GCE Advanced Level (A-Levels) Chemistry Syllabus

Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Chemistry in GCE Advanced Level (A-Levels), not a summary of it.

1. 1

   Atomic Structure and Bonding

   3 topics

   • Atomic Structure
     • Subatomic particles and isotopes
     • Electron configuration and sub-shells
     • Mass spectrometry and ionisation energies
   • Bonding and Structure
     • Ionic, covalent and metallic bonding
     • Electronegativity and polarity
     • Intermolecular forces and hydrogen bonding
   • Shapes of Molecules
     • Electron pair repulsion theory
     • Predicting bond angles and geometry
2. 2

   Physical Chemistry

   6 topics

   • Amount of Substance
     • The mole, Avogadro constant and molar mass
     • Empirical and molecular formulae
     • Reacting masses, gas volumes and concentrations
     • Atom economy and percentage yield
   • Energetics and Thermodynamics
     • Enthalpy changes and calorimetry
     • Hess's law and Born-Haber cycles
     • Entropy and Gibbs free energy
   • Kinetics
     • Collision theory and the Maxwell-Boltzmann distribution
     • Rate equations and orders of reaction
     • Rate-determining step and the Arrhenius equation
   • Chemical Equilibria
     • Le Chatelier's principle
     • Equilibrium constants Kc and Kp
   • Acids, Bases and pH
     • Bronsted-Lowry theory and pH calculations
     • Weak acids, Ka and buffers
     • Titration curves and indicators
   • Electrochemistry
     • Redox and oxidation states
     • Electrode potentials and electrochemical cells
3. 3

   Inorganic Chemistry

   3 topics

   • Periodicity
     • Trends across Period 3
     • Atomic radius, ionisation energy and melting points
   • Group 2 and Group 7
     • Reactions and trends of the alkaline earth metals
     • Halogens, halides and reactivity trends
   • Transition Metals
     • Variable oxidation states and coloured ions
     • Complex ions, ligands and isomerism
     • Catalysis and substitution reactions
4. 4

   Organic Chemistry

   5 topics

   • Introduction to Organic Chemistry
     • Nomenclature and functional groups
     • Structural and stereoisomerism
     • Reaction mechanisms and curly arrows
   • Hydrocarbons
     • Alkanes and combustion
     • Alkenes and electrophilic addition
   • Oxygen-Containing Compounds
     • Alcohols, oxidation and dehydration
     • Aldehydes, ketones and carboxylic acids
     • Esters, acyl chlorides and amides
   • Aromatic Chemistry and Nitrogen Compounds
     • Benzene and electrophilic substitution
     • Amines and amino acids
   • Polymers and Synthesis
     • Addition and condensation polymers
     • Multi-step synthetic routes
5. 5

   Analysis and Practical Skills

   3 topics

   • Spectroscopy
     • Infrared spectroscopy
     • Mass spectrometry of organic molecules
     • Proton and carbon-13 NMR
   • Chromatography
     • Thin-layer and paper chromatography
     • Gas and high-performance liquid chromatography
   • Qualitative Analysis and Practical Techniques
     • Tests for ions and functional groups
     • Measurement uncertainty and error analysis

Chemistry flashcards for GCE Advanced Level (A-Levels)

21 of 71 cards from the Chemistry deck — real questions with worked answers.

1. Define relative atomic mass ($A_r$).

   The weighted mean mass of an atom of an element relative to $\frac{1}{12}$ of the mass of a carbon-12 atom.

2. State the number of protons, neutrons and electrons in $\ce{^{27}_{13}Al^{3+}}$.

   13 protons, 14 neutrons ($27-13$), and 10 electrons ($13-3$).

3. Write the full electron configuration of a chromium atom ($Z=24$).

   $\ce{1s^2 2s^2 2p^6 3s^2 3p^6 3d^5 4s^1}$ (one electron in 4s gives a stable half-filled 3d).

4. What is first ionisation energy?

   The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous +1 ions: $\ce{X(g) -> X+(g) + e-}$.

5. Why is there a general increase in first ionisation energy across a period?

   Nuclear charge increases while electrons are added to the same shell (similar shielding), so atomic radius decreases and the outer electron is held more strongly.

6. Define electronegativity.

   The ability of an atom to attract the bonding pair of electrons in a covalent bond (commonly measured on the Pauling scale).

7. What type of bonding holds particles together in a metal, and how is it described?

   Metallic bonding: the electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons.

8. Distinguish ionic from covalent bonding.

   Ionic bonding is electrostatic attraction between oppositely charged ions (electron transfer); covalent bonding is a shared pair of electrons between two atoms (electron sharing).

9. What is a dative (coordinate) covalent bond? Give an example.

   A covalent bond in which both shared electrons come from the same atom, e.g. the N$\to$H bond in $\ce{NH4+}$ or B$\leftarrow$N in $\ce{NH3.BF3}$.

10. List the three types of intermolecular force in order of increasing strength.

    London (induced dipole-dipole / van der Waals) < permanent dipole-dipole < hydrogen bonding.

11. Using VSEPR, give the shape and bond angle of $\ce{NH3}$.

    Trigonal pyramidal with a bond angle of about $107^\circ$ (one lone pair compresses the angle from the tetrahedral $109.5^\circ$).

12. State the shape and bond angle of a molecule with 6 bonding pairs and no lone pairs, e.g. $\ce{SF6}$.

    Octahedral, with bond angles of $90^\circ$.

13. Why is the bond angle in $\ce{H2O}$ ($104.5^\circ$) smaller than in $\ce{NH3}$ ($107^\circ$)?

    Water has two lone pairs versus ammonia's one; lone pair-lone pair repulsion is greater than lone pair-bond pair, pushing the bonding pairs closer together.

14. State Avogadro's constant and what it represents.

    $L = 6.022\times10^{23}\ \text{mol}^{-1}$, the number of particles in one mole of a substance.

15. Write the ideal gas equation, naming each term with SI units.

    $pV = nRT$, where $p$ is pressure (Pa), $V$ volume (m$^3$), $n$ moles, $R = 8.31\ \text{J K}^{-1}\text{mol}^{-1}$, and $T$ temperature (K).

16. Give the formula linking moles, concentration and volume of a solution.

    $n = c \times V$, where $c$ is concentration in $\text{mol dm}^{-3}$ and $V$ is volume in $\text{dm}^3$.

17. Define percentage atom economy.

    $\text{atom economy} = \dfrac{\text{molar mass of desired product}}{\text{sum of molar masses of all products}} \times 100\%$.

18. State Hess's law.

    The total enthalpy change of a reaction is independent of the route taken, provided the initial and final conditions are the same.

19. Define standard enthalpy of formation, $\Delta_f H^{\ominus}$.

    The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions ($100\ \text{kPa},\ 298\ \text{K}$).

20. Write the equation linking enthalpy change to heat in calorimetry.

    $q = mc\Delta T$, where $m$ is mass of solution (g), $c$ the specific heat capacity ($4.18\ \text{J g}^{-1}\text{K}^{-1}$ for water) and $\Delta T$ the temperature change.

21. Give the Gibbs free energy equation and the condition for a feasible (spontaneous) reaction.

    $\Delta G = \Delta H - T\Delta S$; the reaction is feasible when $\Delta G \leq 0$.

See more Chemistry flashcards →

Planning Chemistry for GCE Advanced Level (A-Levels)

Chemistry (GCE Advanced Level (A-Levels)) FAQ