Electron Affinity Explained: H2 Chemistry Guide

Quick answer: First electron affinity is the enthalpy change when one mole of gaseous atoms each gains one electron to form one mole of gaseous 1− ions. It is usually negative (energy released). The second electron affinity is always positive, because the incoming electron is repelled by the ion’s existing negative charge.

Electron affinity explains why some atoms form negative ions readily, and others do not. It also supplies two of the energy terms in a Born-Haber cycle. This guide covers the definition and sign convention, the periodic trends and their exceptions, the second electron affinity, how electron affinity differs from electronegativity and ionisation energy, and the questions that come up in exams. It sits inside the Chemical Energetics and Periodic Table topics of the H2 Chemistry Syllabus Explained.

Table of Contents

What Is Electron Affinity?

First electron affinity is the enthalpy change when one mole of gaseous atoms each gains one electron to form one mole of gaseous 1− ions.

Cl(g) + e⁻ → Cl⁻(g) ΔH = −349 kJ mol⁻¹

If the nucleus attracts the incoming electron strongly enough, the new ion is more stable than the atom and energy is released, so the value is negative. If repulsion is significant, little energy is released, or energy has to be supplied.

Check the sign convention. H2 Chemistry uses the enthalpy-change convention, so a favourable electron gain has a negative value. Many UK and US textbooks instead quote electron affinity as the energy released, which makes the same process a positive number. Always check which convention a question or data source is using before you compare values.

Factors Affecting Electron Affinity

• Nuclear charge: a larger positive charge increases attraction for the incoming electron.
• Atomic radius: a smaller atom puts the incoming electron closer to the nucleus, increasing attraction.
• Shielding: inner-shell electrons reduce the attraction the incoming electron feels. See Shielding Effect Explained and Effective Nuclear Charge.
• Electron–electron repulsion: crowding in a small subshell makes adding another electron less favourable.

These factors compete, which is why electron affinity trends are useful but not perfectly smooth. For the size effects behind them, see Atomic Radius Explained.

Electron Affinity Trend Across a Period

Across a period, nuclear charge increases while the added electron goes into the same shell. Attraction for the incoming electron grows, so first electron affinity generally becomes more negative from left to right.

The trend has clear exceptions, and each can be explained with electron arrangements: – Group 2 (for example Be, Mg): the outer s subshell is full, so the incoming electron must go into a higher-energy p subshell, and little or no energy is released. – Group 15 (for example N): the p subshell is half-filled, so the incoming electron must pair up with an existing electron and is repelled, making electron affinity much less negative than the neighbours on either side. – Group 18: the shell is full, so the incoming electron must start a new, higher-energy shell. Electron gain is not favourable.

Direction	Main reason	Typical effect
Across a period	Nuclear charge increases	Generally more negative
Down a group	Radius and shielding increase	Generally less negative

For a quick summary of all the periodic trends side by side, see the A Level Periodic Table guide.

Electron Affinity Trend Down a Group

Down a group, atoms have more occupied shells. The incoming electron is added further from the nucleus and feels more shielding, so attraction weakens and first electron affinity generally becomes less negative.

This is easiest to see in Group 17. Halogens have seven outer electrons, so gaining one gives a stable noble-gas arrangement and all the first electron affinities are strongly negative. From Cl to ,I, the values become less negative as expected: −349, −325, −295 kJ mol⁻¹. Fluorine is the exception that students are asked about.

Why Is Fluorine’s Electron Affinity Lower Than Chlorine’s?

Fluorine’s first electron affinity (−328 kJ mol⁻¹) is less negative than chlorine’s (−349 kJ mol⁻¹), even though fluorine is higher up the group.

Bar chart of first electron affinity for F, Cl, Br and I: chlorine is the most negative at −349, fluorine is −328, bromine −325 and iodine −295 kJ per mole.

Fluorine’s 2p subshell is very small. The incoming electron enters a compact region already holding seven electrons, so electron–electron repulsion is strong and cuts the energy released. Chlorine adds the electron to the larger 3p subshell, where there is more space and less repulsion, so more energy is released.

Exam tip: do not explain this exception only by saying “fluorine is smaller”. A complete answer states that the small size gives strong attraction, but it also gives greater electron–electron repulsion in the compact 2p subshell, and the repulsion wins. The same idea explains why oxygen (−141) is less negative than sulfur (−200).

Second Electron Affinity

Second electron affinity is the enthalpy change when one mole of gaseous 1− ions each gains one electron to form one mole of gaseous 2− ions.

X⁻(g) + e⁻ → X²⁻(g)

It is endothermic (positive) for any 1− ion, because the incoming electron is being added to a species that is already negatively charged. The strong electron–electron repulsion means energy must be supplied to force the electron on.

Element	1st EA / kJ mol⁻¹	2nd EA / kJ mol⁻¹
O	−141	about +798
S	−200	about +640

Energy-level diagram for oxygen: O to O⁻ releases 141 kJ per mole, then O⁻ to O²⁻ absorbs 798 kJ per mole, so the net change to form O²⁻ is +657 kJ per mole.

Notice that the total for forming O²⁻ from O is positive (−141 + 798 = +657 kJ mol⁻¹). Forming the oxide ion from a gaseous oxygen atom costs energy overall. Oxides such as MgO are still stable because the lattice energy released when the ions come together is so large.

Electron Affinity vs Electronegativity vs Ionisation Energy

These three terms are related, but each describes something different. Using them interchangeably is a common exam error.

Property	What it describes	What is involved	Typical trend
Electron affinity	Enthalpy change when a gaseous atom gains an electron	An isolated gaseous atom	Generally more negative across a period; less negative down a group
Electronegativity	Relative ability of an atom to attract the bonding pair of electrons	An atom within a covalent bond	Increases across a period; decreases down a group
Ionisation energy	Enthalpy change when a gaseous atom loses an electron	An isolated gaseous atom	Generally increases across a period; decreases down a group

The key differences: – Electron affinity and ionisation energy are measurable energy changes with units of kJ mol⁻¹ and a sign. Electronegativity is a relative scale with no units and no sign. – Electron affinity and ionisation energy apply to isolated gaseous atoms. Electronegativity applies to atoms in a bond. – Ionisation energy removes an electron; electron affinity adds one.

For the full treatment, see the Electronegativity guide linked above and Ionisation Energy Explained.

Useful Electron Affinity Values

Exact values depend on the data source. Use this table to interpret trends, not to memorise numbers. In an exam, use the values given in the question or in the H2 Chemistry Data Booklet.

Element	Approx. first EA / kJ mol⁻¹	What to notice
F	−328	Less negative than Cl despite being above it
Cl	−349	The most negative: the electron enters the larger 3p subshell
Br	−325	Less negative than Cl as radius and shielding increase
I	−295	Less negative than Br; the electron is further from the nucleus
O	−141	Less negative than S because of repulsion in the compact 2p subshell
S	−200	More negative than O despite being below it

Electron Affinity in Born-Haber Cycles

Electron affinity is one of the energy terms in a Born-Haber cycle. When a non-metal forms an anion, the first electron affinity is an energy-releasing step. If the anion has a 2− charge, the second electron affinity is an energy-absorbing step that must be included with the correct (positive) sign.

The step-by-step calculation, with worked examples for NaCl and MgCl₂, is in Lattice Energy Explained, which is the page to use for Born-Haber calculations.

Exam-Style Questions With Model Answers

Question 1 — Group 17 trend (2 marks). Explain why the first electron affinity becomes less negative from chlorine to iodine. Model answer: Down the group, atomic radius and shielding increase, so the incoming electron is further from the nucleus and experiences weaker attraction. Less energy is released when the electron is added.

Question 2 — The F and Cl exception (3 marks). The first electron affinity of chlorine is more negative than that of fluorine. Explain why. Model answer: In fluorine, the incoming electron enters the small 2p subshell, where electron–electron repulsion is strong. In chlorine, it enters the larger 3p subshell, where there is more space and less repulsion. Less energy is released for fluorine, so its first electron affinity is less negative.

Question 3 — The second electron affinity (2 marks). Explain why the second electron affinity of oxygen is positive. Model answer: The second electron is added to the O⁻ ion, which is already negatively charged. The incoming electron is repelled by the existing negative charge, so energy must be supplied.

Question 4 — Period trend exception (3 marks). The first electron affinity of nitrogen is much less negative than that of carbon or oxygen. Explain why. Model answer: In nitrogen, the 2p subshell is half-filled, so the incoming electron must pair up with an existing 2p electron. The electron–electron repulsion from pairing makes the process much less favourable than in carbon, where the electron enters an empty p orbital, or in oxygen, where the nuclear charge is higher.

Common Electron Affinity Mistakes

• Confusing electron affinity with electronegativity (an energy change versus a relative scale).
• Forgetting that the first electron affinity is usually negative while the second is always positive.
• Explaining the F versus Cl exception only by “fluorine is smaller”, without mentioning electron–electron repulsion.
• Saying the second electron is “attracted to” a negative ion, instead of repelled by it.
• Explaining trends using atomic radius alone, without nuclear charge and shielding.
• Mixing up sign conventions between textbooks and the data given in the question.
• Using the first electron affinity only for a 2− anion in a Born-Haber cycle and leaving out the second.

How to Revise Electron Affinity

Revise the reason for each energy change, not the individual numbers. 1. Learn both definitions, including state symbols and “gaseous 1− ions”. 2. Practise explaining trends using nuclear charge, radius, shielding and repulsion, in that order. 3. Place first and second electron affinities side by side and rehearse why the second is positive. 4. Use the F versus Cl and N versus C/O cases as models for explaining exceptions. 5. Practise the comparison between electron affinity, electronegativity and ionisation energy until you can do it from memory.

For wider H2 Chemistry revision structure, see the H2 Chemistry revision guide.

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Electron affinity connects to every other periodic trend. Continue with Electronegativity Explained, or explore our A-Level Chemistry tuition and wider Chemistry programmes, taught in small groups capped at 7 students near Beauty World MRT.