Today we're looking at one of the most powerful ideas in chemistry: periodicity. This is the recurring pattern in properties of elements as you move across a period or down a group in the periodic table. Once you understand why these patterns exist, you can predict properties of elements you've never even studied before.
Why does periodicity happen?
The periodic table is arranged so that elements with similar electron configurations line up in the same group (Talbot, p. 233). Since chemical and physical properties are governed largely by the arrangement of electrons — especially the valence electrons — elements in the same group tend to behave in similar ways, and properties change in a predictable, repeating pattern as you move across each period.
- The period number tells you which outer energy level (shell) is being filled with electrons.
- Elements in the same group share the same number of valence electrons (Talbot, p. 233).
This is why, for example, all group 1 elements are soft, reactive metals that form 1+ ions, while all group 17 elements are reactive non-metals that form 1− ions.
Trends across a period
As you move left to right across a period:
- Nuclear charge increases (more protons), pulling electrons closer to the nucleus.
- Atomic radius decreases, since electrons are added to the same shell but are pulled in more strongly.
- Ionization energy increases, because it becomes harder to remove an electron held tightly by a stronger nuclear pull.
- Electronegativity increases, as atoms attract bonding electrons more strongly.
- Metallic character decreases while non-metallic character increases.
Trends down a group
As you move down a group:
- Atomic radius increases, because each new period adds an extra electron shell, moving electrons further from the nucleus.
- Ionization energy decreases, since outer electrons are further away and shielded by inner shells, making them easier to remove.
- In group 1, metallic character increases down the group (Talbot, p. 233) — elements become more reactive as it's easier to lose their single valence electron.
- In group 17, non-metallic character decreases down the group (Talbot, p. 233) — elements become less reactive as it's harder to attract an extra electron.
From metals to non-metals: a continuum
Talbot (p. 233) highlights that metallic and non-metallic character isn't a strict on/off switch — it's a continuum. This shows up nicely in the oxides elements form:
- Basic oxides — formed by metals (e.g. Na₂O, MgO)
- Amphoteric oxides — behave as both acidic and basic (e.g. Al₂O₃)
- Acidic oxides — formed by non-metals (e.g. SO₃, CO₂)
Worked examples
Example 1: Predict how atomic radius compares between sodium (Na) and chlorine (Cl), both in period 3. Since Cl has a higher nuclear charge and both add electrons to the same shell, Cl has a smaller atomic radius than Na.
Example 2: Predict which is more reactive: potassium (K) or sodium (Na), both group 1 metals. K is below Na, so its valence electron is further from the nucleus and less strongly held → K is more reactive.
Example 3: Classify the oxide Al₂O₃ as basic, acidic, or amphoteric. Aluminium sits near the metal/non-metal boundary, so Al₂O₃ is amphoteric — it reacts with both acids and bases.
Key takeaways
- Periodicity describes repeating trends in properties across periods and down groups, caused by patterns in electron configuration.
- Across a period: atomic radius decreases, ionization energy and electronegativity increase, metallic character decreases.
- Down a group: atomic radius increases, ionization energy decreases; group 1 metallic character increases while group 17 non-metallic character decreases.
- Metal/non-metal character is a continuum, reflected in basic → amphoteric → acidic oxides.
- Understanding these trends lets you predict properties of unfamiliar elements just from their position in the table.