The Periodic Table: Groups and Trends — Revision Notes

Condensed for the final weeks. For the full explanation, use the Periodic Table: Groups and Trends study guide.

How the table is arranged

• Groups (columns) = number of outer-shell electrons. Same group → similar chemical properties, since chemical behaviour is driven almost entirely by the outer shell.
• Periods (rows) = number of occupied shells.
• Metals on the left, non-metals on the right, with a staircase boundary.

Arranged in order of proton (atomic) number — this single ordering principle is what makes the whole table’s structure, and every trend within it, predictable rather than arbitrary.

Group I — the alkali metals

Li, Na, K, Rb, Cs

Going down the group:

Property	Trend	Why
Reactivity	Increases	Outer electron further from nucleus, more shielding → lost more easily
Melting point	Decreases	Weaker metallic bonding
Density	Increases (generally)

2Na + 2H2O -> 2NaOH + H2

Lithium  fizzes steadily
Sodium   melts into a ball, moves rapidly
Potassium ignites with a LILAC flame

All are soft, low-density metals stored under oil; all form 1+ ions and alkaline hydroxides, since each loses its single outer electron in exactly the same way, differing only in how easily.

Group VII — the halogens

F, Cl, Br, I

Going down the group:

Property	Trend
Reactivity	Decreases (opposite of Group I)
Melting/boiling point	Increases
Colour	Darkens: pale yellow → green → red-brown → grey-black
State at rtp	Gas → gas → liquid → solid

Reactivity falls because the outer shell is further from the nucleus, so an electron is gained less easily — the exact mirror image of Group I’s explanation, but for gaining rather than losing an electron.

Displacement: a more reactive halogen displaces a less reactive halide from solution. Cl₂ + 2KBr → 2KCl + Br₂ (solution turns orange)

All exist as diatomic molecules and form 1− ions.

The same reactivity order (Cl > Br > I) is exactly why chloride, bromide and iodide give progressively less soluble, differently coloured precipitates with silver nitrate — a direct link to the halide ion tests in qualitative analysis, and a reminder that the two topics describe the same underlying trend from different angles.

Transition elements vs Group I

	Transition elements	Group I
Density	High	Low
Melting point	High	Low
Reactivity	Low	Very high
Ion charges	Variable (Fe²⁺/Fe³⁺)	Always 1+
Compounds	Coloured	White/colourless
Catalysts	Often	No

Iron is 0620’s named example of variable oxidation numbers (5070 states the general property without naming a specific element): iron(II) compounds (e.g. iron(II) sulfate) and iron(III) compounds (e.g. iron(III) chloride) are both stable, ordinary compounds — the same metal simply forms two different, equally valid ions. This is exactly the property behind the iron(II)/iron(III) hydroxide precipitate colours (green vs red-brown) used to distinguish them in qualitative analysis.

Noble gases

Group VIII (not Group 0) — full outer shell, therefore unreactive (monatomic). Uses: helium in balloons (low density, non-flammable), argon in lamps and welding (inert atmosphere), neon in lighting. “Monatomic” means existing as single, separate atoms rather than as molecules — unlike Group VII, which is diatomic — since there is no bonding pulling their atoms together. See the Periodic Table: Groups and Trends study guide for the full syllabus coverage and worked reasoning behind every trend above.

Exam traps

• Group I reactivity increases down; Group VII decreases down. Opposite directions.
• Explain reactivity by distance from the nucleus and shielding, not “it just is”.
• Halogens are diatomic — write Cl₂, not Cl.
• Transition metals are not in a numbered group in this syllabus.
• Noble gases have a full outer shell — helium’s is 2, not 8.
• Claiming halogen reactivity explains silver halide solubility — the two trends happen to run in the same order (Cl > Br > I) as a memory aid, but they arise from different underlying causes, not a cause-and-effect relationship.
• Treating iron(II) and iron(III) as if one were “more correct” than the other — both are equally valid, stable oxidation states of the same element.

Self-test

1. Why is potassium more reactive than sodium?
2. Predict the state and colour of astatine at room temperature.
3. Will chlorine displace iodine from potassium iodide? Write the equation.
4. Give three ways transition elements differ from Group I metals.
5. Why are noble gases unreactive?
6. Describe how the order of halogen reactivity compares with the order of silver halide solubility, and how this can be used as a memory aid.
7. Name 0620’s example element for variable oxidation numbers, and give its two named ion charges.

Answers: 1. Its outer electron is in a shell further from the nucleus with more shielding, so it is lost more easily. 2. Below iodine → solid, very dark/black. 3. Yes — chlorine is more reactive: Cl₂ + 2KI → 2KCl + I₂. 4. Any three: higher density, higher melting point, lower reactivity, variable ion charge, coloured compounds, act as catalysts. 5. They have a full outer shell, so they have no tendency to lose, gain or share electrons. 6. Halogen reactivity decreases Cl > Br > I, and silver halide solubility happens to decrease in the same order — silver iodide is the least soluble and most strongly coloured precipitate. The two trends follow the same order, which can help you remember one from the other, but reactivity does not explain the solubility trend; they arise from different underlying causes. 7. Iron; Fe²⁺ (iron(II)) and Fe³⁺ (iron(III)).