What are periodic trends in the context of the Periodic Table?

Periodic trends refer to the patterns observed in the properties of elements as you move across a period or down a group in the Periodic Table. These trends include variations in atomic radius, ionization energy, electronegativity, and electron affinity. Understanding these trends helps predict the chemical behavior of elements, which is crucial for Cambridge IGCSE Chemistry students.

How does atomic radius change across a period and down a group?

As you move across a period from left to right in the Periodic Table, the atomic radius generally decreases. This is because the increasing number of protons in the nucleus pulls the electron cloud closer. Conversely, as you move down a group, the atomic radius increases due to the addition of electron shells, which outweighs the increased nuclear charge. This concept is fundamental for students studying Periodic Trends in Cambridge IGCSE Chemistry.

What is the trend in ionization energy across a period and down a group?

Ionization energy tends to increase across a period from left to right. This is due to the increased nuclear charge, which holds the electrons more tightly, making them harder to remove. Conversely, ionization energy decreases as you move down a group because the outer electrons are further from the nucleus and are more easily removed. Understanding ionization energy is essential for mastering Periodic Trends in the Cambridge IGCSE Chemistry syllabus.

Why does electronegativity vary across periods and groups in the Periodic Table?

Electronegativity increases across a period from left to right due to the increasing nuclear charge, which attracts bonding electrons more strongly. However, it decreases down a group because the additional electron shells reduce the nucleus's pull on the bonding electrons. This trend is crucial for predicting how elements will interact in chemical reactions, a key aspect of the Cambridge IGCSE Chemistry curriculum.

How does electron affinity change across a period and down a group?

Electron affinity generally becomes more negative across a period, indicating a stronger attraction for additional electrons due to increased nuclear charge. However, it becomes less negative down a group as the added electron shells reduce the nucleus's effective pull on additional electrons. This trend helps explain the reactivity of nonmetals and is an important concept for students studying Periodic Trends in Cambridge IGCSE Chemistry.

Why is "Saying reactivity increases down BOTH groups" a common mistake in Periodic Trends?

Why it happens: Generalising the rule. How to avoid it: Group I: increases DOWN (lose more easily). Group VII: DECREASES down (gain less easily). Source: 0620/42 — every series

Why is "Skipping shielding in the explanation" a common mistake in Periodic Trends?

Why it happens: Forgetting the structural reason. How to avoid it: Always include: atomic radius CHANGES + SHIELDING + ease of electron transfer.

Why is "Saying bromine displaces chlorine" a common mistake in Periodic Trends?

Why it happens: Inverting reactivity. How to avoid it: More reactive halogen displaces less reactive: F > Cl > Br > I.

The Periodic TableCambridge IGCSE Chemistry 0620 Extended

Periodic Trends

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Periodic Trends — Cambridge IGCSE 0620 Chemistry Extended (2026)

How atomic size, metallic character and reactivity change across periods and down groups. Explained by electron shells and shielding.

What you’ll learn

Mapped to the Cambridge IGCSE 0620 syllabus (2026-2028).

10.2 — Describe trends in atomic radius and reactivity across periods and down groups.
10.2 — Explain trends in terms of electronic structure.

Atomic size (radius)

Across period: smaller. Down group: bigger.

Across a period. Atomic radius DECREASES.

Same number of shells.
Increasing number of protons.
Stronger nuclear pull on the outer electrons.
Outer electrons drawn closer → smaller atom.

Worked. Period 3: Na ( $\sim 186\,\text{pm}$ ) > Mg ( $160$ ) > Al ( $143$ ) > Si ( $117$ ) > P ( $110$ ) > S ( $104$ ) > Cl ( $99$ ) > Ar ( $71$ ).

Down a group. Atomic radius INCREASES.

Adding a new shell each row.
Outer electrons further from the nucleus.
More inner shells SHIELD the outer electrons from the nucleus.
Net pull on outer electrons WEAKER → atom larger.

Worked. Group 1: Li ( $\sim 152\,\text{pm}$ ) < Na ( $186$ ) < K ( $227$ ) < Rb ( $248$ ) < Cs ( $265$ ).

Across a period the nucleus pulls outer electrons closer; down a group each new shell pushes them further out.

Across: more protons → smaller.
Down: more shells + shielding → bigger.
Period 3: Na largest, Ar smallest.
Group 1: Li smallest, Cs largest.

Metallic character

How easily an element loses electrons. Across: less metallic. Down: more metallic.

Metallic character = ease of losing outer-shell electrons (forming positive ions).

Across a period: DECREASES.

Atoms get smaller; outer electrons closer to nucleus.
Pull on outer e⁻ stronger → harder to lose.
More electronegative — easier to GAIN electrons (non-metal behaviour).
Period 3: Na, Mg, Al — metals (lose e⁻). Si — metalloid. P, S, Cl — non-metals (gain e⁻).

Down a group: INCREASES.

Atoms get bigger; outer electrons further from nucleus.
Shielding by inner shells weakens the pull.
Outer e⁻ easier to lose → more metallic.
Group 14: C (non-metal) → Si (metalloid) → Ge (metalloid) → Sn, Pb (metals).

Cambridge tip. When asked "explain why potassium is more metallic than sodium" — say:

K has MORE shells than Na.
Outer e⁻ further from nucleus, more shielded.
Easier to lose.
More metallic.

Always include both shells/shielding AND the consequence (easier to lose).

Metallic = easy to lose e⁻.
Across period: less metallic (smaller atoms).
Down group: more metallic (more shells + shielding).
K more metallic than Na due to more shells.

Group 1 vs Group 17 reactivity

Group 1 reactivity: ↑ down. Group 17: ↓ down. Both explained by atom size.

Group 1 (alkali metals): reactivity INCREASES down the group. Reactivity = how readily the metal LOSES its 1 outer electron.

Atoms get bigger going down.
Outer e⁻ further from nucleus → weaker pull.
Easier to lose → more reactive.

Order of reactivity: Li < Na < K < Rb < Cs.

Worked qualitative. Caesium reacts EXPLOSIVELY with water (often dropped through the surface). Sodium reacts vigorously (fizzes, sometimes ignites). Lithium reacts more sedately (slow fizz). Same chemistry ( $2\text{M} + 2\text{H}_{2}\text{O} \to 2\text{MOH} + \text{H}_{2}$ ) — just different vigour.

Group 17 (halogens): reactivity DECREASES down the group. Reactivity = how readily the non-metal GAINS 1 electron.

Atoms get bigger going down.
Incoming e⁻ would join an outer shell that's farther from the nucleus, more shielded.
Pull on incoming e⁻ weaker → harder to gain → less reactive.

Order of reactivity: F > Cl > Br > I.

Why opposite directions? For metals (electron LOSS), bigger = better. For non-metals (electron GAIN), bigger = worse. Same physics; different role for the outer shell.

Bigger atoms lose electrons more easily but gain them less easily — so the two groups trend in opposite directions.

Halogen displacement. A more reactive halogen displaces a less reactive halogen from a salt: $\text{Cl}_{2} + 2\text{KBr} \to 2\text{KCl} + \text{Br}_{2}.$ Cl displaces Br. The brown colour of Br₂ appears.

G1 reactivity ↑ down: bigger atom, easier to lose e⁻.
G17 reactivity ↓ down: bigger atom, harder to gain e⁻.
Same explanation from atom size — opposite consequences.
Halogen displacement reactions follow the trend.

How it’s examined

Periodic trends are examined every Paper 4 (5-7 marks): explain a trend using shell structure. Examiner reports flag students stopping at 'atom is bigger' without linking to easier electron loss / harder electron gain.

Sources: Cambridge IGCSE Chemistry 0620 syllabus 2026-2028 (10.2); 0620/42 Oct/Nov 2024 — Q16 (periodic trends); 0620 Examiner Reports 2022-2024. Last reviewed 2026-05-07.

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Step-by-step worked examples — Periodic Trends

Step-by-step solutions to past-paper-style questions on periodic trends, written exactly the way a tutor would explain them at the board.

1Reactivity trend in Group I
Extended• Adapted from 0620/42 May/Jun 2024 Q11• Group I
Question
Explain why reactivity INCREASES down Group I.
Step-by-step solution
1. Step 1
  Atoms get LARGER down the group (more shells).
2. Step 2
  Outer electron is FURTHER from the nucleus and SHIELDED by inner shells.
3. Step 3
  Easier to LOSE the outer electron → more reactive.
Answer
Bigger atoms + more shielding → outer electron lost more easily → reactivity increases.
2Reactivity trend in Group VII
Extended• Group VII
Question
Explain why reactivity DECREASES down Group VII.
Step-by-step solution
1. Step 1
  Atoms get larger down the group → outer shell further from nucleus.
2. Step 2
  Group VII atoms GAIN an electron. Bigger atoms → harder to attract an electron.
Answer
Larger atoms attract incoming electrons less strongly → reactivity decreases.
Examiner tip
Direction of trend depends on whether atoms LOSE (G-I) or GAIN (G-VII) electrons.
3Halogen displacement
Extended• displacement
Question
What happens when chlorine is bubbled into potassium bromide solution?
Step-by-step solution
1. Step 1
  Cl is more reactive than Br → Cl displaces Br.
2. Step 2
  Solution turns from colourless to brown/orange (Br₂).
  $\text{Cl}_{2} + 2\text{KBr} \to 2\text{KCl} + \text{Br}_{2}$
Answer
Solution turns brown/orange as $\text{Br}_{2}$ is displaced.

Key Definitions and Keywords — Periodic Trends

Definitions to memorise and the exact keywords mark schemes credit for periodic trends answers — sharpened from recent examiner reports for the 2026 0620 sitting.

Shielding
Examiner keyword
Inner-shell electrons reduce the effective attraction of the nucleus on outer-shell electrons.
Displacement reaction
Examiner keyword
More reactive element takes the place of a less reactive one in a compound.

Common Mistakes and Misconceptions — Periodic Trends

The traps other students keep falling into on periodic trends questions — taken from recent Cambridge IGCSE 0620 examiner reports and mark schemes — and how to avoid them.

✕Saying reactivity increases down BOTH groups
0620/42 — every series
Why it happens
Generalising the rule.
How to avoid it
Group I: increases DOWN (lose more easily). Group VII: DECREASES down (gain less easily).
✕Skipping shielding in the explanation
Why it happens
Forgetting the structural reason.
How to avoid it
Always include: atomic radius CHANGES + SHIELDING + ease of electron transfer.
✕Saying bromine displaces chlorine
Why it happens
Inverting reactivity.
How to avoid it
More reactive halogen displaces less reactive: F > Cl > Br > I.

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Periodic Trends — frequently asked questions

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Periodic Trends

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Detailed Notes

What you’ll learn

How it’s examined

1Reactivity trend in Group I

2Reactivity trend in Group VII

3Halogen displacement

Shielding

Displacement reaction

✕Saying reactivity increases down BOTH groups

✕Skipping shielding in the explanation

✕Saying bromine displaces chlorine

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Assess your understanding

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Periodic Trends — frequently asked questions