Chemistry Paper 1

Atomic radius trend

Atomic radius decreases across a period due to increasing proton number pulling electrons closer.

Element classification

An element is classified as s, p, d or f block based on its Periodic Table position.

Ionisation energy trend

Ionisation energy generally increases across a period due to stronger nuclear attraction.

Periodicity

Trends in element properties with increasing atomic number, caused by atomic structure changes.

Proton number

Number of protons in an atom’s nucleus; determines periodic table order.

Barium meals

Barium sulfate, opaque to x-rays, used to diagnose stomach or intestinal issues.

Extraction of titanium

Titanium is extracted by converting TiO₂ to TiCl₄, then reducing with magnesium.

Flue gases

Gases emitted from chimneys and industrial exhausts.

Solubility

Ability of a substance to dissolve; Group 2 hydroxides increase and sulfates decrease in solubility down the group.

Sparingly soluble

Describes compounds with very low solubility, e.g., magnesium hydroxide.

Test for sulfates

Add acidified barium chloride; white precipitate of barium sulfate indicates presence.

Wet scrubbing

Removes sulfur dioxide from flue gases using an alkali.

Displacement

A more reactive halogen displaces a halide lower in the periodic table.

Disproportionation

A reaction where a substance is both oxidised and reduced.

Electronegativity

Tendency of an atom to attract bonding electrons.

Oxidising ability

Decreases down Group 7 due to increased atomic size and shielding.

Redox reaction

Reaction involving simultaneous oxidation and reduction.

Reducing ability

Increases down Group 7 due to easier electron loss from larger halide ions.

Acceleration

TOF stage where ions are accelerated so they have equal kinetic energy.

Atom

Smallest part of an element that can exist.

Atomic nucleus

Central part of atom with protons and neutrons; positively charged.

Atomic number

Number of protons in an atom's nucleus.

Electron

Negatively charged subatomic particle, relative mass 1/2000.

Electron configuration

Arrangement of electrons in atomic orbitals.

Electron impact ionisation

Ionisation method where high-energy electrons knock off electrons from particles.

Electrospray ionisation

Ionisation method where particles gain H⁺ after being pushed through a nozzle with high voltage.

First ionisation energy

Energy to remove 1 mole of electrons from 1 mole of gaseous atoms.

Ion detection

Final TOF stage where ions hit a detector to generate a mass spectrum.

Ion drift

TOF stage where ions move with constant energy; lighter ions drift faster.

Ionisation

Process of converting particles into ions.

Isotope

Same element, same protons, different neutrons (e.g., ³⁵Cl and ³⁷Cl).

Mass number

Total number of protons and neutrons.

Mass spectrometer

Instrument providing isotope mass and abundance data.

Mass spectrometry

Technique to identify elements and calculate molecular masses.

Neutron

Neutral subatomic particle with relative mass 1.

Nuclear charge

Total proton charge in nucleus; increases across the periodic table.

Proton

Positively charged subatomic particle, relative mass 1.

Second ionisation energy

Energy to remove 1 mole of electrons from 1 mole of gaseous 1+ ions.

Sub-shells

Subdivisions of electron shells (s, p, d, f) with distinct energy levels.

Time of Flight (TOF) spectrometer

Mass spectrometry method with ionisation, acceleration, drift, and detection.

Atom economy

Measure of useful product formed from starting materials.

Avogadro’s constant

Number of particles in one mole.

Concentration

Amount of substance per unit volume (g/dm³ or mol/dm³).

Empirical formula

Simplest ratio of atoms in a compound.

Limiting reactant

Reactant fully used up, limiting product formation.

Mole

Amount of substance with same particles as 12g of carbon-12.

Molecular formula

Actual number of atoms of each element in a molecule.

Percentage by mass

Element’s proportion in a compound or mixture.

Percentage yield

Actual yield as a percentage of theoretical yield.

Relative atomic mass

Average mass of atoms compared to 1/12 of carbon-12.

Relative molecular mass

Average mass of a molecule compared to 1/12 of carbon-12.

Co-ordinate bond

Shared electron pair donated by one atom.

Covalent bond

Shared electron pair between non-metals.

Dipole

Charge difference due to unequal electron sharing in a bond.

Electron pair repulsion

Negatively charged electron pairs repel and arrange far apart.

Electronegativity

Ability of atom to attract electrons in a bond.

Electrostatic forces

Attraction between oppositely charged ions.

Hydrogen bonding

Attraction between H⁺ and lone pair on electronegative atoms (e.g., O, N, F).

Intermolecular forces

Forces between molecules affecting physical properties.

Ion

Charged atom or molecule from electron loss/gain.

Ionic bond

Metal loses and non-metal gains electrons; oppositely charged ions attract.

Ionic compound

Compound of ions held by electrostatic forces.

Lattice

Regular, repeating arrangement in a crystal.

Macromolecular structure

Giant covalent network (e.g., diamond) with high melting points.

Metallic bond

Attraction between metal ions and delocalised electrons.

Permanent dipole-dipole force

Intermolecular force between polar molecules.

Polar bond

Covalent bond with unequal electron sharing, causing charge difference.

Simple molecular structure

Atoms joined by covalent bonds; low melting/boiling points.

Van der Waals

Temporary induced dipoles causing attraction between molecules.

VSEPR theory

Explains molecular shape by minimizing electron pair repulsion.

Calorimetry

Measurement of energy in chemical reactions.

Endothermic reaction

Takes in energy, temperature drops.

Enthalpy change (ΔH)

Heat change at constant pressure.

Exothermic reaction

Releases energy, temperature rises.

Hess’s law

Total enthalpy change is independent of the reaction path.

Mean bond enthalpy

Average bond-breaking energy across compounds.

Molar enthalpy change

Enthalpy change per mole of substance.

Standard conditions

100 kPa and 298 K.

Standard enthalpy of combustion

ΔH when 1 mole is burned in excess O₂.

Standard enthalpy of formation

ΔH when 1 mole is formed from elements in standard states.

Activation energy

Minimum energy required for a successful collision.

Catalyst

Substance increasing rate by lowering activation energy.

Collision theory

Reaction needs sufficient energy collisions to occur.

Concentration effect on rate

Higher concentration = more collisions = faster rate.

Pressure effect on rate

Higher pressure = more collisions = faster rate.

Temperature effect on rate

Higher temp = more kinetic energy = faster reaction.

Maxwell-Boltzmann distribution

Shows energy distribution of molecules at constant temp.

Rate of reaction

Amount of product formed/reactant used over time.

Closed system

System with energy exchange but no matter exchange.

Dynamic equilibrium

Forward and reverse reaction rates equal; concentrations constant.

Le Chatelier’s principle

System at equilibrium shifts to oppose changes.

Equilibrium constant (Kc)

Relationship between concentrations of reactants and products at equilibrium.

Heterogeneous system

Chemicals in different phases.

Homogeneous system

Chemicals in same phase.

Redox reaction

Reaction with simultaneous oxidation and reduction.

Oxidation

Loss of electrons; increase in oxidation number.

Reduction

Gain of electrons; decrease in oxidation number.

Oxidising agent

Accepts electrons; gets reduced.

Reducing agent

Donates electrons; gets oxidised.

Half equation

Shows oxidation or reduction part of redox reaction.