Periodic Trends: Electron Affinity and Electronegativity

Electron affinity is the energy change when a gaseous atom gains an electron, and electronegativity is a measure of how strongly an atom in a bond attracts the shared electrons; both generally increase from left to right across a period and electronegativity decreases down a group. In brief: fluorine is the most electronegative element (3.98 on the Pauling scale), chlorine has the most negative electron affinity (it releases 349 kJ/mol), and the electronegativity difference between two bonded atoms predicts whether the bond is nonpolar, polar or ionic.

These two properties describe how atoms deal with electrons from the nonmetal side of the periodic table. Electron affinity is a property of an isolated atom, and electronegativity is a property of an atom in a molecule. This article defines both, shows the trends and exceptions, uses a table of values, explains how to use electronegativity differences to classify bonds, works through examples and describes the different electronegativity scales and their history.

What is electron affinity?

The electron affinity (EA) is the energy change when one mole of gaseous atoms each gains one electron to form 1− ions: X(g) + e⁻ → X⁻(g). For most elements the process releases energy. By one common convention, the electron affinity is given as the energy released and so it has a positive value, for example 349 kJ/mol for chlorine. By another convention, the enthalpy change is quoted, and it is negative, −349 kJ/mol. Always check which convention is used. Atoms with a large electron affinity form stable anions easily. Atoms with filled subshells, such as the noble gases and the alkaline earth metals, have little or no affinity for an extra electron.

What is electronegativity?

Electronegativity (symbol χ) is the ability of an atom in a chemical bond to attract the bonding electrons toward itself. It is a relative number without units. In a bond between atoms of different electronegativity, the electrons are shared unequally, and the more electronegative atom becomes partially negative (δ−) while the other becomes partially positive (δ+). The most widely used scale was devised by Linus Pauling, who assigned fluorine, the most electronegative element, the value of 3.98.

Data for selected elements

Values are rounded. The electron affinity of nitrogen is close to zero (a small negative value of about −7 kJ/mol is sometimes quoted for the energy released), and those of Be, Mg and the noble gases are negative or nearly zero, which means that no stable 1− ion is formed in the gas phase.

What are the trends?

• Electronegativity increases from left to right across a period, as the effective nuclear charge increases and the atomic radius decreases, and it decreases down a group, as the bonding electrons are farther from the nucleus and more shielded. The most electronegative elements are fluorine, oxygen, nitrogen and chlorine, in the top right corner (excluding the noble gases), and the least are cesium and francium at the bottom left.
• Electron affinity also generally becomes larger (more energy is released) across a period, from the alkali metals to the halogens, and the values for the halogens are the largest in each period.
• Exceptions in electron affinity: chlorine has a higher value than fluorine, even though fluorine is above it. The small 2p orbital of fluorine is crowded, and the incoming electron is repelled by the other electrons. Similarly, sulfur’s value exceeds oxygen’s. Nitrogen has a very low value because of the stable half-filled 2p subshell.

How does electronegativity predict bond type?

The difference in electronegativity (ΔEN) between two bonded atoms gives a guide to the character of the bond. The boundaries are approximate.

Worked examples

1. Hydrogen chloride. ΔEN = 3.16 − 2.20 = 0.96. This falls in the polar covalent range, so the H–Cl bond is polar, with chlorine δ− and hydrogen δ+.
2. Sodium chloride. ΔEN = 3.16 − 0.93 = 2.23. This is above 1.7, so the bond is mainly ionic.
3. Methane. The C–H bond has ΔEN = 2.55 − 2.20 = 0.35, which is nonpolar, or only weakly polar.
4. Water. The O–H bond has ΔEN = 3.44 − 2.20 = 1.24, which is polar covalent, with oxygen δ−. The bent shape means that the molecule is polar overall.
5. Rank the bonds. Rank H–F, H–Cl, H–Br and H–I in order of increasing polarity. The differences are 3.98 − 2.20 = 1.78, 0.96, 2.96 − 2.20 = 0.76 and 2.66 − 2.20 = 0.46. From least to most polar: H–I < H–Br < H–Cl < H–F.
6. Energy of an electron gain. The energy released when 1.00 g of gaseous chlorine atoms (35.45 g/mol) each gain an electron is (1.00 ÷ 35.45) × 349 = 9.84 kJ.
7. Which atom is δ−? In C–O the atom with the higher electronegativity (oxygen, 3.44 versus carbon, 2.55) carries the partial negative charge.

What scales of electronegativity exist?

Who defined electronegativity?

The term was introduced by Jöns Jacob Berzelius in the early 19th century, in the sense of electropositive and electronegative elements. Linus Pauling gave it a quantitative form in 1932 in his work on the nature of the chemical bond, which was a major part of the research for which he received the Nobel Prize in Chemistry in 1954. Robert Mulliken proposed a scale based on atomic properties in 1934. Electronegativity values are still used in school and in research to estimate bond polarity, although the scales are not exact and do not always agree.

Which mistakes are common?

Mistake 1: Mixing up electron affinity and electronegativity

Electron affinity is an energy for an isolated atom gaining an electron. Electronegativity is the tendency of a bonded atom to attract shared electrons.

Mistake 2: Assuming that fluorine has the highest electron affinity

Chlorine has a higher value, though fluorine has the highest electronegativity.

Mistake 3: Using the sign of electron affinity carelessly

State clearly whether you quote energy released (positive) or the enthalpy change (negative).

Mistake 4: Treating the ΔEN limits as strict

The boundaries between bond types are approximate, and many bonds are intermediate.