Cambridge Pre-U Chemistry (Principal Subject) Syllabus

Chemistry (Principal Subject) syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Chemistry (Principal Subject) in Cambridge Pre-U, not a summary of it.

1. 1

   Physical Chemistry

   5 topics

   • Atomic Structure
     • Sub-atomic particles and isotopes
     • Electron configuration and orbitals
     • Ionisation energies and periodic trends
   • Bonding and Structure
     • Ionic, covalent and metallic bonding
     • Shapes of molecules and VSEPR theory
     • Intermolecular forces and hydrogen bonding
   • Energetics and Thermodynamics
     • Enthalpy changes and Hess's law
     • Born-Haber cycles and lattice enthalpy
     • Entropy and Gibbs free energy
   • Chemical Kinetics
     • Rate equations and orders of reaction
     • Rate-determining step and mechanisms
     • The Arrhenius equation and catalysis
   • Equilibria
     • Equilibrium constants Kc and Kp
     • Le Chatelier's principle
     • Acid-base equilibria, pH and buffers
2. 2

   Inorganic Chemistry

   4 topics

   • Periodicity
     • Trends across Period 3
     • Reactions of Period 3 elements and oxides
   • Group Chemistry
     • Group 2 alkaline earth metals
     • Group 17 halogens and reactivity trends
   • Transition Elements
     • Variable oxidation states and colour
     • Complex ions, ligands and shapes
     • Catalytic behaviour
   • Redox and Electrochemistry
     • Oxidation numbers and half-equations
     • Standard electrode potentials and cells
3. 3

   Organic Chemistry

   5 topics

   • Hydrocarbons
     • Alkanes and free-radical substitution
     • Alkenes and electrophilic addition
   • Oxygen-Containing Compounds
     • Alcohols, oxidation and dehydration
     • Aldehydes, ketones and carboxylic acids
     • Esters and esterification
   • Nitrogen Compounds and Aromatics
     • Amines and amides
     • Benzene and electrophilic substitution
   • Isomerism and Mechanisms
     • Structural and stereoisomerism
     • Optical isomerism and chirality
     • Nucleophilic substitution mechanisms
   • Polymers and Synthesis
     • Addition and condensation polymers
     • Multi-step synthetic routes
4. 4

   Analysis and Practical Skills

   4 topics

   • Spectroscopy
     • Infrared spectroscopy
     • Mass spectrometry and fragmentation
     • Nuclear magnetic resonance spectroscopy
   • Qualitative Analysis
     • Tests for cations and anions
     • Tests for organic functional groups
   • Quantitative Techniques
     • Titrations and volumetric analysis
     • Gravimetric analysis
   • Experimental Methodology
     • Measurement uncertainty and error analysis
     • Planning and evaluating investigations

Chemistry (Principal Subject) flashcards for Cambridge Pre-U

25 of 61 cards from the Chemistry (Principal Subject) deck — real questions with worked answers.

1. Define the term 'first ionisation energy' of an element.

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

2. State the order in which the subshells $3d$ and $4s$ are filled, and which is lost first on ionisation.

   $4s$ fills before $3d$ (lower energy when empty), but $4s$ electrons are removed first on ionisation because once $3d$ is occupied, $4s$ becomes higher in energy.

3. Why does first ionisation energy show a small drop from Group 2 to Group 13 (e.g. Mg to Al)?

   The outer electron in Group 13 is in a $p$ subshell, which is higher in energy and slightly shielded by the filled $s$ subshell, so it is more easily removed than the $s$ electron in Group 2.

4. Write the full electron configuration of a $\ce{Cr}$ atom (Z = 24) and explain the anomaly.

   $1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,3d^{5}\,4s^{1}$. A half-filled $3d^{5}$ with $4s^{1}$ is more stable than $3d^{4}\,4s^{2}$ due to the extra exchange stability of a half-filled subshell.

5. Define electronegativity.

   The ability (power) of an atom to attract the bonding pair of electrons in a covalent bond towards itself.

6. Compare the shapes and bond angles of $\ce{NH3}$, $\ce{H2O}$ and $\ce{CH4}$ using VSEPR.

   $\ce{CH4}$: tetrahedral, $109.5^{\circ}$ (4 bond pairs). $\ce{NH3}$: trigonal pyramidal, $107^{\circ}$ (3 bp, 1 lp). $\ce{H2O}$: bent/angular, $104.5^{\circ}$ (2 bp, 2 lp). Lone pairs repel more, reducing the angle.

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

   A covalent bond in which both shared electrons originate from the same atom. Example: the bond from $\ce{N}$ to $\ce{H+}$ in the ammonium ion $\ce{NH4+}$.

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

   London (instantaneous induced dipole–induced dipole) forces $<$ permanent dipole–dipole forces $<$ hydrogen bonding.

9. State Hess's law.

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

10. Write the formula for the standard enthalpy change of reaction using bond enthalpies.

    $$\Delta H_{r} = \sum (\text{bonds broken}) - \sum (\text{bonds formed})$$ where bond breaking is endothermic and bond forming is exothermic.

11. Define the standard enthalpy change of formation, $\Delta H_{f}^{\ominus}$.

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

12. State the relationship between $\Delta G$, $\Delta H$ and $\Delta S$, and the condition for spontaneity.

    $$\Delta G = \Delta H - T\Delta S$$ A reaction is feasible (spontaneous) when $\Delta G < 0$.

13. Define lattice energy.

    The enthalpy change when one mole of an ionic solid is formed from its gaseous ions: $\ce{M+(g) + X-(g) -> MX(s)}$ (an exothermic quantity).

14. How do ionic charge and ionic radius affect the magnitude of lattice energy?

    Lattice energy becomes more exothermic (larger magnitude) with greater ionic charges and smaller ionic radii, since $\text{LE} \propto \dfrac{q_{+}q_{-}}{r_{+}+r_{-}}$.

15. State the two key effects of a catalyst on the rate and energy profile of a reaction.

    It provides an alternative reaction pathway of lower activation energy $E_{a}$; it does not change $\Delta H$ or the position of equilibrium, only the rate.

16. Write the Arrhenius equation and identify each term.

    $$k = A\,e^{-E_{a}/RT}$$ where $k$ = rate constant, $A$ = pre-exponential (frequency) factor, $E_{a}$ = activation energy, $R$ = gas constant, $T$ = absolute temperature.

17. For a reaction with rate $= k[A]^{2}[B]$, give the overall order and the units of $k$.

    Overall order $= 3$. Units of $k$: $\text{mol}^{-2}\,\text{dm}^{6}\,\text{s}^{-1}$.

18. How can the half-life of a reaction be used to identify first order kinetics?

    For a first-order reaction the half-life $t_{1/2}$ is constant (independent of concentration); a constant successive half-life on a concentration–time graph confirms first order.

19. What is the rate-determining step?

    The slowest step in a multi-step reaction mechanism; its molecularity and the species involved determine the experimental rate equation.

20. State Le Chatelier's principle.

    If a change is imposed on a system at equilibrium, the position of equilibrium shifts in the direction that tends to oppose (minimise) that change.

21. Write the expression for $K_{c}$ for the reaction $\ce{N2 + 3H2 <=> 2NH3}$.

    $$K_{c} = \frac{[\ce{NH3}]^{2}}{[\ce{N2}][\ce{H2}]^{3}}$$

22. Define $K_{p}$ and write it for $\ce{2SO2 + O2 <=> 2SO3}$.

    $K_{p}$ is the equilibrium constant in terms of partial pressures: $$K_{p} = \frac{p_{\ce{SO3}}^{2}}{p_{\ce{SO2}}^{2}\,p_{\ce{O2}}}$$

23. How does increasing temperature affect $K_{c}$ for an exothermic reaction?

    For an exothermic forward reaction, increasing temperature shifts equilibrium backwards and decreases $K_{c}$ (only temperature changes the value of $K$).

24. Define a Brønsted–Lowry acid and base.

    An acid is a proton ($\ce{H+}$) donor; a base is a proton acceptor.

25. Give the expression for $\text{pH}$ and for $K_{w}$ at $298\,\text{K}$.

    $$\text{pH} = -\log_{10}[\ce{H+}]$$ and $K_{w} = [\ce{H+}][\ce{OH-}] = 1.0\times10^{-14}\,\text{mol}^{2}\,\text{dm}^{-6}$.

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