Periodicity
What is the approximate charge and relative mass of an electron?
Electron: charge = −1 (−1.602×10−19 C), relative mass ≈ 1/1836 of a proton (0) (~0.00055 u).
Give the type of bonding in sodium chloride and one feature of its structure.
Ionic bonding; crystalline lattice of alternating ions; high melting/boiling point and conducts when molten/dissolved.
Define relative atomic mass (Ar)
Ar is the weighted mean mass of an atom of an element compared with 1/12 of a carbon-12 atom.
Define exothermic and endothermic reaction in terms of energy flow.
Exothermic: releases heat to surroundings (products lower enthalpy than reactants). Endothermic: absorbs heat (products higher enthalpy).
State the functional group and general formula of an alcohol.
Alcohol: –OH group; general formula CnH2n+1OH (or R–OH).
State the meaning of first ionisation energy.
First ionisation energy: energy required to remove one mole of electrons from one mole of gaseous atoms to form 1+ ions.
Draw (or describe) the electron pair geometry around a carbon atom in methane and state the bond angle.
Tetrahedral geometry, bond angle ≈ 109.5
Calculate the number of moles in 20.0 g of oxygen gas, O2.
0.625 mol
Sketch (describe) a reaction profile showing activation energy and ΔH for an exothermic reaction.
Reaction profile: peak representing Ea above reactants; products lower than reactants; ΔH negative. (Label Ea and ΔH.)
Name the following: CH3-CH2-CH2-CH3 (IUPAC)
CH3CH2CH2CH3 = butane
Explain, using shell and sub-shell structure, why potassium has a lower first ionisation energy than neon.
Potassium has an extra shell (4s electron) with greater shielding and larger radius; nuclear attraction on outer electron is weaker than neon’s tightly held electrons.
Explain why diamond conducts heat well but not electricity; contrast with graphite.
Diamond: giant covalent lattice, strong C–C bonds, no delocalised electrons → no electrical conductivity; graphite: layers with delocalised electrons between layers → conducts electricity and layers slide (lubricant).
A 2.00 g sample of a hydrate loses 0.36 g on heating to give anhydrous salt. Calculate the percentage of water in the hydrate.
% water = 18.0%
State two methods to increase the rate of a reaction and explain why each works at the particle level.
Increase concentration (more collisions per unit time); increase temperature (particles have more kinetic energy → more collisions exceed Ea); use catalyst (lowers Ea).
Describe one test to distinguish an alkene from an alkane and state the observation for a positive test.
Bromine water decolourisation or Baeyer's test (cold, dilute KMnO4); alkene decolourises bromine (brown to colourless), alkane does not under same conditions.
Describe how successive ionisation energies provide evidence for the number of electrons in outer shells (illustrate with a likely pattern for magnesium).
Successive IE: relatively small increases until removal from inner shell causes large jump — use magnesium as example: moderate IEs for 1st & 2nd, large jump for 3rd indicating two outer electrons.
Describe how hydrogen bonding affects the boiling point of water compared with hydrogen sulfide.
H2O has hydrogen bonding between molecules raising boiling point; H2S lacks effective H-bonding → much lower bp.
A reaction: 2A + 3B → C. If 0.50 mol A reacts with 0.75 mol B, determine the limiting reagent and the maximum moles of C formed.
2A needs 0.75 mol B for 0.50 mol A. For 0.50 mol A need 0.75 mol B (2:3 ratio). Both exactly used: limiting reagent none (both exactly used) → moles C = 0.25 mol
Define enthalpy change of formation and how it differs from enthalpy change of combustion.
Enthalpy of formation: enthalpy change when 1 mol compound forms from elements in standard states. Combustion: enthalpy when 1 mol substance reacts completely with O2 to form oxides.
Outline the mechanism (brief steps) for nucleophilic substitution of a haloalkane by OH− in aqueous conditions (SN1 vs SN2 — indicate which types of substrates favour each).
SN2: single-step bimolecular backside attack, inversion of configuration, favoured by primary substrates and strong nucleophiles. SN1: two-step via carbocation intermediate, racemisation, favoured by tertiary substrates and polar protic solvents.
Predict and explain how atomic radius and electronegativity change across Period 3 (Na → Ar). Include reasons based on nuclear charge and shielding.
Across Period 3: atomic radius decreases (increased nuclear charge, similar shielding), electronegativity increases (stronger attraction for bonding electrons).
Explain the factors that determine lattice enthalpy and predict whether MgO or NaCl has the larger (more exothermic) lattice enthalpy, giving reasons.
Lattice enthalpy depends on ionic charges (greater = larger) and ionic radii (smaller = larger); Mg2+ and O2− vs Na+ and Cl−: MgO has much larger lattice enthalpy.
Given: 0.250 mol of H2 reacts with excess O2. Calculate the mass of water formed. (Equation: 2H2 + O2 → 2H2O)
0.250 mol H2 produces 0.250 mol H2O (1:1). Mass = 0.250 × 18.0 = 4.50 g.
Explain how a catalyst affects activation energy and the reaction mechanism; give one example relevant to industry or laboratory.
Catalyst provides an alternative lower-Ea pathway, often via intermediate steps; example: Haber process uses Fe catalyst to increase rate and lower temperature required.
For the aromatic compound benzene, explain why electrophilic substitution (e.g., nitration) occurs rather than addition, including a comment on stability and resonance.
Addition would break aromaticity; electrophilic substitution preserves aromatic system. Benzene’s delocalised pi electrons create a stable resonance system; EAS proceeds via sigma complex then reformation of aromaticity.