Chemistry remove exam

Solid

Particles closely packed, vibrate in fixed positions, low energy.

Liquid

Particles close but can move past each other, moderate energy.

Gas

Particles far apart, move randomly and quickly, high energy.

Interconversions

Melting, freezing, boiling, condensation, sublimation.

Diffusion

Spreading of particles from high to low concentration.

Gas diffusion rate

Depends on molecular mass — lighter gases diffuse faster.

Element

One type of atom only.

Compound

Two or more elements chemically bonded.

Mixture

Two or more substances not chemically bonded; can be separated physically.

Pure substances

Have fixed melting/boiling points.

Simple distillation

Separates solvent from solution.

Fractional distillation

Separates liquids with different boiling points.

Filtration

Separates insoluble solid from liquid.

Crystallisation

Obtains solid crystals from a solution.

Paper chromatography

Separates substances based on solubility and attraction to paper.

Rf value

Rf = distance moved by substance ÷ distance moved by solvent.

Atom

Smallest particle of an element.

Molecule

Two or more atoms chemically bonded (same or different elements).

Nucleus

Contains protons (+1 charge, mass 1) and neutrons (0 charge, mass 1).

Electron

Negative charge (-1), negligible mass, orbit nucleus in shells.

Atomic number (Z)

Number of protons = number of electrons.

Mass number (A)

Total protons + neutrons.

Isotopes

Atoms of the same element with different numbers of neutrons.

Relative atomic mass (Ar)

Average mass of isotopes weighted by abundance.

Periodic Table

Arranged in order of increasing atomic number.

Electron configuration

Fill shells in order: 2, 8, 8, 2.

Metals

Good conductors, basic oxides.

Non-metals

Poor conductors, acidic oxides.

Similar properties in groups

Same number of outer electrons → similar chemical reactions.

Noble gases

Group 0 → full outer shell → stable → unreactive.

Formation of ions

Atoms gain or lose electrons to achieve a full outer shell.

Cations

Metals lose electrons → positive ions.

Anions

Non-metals gain electrons → negative ions.

Common ion charges

Group 1: +1, Group 2: +2, Group 3: +3; Group 5: -3, Group 6: -2, Group 7: -1.

Transition metals

Ag⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Zn²⁺.

Other ions

H⁺, OH⁻, NH₄⁺, CO₃²⁻, NO₃⁻, SO₄²⁻.

Writing formulae

x

Ionic dot-and-cross diagrams

Show electron transfer between metals and non-metals (Groups 1, 2, 3 with 5, 6, 7). Only outer electrons shown.

Ionic bonding

Electrostatic attraction between oppositely charged ions.

Giant ionic lattices

Strong electrostatic forces → high melting and boiling points.

Electrical conductivity of ionic compounds

Conduct when molten or dissolved (ions free to move). Do not conduct when solid (ions fixed in lattice).

Covalent bonding

Formed when two atoms share a pair of electrons.

Covalent bond explanation

Electrostatic attraction between shared electrons and nuclei.

Covalent dot-and-cross diagrams

Diatomic molecules: H₂, O₂, N₂, Cl₂, HCl; Inorganic molecules: H₂O, NH₃, CO₂; Organic (up to two carbons): CH₄, C₂H₆, C₂H₄, halogenated forms.

Simple molecular substances

Weak intermolecular forces → low melting and boiling points. Usually gases or liquids at room temperature.

Effect of molecular mass

As molecular mass increases → stronger intermolecular forces → higher melting and boiling points.

Giant covalent structures

Strong covalent bonds throughout → very high melting and boiling points.

Diamond, graphite, C₆₀ fullerene

Diamond: each carbon bonds to 4 others → hard, non-conductive. Graphite: each carbon bonds to 3 others → layers, conducts electricity. C₆₀ fullerene: discrete molecules, soft, does not conduct well.

Conductivity of covalent compounds

Covalent compounds do not conduct electricity (no free ions or electrons).

Metallic lattice

Positive metal ions in a sea of delocalised electrons.

Metallic bonding

Electrostatic attraction between positive ions and delocalised electrons.

Properties of metals

Conduct electricity (free electrons). Malleable and ductile (layers of ions can slide).

Reactions with water

All alkali metals react with water → hydrogen + metal hydroxide.

Trends in reactivity (Group 1)

Reactivity increases down the group as outer electron is more easily lost.

Predicting properties (Group 1)

Use patterns to predict reactions, melting points, density of unknown alkali metals.

Halogen properties

Halogens: fluorine (gas, yellow), chlorine (gas, green), bromine (liquid, red-brown), iodine (solid, grey). Become darker and less reactive down the group.

Predicting halogen properties

Use trends to predict colour, state, and reactivity of unknown halogens.

Displacement reactions

More reactive halogen displaces a less reactive halogen from a compound. e.g. Cl₂ + 2KBr → 2KCl + Br₂.