Q and Le Chatelier's Principle (Wk 8)Reversible reactions settle into a dynamic equilibrium where forward and reverse rates are equal and concentrations stay constant.
- Dynamic equilibrium: forward rate = reverse rate; concentrations constant (not necessarily equal)
- Equilibrium constant K<sub>c</sub> = [products]^coeff / [reactants]^coeff
- Pure solids and liquids omitted from K (heterogeneous equilibria)
- Magnitude of K and the extent of reaction (large K favours products)
- Reaction quotient Q: same form as K, evaluated at any instant
- Comparing Q to K: Q < K favours forward, Q > K favours reverse, Q = K is equilibrium
Acids, Bases & Aqueous Equilibria
Bronsted-Lowry acids/bases, pH/pOH, weak-acid equilibria with ICE tables, buffers and solubility (Ksp) (Wks 8-9)Equilibrium applied to water chemistry. You define acids and bases by the Brønsted–Lowry model, identify conjugate pairs, and use the pH/pOH scale with Kw.
- Brønsted–Lowry acids (proton donors) and bases (proton acceptors); conjugate acid–base pairs; amphiprotic species
- Self-ionisation of water: K<sub>w</sub> = [H<sub>3</sub>O<sup>+</sup>][OH<sup>−</sup>] = 1.0 × 10<sup>−14</sup> at 25 °C
- pH = −log<sub>10</sub>[H<sub>3</sub>O<sup>+</sup>], pOH = −log<sub>10</sub>[OH<sup>−</sup>], pH + pOH = 14.0
- Strong acids/bases: pH directly from concentration (full dissociation)
- Weak acids/bases: K<sub>a</sub> / K<sub>b</sub>, pK<sub>a</sub> = −log K<sub>a</sub>, pK<sub>a</sub> + pK<sub>b</sub> = 14.0
- ICE-table method for weak-acid pH and the small-x approximation; polyprotic acids
Organic Chemistry I: Alkanes, Isomerism & Stereochemistry
Alkane structure and IUPAC naming, structural isomers, Newman conformations, chirality and R/S, and the functional-group map (Wk 10)The first organic week builds the structural toolkit.
- Alkanes C<sub>n</sub>H<sub>2n+2</sub>; tetrahedral sp<sup>3</sup> bonding in CH<sub>4</sub> and C<sub>2</sub>H<sub>6</sub>; condensed and skeletal drawing
- Structural (constitutional) isomers: same formula, different connectivity
- IUPAC nomenclature: longest chain as parent, lowest locants, alphabetical substituents
- Newman projections: staggered vs eclipsed conformations (interconvert by rotation)
- Stereoisomerism vs conformation: configurational isomers need bond breaking
- Chirality, stereocentre, enantiomers and racemates
Organic Chemistry II: Alkenes, Alkynes & Aromatics
Cycloalkanes, cis-trans and E/Z isomerism, alkenes and alkynes, polyenes and benzene (Wk 11)The final week extends organic structure to unsaturated and cyclic systems.
- Cycloalkanes: 'cyclo' nomenclature
- Geometrical (cis–trans) isomerism in rings and across C=C double bonds
- Restricted rotation about a C=C double bond as the cause of cis–trans isomerism
- Alkenes C<sub>n</sub>H<sub>2n</sub>: sp<sup>2</sup> bonding and the π bond in ethene; −ene nomenclature with lowest locants
- E/Z isomerism assigned by CIP priorities on each sp<sup>2</sup> carbon
- Polyenes: conjugation and colour (e.g. carotenoids)
The whole subject is examined in one closed-book exam worth 60% — Section A multiple choice (60 marks) plus Section B extended written answers (60 marks), 2 hours of writing plus 15 minutes reading time — with an official appendices booklet (constants, formula sheet, Periodic Table) supplied, so marks come from choosing the right relationship and substituting correctly with significant figures, not from memorising formulae.
Note the hurdle: you must satisfactorily complete the practical work and attend a minimum of 4 of the 6 lab classes to pass the subject.
Worked example · free
Section B archetype — Redox titration: percentage of iron in an ore
- 1 mark — correct reductant identifiedFe2+ is oxidised to Fe3+ (loses an electron), so Fe2+ is the reductant; MnO4− is reduced and is the oxidant.
- 1 mark — n = cV with volume in litresn(MnO4−) = cV = 0.0200 mol L−1 × 0.02400 L = 4.80 × 10−4 mol.
- 1 mark — applies the 5:1 ratioMole ratio Fe2+ : MnO4− = 5 : 1, so n(Fe2+) = 5 × 4.80 × 10−4 = 2.40 × 10−3 mol.
- 1 mark — mass of iron via m = nMm(Fe) = n × M = 2.40 × 10−3 mol × 55.85 g mol−1 = 0.134 g.
- 1 mark — correct percentage%Fe = (mass of Fe ÷ mass of sample) × 100 = (0.134 ÷ 0.450) × 100 = 29.8 %.
- 1 mark — significant figures + sanity checkReport to 3 significant figures: the inputs (0.0200 M, 24.00 mL, 0.450 g) carry 3 sig figs, so the answer is 29.8 % Fe — a sensible value for an iron ore.
Key terms
- Mole (n)
- The amount of substance containing Avogadro's number (6.0221 × 1023) of particles. Linked to mass by n = m/M and to solution concentration by n = cV — the hinge of every stoichiometry calculation.
- Limiting reagent
- The reactant that runs out first and so caps the amount of product. Found by dividing each reactant's moles by its balancing coefficient; the smallest value is limiting.
- Enthalpy change (ΔH)
- The heat exchanged at constant pressure. ΔH < 0 is exothermic (heat released); ΔH > 0 is endothermic. Measured by calorimetry via q = mcΔT and combined with Hess's Law.
- Standard cell potential (E°cell)
- The voltage of a galvanic cell under standard conditions, E°cell = E°reduction + E°oxidation. A positive E°cell means the redox reaction is spontaneous.
- pH
- A measure of acidity, pH = −log10[H3O+]. With pH + pOH = 14.0 at 25 °C, it links to pOH, Kw and (for weak acids) Ka.
- Significant figures
- The reliable digits in a measurement. For × and ÷ the answer takes the fewest sig figs of the inputs; for + and − it takes the fewest decimal places. Explicitly marked in Section B.
CHEM10007 FAQ
Is the CHEM10007 exam open or closed book?
Closed-book. You may not bring cheat sheets, notes or textbooks, but an appendices booklet (physical constants, a chemical-relationships formula sheet, selected Ka/Kb/Ksp and standard reduction potentials, and a Periodic Table) is supplied. An FX82 calculator and an unassembled molecular model kit are allowed.
How long is the exam and how is it structured?
It is a 2-hour written exam with 15 minutes of reading time (no writing during reading). Section A is multiple choice (60 marks, write the capital letter in the box) and Section B is extended written answers (60 marks). The Subject Overview table lists 2.5 hours; the dedicated Exam Information page states 2 hours writing + 15 minutes reading — check the official timetable for your offering.
Do I have to pass the labs to pass the subject?
Yes. Practical work is a hurdle: you must attend a minimum of 4 of the 6 practical classes and satisfactorily complete (pass) the practical reports overall. A lab report cannot be submitted without doing the experiment, and a pre-lab quiz must be completed before each class to gain entry.
Do I need prior chemistry to take CHEM10007?
No. The welcome page states the subject requires no prior knowledge of chemistry — it starts from the nature of matter and builds the full first-year toolkit over 11 weeks of dual-delivery (online plus on-campus) teaching.
What textbook does CHEM10007 use?
Chemistry: Core Concepts, 3rd Edition (Blackman, Southam, Lawrie, Williamson & Thompson, Wiley). All weekly readings reference 3rd-edition section numbers; earlier editions are not referenced.
Questions students ask in each CHEM10007 chapter
- What are the main limitations of the Bohr model?
- It cannot account for electron–electron repulsion in multi-electron atoms, cannot predict the spectra of atoms larger than hydrogen or magnetic behaviour, and (classically) a circulating electron should radiate energy and spiral into the nucleus. The quantum-mechanical model replaces fixed orbits with probability orbitals. Chapter 01 · Atoms, the Periodic Table & Electronic Structure
- How do I count valence electrons for an ion?
- Sum the valence electrons of every atom, then add one electron for each unit of negative charge or subtract one for each unit of positive charge. For NO3− that is 5 + 18 + 1 = 24. Chapter 02 · Chemical Bonding, Lewis Structures & Shape
- How do I find the strongest intermolecular force in a substance?
- First decide if the molecule is polar. If it has H bonded to N, O or F it can hydrogen bond; if it is polar but cannot, the strongest force is dipole–dipole; if it is non-polar, only dispersion forces act. All molecules have dispersion forces underneath. Chapter 03 · Intermolecular Forces, Chemical Reactions & the Mole
- How do I decide which reactant is limiting?
- Convert each reactant's mass to moles, then divide each by its balancing coefficient. The reactant giving the smallest value is limiting; use its moles for all product calculations. The other reactant is in excess. Chapter 04 · Stoichiometry, Solutions & Gases
- What is the difference between heat and temperature?
- Temperature measures the average kinetic energy of particles; heat (q) is the energy transferred because of a temperature difference. A small mass at high temperature can carry less heat than a large mass at lower temperature. Chapter 05 · Thermochemistry, Calorimetry & Hess's Law
- How do I balance a redox reaction in acidic solution?
- Split into oxidation and reduction half-equations. In each, balance the main atoms, then balance O with H2O, balance H with H+, and balance charge with electrons. Scale so electrons lost equal electrons gained, then add the halves and cancel. Chapter 06 · Redox & Electrochemistry
- Why does raising the temperature speed up a reaction?
- Higher temperature widens the energy distribution so a larger fraction of molecules have energy at or above Ea. There are more frequent and more energetic effective collisions per second, increasing the rate. Chapter 07 · Reaction Kinetics
- What is the difference between Q and K?
- They have the identical algebraic form, but K is the fixed value at equilibrium (at a given temperature) while Q is calculated for the actual concentrations at any instant. Comparing Q to K tells you which way the reaction must shift. Chapter 08 · Chemical Equilibrium
- When can I use the small-x approximation?
- When the acid is weak and reasonably concentrated, x (the amount ionised) is tiny compared with the initial concentration, so 0.200 − x ≈ 0.200. Always check afterwards: if x exceeds about 5 % of the initial value, redo the calculation with the full quadratic. Chapter 09 · Acids, Bases & Aqueous Equilibria
- What is the difference between a conformation and a configurational isomer?
- Conformations (staggered, eclipsed) interconvert freely by rotation about a single bond, so they are the same molecule. Configurational isomers (such as enantiomers or cis/trans) can only interconvert by breaking and remaking bonds, so they are genuinely different compounds. Chapter 10 · Organic Chemistry I: Alkanes, Isomerism & Stereochemistry
- When does cis–trans isomerism occur?
- It occurs when there is restricted rotation — across a C=C double bond or within a ring — and each of the relevant carbons carries two different groups. Free rotation about single bonds (as in alkanes) prevents it. Chapter 11 · Organic Chemistry II: Alkenes, Alkynes & Aromatics
How to study for the exam
Treat the appendices booklet as a working tool, not a crutch: print the formula sheet early and practise choosing the right relationship for each question type, because the marks live in method and significant figures, not recall.
Build fluency in the high-yield Section B archetypes first — limiting reagent / % yield, redox titration, weak-acid pH via an ICE table, calorimetry (q = mcΔT linked to moles and ΔH), and gas-law plus pV = nRT manipulation — since these recur across the mock and practice exams.
Use the 5 Feedback Quizzes as low-stakes diagnostics across the semester rather than cramming them, and protect the lab hurdle by completing every pre-lab quiz and attending well above the 4-class minimum. In the final fortnight, rehearse the set up → substitute → sig figs → sanity-check rhythm on past-style questions, and keep T in kelvin and volumes in litres as automatic reflexes.
