Module 2
Periodic table & separation techniques
Quick-reference revision notes for parents.
2.1 Three elements (atoms, elements, compounds)
- Element: made of only one type of atom (e.g. iron, oxygen).
- Compound: two or more elements chemically bonded together (e.g. water H₂O, salt NaCl).
- Atom: the smallest particle of an element. Each element has its own symbol (Fe, O, Na).
2.2 Physical properties of metals and non-metals
| Metal | Non-metal |
|---|---|
| Shiny when polished | Dull |
| Good conductor (heat & electricity) | Poor conductor |
| Malleable (can be hammered into shape) | Brittle (snap) |
| Ductile (can be pulled into wires) | Not ductile |
| Usually solid at room temp (except mercury) | Often gas or low-melting solid |
| High density | Low density |
2.3 Chemical properties of metals and non-metals
- Metals form positive ions; react with acids → salt + hydrogen.
- Non-metals form negative ions (or share); often react with metals to form ionic compounds.
- Metal oxides are basic; non-metal oxides are acidic.
2.4 Groups and periods
- Group = column. Elements in the same group have similar properties (same number of outer electrons).
- Period = row. Properties change gradually across a period.
- Metals on the left and middle; non-metals on the right.
2.5 Group 1 — Alkali metals
Lithium (Li), sodium (Na), potassium (K), rubidium (Rb)…
- Soft (cut with a knife), low density, shiny when freshly cut, tarnish quickly in air.
- React with water → metal hydroxide + hydrogen gas. Solution is alkaline.
- Reactivity increases down the group (Li < Na < K — increasingly violent).
2.6 Group 7 — Halogens
Fluorine (F), chlorine (Cl), bromine (Br), iodine (I).
- Diatomic molecules (Cl₂, Br₂, …).
- Coloured non-metals — pale yellow gas (F) → orange liquid (Br) → grey solid (I).
- Reactivity decreases down the group (opposite of Group 1).
- A more reactive halogen will displace a less reactive one from a salt.
2.7 Group 0 — Noble gases
- Helium (He), neon (Ne), argon (Ar)…
- Full outer electron shell → inert (don't react).
- Used where unreactivity matters (helium balloons, argon in light bulbs).
2.8 Pure substances
A pure substance contains only one element or one compound. It melts and boils at sharp, fixed temperatures.
2.9 Mixtures
Two or more substances not chemically joined — they keep their own properties and can be separated physically.
2.10 Solutions
- Solute: the substance that dissolves (e.g. salt).
- Solvent: the substance it dissolves in (e.g. water).
- Solution: the mixture (salty water).
- Concentration: how much solute per amount of solvent. Concentrated = lots; dilute = little.
2.11 Solubility
The maximum mass of solute that can dissolve in a given mass of solvent at a given temperature. For most solid solutes, solubility increases with temperature.
2.12 Filtration
Separates an insoluble solid from a liquid. Solid stays on the filter paper (residue); liquid passes through (filtrate).
2.13 Evaporation and distillation
- Evaporation: gentle heating drives off the solvent, leaving the solid solute behind. Used to recover salt from salt water.
- Distillation: heat to boiling, condense the vapour back to liquid in a different container. Used to recover the solvent (e.g. pure water from sea water) — or to separate two liquids with different boiling points (fractional distillation).
2.14 Chromatography
Separates substances dissolved in the same solvent (e.g. inks). Each substance travels at its own rate up the paper.
Rf = distance moved by spot ÷ distance moved by solvent. Each pure substance has its own Rf in a given solvent — useful for identification.
Quick reference
- Metals: shiny, conducts, malleable. Non-metals: dull, brittle, poor conductor
- Group 1 reactivity ↑ down; Group 7 reactivity ↓ down
- Group 0 = inert (full outer shell)
- Filtration: insoluble solid from liquid
- Evaporation: keep the solute. Distillation: keep the solvent
- Chromatography: separates dissolved substances on paper