IMAT Chemistry · Periodic Table

Ionisation Energy, Electron Affinity and Electronegativity

How strongly an atom holds its own electrons (ionisation energy), welcomes an extra one (electron affinity) and pulls shared ones (electronegativity) all rise across a period and fall down a group.

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Why this matters

First ionisation energy is the trend asked most often: in 2018, 2022 and the 2023 ministry paper, each time as which element has the highest value or which order is right. 2021 asked which group 17 properties fall down the group and which equation is the first electron affinity of chlorine, and 2012 asked the signs of oxygen's first and second electron affinities.

Concept 1 of 4: First and successive ionisation energies

Ionisation energy is the price of pulling an electron off an atom. It is high when the electron is close to a strongly charged nucleus and low when it is far away and shielded by inner shells. That makes it the mirror image of atomic radius: small atoms hold their electrons tightly.

Definition

The first ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms. It is always positive (endothermic).

Across a period it increases: nuclear charge rises and the atom shrinks. Each noble gas has the highest value in its period; helium has the highest of all.
Down a group it decreases: the outer electron is further out and more shielded. Group 1 metals have the lowest values in each period.
Two small dips: group 13 is lower than group 2 (Be > B, Mg > Al), and group 16 lower than group 15 (N > O).
Successive ionisation energies always rise. A very large jump shows that the next electron must come from a full inner shell, so the number of electrons removed before the jump is the group number.

First ionisation energy

\mathrm{X(g)} \rightarrow \mathrm{X^{+}(g)} + e^- \qquad \Delta H > 0

$\mathrm{X(g)}$ a gaseous atom of the element
$\Delta H$ energy change, in kJ/mol; positive because energy is taken in

Worked example

An element has successive ionisation energies of 578, 1817, 2745 and 11 577 kJ/mol. In which group is it?

Practice this conceptself-check · 3 quick reps

The noble gas, not the halogen, has the highest ionisation energy in a period

Ionisation energy keeps rising across a period all the way to group 18. Halogens are high but each noble gas is higher. Down a group it falls, so an element lower in the same group always has the smaller value.

Concept 2 of 4: First and second electron affinity

Electron affinity is the energy change when a neutral gaseous atom accepts an electron. A non-metal atom attracts the incoming electron towards its nucleus, so energy is released. But a negative ion already repels electrons, so forcing a second electron onto it costs energy.

Definition

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

For most non-metals it is negative (exothermic). Halogens have the most exothermic values in each period.
Chlorine has the most exothermic first electron affinity of all, slightly more than fluorine: fluorine is so small that its own electrons repel the newcomer.
The second electron affinity (adding an electron to a 1− ion) is always positive (endothermic), because the ion repels the electron.
Oxygen: the first electron affinity is about −141 kJ/mol (mildly exothermic); the second is large and positive. Forming $\mathrm{O^{2-}}$ from O is endothermic overall.
Do not confuse it with ionisation energy, which removes an electron.

First and second electron affinity

\mathrm{X(g)} + e^- \rightarrow \mathrm{X^{-}(g)} \qquad \mathrm{X^{-}(g)} + e^- \rightarrow \mathrm{X^{2-}(g)}

first equationfirst electron affinity: one gaseous atom gains one electron
second equationsecond electron affinity: the gaseous 1− ion gains another

Worked example

Write the equations for the first and second electron affinities of sulfur, and give the sign of each.

Practice this conceptself-check · 3 quick reps

Second electron affinities are endothermic

The first electron affinity of oxygen or sulfur is exothermic, but the second is strongly endothermic. Options giving both as exothermic, or the first as very exothermic and the second as exothermic too, are wrong.

Concept 3 of 4: Electronegativity and its values

In a covalent bond two atoms share a pair of electrons, but not always equally. The atom whose nucleus pulls harder on that shared pair is the more electronegative one. Like ionisation energy, the pull is strongest for small atoms with a high nuclear charge, so fluorine wins.

Definition

Electronegativity is the ability of an atom in a covalent bond to attract the shared pair of electrons. It is measured on the Pauling scale, which has no units.

It increases across a period and decreases down a group.
Fluorine is the most electronegative element (about 4.0), then oxygen, then chlorine and nitrogen.
The least electronegative elements are the group 1 metals at the bottom of the table.
A large difference in electronegativity gives a polar or ionic bond; equal values give a non-polar bond.

Element	Pauling value	Where it sits
Fluorine	3.98	Top of group 17: the highest of all
Oxygen	3.44	Group 16, period 2: second highest
Chlorine	3.16	Group 17, period 3
Nitrogen	3.04	Group 15, period 2
Carbon	2.55	Group 14, period 2
Hydrogen	2.20	Slightly below carbon
Sodium	0.93	Group 1, period 3
Caesium	0.79	Group 1, period 6: among the lowest

Practice this conceptself-check · 3 quick reps

Oxygen is more electronegative than chlorine

Fluorine is first and oxygen second. Chlorine and nitrogen come after oxygen. Picking chlorine because it is a halogen is the usual slip.

Concept 4 of 4: Summary of the periodic trends across a period and down a group

Every trend in this chapter comes from two causes. Across a period, the nuclear charge rises while the shell stays the same, so electrons are held more tightly. Down a group, a new shell is added each time, so the outer electrons are further away and more shielded. Work out which cause applies, and the direction of any trend follows.

Definition

Metallic character (the ease of losing electrons) runs opposite to ionisation energy and electronegativity.
Metals get more reactive down groups 1 and 2; halogens get less reactive down group 17.
The most vigorous reactions are between a metal at the bottom left and a non-metal at the top right, such as caesium with fluorine. Noble gases react with almost nothing.
Melting and boiling points of the halogens and the noble gases rise down the group (bigger molecules or atoms, stronger London forces); those of the group 1 metals fall.

Property	Across a period (left to right)	Down a group	Main reason
Atomic radius	Decreases	Increases	Nuclear charge across; extra shells down
First ionisation energy	Increases (small dips at groups 13 and 16)	Decreases	How tightly the outer electron is held
Electronegativity	Increases	Decreases	Pull of the nucleus on shared electrons
First electron affinity	Generally more exothermic	Less exothermic (from chlorine down)	Pull of the nucleus on an added electron
Metallic character	Decreases	Increases	Ease of losing electrons
Reactivity of metals (groups 1, 2)	Group 1 more reactive than group 2	Increases	Ease of losing electrons
Reactivity of halogens (group 17)	Halogens most reactive in their period	Decreases	Ease of gaining an electron

Practice this conceptself-check · 3 quick reps

Metal reactivity rises down a group; halogen reactivity falls

Metals react by losing electrons, which gets easier further down, so caesium is more reactive than lithium. Halogens react by gaining electrons, which gets harder further down, so fluorine is more reactive than iodine. Melting points of the halogens rise down the group while their electronegativity and ionisation energy fall.

Summary — formulas & gotchas at a glance

A revision cheat-sheet for the formulas and gotchas above. Click any concept name to jump back to its full explanation.

Formulas (2)

First and successive ionisation energies
First ionisation energy
$\mathrm{X(g)} \rightarrow \mathrm{X^{+}(g)} + e^- \qquad \Delta H > 0$
First and second electron affinity
First and second electron affinity
$\mathrm{X(g)} + e^- \rightarrow \mathrm{X^{-}(g)} \qquad \mathrm{X^{-}(g)} + e^- \rightarrow \mathrm{X^{2-}(g)}$

Reference tables (2)

Electronegativity and its values8 rows

Element	Pauling value	Where it sits
Fluorine	3.98	Top of group 17: the highest of all
Oxygen	3.44	Group 16, period 2: second highest
Chlorine	3.16	Group 17, period 3
Nitrogen	3.04	Group 15, period 2
Carbon	2.55	Group 14, period 2
Hydrogen	2.20	Slightly below carbon
Sodium	0.93	Group 1, period 3
Caesium	0.79	Group 1, period 6: among the lowest

Summary of the periodic trends across a period and down a group7 rows

Property	Across a period (left to right)	Down a group	Main reason
Atomic radius	Decreases	Increases	Nuclear charge across; extra shells down
First ionisation energy	Increases (small dips at groups 13 and 16)	Decreases	How tightly the outer electron is held
Electronegativity	Increases	Decreases	Pull of the nucleus on shared electrons
First electron affinity	Generally more exothermic	Less exothermic (from chlorine down)	Pull of the nucleus on an added electron
Metallic character	Decreases	Increases	Ease of losing electrons
Reactivity of metals (groups 1, 2)	Group 1 more reactive than group 2	Increases	Ease of losing electrons
Reactivity of halogens (group 17)	Halogens most reactive in their period	Decreases	Ease of gaining an electron

Watch out for (4)

The noble gas, not the halogen, has the highest ionisation energy in a period
→ First and successive ionisation energies
Second electron affinities are endothermic
→ First and second electron affinity
Oxygen is more electronegative than chlorine
→ Electronegativity and its values
Metal reactivity rises down a group; halogen reactivity falls
→ Summary of the periodic trends across a period and down a group

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