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Module 1 · L17 of 21 25 min ⚡ +50 XP in Learn · +25 to complete Year 11 · Module 1 · IQ3

Periodic Trends: Atomic Radius

Today's hook, Cram more protons and more electrons into an atom and it gets smaller, not bigger. Across a period, atoms shrink even as they gain particles.

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Warm up and recall

Warm up first

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

Worksheets

Practise this lesson

Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

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Lithium and fluorine are in the same period (Period 2). Lithium has a larger atomic radius than fluorine. Sodium and chlorine are in the same period (Period 3), and sodium also has a larger atomic radius than chlorine. Yet sodium is larger than lithium, and chlorine is larger than fluorine. What is causing the atomic radius to decrease across a period, and what is causing it to increase down a group?

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What you'll master, and the words for it

What you'll master

Know

Key facts

The direction of the atomic-radius trend: decreases across a period, increases down a group
Definitions of effective nuclear charge and shielding
How forming a cation or anion changes the radius (cations smaller, anions larger)

Understand

Concepts

Why across a period stronger nuclear attraction wins over electron repulsion
Why down a group extra shells make the atom larger despite increasing proton number
Why cations lose a shell (often) and anions expand due to repulsion

Can do

Skills

Rank elements by atomic radius using period and group position
Compare an atom with its ion (e.g. Na vs Na⁺ vs Cl⁻) and justify the order
Write a model exam sentence using shielding and effective nuclear charge to explain a radius comparison

The Main Atomic Radius Trend

core concept

Direction	Trend	Main reason
Across a period (left to right)	Atomic radius decreases	More protons are added, but electrons go into the same main shell. Shielding changes only slightly, so the stronger nuclear attraction pulls the outer electrons closer.
Down a group	Atomic radius increases	A new electron shell is added each step down the group. Outer electrons are further from the nucleus and more shielded, so the atom is larger.

Fast rule: across a period, stronger nuclear pull wins, so atoms get smaller. Down a group, extra occupied shells win, so atoms get larger.

Which radius, and where the simple rule bends. These trends use the covalent (empirical) atomic radius, roughly half the distance between the nuclei of two bonded atoms of the element, so every value measures the same thing and the comparison is fair. Two honest complications: (1) a noble gas forms no ordinary covalent bond, so its radius is quoted on a different basis (van der Waals) and is not directly comparable with the others; (2) across the transition (d-block) metals the radius stays almost flat, because the added electrons go into an inner d subshell, so 'decreases across a period' is a main-group pattern, not a law that holds atom-by-atom everywhere.

Atomic radius across Period 3 as a comparison table

Comparison of Li, Na, K down Group 1 showing atomic radius increasing with proton count and added electron shells.

Atomic radius decreases across a period because protons are added (increasing nuclear charge) while electrons stay in the same shell, effective nuclear charge (Z_eff) increases, pulling outer electrons closer. Atomic radius increases down a group because each step adds an extra occupied electron shell, moving outer electrons further from the nucleus despite more shielding.

Pause, copy the highlighted atomic radius trends into your book before moving on.

Two truths and a lie: pick the false statement about the atomic radius trend.

How to Explain the Trend in Exam Language

core concept

For full-mark responses, avoid saying only “it gets bigger” or “it gets smaller.” State what happens to shells, shielding, and nuclear attraction.

Model sentence, across a period: Atomic radius decreases because proton number increases while electrons are added to the same shell, so shielding changes little and the stronger effective nuclear charge pulls the outer electrons closer to the nucleus.

Model sentence, down a group: Atomic radius increases because each element has an extra occupied shell, so the outer electrons are further from the nucleus and more shielded by inner electrons.

This trend sets up the next lesson. Smaller atoms generally hold outer electrons more tightly, which helps explain why ionisation energy generally increases across a period. The word 'generally' matters: the rise is not perfectly smooth, there are small, predictable dips (for example between Mg and Al, and between P and S) caused by which subshell the electron leaves and by electron pairing, which you will explain when you study ionisation energy in detail.

We just saw the two atomic radius trends. That raises a question: examiners specifically award marks for mechanistic explanations, what precise language is required for across-period and down-group answers? This card answers it → every atomic radius answer must explicitly link Z_eff, shielding, and shell count to the observed change in radius.

Exam language for atomic radius: across a period, "increasing proton number raises Z_eff while electrons enter the same shell, pulling outer electrons closer." Down a group, "each element has one more occupied electron shell, so outer electrons are further from the nucleus and more shielded." Always mention: protons (nuclear charge), shielding (inner electrons), and shell number (distance).

Add the highlighted exam language to your notes before the check below.

Match each comparison to the exam-language explanation it requires.

Across a period: radius decreases
Down a group: radius increases
Atom → cation: radius decreases
Atom → anion: radius increases

Extra electrons added to the same outer shell increase electron-electron repulsion, so the electron cloud expands.
Proton number increases while electrons are added to the same shell; shielding changes little, so effective nuclear charge pulls electrons closer.
Loss of electrons (often an entire outer shell) reduces repulsion, so the remaining electrons are pulled in more tightly.
Each element has an extra occupied shell, so outer electrons are further from the nucleus and more shielded by inner electrons.

Atoms Compared with Ions

core concept

Atomic radius questions often connect to ions. When an atom loses electrons to form a cation, it becomes smaller. When an atom gains electrons to form an anion, it becomes larger.

Species change	Size effect	Reason
Atom → cation	Smaller	There are fewer electrons and often one fewer occupied shell. Electron-electron repulsion drops, so the remaining electrons are pulled in more strongly.
Atom → anion	Larger	Extra electrons increase repulsion in the outer shell, so the electron cloud expands.

We just saw the exact language for atomic radius trend answers. That raises a question: when atoms form ions, how do their sizes change, and why? This card answers it → cations shrink (fewer electrons, often an entire shell removed); anions grow (extra electrons increase repulsion and expand the cloud); isoelectronic series show size decreasing with increasing proton number.

Cation (positive ion): fewer electrons, often loses a whole shell → smaller than parent atom (e.g. Na⁺ < Na). Anion (negative ion): extra electrons in outer shell increase repulsion → cloud expands → larger than parent atom (e.g. Cl⁻ > Cl). For isoelectronic species (same electron count): more protons → stronger pull on same electrons → smaller ion (e.g. Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻).

Pause, write the highlighted cation/anion size rules into your book.

Quick check: which species has the largest radius?

Quick Check

core concept

1. Which atom is larger: Li or K? Explain using shell number and shielding.

2. Which atom is smaller: P or Cl? Explain using effective nuclear charge.

3. Which is larger: Na or Na⁺? Explain why.

We just saw the cation/anion size rules and the isoelectronic series. That raises a question: when a question gives you a specific pair to compare, what is the fastest systematic way to reason about relative size? This card answers it → apply the two-step pattern check: same shell? (more protons wins); different shell? (extra shell wins); isoelectronic? (more protons wins).

Pattern checks: same shell + more protons → smaller radius. Same protons + extra shell → larger radius. Isoelectronic series rule: same electron count, more protons → smaller. Example: Li vs K, K is larger (4 shells vs 2); Na vs Na⁺, Na is larger (Na⁺ lost an entire outer shell). Always connect periodic trends to Z_eff, shielding, and the number of occupied shells.

Pause, copy the highlighted pattern check rules into your book before moving on.

Two truths and a lie: three claims about size comparisons. Which is false?

Worked examples

Worked examples · reveal as you go

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Rank Na, Mg, and Al in order of decreasing atomic radius.

All three elements are in Period 3, so their outer electrons are in the same main shell.

This is the key observation, we're not comparing atoms with different shell counts, just comparing atoms in the same period.

Across a period, proton number increases from Na to Mg to Al, so effective nuclear charge increases.

Na (11p), Mg (12p), Al (13p), as we move right, the nucleus pulls harder on the outer electrons.

Stronger attraction pulls the outer electrons closer, so radius decreases left to right.

Therefore: Na > Mg > Al in order of decreasing atomic radius.

Put the method in order

Predict, then reveal+8 XP

1 · Predict

2 · Reveal

3 · Compare

Rank these four species from smallest to largest atomic/ionic radius: Mg, Mg²⁺, Ca, K. Justify your order.

Confidence: 50%

Misconception to fix

Common errors · the 3 traps that cost marks

Misconception to fix

Wrong: Atomic radius increases across a period because more electrons mean a larger atom.

Misconception to fix

Right: Atomic radius decreases across a period because the increasing nuclear charge pulls electrons closer to the nucleus. The added electrons enter the same shell and do not shield each other effectively, so the effective nuclear charge increases.

Forgetting that cations lose a whole shell, not just one electron

Students often picture Na → Na⁺ as "Na minus one tiny dot" and assume the radius barely changes. In reality, Na is 2,8,1 and Na⁺ is 2,8, the entire third shell is gone, so Na⁺ is dramatically smaller than Na. Forgetting this leads to wrong rankings on isoelectronic series.

Fix: When forming a main-group cation, check whether the lost electron was the last one in its shell, if so, draw the ion with one fewer shell.

Drill, then revisit

Quick-fire practice · 5 reps +2 XP per reveal

Compare the atomic radius of Na and Cl. Explain the trend across Period 3 in terms of nuclear charge and shielding.

Explain why F has a higher electronegativity than Cl, despite both being in Group 17.

Predict which element, Li or Cs, has the higher first ionisation energy. Justify your prediction.

Which is larger: Cl or Cl⁻? Justify your answer.

Rank the isoelectronic species Na⁺, F⁻, Mg²⁺ and O²⁻ from smallest to largest, and explain the rule you used.

Working · 1

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Working · 2

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Revisit your thinking

Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?

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Interactive Tool, Periodic Table Explorer Open fullscreen ↗

The Periodic Table tool shows elements are arranged in periods and groups. Elements in the SAME GROUP have…

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Multiple choice

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Pick your answer, then rate your confidence. That tells the system what to drill next.

Spot the error+5 XP

A student is asked to compare the radii of N (Z=7) and F (Z=9). One line in their working contains an error, click it.

Both N and F are in Period 2, so they have the same number of occupied electron shells.
From N to F, the nuclear charge (number of protons) increases from 7 to 9.
Shielding by inner electrons changes very little within the same period.
Effective nuclear charge increases, so F is larger than N.

Short answer

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Q1. 1. Which atom is larger: Li or K? Explain using shell number and shielding.

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Q2. 2. Which atom is smaller: P or Cl? Explain using effective nuclear charge.

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Q3. 3. Which is larger: Na or Na⁺? Explain why.

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📖 Comprehensive answers (click to reveal)

1

K is larger. It is lower in Group 1, so it has more occupied shells. The outer electron is further from the nucleus and more shielded by inner electrons.

2

Cl is smaller. P and Cl are in the same period, so the outer electrons are in the same main shell. Cl has more protons, so the effective nuclear charge is greater and the outer electrons are pulled closer.

3

Na is larger than Na⁺. Forming Na⁺ removes the 3s valence electron, leaving one fewer occupied shell and reducing electron-electron repulsion.

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Previous: Electron Configuration: Subshell Notation ← Next: Periodic Trends: Electronegativity and Reactivity →

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