Engineering and Science Admissions Test (ESAT) Chemistry Module Syllabus

Chemistry Module syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Chemistry Module in Engineering and Science Admissions Test (ESAT), not a summary of it.

1. 1

   Atomic Structure and the Periodic Table

   3 topics

   • Atomic Structure
     • Protons, neutrons and electrons
     • Isotopes and relative atomic mass
     • Electronic configuration and energy levels
   • The Periodic Table
     • Periods, groups and trends
     • Group 1 alkali metals and Group 7 halogens
     • Transition metals overview
   • Bonding and Structure
     • Ionic, covalent and metallic bonding
     • Simple molecular and giant structures
     • Intermolecular forces
2. 2

   Quantitative Chemistry

   3 topics

   • The Mole and Amount of Substance
     • Avogadro's constant and molar quantities
     • Empirical and molecular formulae
   • Chemical Calculations
     • Reacting masses and limiting reagents
     • Concentration of solutions
     • Percentage yield and atom economy
   • Gas Volumes and Titrations
     • Molar gas volume calculations
     • Acid-base titration calculations
3. 3

   Chemical Reactions and Energetics

   3 topics

   • Types of Reaction
     • Acids, bases and neutralisation
     • Redox reactions and oxidation states
     • Precipitation and displacement
   • Energetics
     • Exothermic and endothermic reactions
     • Enthalpy change and energy profile diagrams
     • Bond energy calculations
   • Electrolysis
     • Electrolysis of molten and aqueous compounds
     • Products at electrodes
4. 4

   Rates, Equilibria and Reversible Reactions

   3 topics

   • Rates of Reaction
     • Collision theory and activation energy
     • Effect of concentration, temperature and surface area
     • Catalysts
   • Chemical Equilibrium
     • Dynamic equilibrium
     • Le Chatelier's principle
   • Reversible Reactions and Industrial Processes
     • Reversible reactions and conditions
     • The Haber process
5. 5

   Organic and Applied Chemistry

   4 topics

   • Hydrocarbons and Functional Groups
     • Alkanes, alkenes and homologous series
     • Alcohols, carboxylic acids and esters
     • Crude oil, fractional distillation and cracking
   • Polymers and Reactions
     • Addition and condensation polymerisation
     • Combustion of fuels
   • Chemical Analysis
     • Tests for gases and ions
     • Flame tests and identification
   • Earth and Atmospheric Chemistry
     • Composition of the atmosphere
     • Greenhouse gases and pollution

Chemistry Module flashcards for Engineering and Science Admissions Test (ESAT)

18 of 54 cards from the Chemistry Module deck — real questions with worked answers.

1. What are the relative masses and charges of the three subatomic particles?

   Proton: relative mass $1$, charge $+1$. Neutron: relative mass $1$, charge $0$. Electron: relative mass $\frac{1}{1836}$ (negligible), charge $-1$.

2. Define the terms atomic number ($Z$) and mass number ($A$).

   Atomic number $Z$ = number of protons in the nucleus (defines the element). Mass number $A$ = total number of protons + neutrons. Number of neutrons $= A - Z$.

3. What are isotopes?

   Atoms of the same element with the same number of protons (same $Z$) but different numbers of neutrons, hence different mass numbers $A$. They have identical chemical properties but different physical properties (e.g. density).

4. How is relative atomic mass ($A_r$) calculated from isotopic abundances?

   $$A_r = \frac{\sum (\text{isotope mass} \times \%\text{ abundance})}{100}$$ It is the weighted mean mass of an atom relative to $\tfrac{1}{12}$ of a carbon-12 atom.

5. State the maximum number of electrons in the first four electron shells (energy levels).

   Using $2n^{2}$: shell 1 holds $2$, shell 2 holds $8$, shell 3 holds $18$ (commonly $8$ before filling), shell 4 holds $32$.

6. Write the sub-shell (orbital) electron configuration of iron, $\ce{Fe}$ ($Z=26$).

   $1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,3d^{6}\,4s^{2}$. Note $4s$ fills before $3d$ but is written/removed first on ionisation.

7. How are elements arranged in the modern periodic table, and what do groups and periods represent?

   Arranged by increasing atomic number $Z$. A group (vertical column) = elements with the same number of outer-shell electrons (similar chemistry); a period (horizontal row) = same number of occupied electron shells.

8. Describe and explain the trend in atomic radius across a period and down a group.

   Across a period: decreases, because nuclear charge increases while electrons add to the same shell, pulling them closer. Down a group: increases, because additional shells and greater shielding outweigh increased nuclear charge.

9. Define first ionisation energy and write its equation for a general element X.

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

10. Why does first ionisation energy generally increase across a period?

    Nuclear charge increases, atomic radius decreases, and shielding stays roughly constant, so the outer electron is held more strongly and harder to remove.

11. State the typical properties of Group 1 alkali metals and their reaction with water.

    Soft, low-density, low melting metals; reactivity increases down the group. They react with water to give a hydroxide and hydrogen: $$\ce{2Na + 2H2O -> 2NaOH + H2}$$

12. State the trend in reactivity of Group 7 halogens and write a displacement example.

    Reactivity decreases down the group (oxidising power falls). A more reactive halogen displaces a less reactive one: $$\ce{Cl2 + 2KBr -> 2KCl + Br2}$$

13. What is ionic bonding and what structure does it form?

    Electrostatic attraction between oppositely charged ions formed by electron transfer from metal to non-metal. It forms a giant ionic lattice (e.g. $\ce{NaCl}$).

14. Explain why ionic compounds have high melting points and conduct only when molten or aqueous.

    Strong electrostatic forces throughout the giant lattice require much energy to break (high melting point). Ions are fixed in the solid so cannot move; when molten or dissolved the ions are free to move and carry charge.

15. What is a covalent bond, and how do single, double and triple bonds differ?

    A covalent bond is a shared pair of electrons between two non-metal atoms. Single = 1 shared pair, double = 2 shared pairs, triple = 3 shared pairs; bond strength and shortness increase from single to triple.

16. Define a dative (coordinate) covalent bond and give an example.

    A covalent bond where both shared electrons come from the same atom. Example: the ammonium ion $\ce{NH4+}$, where the lone pair on nitrogen forms the fourth $\ce{N-H}$ bond.

17. List the three types of intermolecular force from weakest to strongest.

    London (induced dipole–dipole / van der Waals) forces < permanent dipole–dipole forces < hydrogen bonding. All are much weaker than covalent or ionic bonds.

18. Compare the structure and properties of diamond and graphite.

    Both are giant covalent carbon. Diamond: each C bonded to 4 others, very hard, non-conductor. Graphite: each C bonded to 3 in layers with delocalised electrons, soft/slippery, conducts electricity.

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