CHEM1011 Fundamentals of Chemistry 1A

How CHEM1011 is assessed

What CHEM1011 covers

Thirteen weeks of foundational chemistry → one exam-ready map. Each topic links to its free chapter guide.

The final exam is 55% of your grade and is closed-book, and the laboratory is a separate hurdle you must pass to pass the unit, so this guide teaches each topic to exam standard: the method examiners reward, the formula they expect, and the calculation steps that earn the marks.

Worked example · free

Limiting reagent & mass of product — the workhorse calculation

Key terms

Mole (n)

The amount of substance containing Avogadro's number (6.022 × 10²³) of particles; linked to mass by n = m/M, where M is the molar mass in g mol⁻¹.

Limiting reagent

The reactant that runs out first; it caps the amount of product. Found by comparing each reactant's moles divided by its stoichiometric coefficient.

Electronegativity

The ability of a bonded atom to attract the shared electrons; it rises across a period and falls down a group, and a large difference (ΔEN) signals ionic rather than covalent bonding.

Enthalpy change (ΔH)

The heat absorbed or released at constant pressure; a negative ΔH is exothermic, a positive ΔH endothermic. It is path-independent (Hess's Law).

Equilibrium constant (K)

The ratio of product to reactant activities at equilibrium for aA + bB ⇌ cC + dD: K = [C]^c[D]^d / [A]^a[B]^b; pure solids and liquids are omitted.

pH

A measure of acidity: pH = −log₁₀[H₃O⁺]; pH + pOH = 14 at 25 °C, and at the half-equivalence point of a weak-acid titration pH = pKa.

CHEM1011 FAQ

Questions students ask in each CHEM1011 chapter

What's the difference between mass number and average atomic mass?

Mass number A is a whole-number count of protons + neutrons for one specific isotope (e.g. ³⁷Cl has A = 37). Average atomic mass is the abundance-weighted average over all of an element's natural isotopes and is usually not a whole number (chlorine ≈ 35.5). The periodic-table value is the average; a superscript on a symbol is the mass number of that one isotope. Chapter 01 · Atoms, Light and the Mole

How do I tell whether a bond is ionic, polar covalent, or non-polar covalent?

Use the electronegativity difference (ΔEN) between the two atoms. Roughly: ΔEN below ~0.4 is non-polar covalent (electrons shared evenly); ~0.4 up to ~2 is polar covalent (shared but unevenly, giving δ+/δ− partial charges); and above ~2 is ionic (electrons effectively transferred). A quick sanity check is metal + non-metal → ionic, non-metal + non-metal → covalent. Chapter 02 · Bonding, Equations and Stoichiometry

How do I find the mass of solid to make a solution of a given molarity?

Convert the target volume to litres, find moles with n = cV, then convert to mass with m = nM. In one line, m = c × V × M. Read the molar mass off the periodic table provided and report to 3 significant figures with the unit (g). Chapter 03 · Solutions and Periodic Trends

What's the 5-step method for drawing any Lewis structure?

(1) Pick the central atom — the lowest-electronegativity / highest-valence atom (never H or F). (2) Count total valence electrons by summing group numbers, then subtract for a positive charge or add for a negative one. (3) Draw a single bond from the centre to each outer atom. (4) Add lone pairs to the outer atoms first, then any leftover to the centre. (5) If the central atom is short of an octet, convert adjacent lone pairs into double or triple bonds until it reaches 8 (stop at 6 if the centre is boron). Do the same five steps every time — the marks are in the order. Chapter 04 · Lewis Structures and VSEPR

How do I know which way to number the carbon chain in IUPAC naming?

Number from the end that gives the principal functional group (or the multiple bond) the lowest locant first. Only if there is a tie do you then minimise the substituent locants. So a four-carbon chain with a terminal double bond is but-1-ene, not but-3-ene — always check both directions before committing. Chapter 05 · Organic Chemistry: Groups, Naming and Isomers

What is the difference between calculating ΔH from formation enthalpies and from bond enthalpies?

Both give a reaction enthalpy but set up oppositely. From standard enthalpies of formation you do Σn·ΔHf°(products) − Σn·ΔHf°(reactants), which gives an essentially exact value. From average bond enthalpies you do Σ(bonds broken) − Σ(bonds formed), where every bond term is a positive look-up value; because the values are averages, this route only gives an estimate. If a question supplies ΔHf° use the first method; if it supplies bond energies use the second. Chapter 06 · Thermodynamics

Why do pure solids and pure liquids get left out of K and Q?

Their activity is taken as 1, so they make no contribution to the ratio. For CaCO₃(s) ⇌ CaO(s) + CO₂(g) the expression is simply K = [CO₂] — both solids drop out. Water acting as a pure solvent is omitted too; dissolved [aq] species and gases [g] always stay in. Chapter 07 · Chemical Equilibrium

When do I use Henderson–Hasselbalch instead of an ICE table?

Use Henderson–Hasselbalch (pH = pKₐ + log([A⁻]/[HA])) whenever the flask holds appreciable amounts of BOTH a weak acid and its conjugate base — a prepared buffer, or anywhere on the buffer plateau of a weak-acid titration before equivalence. Use a full ICE table on Kₐ for the initial pH of a pure weak acid, and ICE on K_b of the conjugate base at the equivalence point. Chapter 08 · Acids, Bases and Titrations

Why does an alcohol boil higher than an aldehyde or alkane of the same mass?

Because boiling has to overcome intermolecular forces, and the alcohol's O–H lets it hydrogen-bond — the strongest of the three forces. The aldehyde only has dipole–dipole (its H is on carbon, not O), and the alkane has London dispersion alone. At similar mass the dispersion is comparable, so the force type decides: H-bond > dipole–dipole > London, and the boiling points follow the same order. Chapter 09 · Functional Groups, Fats and Polymers

How to study for the exam

Treat chemistry as a set of repeatable calculation workflows. For every topic, practise the method from a blank page: the mole map (n = m/M), the limiting-reagent comparison, the five-step Lewis build, the VSEPR angle table, the two routes to ΔH, the ICE table, and the titration region map. Memorise the methods, not the constants — the datasheet and periodic table are provided.

Because the final is 55% and closed-book and the laboratory is a hurdle, banked, well-practised calculation marks are the safest marks in the unit.