DPST1031 Chemistry A: Atoms, Molecules and Energy
The Complete Exam Bible · T3 2026DPST1031 Overview
- 6 units of credit
- UNSW Sydney
- Term 3 2026
- Level 1 diploma course
- In person final exam
Chemistry A: Atoms, Molecules and Energy is the first chemistry course in the UNSW College Diploma of Science and Diploma of Engineering pathways, delivered on behalf of UNSW Sydney and taught in Term 3 2026 over twelve teaching weeks plus a study week.
- Assessed by Labs, eight quizzes, two in-term tests and a 40% final
- Hardest step Choosing the right equilibrium or thermodynamic relation before substituting
- How to prepare Solve mastery-style problems with only the data sheet beside you
What you get in the full DPST1031 Exam Bible
- 13 chapters, 72 pages — every examinable topic, week by week
- 32 original practice questions, each with a model answer in the Final Examination format
- 13 worked examples, solved the way DPST1031 marks
- 22 computed figures, redrawn from the course’s own numbers
- Exam-morning plan for the Final Examination (40%) — what to revise last, where marks are lost
Free on this page: assessment table, chapter map, one worked example, key terms, FAQ. The rest is in the full exam bible.
How DPST1031 is assessed
| Component | Weight | Format |
|---|---|---|
| Laboratory Work: Hurdle Tasks · hurdle | 10% | Core laboratory skills demonstrated during lab classes; every core skill must be awarded to pass the course |
| Laboratory Work: Graded Tasks | 10% | Marks for results and task questions in the lab experiments |
| Online Quizzes | 8% | Eight threshold quizzes in Moodle; 1% each once a 10/10 attempt is recorded within the two-week window |
| In-term Tests (two tests) | 32% | 16% each; 45-minute paper tests in person on threshold content; best of two attempts, second attempt capped at 75% |
| Final Examination | 40% | In person in the exam period; focuses on mastery content; multiple choice and written response |
These weights come from the assessment table on the Term 3 2026 Moodle course page, and they match the course outline (laboratory 20%, quizzes 8%, in-term tests 32%, final 40%). The UNSW Handbook entry for DPST1031 lists a different split (final examination 60%, laboratory work 20%, mid-session test 10%, computer assignment 10%); follow the Moodle table and course outline for this offering and check Moodle if the two disagree. To pass, the outline also requires at least 80% attendance at lectures and tutorials, at least six laboratory classes, every core laboratory skill, and a course mark of at least 50%.
What other students are asking about DPST1031
- Where can I find lecture notes or study guides for DPST1031 at UNSW?
- The DPST1031 study guide in AskSia Library covers Chemistry A: Atoms, Molecules and Energy in 13 chapters, with the assessment table, a chapter-by-chapter topic map, key terms and a worked example, free to read on this page.
- Is there a final exam in DPST1031, and how much is it worth?
- Yes. In the offering this AskSia Library guide is built from, Final Examination carries 40% of the DPST1031 mark; the full split is Laboratory Work: Hurdle Tasks 10%, Laboratory Work: Graded Tasks 10%, Online Quizzes 8%, In-term Tests (two tests) 32%, Final Examination 40%.
- What does DPST1031 cover on quantum numbers and electron configurations?
- In AskSia Library's DPST1031 guide, the Quantum Numbers and Electron Configurations chapter explains orbitals and nodes, the four quantum numbers, Pauli, Hund and Aufbau, configurations of atoms and ions, paramagnetism.
- What does DPST1031 cover on periodic trends and effective nuclear charge?
- In AskSia Library's DPST1031 guide, the Periodic Trends and Effective Nuclear Charge chapter explains shielding and effective charge, atomic and ionic radius, ionisation energy, electron affinity, electronegativity and the anomalies.
- What does DPST1031 cover on lewis structures, formal charge and vsepr?
- In AskSia Library's DPST1031 guide, the Lewis Structures, Formal Charge and VSEPR chapter explains valence electron counts, octet exceptions, formal charge, resonance, electron domains, molecular shape and bond angles.
Current DPST1031 dates
| Date | Item | Control |
|---|---|---|
| 5 October 2026 | In-term Test 1, first attempt | Week 6 of the Term 3 2026 schedule; threshold topics 1 to 4 |
| 19 October 2026 | In-term Test 1, second attempt | Week 8 of the Term 3 2026 schedule; mark capped at 75% |
| 9 November 2026 | In-term Test 2, first attempt | Week 11 of the Term 3 2026 schedule; threshold topics 5 to 8 |
| 16 November 2026 | In-term Test 2, second attempt | Week 12, which is also the exam revision week |
| 23 November 2026 | Study week | Week 13; the final exam time and place are released on Moodle |
Dates are as published in the Term 3 2026 course schedule on Moodle (dates are the Monday that starts each week). Confirm exact deadlines and submission settings in the live LMS.
What DPST1031 covers
The course climbs from the quantum structure of single atoms to bonding, intermolecular forces and gases, then applies equilibrium to acids and bases, thermodynamics to spontaneity, and both to electrochemical cells.
Light, Quantisation and the Hydrogen Atom
Photon energy, wavelength and frequency, hydrogen energy levels, emission and absorption lines, the Rydberg equation and ionisationThis chapter builds the quantum picture that the rest of the course depends on. Light is treated both as a wave, with wavelength, frequency and speed linked by c = λν, and as a stream of photons carrying E = hν.
- Wavelength, frequency and the speed of light
- Photon energy and Planck's constant
- The photoelectric effect and work function
- De Broglie wavelength of moving particles
- Hydrogen energy levels and the principal quantum number
- Emission, absorption and ionisation
Quantum Numbers and Electron Configurations
Orbitals and nodes, the four quantum numbers, Pauli, Hund and Aufbau, configurations of atoms and ions, paramagnetismAtomic orbitals emerge from the Schrödinger equation, and four quantum numbers describe every electron: n sets the level, l the subshell shape, the magnetic number the orientation and the fourth number the spin.
- Atomic orbitals and electron probability
- The four quantum numbers and allowed values
- Counting orbitals and nodes
- Aufbau principle and filling order
- Pauli exclusion principle
- Hund's rule and arrows in boxes
Periodic Trends and Effective Nuclear Charge
Shielding and effective charge, atomic and ionic radius, ionisation energy, electron affinity, electronegativity and the anomaliesPeriodic trends are the payoff of electron configurations.
- Shielding and effective nuclear charge
- Atomic radius across periods and down groups
- Cation and anion radii
- Isoelectronic series
- First ionisation energy and its dips
- Successive ionisation energies
Lewis Structures, Formal Charge and VSEPR
Valence electron counts, octet exceptions, formal charge, resonance, electron domains, molecular shape and bond anglesLewis diagrams turn a formula into a picture of shared and unshared valence electrons, and VSEPR theory turns that picture into a three-dimensional shape.
- Counting valence electrons for molecules and ions
- Lewis diagrams of organic functional groups
- Octet exceptions and expanded valence shells
- Formal charge and the preferred structure
- Resonance and delocalisation
- Electron domains and electron geometry
Valence Bond, Hybrid and Molecular Orbitals
Orbital overlap, sigma and pi bonds, sp, sp2 and sp3 hybrids, molecular orbital diagrams, bond order and magnetismLewis structures say where electron pairs are; this chapter explains why the bonds form and how strong they are.
- Bond formation and the potential energy curve
- Sigma and pi bonds from orbital overlap
- Hybrid orbitals: sp, sp² and sp³
- Hybridisation in organic molecules
- Bonding and antibonding molecular orbitals
- MO diagrams for second-row diatomics
Intermolecular Forces and Liquid Properties
Bond and molecular polarity, dispersion, dipole and hydrogen bonding, boiling point, viscosity, surface tension, miscibilityThis chapter connects the structure of a molecule to the way a bulk substance behaves. It begins with bond polarity from electronegativity difference and uses VSEPR shapes to decide whether bond dipoles cancel.
- Electronegativity and bond polarity
- Molecular polarity from bond dipoles
- Polarizability and dispersion forces
- Dipole-dipole forces
- Hydrogen bonding requirements
- Ranking boiling and melting points
Ideal Gases, Real Gases and Raoult's Law
pV = nRT and Dalton's law, van der Waals corrections, vapour pressure, Raoult's law, deviations and Henry's lawGases and liquid mixtures are where intermolecular forces become numbers.
- The ideal gas law and the two values of R
- Pressure, volume and temperature conversions
- Dalton's law and mole fractions
- Real gases and the van der Waals equation
- Vapour pressure and liquid-vapour equilibrium
- Raoult's law for binary mixtures
Equilibrium Constants, Reaction Quotients and ICE Tables
K expressions, Q versus K, ICE tables and the small x test, Kc and Kp, combining reactions, Le Chatelier and van 't HoffChemical equilibrium is the quantitative core of the second half of the course.
- Dynamic equilibrium and equal rates
- Writing Kc for homogeneous and heterogeneous reactions
- The reaction quotient Q
- K for reversed, scaled and combined reactions
- Converting Kc and Kp
- ICE tables
Acids, Bases, Buffers and Titration Curves
Brønsted-Lowry and Lewis acids, pH and Kw, weak acid and base pH, salt solutions, Henderson-Hasselbalch buffers, titration curvesAcids and bases apply the equilibrium methods of the previous chapter to the reactions students meet most often.
- Brønsted-Lowry and Lewis definitions
- Conjugate acid and base pairs
- Strong versus weak acids and bases
- Kw, pH and pOH
- Weak acid and weak base pH
- Acidic, basic and neutral salts
Heat, Enthalpy and Hess's Law
System and surroundings, q, w and the first law, calorimetry, enthalpy and ΔU, formation enthalpies, Hess's law and Born-HaberThermochemistry measures the energy that chemical change exchanges with its surroundings.
- System, surroundings and universe
- The first law and the sign of q and w
- Heat capacity and calorimetry
- Constant-volume and constant-pressure heat
- Enthalpy versus internal energy
- Standard enthalpy changes and standard states
Entropy, Gibbs Energy and Spontaneity
Entropy as energy dispersal, standard entropies, the second and third laws, ΔG = ΔH − TΔS, ΔG° and K, temperature dependenceThis chapter answers the question enthalpy cannot: will a change happen by itself?
- Entropy as energy dispersal
- The second and third laws
- Predicting the sign of ΔS
- Standard reaction entropy from tables
- Entropy of the surroundings and universe
- Gibbs energy and spontaneity
Redox, Galvanic Cells and Electrolysis
Oxidation numbers, half-equation balancing, cell diagrams, standard potentials, E° with ΔG° and K, Nernst, corrosion, FaradayElectrochemistry brings the course together: electron transfer, equilibrium and Gibbs energy in one system.
- Oxidation numbers and their rules
- Oxidising and reducing agents
- Balancing half-equations in acid and base
- Galvanic cells, anode and cathode
- Cell diagrams and the salt bridge
- Standard reduction potentials and E°cell
Laboratory Core Skills and Chemical Analysis
Weighing by difference, standard solutions, pipette and titration technique, titre accuracy, data sheets and report skillsThe laboratory program is worth 20% of the course and carries the one hurdle that can fail a student outright: every core laboratory skill must be awarded by the end of term.
- Core and non-core laboratory skills
- How the hurdle and graded marks combine
- Attendance, make-up labs and pre-lab work
- Weighing by difference
- Preparing a standard solution
- Pipette and burette technique
It assumes Year 12 chemistry and starts by asking why atoms absorb and emit light only at particular wavelengths. That quantum picture of hydrogen leads to orbitals, quantum numbers and electron configurations, and from there to the periodic trends in radius, ionisation energy, electron affinity and electronegativity. The middle of the course turns structure into behaviour.
Lewis structures, formal charge and VSEPR predict molecular shape; valence bond theory, hybrid orbitals and molecular orbital diagrams explain why bonds form and how strong they are; polarity and intermolecular forces then account for boiling points, viscosity, surface tension and the vapour pressure of mixtures through Raoult's law. Gases are handled with the ideal gas law and the van der Waals correction.
The second half is quantitative. Chemical equilibrium introduces K, Q and ICE tables, which are then applied to weak acids, bases, buffers and titration curves. Thermochemistry builds enthalpy, entropy and Gibbs energy, links the sign of ΔG° to the size of K, and finishes with redox chemistry, galvanic cells, the Nernst equation, batteries and electrolysis. Every topic is written in two layers.
Threshold outcomes are the pass-level skills checked by the online quizzes and the two in-term tests; Mastery outcomes combine ideas across topics and are the focus of the final examination. This guide follows that split chapter by chapter, works each standard calculation from the data sheet constants, and flags the reasoning steps where marks are usually lost.
Worked example · free
Buffer and equivalence pH in an acetic acid titration
- +1Moles: acid 2.50 mmol; 12.5 mL of base adds 1.25 mmol, converting half the acid to acetate.
- +1With equal moles of acid and acetate, Henderson-Hasselbalch gives pH = pKa + log(1) = pKa = −log(1.8 × 10⁻⁵) = 4.74.
- +1At equivalence, 25.0 mL of base has been added: 2.50 mmol of acetate in 50.0 mL, so [acetate] = 0.0500 M.
- +1Kb = Kw ÷ Ka = 5.6 × 10⁻¹⁰, so [OH⁻] = √(0.0500 × 5.6 × 10⁻¹⁰) = 5.3 × 10⁻⁶ M and pOH = 5.28.
- +1pH = 14.00 − 5.28 = 8.72, above 7 because acetate is a weak base.
Key terms
- Photon
- A packet of electromagnetic energy whose energy equals Planck's constant multiplied by its frequency.
- Principal quantum number
- The integer n that sets the energy level and overall size of an orbital in a hydrogen-like atom.
- Effective nuclear charge
- The net positive charge felt by an electron after shielding by the electrons between it and the nucleus.
- Formal charge
- The charge an atom would carry in a Lewis structure if bonding electrons were shared exactly equally.
- Electron domain
- A region of electron density around a central atom, counting a lone pair or any bond as one domain.
- Hybrid orbital
- A mixed atomic orbital, such as sp3, used in valence bond theory to explain observed molecular shapes.
- Bond order
- Half the difference between bonding and antibonding electrons, used to compare bond strength and length.
- Hydrogen bond
- A strong dipole attraction between an H atom bonded to N, O or F and a lone pair on another N, O or F.
- Vapour pressure
- The pressure exerted by a vapour in dynamic equilibrium with its liquid at a given temperature.
- Reaction quotient
- The ratio Q built like the equilibrium expression but from current concentrations, showing which way a reaction will shift.
- Buffer
- A solution that resists pH change when small amounts of acid or base are added, made from a weak acid and the base it forms on losing a proton.
- Standard enthalpy of formation
- The enthalpy change when one mole of a compound forms from its elements in their standard states.
- Gibbs energy
- The state function G = H − TS whose change at constant temperature and pressure decides spontaneity.
- Standard reduction potential
- The voltage of a half-cell written as a reduction, measured against the standard hydrogen electrode.
DPST1031 FAQ
How is this chemistry course assessed in Term 3 2026?
The Moodle assessment table gives laboratory work 20%, split between hurdle tasks and graded tasks, eight online quizzes 8%, two in-term tests 32% and a final examination 40%. The UNSW Handbook shows a different split, so treat the Moodle table and outline as the operative version and confirm on Moodle if anything changes before the exam period.
What is the difference between Threshold and Mastery content?
Threshold outcomes are the pass-level skills, such as calculating a photon energy or drawing a Lewis structure, and they are checked by the quizzes and in-term tests. Mastery outcomes go deeper or combine topics, for example using the Rydberg equation with ionisation, or linking ΔG° to K and E°. The final examination focuses on Mastery content but still assumes every Threshold skill.
Can I resit an in-term test?
Yes. Each in-term test has two sittings and the better mark counts, but the second attempt is capped at 75% of the marks. In Term 3 2026 the first sittings fall in Weeks 6 and 11 and the second sittings in Weeks 8 and 12, so plan to be fully prepared for the first attempt.
What does the final exam focus on?
The outline describes a two-hour paper with ten minutes of reading time worth 40% of the course, mixing multiple-choice and written-response questions and focusing on Mastery content. The Term 3 Moodle page says exact time, location and format details are released later in term, so check that notice once it is posted.
Why do the laboratory core skills matter so much?
Laboratory core skills are all-or-nothing. Once every core skill has been awarded you receive the hurdle 10% in one lump, and the graded skills add up to another 10%. Missing even one core skill means failing the laboratory component and therefore the course, so use the make-up labs if a skill is still outstanding.
Which formulas are on the data sheet?
The course data sheet lists the gas constant in two units, Avogadro's number, the Faraday constant and equation, the Nernst equation, the Planck and Rydberg equations, the Gibbs relation with ΔH and TΔS, the link between ΔG°, K and E°, van 't Hoff, Henderson-Hasselbalch, van der Waals and Kp = Kc(RT)Δn. You still need to know when each one applies.
Which calculations cost students the most marks?
The recurring trouble spots are ICE tables where the small x approximation is not justified, buffer problems that skip the mole stoichiometry step, sign errors in ΔH, ΔS and ΔG, unit slips between kJ and J or °C and K, and choosing the wrong half-equation direction in a cell. Each chapter of this guide drills one of these with fresh numbers.
Questions students ask in each DPST1031 chapter
- Why does brighter light not eject electrons below the threshold frequency?
- Brightness only increases the number of photons. Each photon still carries E = hν, and one photon gives its energy to one electron, so if hν is below the work function no single electron ever receives enough energy to leave the metal. Chapter 01 · Light, Quantisation and the Hydrogen Atom
- Why does 4s fill before 3d but lose electrons first?
- In a neutral atom the 4s orbital is lower in energy while it fills, because s electrons penetrate close to the nucleus. Once 3d electrons are present they shield and reorder the energies, so the 4s electrons become the outermost and are the first removed when a cation forms. Chapter 02 · Quantum Numbers and Electron Configurations
- Why does atomic radius decrease across a period?
- Across a period electrons are added to the same shell while protons are added to the nucleus. Electrons in one shell shield each other poorly, so the effective nuclear charge rises and the whole shell is pulled closer. The radius therefore shrinks from the group 1 metal to the noble gas. Chapter 03 · Periodic Trends and Effective Nuclear Charge
- How do I decide which Lewis structure is best?
- Calculate the formal charge of every atom in each candidate. Prefer the structure with charges closest to zero, and if a negative formal charge is unavoidable, place it on the most electronegative atom. The charges must always add up to the overall charge of the species, which is a quick arithmetic check. Chapter 04 · Lewis Structures, Formal Charge and VSEPR
- How do I find the hybridisation of an atom?
- Count the electron domains around it using VSEPR, counting lone pairs and each bond once. Two domains means sp, three means sp² and four means sp³. The number of unhybridised p orbitals left over tells you how many π bonds that atom can form. Chapter 05 · Valence Bond, Hybrid and Molecular Orbitals
- Is a polar molecule always the one with the higher boiling point?
- No. For molecules of similar size, polarity raises the boiling point, but a large non-polar molecule can have dispersion forces strong enough to beat a small polar one. Octane boils above water's boiling point, so always compare sizes as well as polarity. Chapter 06 · Intermolecular Forces and Liquid Properties
- Which value of R should I use?
- Use 0.08206 L atm mol⁻¹ K⁻¹ when pressure is in atmospheres and volume in litres. Use 8.314 J mol⁻¹ K⁻¹ when pressure is in pascals and volume in cubic metres, or whenever an energy is involved. Temperature must always be in kelvin with either value. Chapter 07 · Ideal Gases, Real Gases and Raoult's Law
- What is the difference between Q and K?
- Both use the same expression. K uses equilibrium concentrations and is a constant at a given temperature; Q uses the concentrations present at any moment. If Q is smaller than K the reaction moves forward, if larger it moves backward, and if equal the system is already at equilibrium. Chapter 08 · Equilibrium Constants, Reaction Quotients and ICE Tables
- What is the difference between a strong acid and a concentrated acid?
- Strength describes how completely an acid ionises in water; concentration describes how much acid is dissolved per litre. A strong acid ionises completely at any concentration, while a weak acid ionises only partly even when concentrated, so the two ideas must be judged separately. Chapter 09 · Acids, Bases, Buffers and Titration Curves
- What sign convention does the course use for work?
- Signs are taken from the system's point of view. Heat absorbed by the system and work done on the system are positive; heat released and work done by the system are negative. With that convention the first law is written ΔU = q + w. Chapter 10 · Heat, Enthalpy and Hess's Law
- Why can an endothermic reaction be spontaneous?
- Spontaneity depends on ΔG = ΔH − TΔS, not on ΔH alone. If an endothermic reaction has a large enough positive ΔS, the −TΔS term outweighs ΔH at high temperature and ΔG becomes negative. Melting ice above 0 °C and decomposing limestone in a kiln are examples. Chapter 11 · Entropy, Gibbs Energy and Spontaneity
- How do I know which electrode is the anode?
- In a galvanic cell, the half-cell with the lower standard reduction potential is reversed and becomes the anode, where oxidation happens. The half-cell with the higher potential is the cathode, where reduction happens. Oxidation is always at the anode in every type of cell. Chapter 12 · Redox, Galvanic Cells and Electrolysis
- What happens if I miss a core laboratory skill?
- Core skills are all or nothing. If any core skill has not been awarded by the end of term, you fail the laboratory component and therefore the course. There are several chances to show each skill, and make-up laboratory classes give a final opportunity, so use them for any skill still outstanding. Chapter 13 · Laboratory Core Skills and Chemical Analysis
How to study for the exam
Work through the course in the order it is examined. Before In-term Test 1, make sure every Threshold skill in the first four topics is automatic: photon calculations, configurations, Lewis structures and shapes, and intermolecular force identification. Do the online quiz until it reaches 10/10, then attempt the tutorial questions without notes.
Before In-term Test 2, repeat the routine for equilibrium, acids and bases, thermochemistry and electrochemistry, where most marks come from setting up the calculation correctly. For the final, switch to Mastery practice. Sit with only the data sheet and a calculator, choose a mixed set of problems, and write the relationship you are using before you substitute any number.
After each problem, check units, significant figures and signs, then write one sentence explaining the chemistry behind the answer. Finally, connect topics: practise questions that join bonding to intermolecular forces, equilibrium to Gibbs energy, and Gibbs energy to cell potential, because the final rewards linking ideas rather than repeating isolated steps.
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