# AHL: Extended periodic trends

> IB Chemistry HL · IB Chemistry HL (2025 syllabus)
> Source: https://www.owlsprep.com/study/ib-chemistry-hl-u3-ahl-extended-periodic-trends/

This subtopic extends core periodic trends to cover electron affinity and melting/boiling point trends across periods and down groups. You will connect these trends to atomic structure, bonding, and intermolecular forces, a common exam topic for Paper 1 and 2.

**Prerequisites:** [Core periodic trends (atomic radius, ionization energy)](https://www.owlsprep.com/study/ib-chemistry-hl-u3-core-periodic-trends/); [Bonding types and intermolecular forces](https://www.owlsprep.com/study/ib-chemistry-hl-u4-bonding-and-structure/)

## Learning objectives

- Explain periodic trends in first electron affinity across periods and down groups
- Compare and explain trends in melting/boiling points across period 3 and down common groups
- Relate property trends to atomic structure, bonding and intermolecular forces
- Predict properties of unknown elements based on their position in the periodic table

## Electron Affinity: Definition and Periodic Trends

**First Electron Affinity** — The energy released (or absorbed) when one mole of gaseous neutral atoms gains one mole of electrons to form one mole of gaseous 1- anions. Most first electron affinities are negative (exothermic), while all second electron affinities are positive (endothermic) due to electrostatic repulsion between the negative ion and incoming electron.

*Notation:* E_{ea}

*Example:* First $E_{ea}$ of chlorine is -349 kJ mol⁻¹, meaning 349 kJ of energy is released per mole of Cl atoms gaining electrons.

The trend in first electron affinity across a period generally increases in magnitude (becomes more negative) from left to right. This is because atomic radius decreases across a period, so the incoming electron is attracted more strongly to the higher effective nuclear charge, resulting in a more exothermic process. Down a group, first electron affinity generally decreases in magnitude (becomes less negative) as atomic radius increases, so the incoming electron is further from the nucleus and experiences less attraction.

> **info**
>
> Fluorine is an exception to the group 17 trend. Its small atomic radius causes extra electron-electron repulsion, so its first $E_{ea}$ is less negative than chlorine's.

**Worked example:** Explain why the first electron affinity of sulfur is less negative than that of chlorine.

1. Step 1: Identify the position of both elements: both are in period 3, sulfur in group 16, chlorine in group 17.
2. Step 2: Compare effective nuclear charge and atomic radius: higher atomic number for Cl means higher effective nuclear charge, smaller atomic radius.
3. $$Z_{eff}(Cl) > Z_{eff}(S), \quad r(Cl) < r(S)$$
4. Step 3: The nucleus of Cl attracts the incoming electron more strongly than sulfur, so more energy is released when Cl gains an electron.
5. Step 4: Conclusion: First $E_{ea}$ of Cl is more negative than sulfur, matching the observation.

## Melting Point Trends Across Period 3

Melting point depends on the strength of the interactions holding the element's particles together. Across period 3, the bonding type changes from metallic to giant covalent to simple molecular, which drives a characteristic trend in melting point.

**Worked example:** Arrange Na, Mg, Si, Cl₂ and S₈ in order of increasing melting point, and explain your answer.

1. Step 1: Identify the bonding and structure for each element: Na and Mg are giant metallic, Si is giant covalent, S₈ and Cl₂ are simple molecular.
2. Step 2: Compare overall interaction strength: Giant covalent bonds > metallic bonds > intermolecular forces in simple molecular structures.
3. Step 3: Compare strengths within each category: Mg has higher charge density than Na, so metallic bonding is stronger in Mg. S₈ has larger molecular mass than Cl₂, so London dispersion forces are stronger in S₈.
4. Step 4: Final order from lowest to highest melting point:
5. $$Cl_2 < S_8 < Na < Mg < Si$$

| Element | Structure/Bonding | Melting Point (°C) |
| --- | --- | --- |
| Na | Giant metallic | 98 |
| Mg | Giant metallic | 650 |
| Al | Giant metallic | 660 |
| Si | Giant covalent | 1410 |
| P₄ | Simple molecular | 44 |
| S₈ | Simple molecular | 115 |
| Cl₂ | Simple molecular | -101 |
| Ar | Simple molecular | -189 |

## Melting Point Trends Down Groups

Trends in melting point down a group depend entirely on the type of bonding and structure of the group's elements, so different groups have opposite trends.

**Worked example:** Explain why melting point increases down group 17 but decreases down group 1.

1. Step 1: Analyze group 1: All elements are giant metallic structures with metallic bonding between cations and delocalized electrons.
2. Step 2: Down group 1, cation radius increases, charge density of the cation decreases, so the strength of metallic bonding decreases.
3. Less energy is needed to overcome weaker metallic bonding, so melting point decreases down group 1.
4. Step 3: Analyze group 17: All elements are diatomic simple molecular structures, with London dispersion forces between molecules.
5. Step 4: Down group 17, molecular mass and number of electrons increase, so the strength of London dispersion forces increases.
6. More energy is needed to overcome stronger intermolecular forces, so melting point increases down group 17.

> **tip**
>
> Always check the structure of the element before explaining a melting point trend: bonding type is always the root cause.

## Common pitfalls

- **Wrong:** Confusing electron affinity with electronegativity
  - Why it fails: Both describe attraction for electrons, but they describe different processes: electron affinity is energy change for a gaseous atom gaining a free electron, while electronegativity describes attraction for bonding electrons in a covalent bond
  - Correct: Remember: electron affinity = free electron gain (energy change), electronegativity = bonding electron attraction
- **Wrong:** Stating fluorine has the most negative electron affinity in group 17
  - Why it fails: Fluorine's small atomic size causes extra electron-electron repulsion that offsets higher effective nuclear charge, so chlorine is more negative
  - Correct: Remember chlorine has the most negative first electron affinity in the periodic table
- **Wrong:** Claiming melting point increases across all of period 3
  - Why it fails: Melting point increases from Na to Si, then drops sharply for non-metals with simple molecular structures
  - Correct: The trend across period 3 is increasing melting point up to group 14, then decreasing for non-metals
- **Wrong:** Saying all electron affinities are negative (exothermic)
  - Why it fails: Second electron affinities (adding an electron to a 1- ion) are always positive (endothermic) due to electrostatic repulsion
  - Correct: Only most first electron affinities are negative; all second and higher electron affinities are positive
- **Wrong:** Explaining melting point trend down group 17 in terms of atomic radius
  - Why it fails: Melting point depends on intermolecular force strength between molecules, not attraction within the molecule
  - Correct: Down group 17, increasing molecular mass leads to stronger London dispersion forces, increasing melting point

## Cheatsheet

| Property | Trend Across Period (left → right) | Trend Down Group (top → bottom) | Key Exception |
| --- | --- | --- | --- |
| First $E_{ea}$ | More negative (more exothermic) | Less negative | F: less negative than Cl |
| Period 3 Melting Point | Increases to Si, then decreases | - | - |
| Group 1 Melting Point | - | Decreases | None |
| Group 17 Melting Point | - | Increases | None |

## What's next

Extended periodic trends build on core periodicity concepts and connect directly to bonding and structure topics, which make up ~20% of the IB Chemistry HL exam. Understanding how structure and bonding drive property trends is critical for answering data-based questions in Paper 2 and multiple-choice questions in Paper 1, where you will often be asked to predict properties of unfamiliar elements or explain deviations from general trends. This subtopic also provides a foundation for understanding element reactivity, which is explored in redox and organic chemistry topics. Mastering these trends will help you connect atomic structure to bulk material properties across the entire syllabus.

- [AHL: Advanced Lewis acid-base concepts](https://www.owlsprep.com/study/ib-chemistry-hl-u3-ahl-advanced-lewis-acid-base/)
- [AHL: Acid-base titration curves](https://www.owlsprep.com/study/ib-chemistry-hl-u3-ahl-acid-base-titration-curves/)

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