Periodic Table
Concise revision notes aligned to the O Level syllabus.
Structure of the Periodic Table
Elements are arranged in order of increasing atomic number (number of protons). The table has two axes:
- Groups (vertical columns, 1–7 then 0) — elements in the same group have the same number of outer electrons and share similar chemical properties.
- Periods (horizontal rows, 1–7) — each period represents one filled electron shell.
A diagonal "staircase" separates metals (left) from non-metals (right). Elements on the boundary — silicon, germanium, arsenic — are metalloids with intermediate properties.
The atomic number equals the number of protons, which equals the number of electrons in a neutral atom. The period number tells you how many electron shells the atom has.
Group 1 — Alkali Metals
Li · Na · K · Rb · Cs · Fr. All have one outer electron, making them highly reactive.
| Element | Symbol | Period | Reaction with water |
|---|---|---|---|
| Lithium | Li | 2 | Slow fizzing, floats |
| Sodium | Na | 3 | Fast, melts into ball, moves on surface |
| Potassium | K | 4 | Violent, lilac flame, may ignite hydrogen |
| Rubidium | Rb | 5 | Very violent |
| Caesium | Cs | 6 | Explosively violent |
All reactions with water produce a metal hydroxide + hydrogen:
2Na + 2H₂O → 2NaOH + H₂
Reactivity increases down Group 1. The outer electron is further from the nucleus and more easily lost. This is the key explanatory principle for all Group 1 reactivity trends.
They also react with:
- Oxygen → metal oxide (e.g. 4Na + O₂ → 2Na₂O)
- Chlorine → metal chloride (e.g. 2Na + Cl₂ → 2NaCl)
Group 7 — Halogens
F · Cl · Br · I · At. All have seven outer electrons and exist as diatomic molecules.
| Halogen | Symbol | State at r.t.p. | Colour / appearance |
|---|---|---|---|
| Fluorine | F₂ | Gas | Pale yellow |
| Chlorine | Cl₂ | Gas | Yellow-green |
| Bromine | Br₂ | Liquid | Orange-brown, dense vapour |
| Iodine | I₂ | Solid | Grey-black, purple vapour |
What is a Halide?
A halide is a compound formed when a halogen gains one electron to become a negatively charged ion (X⁻). The ion is named by adding -ide to the halogen's root:
| Halogen | Ion formed | Ion name |
|---|---|---|
| Fluorine | F⁻ | Fluoride |
| Chlorine | Cl⁻ | Chloride |
| Bromine | Br⁻ | Bromide |
| Iodine | I⁻ | Iodide |
Halides occur in ionic salts (e.g. NaCl, KBr, NaI) and in molecular compounds like HCl. The halide ion has a full outer shell of 8 electrons and is stable.
Reactions with Hydrogen
Each halogen reacts with hydrogen to form a hydrogen halide gas, which dissolves in water to give a strong acid (except HF, which is weak):
| Reaction | Product | Acid formed |
|---|---|---|
| H₂ + F₂ → 2HF | Hydrogen fluoride | Hydrofluoric acid (weak) |
| H₂ + Cl₂ → 2HCl | Hydrogen chloride | Hydrochloric acid (strong) |
| H₂ + Br₂ → 2HBr | Hydrogen bromide | Hydrobromic acid (strong) |
| H₂ + I₂ ⇌ 2HI | Hydrogen iodide | Hydroiodic acid (strong) |
The reaction with iodine is reversible (⇌) and less vigorous — consistent with iodine being the least reactive of the four.
Displacement Reactions
A more reactive halogen displaces a less reactive halide ion from solution. The colour change is the observable evidence.
| Halogen added | Halide solution | Observation | Equation |
|---|---|---|---|
| Cl₂ | KBr (aq) | Colourless → orange-brown | Cl₂ + 2KBr → 2KCl + Br₂ |
| Cl₂ | KI (aq) | Colourless → brown (iodine released) | Cl₂ + 2KI → 2KCl + I₂ |
| Br₂ | KI (aq) | Orange → darker brown | Br₂ + 2KI → 2KBr + I₂ |
| Br₂ | KCl (aq) | No change | No reaction — Br is less reactive than Cl |
| I₂ | KCl (aq) | No change | No reaction |
| I₂ | KBr (aq) | No change | No reaction |
Reactivity decreases down Group 7. Each successive halogen needs to gain one electron to complete its outer shell, but the increasing atomic radius makes the nucleus less able to attract that electron.
Testing for Halide Ions (Silver Nitrate Test)
To identify which halide ion is present in a solution, add dilute nitric acid (to remove interfering ions), then add silver nitrate solution (AgNO₃):
| Halide ion | Precipitate formed | Colour | Solubility in ammonia |
|---|---|---|---|
| Cl⁻ | AgCl | White | Dissolves in dilute ammonia |
| Br⁻ | AgBr | Cream | Dissolves only in concentrated ammonia |
| I⁻ | AgI | Yellow | Insoluble in ammonia |
The ionic equation is the same for all three: Ag⁺ + X⁻ → AgX↓
The ammonia solubility test distinguishes white AgCl from cream AgBr when the colours are hard to tell apart.
White–Cream–Yellow is the order for AgCl, AgBr, AgI. Going down the group the precipitate gets darker and less soluble in ammonia — a direct parallel to decreasing reactivity.
Group 0 — Noble Gases
He · Ne · Ar · Kr · Xe · Rn. Full outer shells make them chemically inert — they form virtually no compounds under normal conditions.
| Gas | Symbol | Key use |
|---|---|---|
| Helium | He | Balloons, airships (non-flammable) |
| Neon | Ne | Advertising signs (glows red-orange) |
| Argon | Ar | Filling light bulbs, welding shield gas |
| Krypton | Kr | High-speed flash photography |
| Xenon | Xe | Anaesthesia, high-intensity lamps |
Transition Metals
The block of elements between Groups 2 and 3 (periods 4–7). Key properties:
- High melting points and densities — strong metallic bonding
- Variable oxidation states — e.g. iron is Fe²⁺ or Fe³⁺; copper is Cu⁺ or Cu²⁺
- Coloured compounds — CuSO₄ (blue), Fe₂O₃ (red-brown), K₂Cr₂O₇ (orange)
- Good catalysts — iron in the Haber process, platinum in catalytic converters, MnO₂ as a catalyst for H₂O₂ decomposition
Common examples to know: Fe (iron), Cu (copper), Zn (zinc), Ni (nickel), Cr (chromium), Mn (manganese), Co (cobalt).
Trends Across the Table
More protons pull electrons closer across a period; more shells add size down a group.
Harder to remove electrons when nucleus pull is stronger; easier when outer shell is further away.
More protons attract bonding electrons more strongly; larger atoms attract them less.
Elements become less metallic across a period; outer electrons are more easily lost down a group.
Period 3 Elements
Na through Ar — a useful cross-section showing the full range from reactive metal to inert gas within a single period.
Very reactive, reacts vigorously with water
Moderately reactive, burns with bright white flame
Forms protective oxide layer, resists corrosion
Semiconductor, used in electronics
Two allotropes: white (reactive) and red
Yellow solid, used to make sulfuric acid
Yellow-green gas, Group 7 halogen
Colourless, odourless, completely inert
Common Symbols to Memorise
These symbols come from Latin or German names and don't follow the obvious pattern:
| Element | Symbol | Origin |
|---|---|---|
| Sodium | Na | Natrium (Latin) |
| Potassium | K | Kalium (Latin) |
| Iron | Fe | Ferrum (Latin) |
| Copper | Cu | Cuprum (Latin) |
| Silver | Ag | Argentum (Latin) |
| Gold | Au | Aurum (Latin) |
| Lead | Pb | Plumbum (Latin) |
| Mercury | Hg | Hydrargyrum (Greek/Latin) |
| Tungsten | W | Wolfram (German) |
| Tin | Sn | Stannum (Latin) |
These are the highest-frequency symbols in O Level exams. If you know all ten cold, you've neutralised the easiest marks in any periodic table question.
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