Secondary 4 Chemistry

A method that separates the components of a liquid mixture using their different boiling points.

Air compressed and cooled to about −200°C into liquid form; the raw material for fractional distillation.

The temperature at which a liquid boils; nitrogen boils at −196°C and oxygen at −183°C.

A change in which no new substance forms, such as the fractional distillation of liquid air.

A colourless, unreactive gas making up about 78% of air, formula N₂, used to make fertilisers.

A noble gas making up about 0.9% of air; very unreactive, used to fill light bulbs.

A glowing splint relights in the gas, proving that it is oxygen.

A splint that has stopped burning but still glows — the standard tool for testing oxygen.

Substances that lower air quality, such as sulfur dioxide, nitrogen dioxide and suspended particulates.

Matter made of two or more substances in variable proportions, like air.

The substance that dissolves in a solvent, such as salt in seawater.

The liquid that dissolves a solute, such as water.

A solution that cannot dissolve any more solute at that temperature.

The slow change of a liquid into gas at its surface; heating speeds it up, and it is used to obtain salt from seawater.

The process in which solute separates from a saturated solution as crystals.

Boiling a liquid and condensing the vapour to collect the pure liquid.

The apparatus that cools hot vapour back into liquid; cooling water flows in at the bottom and out at the top.

Identifying a metal element from the colour it gives a flame: sodium is golden yellow, calcium brick red.

Adding silver nitrate solution; a white precipitate shows that chloride is present.

A white powder that turns blue when it meets water, used to test for water.

The solid outermost layer of the Earth, made of oxygen, silicon, aluminium, iron and other elements.

A naturally occurring solid of fairly definite composition — the building block of rocks.

Rock containing minerals from which metals can be extracted profitably, such as haematite.

Formula CaCO₃ — the main substance of limestone, chalk and marble.

A common natural form of calcium carbonate and the raw material for cement and glass.

A reaction in which heat breaks a substance into two or more new substances, such as the thermal decomposition of calcium carbonate.

An equation written in words, e.g. calcium carbonate → calcium oxide + carbon dioxide.

Calcium hydroxide solution; it turns milky with carbon dioxide and is used to test for CO₂.

The gradual wearing away and changing of rock by heat, water and acid.

Obtaining a metal from its ore by physical means, by heating alone or by heating with carbon.

The smaller particles that make up an atom: protons, neutrons and electrons.

A subatomic particle of relative charge +1 and relative mass about 1; its number fixes which element the atom is.

An uncharged particle of relative mass about 1, found in the nucleus.

A particle of relative charge −1 and relative mass about 11836, moving in shells outside the nucleus.

The central core of protons and neutrons, holding nearly all of the atom's mass.

The proton number Z of an element, fixing its position in the periodic table.

The sum of the proton and neutron numbers: A = Z + N.

Atoms of the same element with the same proton number but different neutron numbers, e.g. ¹²C and ¹³C.

The abundance-weighted average mass of an element's isotopes on the ¹²C = 12.00 scale.

An element with properties between metals and non-metals, such as silicon and boron.

A horizontal row of the periodic table; the period number equals the number of occupied shells.

A vertical column; members share the same outermost electron number and similar chemistry.

The regions around the nucleus where electrons reside, holding up to 2, 8 and 8 (up to Z = 20).

How electrons are written shell by shell, e.g. 2,8,1 for sodium.

A Group I metal — soft, reactive, and reacts with water to give hydrogen.

A Group VII non-metal, existing as diatomic molecules, reactivity decreasing down the group.

A Group 0 element with a full outermost shell; very unreactive, e.g. helium and argon.

The rule that a full outermost shell of 8 electrons is stable, explaining bonding tendencies.

A molecule of two atoms, such as H₂, O₂ and Cl₂.

The electrostatic attraction between cations and anions that holds an ionic compound together.

An ion formed by losing electrons and carrying positive charge, e.g. Na⁺.

An ion formed by gaining electrons and carrying negative charge, e.g. Cl⁻.

The handing-over of electrons from metal atoms to non-metal atoms — the key step in forming ionic bonds.

The attraction between opposite charges — the nature of the ionic bond.

The huge regular lattice of alternating cations and anions, as in sodium chloride.

The temperature at which a solid melts; ionic compounds have high values, e.g. NaCl at 801°C.

The ability to carry current; ionic compounds conduct only when molten or dissolved.

A diagram using dots and crosses to show outermost electrons and how bonding occurs.

The regular, repeating three-dimensional arrangement of ions.