What Is Electron Affinity?

Electron affinity is defined as the energy change when a gaseous atom gains an electron to form a gaseous anion. It is the exact opposite of ionization energy: while ionization energy measures the energy required to remove an electron from an atom, electron affinity describes the energy released or absorbed when an atom gains an electron. Think of it this way — electrons are attracted to the nucleus, so removing them costs energy. But adding an electron? That is where electrons want to go — at least for the first one.

The process occurs entirely in the gaseous phase, just like ionization energy. Instead of producing a cation, the atom gains an electron to form an anion:

X(g) + e− → X−(g)

Most first electron affinities are exothermic (negative values), meaning the atom releases energy upon gaining an electron. There are notable exceptions, and understanding where those exceptions occur — and why — is a major focus of this lesson.

Sign Convention: Chlorine vs. Argon

Chlorine holds the title for the most negative electron affinity of all the elements at −349 kJ/mol. Hand chlorine an electron and it showers you with energy — it is absolutely delighted to receive that extra electron and complete its octet.

Feeling generous, you might try offering an electron to argon. Bad idea. Argon does not want your electron. With a completely filled octet, argon is perfectly content as it is. It did not want to lose an electron (hence its sky-high ionization energy), and it certainly does not want to gain one either. Argon has a positive electron affinity, meaning the process is endothermic — you would have to force that electron onto argon.

This pattern is typical for the noble gases. Their electron configurations are so stable that forcing an additional electron into the next energy level is energetically unfavorable.

3 Exceptions to the Electron Affinity Periodic Trend

The noble gases are not alone in resisting extra electrons. The same electron configurations that created exceptions in ionization energy trends also produce exceptions in electron affinity. There are three categories of elements that consistently break the general electron affinity trend.

Filled s Subshells (Group 2 Elements)

Elements with a filled s subshell — beryllium, magnesium, calcium, and the rest of Group 2 — tend to have positive or only very slightly negative electron affinities. Magnesium, for instance, does not want your electron. It already has a filled 3s subshell and would prefer to stay that way rather than begin populating the higher-energy 3p orbital. Screening and shielding effects reinforce this resistance.

Half-Filled p Subshells (Group 15 Elements)

The half-filled p subshells of Group 15 also violate the general electron affinity trend. Consider the progression across Period 3: electron affinity values become progressively more negative as you move to the right — from −43 kJ/mol to −134 kJ/mol — but then phosphorus breaks the pattern at only −72 kJ/mol. While still technically negative, it falls well short of the expected value around −200 kJ/mol. For nitrogen, the value is even slightly positive.

The half-filled p subshell enjoys a special stability due to exchange energy. Adding one more electron disrupts this favorable arrangement, making the process less exothermic than expected.

What Is the Periodic Trend for Electron Affinity?

There is no well-defined vertical trend for electron affinity — and that makes it unique among the periodic trends you have studied so far. The only reliable general rule is that electron affinity increases (becomes more negative) moving from left to right across a period, excluding the three exception groups.

Why Is There No Vertical Trend for Electron Affinity?

The Period 3 elements generally possess the most negative electron affinities on the periodic table. Chlorine in Period 3 holds the overall record at −349 kJ/mol. Values tend to become more negative as you approach Period 3 from either direction — whether coming down from Period 2 or up from lower periods. This non-linear behavior is why no simple "up = more negative" vertical trend exists for electron affinity.

Because the vertical trend is so inconsistent, most general chemistry courses teach one rule: electron affinity becomes more negative moving to the right across a period, with the usual exceptions at noble gases, filled s subshells, and half-filled p subshells.

Successive Electron Affinities: Why the 2nd Is Always Endothermic

Just as we discussed successive ionization energies, atoms can also have successive electron affinities. The general trend we just covered applies only to the first electron affinity. What happens when an atom needs to gain a second electron? That is where things get interesting — and where a common textbook simplification gets exposed.

The Sodium Oxide Example

Consider the formation of sodium oxide (Na₂O). Sodium loses an electron to become Na⁺, and oxygen gains two electrons to become O²−. The story students are usually told: sodium wants to lose an electron to achieve a noble gas configuration, and oxygen wants to gain two electrons to complete its octet.

There are two problems with that story.

Problem 1: Sodium does not actually want to lose an electron. Ionization always costs energy. The alkali metals have the lowest ionization energies, making them the best candidates to lose electrons — but it still requires an energy input.

Problem 2: Oxygen does not want to gain two electrons. The first electron affinity of oxygen is exothermic at −141 kJ/mol — no issues there. But once oxygen gains that first electron and becomes O−, it is now a negatively charged ion trying to attract another negatively charged electron. Like charges repel.

What Really Drives Ionic Compound Formation?

If sodium does not want to lose electrons and oxygen does not want to gain a second electron, what actually drives the formation of sodium oxide and other ionic compounds?

The answer lies in lattice energy — the powerful electrostatic attraction between the resulting cations and anions in the crystal lattice. Once you have Na⁺ and O²− ions, their mutual attraction releases an enormous amount of energy, more than enough to compensate for the unfavorable steps along the way. This cation-anion attraction is the true driving force behind ionic compound formation, a topic explored in depth in the next chapter on chemical bonding.

What Is Electronegativity?

Electronegativity is defined as an atom's ability to attract shared electrons within a chemical bond. This is a crucial distinction from electron affinity: electron affinity measures the energy change when an isolated gaseous atom gains an electron, while electronegativity describes how strongly an atom pulls shared electrons toward itself when it is bonded to another atom.

The most widely used measure is the Pauling electronegativity scale, which assigns numerical values to elements based on bond energies. On this scale, the undisputed champion is fluorine — the most electronegative element on the entire periodic table, with a value of 4.0.

In covalent bonds, two atoms (typically non-metals) share electrons. Electronegativity determines how evenly those electrons are shared. The more electronegative atom pulls the shared electrons closer to itself, creating an uneven distribution of electron density that we call a polar covalent bond.

Polar Bonds and Unequal Electron Sharing

Consider a carbon-fluorine bond. If the shared electrons were distributed equally, they would spend the same amount of time near each atom. In reality, fluorine's superior electronegativity pulls the electron density closer to itself, and the electrons spend more time near fluorine than near carbon.

If the electrons were completely transferred to fluorine, we would have full ionic charges: F− and C⁺. But in a polar covalent bond, the transfer is incomplete — the atoms share the electrons, just not equally. The result is partial charges:

• Fluorine acquires a partial negative charge (δ−)
• Carbon acquires a partial positive charge (δ⁺)

The Greek letter delta (δ) denotes "partial" — the same symbol used for partial derivatives in calculus. These partial charges indicate an incomplete separation of charge, not a full +1 or −1. The greater the electronegativity difference between two bonded atoms, the more polar the bond.

What Is the Periodic Trend for Electronegativity?

Unlike electron affinity, electronegativity has a clear two-dimensional periodic trend: electronegativity increases toward fluorine from both directions. It increases as you move up a group, and it increases as you move from left to right across a period. This makes electronegativity one of the most intuitive periodic trends to memorize.

Notice that the trend points toward fluorine, not helium. The noble gases are excluded from electronegativity discussions because they generally do not form bonds, making the concept inapplicable to them.

Applying the Electronegativity Trend

Within any period, the closer an element sits to fluorine, the more electronegative it is. For example, across Period 2: boron < carbon < nitrogen < oxygen < fluorine. Within any group, electronegativity increases going up: iodine < bromine < chlorine < fluorine.

The underlying logic is intuitive: smaller atoms with more protons hold onto shared electrons more tightly. Fluorine is both small and highly charged, making it the ultimate electron hog on the periodic table.