Chemistry
Periodic table trends — quick study summary
The periodic table is organised so elements with similar properties fall in the same group (column). Four trends explain almost every exam question: atomic radius decreases left-to-right and increases top-to-bottom; ionisation energy and electronegativity rise toward the top-right (fluorine area); metallic character rises toward the bottom-left (caesium area). All trends come from nuclear charge and shielding.
Key points
- Atomic radius: ↓ across a period (more protons pulling electrons in), ↑ down a group (more shells)
- Ionisation energy: ↑ across (harder to remove from smaller atom), ↓ down a group
- Electronegativity: ↑ toward fluorine (top-right, excluding noble gases)
- Metallic character: ↑ down-left toward caesium; non-metals dominate top-right
- Group 1 (alkali metals): one outer electron, very reactive, increasingly so down the group
- Group 7 (halogens): seven outer electrons, very reactive non-metals, reactivity DECREASES down the group
Practice quiz
Click each question to reveal the answer.
1. Which element has a larger atomic radius — lithium or caesium?
- Lithium
- Caesium
- They are equal
- Hydrogen
Answer: Caesium
Atomic radius increases down a group because each new period adds an electron shell.
2. Does ionisation energy increase or decrease as you move down Group 1?
Answer: Decrease
Outer electrons are further from the nucleus and more shielded, so they're easier to remove.
3. Which is more electronegative — fluorine or chlorine?
Answer: Fluorine
Fluorine is smaller, so its nucleus pulls bonding electrons more strongly.
4. Why does reactivity DECREASE down Group 7 (halogens)?
Answer: Halogens gain electrons; larger atoms attract them less strongly so reactivity falls
The opposite of Group 1 metals, which give up electrons more easily as size increases.
Last reviewed: May 2026