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 "Giving metals a negative charge or non-metals a positive charge." a common mistake in Periodic Trends?

Why it happens: Confusing which type of element loses and which gains electrons. How to avoid it: Metals LOSE electrons → POSITIVE ions; non-metals GAIN electrons → NEGATIVE ions. Group I → +1, VII → −1. Source: 0620/42 — recurring

Why is "Saying a Group VII element forms a +7 ion (or Group VI a +6)." a common mistake in Periodic Trends?

Why it happens: Wrongly applying 'charge = group number' to non-metals. How to avoid it: For non-metals use charge = (group number − 8): Group VII → −1, Group VI → −2. 'Charge = group number' only works for metals in Groups I–III.

Why is "Saying reactivity simply 'increases' across a period." a common mistake in Periodic Trends?

Why it happens: Trying to force a single-direction trend like the ones down groups. How to avoid it: Across a period the key change is in CHARACTER: metallic → non-metallic. Describe the change in character, not a blanket reactivity direction.

Why is "Misreading position when classifying an element as metal or non-metal." a common mistake in Periodic Trends?

Why it happens: Not relating the element to the dividing line. How to avoid it: Picture the staircase dividing line: left of it = metal, top-right of it = non-metal. The further right, the more non-metallic.

Cambridge IGCSE Chemistry (0620)

Periodic Trends

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Short Notes - Periodic Trends

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Detailed Study Notes

Detailed notes on The Periodic Table for Cambridge IGCSE Chemistry, covering key concepts, explanations, examples, and exam-focused revision points.

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.

At a glance

Atomic size: DECREASES across a period; INCREASES down a group.
Across a period: same shells, more protons → stronger pull → smaller atom.
Down a group: more shells, more shielding → outer electrons further out, weaker pull.
Metallic character: DECREASES across a period; INCREASES down a group.
Group 1 reactivity: increases DOWN the group (easier to lose outer e⁻).
Group 17 reactivity: decreases DOWN the group (harder to attract incoming e⁻).
Two-step explanation: (1) shell structure / shielding, (2) effect on electron loss/gain.

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.

Quick recap

Atomic radius: across ↓, down ↑.
Metallic character: across ↓, down ↑.
G1 reactivity: ↑ down (easier to lose e⁻).
G17 reactivity: ↓ down (harder to gain e⁻).
Always explain via shells + shielding.

Memorise this

Verbatim phrases and definitions Cambridge mark schemes credit.

Atomic radius — distance from nucleus to outer-shell electrons.
Shielding — reduction in nuclear attraction on outer electrons by inner-shell electrons.
Metallic character — tendency of an element to lose electrons and form positive ions.

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.

Take this whole topic with you

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.

1Metallic to non-metallic character across a period
Getting started• period trend, metallic character
Question
State how the character of the elements changes as you move from left to right across a period (e.g. Period 3, Na → Ar).
Step-by-step solution
1. Step 1
  On the LEFT of a period the elements are metals (e.g. Na, Mg, Al).
2. Step 2
  On the RIGHT of a period the elements are non-metals (e.g. P, S, Cl), ending with a noble gas (Ar).
3. Step 3
  So across a period the elements change from metallic to non-metallic character.
Answer
Across a period (left → right) the elements change from metallic to non-metallic character.
Examiner tip
The reliable phrase Cambridge rewards is 'metallic to non-metallic across a period'. Do not say reactivity simply 'increases' across a period — the trend is about character, not a single direction of reactivity.
2Predicting ionic charge from group number
Getting started• ionic charge, group number
Question
Give the charge of the ion formed by each element: sodium (Group I), magnesium (Group II), aluminium (Group III), oxygen (Group VI) and chlorine (Group VII).
Step-by-step solution
1. Step 1
  Groups I–III are metals — they LOSE their outer electrons, forming positive ions. Charge = group number.
  $\text{Na}^{+}, \quad \text{Mg}^{2+}, \quad \text{Al}^{3+}$
2. Step 2
  Groups VI–VII are non-metals — they GAIN electrons to fill the outer shell. Charge = (group number − 8).
  $\text{O}^{2-}, \quad \text{Cl}^{-}$
Answer
$\text{Na}^{+}$ , $\text{Mg}^{2+}$ , $\text{Al}^{3+}$ , $\text{O}^{2-}$ , $\text{Cl}^{-}$ .
Examiner tip
Group I → +1, II → +2, III → +3; Group VII → −1, VI → −2. Metals lose electrons (positive); non-metals gain electrons (negative).
3Metal or non-metal from position
Building confidence• classification, position
Question
Element X is in Group II and element Y is in Group VI. Using their positions, state whether each is a metal or a non-metal and the charge of the ion it forms.
Step-by-step solution
1. Step 1
  Group II is on the left of the table → X is a metal. It has 2 outer electrons, which it loses.
2. Step 2
  Losing 2 electrons gives a 2+ ion: $\text{X}^{2+}$ .
3. Step 3
  Group VI is on the right → Y is a non-metal. It has 6 outer electrons and gains 2 to fill the shell.
4. Step 4
  Gaining 2 electrons gives a 2− ion: $\text{Y}^{2-}$ .
Answer
X is a metal forming $\text{X}^{2+}$ ; Y is a non-metal forming $\text{Y}^{2-}$ .
Examiner tip
Position decides the class: left = metal (loses e⁻, positive ion), right = non-metal (gains e⁻, negative ion). The dividing line runs as a 'staircase' from around B/Al down to Po/At.
4The metal / non-metal dividing line
Building confidence• dividing line, classification
Question
Describe where the dividing line between metals and non-metals lies in the Periodic Table, and where most metals and most non-metals are found relative to it.
Step-by-step solution
1. Step 1
  There is a 'staircase' dividing line running diagonally down the right-hand side of the table (from around boron/aluminium towards astatine).
2. Step 2
  Metals are to the left of this line — the large majority of elements.
3. Step 3
  Non-metals are to the right of (and above) the line, in the top-right region.
Answer
A diagonal 'staircase' line on the right separates metals (left, the majority) from non-metals (top-right). The closer an element is to the right, the more non-metallic it is.
Examiner tip
Use this line to predict class from position: an element in the top-right region (e.g. Group VI–VII) is a non-metal; one on the far left (Group I–II) is a metal.
5Explaining the charge trend by electron arrangement
Stretch• electron arrangement, ionic charge, explanation
Question
Using electron arrangement, explain why a Group I element forms a +1 ion while a Group VII element forms a −1 ion.
Step-by-step solution
1. Step 1
  A Group I atom has 1 electron in its outer shell. Losing this one electron leaves a full outer shell (a stable noble-gas arrangement).
2. Step 2
  Losing 1 negative electron leaves the ion with one more proton than electron, so the ion has a +1 charge (e.g. $\text{Na} \to \text{Na}^{+} + \text{e}^{-}$ ).
3. Step 3
  A Group VII atom has 7 outer electrons — just one short of a full shell. It is easier to gain 1 electron than to lose seven.
4. Step 4
  Gaining 1 negative electron gives one more electron than protons, so the ion has a −1 charge (e.g. $\text{Cl} + \text{e}^{-} \to \text{Cl}^{-}$ ).
Answer
Group I atoms lose their single outer electron to reach a full shell → +1 ion; Group VII atoms gain one electron to complete their outer shell → −1 ion.
Examiner tip
The reasoning chain Cambridge wants: number of outer electrons → lose or gain to reach a full (noble-gas) shell → resulting charge. Always state whether electrons are LOST or GAINED.
6How physical properties change across a period
Stretch• properties, period trend
Question
Compare two elements in the same period — one on the far left and one on the far right — in terms of conductivity, state and how they form ions. Use sodium and chlorine in Period 3 as your example.
Step-by-step solution
1. Step 1
  Sodium (left, metal): a solid that conducts electricity, shiny when fresh, and forms a positive ion ( $\text{Na}^{+}$ ) by losing its 1 outer electron.
2. Step 2
  Chlorine (right, non-metal): a gas at room temperature that does not conduct electricity, and forms a negative ion ( $\text{Cl}^{-}$ ) by gaining 1 electron.
3. Step 3
  Moving left → right, character changes from metallic to non-metallic: conductivity is lost, and the element changes from electron-loser (positive ions) to electron-gainer (negative ions).
Answer
Sodium (metal, left) conducts and forms $\text{Na}^{+}$ ; chlorine (non-metal, right) does not conduct and forms $\text{Cl}^{-}$ — showing the metallic-to-non-metallic change across the period.
Examiner tip
Strong answers link the property change to the underlying change in character and electron behaviour (lose vs gain), not just listing differences.

Model Answers — Periodic Trends

High-scoring sample answers for periodic trends on the Cambridge IGCSE 0620 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.

Question 1
Paper 4 short-answer style1 mark
State the charge of the ion formed by an element in Group II. (1 mark)
Model answer
A 2+ charge (e.g. $\text{Mg}^{2+}$ , $\text{Ca}^{2+}$ ).
Why this scores
One mark for 2+ (the charge equals the group number for Groups I–III). Accept '+2'.
Question 2
Paper 4 structured style2 marks
Describe how the character of the elements changes across Period 3 from sodium to argon. (2 marks)
Model answer
The elements change from metallic on the left (sodium, magnesium, aluminium) to non-metallic on the right (such as phosphorus, sulfur and chlorine). In other words, metallic character decreases and non-metallic character increases from left to right across the period.
Why this scores
Two marks: (1) metals on the left / non-metals on the right; (2) the direction of the change (metallic → non-metallic across the period).
Question 3
Paper 4 structured style3 marks
Explain, in terms of electrons, why metals form positive ions but non-metals form negative ions. (3 marks)
Model answer
Metals have few electrons in their outer shell, so they lose these electrons to achieve a full outer shell. Losing negative electrons leaves the ion with more protons than electrons, giving a positive charge. Non-metals have almost full outer shells, so it is easier for them to gain electrons to complete the shell. Gaining negative electrons gives the ion more electrons than protons, so it has a negative charge.
Why this scores
Three marks: (1) metals lose electrons → positive ions; (2) non-metals gain electrons → negative ions; (3) link to reaching a full (stable) outer shell. The words 'lose' and 'gain' must be correct.
Question 4
Paper 4 (Extended) structured style4 marks
An element is in Group VI of Period 3. Predict whether it is a metal or a non-metal, the charge of the ion it forms, and explain your answers using its position and electron arrangement. (4 marks)
Model answer
It is a non-metal, because it is on the right-hand side of the Periodic Table, to the right of the dividing line between metals and non-metals. Being in Group VI, it has 6 electrons in its outer shell. It is closer to a full shell to gain 2 electrons than to lose six, so it forms a 2− ion. Gaining 2 negative electrons gives the ion two more electrons than protons, hence the −2 charge (e.g. sulfur forms $\text{S}^{2-}$ ).
Why this scores
Four marks: (1) non-metal; (2) reason from position (right of the table / dividing line); (3) 6 outer electrons → gains 2; (4) charge of 2−. Awarding depends on a correct lose/gain statement.
Question 5
Paper 4 (Extended) structured style5 marks
Explain why the charge of the ion formed by an element can be predicted from its group number, referring to Groups I, II, III, VI and VII. (5 marks)
Model answer
The group number tells you the number of outer-shell electrons. In Groups I, II and III the elements are metals with 1, 2 or 3 outer electrons. They lose these electrons to reach a full outer shell, forming ions with charges of +1, +2 and +3 respectively (e.g. $\text{Na}^{+}$ , $\text{Mg}^{2+}$ , $\text{Al}^{3+}$ ). In Groups VI and VII the elements are non-metals with 6 or 7 outer electrons. It is easier for them to gain electrons to complete the shell: a Group VII atom gains 1 electron to form a −1 ion (e.g. $\text{Cl}^{-}$ ) and a Group VI atom gains 2 electrons to form a −2 ion (e.g. $\text{O}^{2-}$ ). So the charge equals the group number for metals (Groups I–III) and equals the group number minus 8 for non-metals (Groups VI–VII).
Why this scores
Five marks across: group number = outer electrons; Groups I–III lose electrons → +1/+2/+3; Groups VI/VII gain electrons → −2/−1; correct example ions; link to achieving a full outer shell.
Question 6
Paper 4 (Extended) structured style6 marks
Using Period 3, describe and explain how the elements change in character and in the type of ion they form as you move from left to right across the period. (6 marks)
Model answer
Across Period 3 the elements change from metallic to non-metallic character. On the left, sodium, magnesium and aluminium are metals: they have only 1, 2 or 3 outer electrons, which they lose to form positive ions ( $\text{Na}^{+}$ , $\text{Mg}^{2+}$ , $\text{Al}^{3+}$ ). As you move right, the number of outer electrons increases, so the elements become non-metals (such as phosphorus, sulfur and chlorine). These have nearly full outer shells, so it is easier for them to gain electrons to complete the shell, forming negative ions (for example sulfur forms $\text{S}^{2-}$ and chlorine forms $\text{Cl}^{-}$ ). The period ends with the noble gas argon, which has a full outer shell and therefore does not normally form ions. Overall, moving left → right, metallic character decreases, non-metallic character increases, and the elements change from electron-losers (positive ions) to electron-gainers (negative ions).
Why this scores
Up to 6 marks: metallic → non-metallic change; metals on the left lose electrons → positive ions with correct charges; non-metals on the right gain electrons → negative ions with correct charges; outer-electron number increases across the period; noble gas (argon) has a full shell / forms no ions; overall summary of the trend.

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.

Metallic character
Examiner keyword
The tendency of an element to behave like a metal — to lose electrons and form positive ions. It DECREASES from left to right across a period.
Non-metallic character
Examiner keyword
The tendency of an element to behave like a non-metal — to gain electrons and form negative ions. It INCREASES from left to right across a period.
Group number
Examiner keyword
The number of the vertical column, equal to the number of electrons in the outer shell — this determines the charge of the ion an element forms.
Metal / non-metal dividing line
Examiner keyword
The diagonal 'staircase' line on the right of the Periodic Table separating metals (to its left) from non-metals (to its right and above).

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.

✕Giving metals a negative charge or non-metals a positive charge.
0620/42 — recurring
Why it happens
Confusing which type of element loses and which gains electrons.
How to avoid it
Metals LOSE electrons → POSITIVE ions; non-metals GAIN electrons → NEGATIVE ions. Group I → +1, VII → −1.
✕Saying a Group VII element forms a +7 ion (or Group VI a +6).
Why it happens
Wrongly applying 'charge = group number' to non-metals.
How to avoid it
For non-metals use charge = (group number − 8): Group VII → −1, Group VI → −2. 'Charge = group number' only works for metals in Groups I–III.
✕Saying reactivity simply 'increases' across a period.
Why it happens
Trying to force a single-direction trend like the ones down groups.
How to avoid it
Across a period the key change is in CHARACTER: metallic → non-metallic. Describe the change in character, not a blanket reactivity direction.
✕Misreading position when classifying an element as metal or non-metal.
Why it happens
Not relating the element to the dividing line.
How to avoid it
Picture the staircase dividing line: left of it = metal, top-right of it = non-metal. The further right, the more non-metallic.

Periodic Trends — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Periodic Trends

Short Notes - Periodic Trends

Detailed Study Notes

At a glance

What you’ll learn

Quick recap

Memorise this

How it’s examined

Take this whole topic with you

1Metallic to non-metallic character across a period

2Predicting ionic charge from group number

3Metal or non-metal from position

4The metal / non-metal dividing line

5Explaining the charge trend by electron arrangement

6How physical properties change across a period

Metallic character

Non-metallic character

Group number

Metal / non-metal dividing line

✕Giving metals a negative charge or non-metals a positive charge.

✕Saying a Group VII element forms a +7 ion (or Group VI a +6).

✕Saying reactivity simply 'increases' across a period.

✕Misreading position when classifying an element as metal or non-metal.

Periodic Trends — frequently asked questions