Chem Paper 1 ocr a

Explain the trend in first ionisation energy across a period.

Increases across period , nuclear charge increases, shielding similar, attraction increases.

Decrease at Group 13 (Al), increase at Group 15 (P), decrease at Group 16 (S) due to sub-shell structure.

Explain why sodium has a larger atomic radius than potassium

Na has fewer shells, less shielding, stronger nuclear attraction. K has more shells, more shielding, weaker attraction.

Explain the factors affecting lattice enthalpy.

Ion charge (higher = more exothermic), ion size (smaller = more exothermic), structure type, electrostatic attraction.

Explain the difference between exothermic and endothermic reactions.

Exothermic: releases heat, products lower energy, negative enthalpy change.

Endothermic: absorbs heat, products higher energy, positive enthalpy change.

Explain entropy and what increases it.

Measure of disorder.

Increases with: more moles of gas,

temperature increase,

phase change solid to liquid to gas, mixing/dissolving.

Explain how catalysts affect reactions without being used up.

Provide alternative pathway with lower activation energy.

Not consumed. Speed up forward and reverse equally.

Do not affect equilibrium position.

Explain the difference between heterogeneous and homogeneous catalysts.

Heterogeneous: different physical state from reactants, surface adsorption.

Homogeneous: same state, forms intermediate with reactants.

Explain Le Chatelier's principle for concentration changes.

Increase reactant concentration shifts right to reduce it.

Increase product shifts left. System opposes the change.

Explain Le Chatelier's principle for pressure changes in gaseous systems.

Increase pressure shifts to side with fewer gas moles.

Decrease pressure shifts to side with more gas moles.

Explain Le Chatelier's principle for temperature changes.

Increase temperature shifts in endothermic direction.

Decrease temperature shifts in exothermic direction.

Explain the difference between strong and weak acids.

Strong: fully dissociate, strong conductor, low pH.

Weak: partially dissociate, equilibrium with undissociated molecules, poor conductor.

Explain the difference between a Bronsted-Lowry acid and base.

Acid: proton donor.

Base: proton acceptor.

Conjugate acid-base pairs differ by one proton.

Explain what is meant by a buffer solution and how it works.

Resists pH change.

Contains weak acid and conjugate base.

Added acid reacts with conjugate base.

Added base reacts with weak acid.

Explain the difference between oxidation and reduction.

Oxidation: loss of electrons, increase oxidation number, gain oxygen, loss hydrogen.

Reduction: gain of electrons, decrease oxidation number, loss oxygen, gain hydrogen.

Explain the trend in atomic radius down a group.

More electron shells added, increased shielding, outer electron further from nucleus, weaker nuclear attraction, atomic radius increases.

Explain why Group 1 metals become more reactive down the group.

Outer electron further from nucleus, more shielding, weaker attraction, electron lost more easily, ionisation energy decreases.

Explain why halogens become less reactive down the group.

Atomic radius increases, shielding increases, nuclear attraction for electrons decreases, harder to gain electron, electron affinity decreases.

Describe the reactions of Group 2 metals with water.

Reactivity increases down group. Forms hydroxide and hydrogen gas. Calcium: slow reaction. Magnesium: reacts with steam only. Barium: vigorous reaction.

Explain the trend in thermal stability of Group 2 carbonates.

Down group: thermal stability increases.

Larger cation polarises carbonate ion less.

Charge density decreases, less distortion of anion.

Explain the trend in thermal stability of Group 2 nitrates.

Down group: thermal stability increases.

Larger cation polarises nitrate less.

Decompose to oxide, nitrogen dioxide, oxygen.

Describe the reactions of halogens with halide ions.

More reactive halogen displaces less reactive from solution.

Cl2 + 2Br- → 2Cl- + Br2.

Cl2 + 2I- → 2Cl- + I2.

Br2 + 2I- → 2Br- + I2.

Explain why hydrogen bonding occurs and its effects on properties.

H bonded to N, O, or F.

Lone pair on electronegative atom attracts delta+ H.

Increases boiling point, increases viscosity, ice less dense than water.

Explain the difference between sigma and pi bonds.

Sigma: head-on overlap, single bond, allows rotation.

Pi: sideways overlap of p orbitals, double/triple bond, restricts rotation, weaker than sigma.

Explain VSEPR theory and molecular shapes.

Electron pairs repel each other.

Arrange to minimise repulsion.

Lone pairs repel more than bonding pairs.

Determines bond angles and molecular geometry.

Explain the difference between ionic and covalent bonding.

Ionic: electrostatic attraction between oppositely charged ions, metal + non-metal, high melting point, conducts when molten/dissolved. Covalent: shared electron pair, non-metals, various melting points, usually non-conducting.

Explain metallic bonding and its properties.

Positive ions in sea of delocalised electrons. Strong electrostatic attraction. High melting point, conducts electricity, malleable, ductile.

Explain the difference between intermolecular forces and intramolecular bonds.

Intramolecular: bonds within molecule (covalent, ionic, metallic) — strong. Intermolecular: forces between molecules (van der Waals, dipole-dipole, hydrogen bonds) — weaker.

Explain how intermolecular forces affect boiling points.

Stronger intermolecular forces → higher boiling point.

Hydrogen bonding > dipole-dipole > van der Waals.

More electrons → stronger van der Waals

Explain the difference between polar and non-polar molecules.

Polar: uneven charge distribution, permanent dipole, electronegativity difference.

Non-polar: even charge distribution, no permanent dipole, symmetrical or same atoms.

How to show reaction has gone to completion?

All limiting reactant used up (e.g., metal dissolved, bubbles stop)

Define relative atomic mass.

Weighted mean mass of an atom of an element / isotope compared with 1/12 of the mass of an atom of carbon-12

Explain why a metal conducts electricity in both solid and liquid states.

Metallic bonding with delocalised electrons that are free to move / mobile in both states

Explain why a molecular substance has a low melting point and does not conduct electricity.

Simple molecular structure with weak induced dipole-dipole forces / London forces between molecules; no mobile charge carriers / ions / delocalised electrons

Explain why an ionic compound has a high melting point, does not conduct when solid, but conducts when molten.

Giant ionic lattice with strong electrostatic forces between oppositely charged ions; ions fixed / cannot move in solid but free to move when molten

Explain the trend in reactivity down Group 7.

Reactivity decreases; atomic radius increases; ability to gain an electron / electron affinity decreases

Explain why transition metal compounds are coloured.

d-d transitions; d-orbitals are split into two energy levels by ligands; electron promoted by absorbing visible light; complementary colour observed

Explain the difference between a d-block element and a transition element.

d-block = highest energy/valence electron in d-orbital; transition element = forms one or more stable ions with incomplete d-subshell

Explain why Scandium is not a transition element.

Sc³⁺ has no d electrons / empty d-subshell

Explain why Zinc is not a transition element.

Zn²⁺ has full d-subshell (3d¹⁰)

Explain the conditions that give maximum yield for an exothermic reaction with fewer moles of gas on the product side

High pressure (shifts right, fewer gas moles) and low temperature (exothermic, produces heat)

Explain why low temperature is not used industrially despite giving maximum yield

Rate too slow; compromise temperature used with catalyst to give reasonable rate

Explain how a catalyst makes a process more sustainable

Lower activation energy → less energy required; provides alternative route → fewer by-products / less waste

Explain why a buffer solution forms when a weak acid is partially neutralised by a strong base.

Weak acid remains in excess AND conjugate base (salt) is formed

Explain why adding a small amount of water to a buffer does not change its pH significantly.

[HA] and [A⁻] dilute equally; ratio [HA]:[A⁻] stays constant

Explain how to determine the direction of a redox reaction using standard electrode potentials.

More positive E° → reduction; more negative E° → oxidation; use "more positive / more negative" NOT "higher / lower"

Define electronegativity

The attraction of a bonded atom for the electrons in a covalent bond

Describe metallic bonding and explain why a metal conducts electricity.

Positive ions/cations in a sea of delocalised electrons; electrons are mobile/free to move and carry charge

Explain why there is a large jump between the 2nd and 3rd ionisation energies of an element.

2nd electron removed from inner shell / closer to nucleus / less shielded / greater nuclear attraction

Explain why there is a large jump between the 10th and 11th ionisation energies.

10th electron from 2nd shell, 11th from 1st shell / much closer to nucleus / much less shielded

Define lattice enthalpy.

Enthalpy change when one mole of ionic solid is formed from its gaseous ions

Define enthalpy change of solution.

Enthalpy change when one mole of solute dissolves in water to form an infinitely dilute solution

Explain why it is difficult to predict whether enthalpy of solution becomes more or less exothermic down a group.

Lattice enthalpy becomes less exothermic (larger ions, weaker attraction) BUT hydration enthalpy also becomes less exothermic (larger ions, weaker attraction to water) — two opposing trends

Define a bidentate ligand.

A ligand that donates two lone pairs to the metal ion / forms two coordinate bonds

Explain why transition metal compounds are coloured.

d-d transitions; d-orbitals split into two energy levels by ligands; electron promoted by absorbing visible light; complementary colour observed

Explain the trend in boiling points down the halogens.

London forces / induced dipole-dipole interactions become stronger; more electrons; larger electron cloud; NOT covalent bond strength

Explain, in terms of oxidation numbers, why disproportionation has taken place.

Same element is both oxidised and reduced / oxidation number increases AND decreases

Explain, in terms of electrode potentials, why a species disproportionates in acid conditions.

E°cell > 0 / overall cell potential is positive / reaction is thermodynamically feasible

Explain why a reaction with ΔH positive and ΔS positive is feasible at high temperatures but not low temperatures.

ΔG = ΔH - TΔS; at low T, ΔH dominates (ΔG > 0); at high T, TΔS dominates (ΔG < 0)

Explain why a catalyst makes a process more sustainable.

Lower activation energy → less energy required; alternative route → fewer by-products / less waste / higher atom economy

Explain why a buffer solution resists pH change on addition of small amounts of acid.

A⁻ reacts with added H⁺ → HA; ratio [HA]:[A⁻] maintained; equilibrium shifts to oppose change

Define enthalpy change of atomisation.

Enthalpy change when one mole of gaseous atoms is formed from the element in its standard state

Define first electron affinity.

Enthalpy change when one mole of gaseous atoms gains one electron to form one mole of gaseous 1- ions

Define enthalpy change of hydration.

Enthalpy change when one mole of gaseous ions dissolves in water to form an infinitely dilute solution

Explain why lattice enthalpy becomes less exothermic down a group.

Ionic radius increases — attraction between ions decreases

Explain why hydration enthalpy becomes less exothermic down a group.

Ionic radius increases — attraction between ion and water molecules decreases

Explain why it is difficult to predict whether enthalpy of solution becomes more or less exothermic down a group.

Lattice enthalpy becomes less exothermic BUT hydration enthalpy also becomes less exothermic — two opposing trends

Explain why Group 2 carbonates become more thermally stable down the group.

Cation radius increases — charge density decreases — less polarising power on carbonate ion

Explain why Group 2 hydroxides become more soluble down the group.

Lattice enthalpy decreases more than hydration enthalpy — enthalpy of solution becomes more exothermic

Explain, in terms of oxidation numbers, why disproportionation has taken place.

Same element is both oxidised and reduced / oxidation number increases AND decreases

Explain, in terms of electrode potentials, why a species disproportionates in acid conditions.

E°cell > 0 / overall cell potential is positive / reaction is thermodynamically feasible

why an alkaline hydrogen-oxygen fuel cell does not produce

No carbon-containing fuel / only H₂ and O₂ used — water is the only product

Explain why an electrochemical cell needs a salt bridge

Completes the circuit / allows ion transfer / maintains electrical neutrality / prevents mixing of solutions

Explain why a reaction with ΔH positive and ΔS positive is feasible at high temperatures but not low temperatures.

ΔG = ΔH - TΔS; at low T, ΔH dominates (ΔG > 0); at high T, TΔS dominates (ΔG < 0)

Explain why a catalyst makes a process more sustainable

Lower activation energy → less energy required; alternative route → fewer by-products / less waste / higher atom economy

Explain why a buffer solution resists pH change on addition of small amounts of acid.

A⁻ reacts with added H⁺ → HA; ratio [HA]:[A⁻] maintained; equilibrium shifts to oppose change