# Periodic trends

> IB Chemistry HL · S3: Classification of matter
> Source: https://www.owlsprep.com/study/ib-chemistry-hl-u3-periodic-trends/

This module explains key periodic trends across periods and down groups of the periodic table, links trends to underlying atomic structure, and covers common exam question patterns and exceptions to general rules.

**Prerequisites:** [Electronic configuration and periodic table arrangement](https://www.owlsprep.com/study/ib-chemistry-hl-u2-electronic-configuration/)

## Learning objectives

- Describe periodic trends in atomic radius, ionic radius, ionization energy, electron affinity and electronegativity
- Explain all trends using concepts of nuclear charge, shielding and effective nuclear charge
- Identify and explain exceptions to general first ionization energy trends
- Predict and compare properties of elements/ions from their periodic position

## Core Concepts: Shielding and Effective Nuclear Charge

All periodic trends can be explained by two core atomic concepts: shielding (or screening) and effective nuclear charge ($Z_{\text{eff}}$). Inner electron shells repel outer electrons, reducing the net pull of the nucleus on outer electrons. This reduction in attractive force is called shielding.

**Effective nuclear charge** — The net positive charge experienced by outer electrons, approximated by $Z_{\text{eff}} = Z - S$, where $Z$ is total nuclear charge (number of protons) and $S$ is the shielding constant from inner electrons.

*Notation:* $Z_{\text{eff}}$

*Example:* For sodium (Z=11), there are 10 inner shielding electrons, so $Z_{\text{eff}} \approx 1$ for the outer 3s electron.

Across a period (row), the number of protons increases by one per element, while the number of inner shielding electrons stays constant. This means $Z_{\text{eff}}$ increases steadily across a period. Down a group (column), nuclear charge increases, but the number of electron shells (and shielding) increases by one per period, so $Z_{\text{eff}}$ changes very little.

**Check your understanding**

Check your understanding of core concepts

1. Which correctly describes the trend of $Z_{\text{eff}}$ across period 3 from Na to Cl?

   - Zeff decreases
   - Zeff stays constant
   - Zeff increases
   - Zeff increases then decreases

   *Answer:* Zeff increases

   *Why:* All period 3 elements have 10 inner shielding electrons, so increasing proton number causes Zeff to increase across the period.

> **Exam tip:** IB exams award separate marks for mentioning both nuclear charge and shielding when explaining trends. Always reference both factors.

## Atomic and Ionic Radius Trends

The size of atoms and ions directly follows the trend in effective nuclear charge. Higher $Z_{\text{eff}}$ pulls electrons closer to the nucleus, reducing radius.

**Atomic radius** — Half the distance between the nuclei of two adjacent covalently bonded atoms of the same element.

- Across a period: Increasing $Z_{\text{eff}}$ pulls outer electrons closer, so atomic radius decreases left to right
- Down a group: Increasing number of electron shells and increased shielding outweigh higher nuclear charge, so atomic radius increases top to bottom

**Worked example:** Explain why the atomic radius of sodium (Z=11) is larger than chlorine (Z=17).

1. Both Na and Cl are in period 3, so both have 10 inner shielding electrons.
2. Sodium has 11 protons, chlorine has 17 protons, so nuclear charge is higher for Cl.
3. Shielding is approximately equal for both elements, so effective nuclear charge $Z_{\text{eff}}$ is higher for Cl.
4. Higher $Z_{\text{eff}}$ pulls outer electrons closer to the nucleus, so Cl has a smaller atomic radius.

Ionic radius follows similar general trends but has key differences for cations and anions: metal cations lose their outer valence shell, so they are much smaller than their parent atom. Non-metal anions gain electrons for the same nuclear charge, so they are larger than their parent atom. For isoelectronic ions (ions with the same number of electrons), radius decreases as nuclear charge increases.

**Worked example:** Explain why $Na^+$ has a smaller ionic radius than $F^-$, even though they are isoelectronic.

1. Both ions have 10 electrons with the same electron configuration, so shielding is identical.
2. $Na^+$ has 11 protons (Z=11), while $F^-$ has 9 protons (Z=9), so $Na^+$ has a higher nuclear charge.
3. Higher nuclear charge gives higher $Z_{\text{eff}}$, which pulls electrons closer to the nucleus, resulting in a smaller radius.

> **Exam tip:** When comparing isoelectronic ions, always state that electron count is equal, so the difference comes from differing proton number.

## Ionization Energy, Electronegativity and Electron Affinity

These properties all depend on the strength of nuclear attraction for electrons, so they follow the opposite trend to atomic radius.

**First ionization energy** — Minimum energy required to remove one mole of outer electrons from one mole of gaseous atoms.

The general trend for first ionization energy is: increases across a period (higher $Z_{\text{eff}}$ means electrons are harder to remove) and decreases down a group (larger atomic radius means electrons are easier to remove). There are two consistent exceptions across any period: between group 2 and 13, and between group 15 and 16. These exceptions come from differences in orbital energy and electron-electron repulsion.

**Worked example:** State and explain the trend in first ionization energy down group 1 (alkali metals).

1. First ionization energy decreases down group 1.
2. Down the group, the number of electron shells increases, so atomic radius increases and shielding from inner electrons increases.
3. The outer electron is further from the nucleus and experiences more shielding, so net attraction to the nucleus decreases.
4. Less energy is required to remove the outer electron, so first ionization energy decreases.

Electronegativity follows the same general trend as ionization energy: it increases across a period and decreases down a group. Fluorine is the most electronegative element, francium the least. Electron affinity (energy change for adding an electron) generally becomes more negative (more energy released) across a period, as higher $Z_{\text{eff}}$ attracts the added electron more strongly.

## Applying Periodic Trends to Exam Questions

Most exam questions on this topic ask you to compare a property of two elements/ions and explain the difference. The mark scheme always rewards linking your answer to atomic structure concepts, not just stating the trend.

**Worked example:** Explain why chlorine has higher electronegativity than bromine.

1. Chlorine and bromine are in the same group (group 17), with chlorine higher up the group.
2. Bromine has one more electron shell than chlorine, so atomic radius is larger and shielding from inner electrons is greater.
3. Even though bromine has a higher nuclear charge, the increased shielding and larger radius reduce net attraction for bonding electrons.
4. Weaker attraction for shared bonding electrons means bromine has lower electronegativity than chlorine.

**Check your understanding**

Test your understanding of exceptions

1. What is the correct order of first ionization energy (highest to lowest) for N, O, F?

   - O > N > F
   - F > N > O
   - F > O > N
   - N > O > F

   *Answer:* F > N > O

   *Why:* F has the highest Zeff so highest IE. N has a stable half-filled p subshell, so higher IE than O, giving order F > N > O.

> **Exam tip:** Never just state 'it follows the trend' as your explanation. Always link back to Zeff, shielding and radius to get full marks.

## Common pitfalls

- **Wrong:** Stating atomic radius increases across a period because nuclear charge increases
  - Why it fails: Forgets that shielding stays constant across a period, so the effect of increasing nuclear charge dominates, pulling electrons closer
  - Correct: State that across a period, nuclear charge increases, shielding is constant, so Zeff increases, pulling electrons closer, decreasing atomic radius
- **Wrong:** Claiming cations are larger than their parent atom
  - Why it fails: Cations lose their outer valence electron shell, so the remaining electrons are pulled closer by the nucleus
  - Correct: State that cations lose their outer valence shell, so they have fewer electron shells, leading to a smaller radius than the parent atom
- **Wrong:** Explaining decreasing ionization energy down a group only by increasing nuclear charge
  - Why it fails: Nuclear charge does increase down a group, but increased shielding and larger atomic radius outweigh this effect
  - Correct: Explain decreasing ionization energy down a group by increasing atomic radius and increased shielding, which reduce net attraction for the outer electron
- **Wrong:** Confusing electronegativity and electron affinity
  - Why it fails: They describe different properties: electronegativity refers to bonded atoms, while electron affinity refers to free gaseous atoms
  - Correct: Distinguish the two: electronegativity is attraction for shared bonding electrons, electron affinity is energy change for adding an electron to a free gaseous atom

## Cheatsheet

| Property | Across period (left → right) | Down group (top → bottom) |
| --- | --- | --- |
| Atomic radius | Decreases | Increases |
| Ionic radius (same ion type) | Decreases | Increases |
| First ionization energy | Generally increases | Decreases |
| Electronegativity | Increases | Decreases |
| Electron affinity (magnitude) | Generally increases | Generally decreases |
| Effective nuclear charge | Increases | Approximately constant |

## What's next

Periodic trends are foundational to all further topics in IB Chemistry. This knowledge is used to explain trends in reactivity of group 1 and group 7 elements, the changing acid-base properties of period 3 oxides, and to predict the type of bonding between two elements based on their electronegativity difference. Mastery of this topic is required for units on chemical bonding, thermodynamics and reactivity, so it is important to be able to both state trends and explain their origins clearly for exam questions.

- [d-block elements and transition metals](https://www.owlsprep.com/study/ib-chemistry-hl-u3-d-block-elements-and-transition/)
- [Acid-base classification: Bronsted-Lowry and Lewis](https://www.owlsprep.com/study/ib-chemistry-hl-u3-acid-base-classification-bronsted-lowry/)
- [Salts and buffer solutions](https://www.owlsprep.com/study/ib-chemistry-hl-u3-salts-and-buffer-solutions/)

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