Electron Affinity and Electronegativity
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You should be comfortable with the concept of effective nuclear charge (\(Z_{\text{eff}}\)) and Slater's rules as established in the previous lecture on Ionisation Energy. Familiarity with basic quantum numbers and orbital filling order is assumed.
Why this matters
Electron affinity and electronegativity quantify how atoms attract and share electrons, forming the basis for understanding chemical bonding and reactivity.
While ionisation energy measures the energy required to remove an electron, electron affinity measures the energy change when an electron is added. Electronegativity, conversely, describes the tendency of an atom to attract shared electrons in a covalent bond. Together, these properties explain why certain elements form ionic bonds while others form polar covalent bonds, and they provide a quantitative framework for predicting molecular polarity and reaction pathways.
The Intuitive Picture
Imagine two atoms approaching each other. If one atom has a strong desire to gain an electron and the other has a weak hold on its own, the electron will shift towards the first atom, creating a dipole. Electron affinity is the "welcome" an atom extends to an extra electron. A high (more negative) electron affinity means the atom releases energy when it accepts an electron, indicating a strong attraction. Electronegativity is the "pull" an atom exerts on shared electrons. It is a relative scale, not an absolute energy value, but it correlates strongly with both ionisation energy and electron affinity. Elements with high electronegativity, like fluorine, are electron-hungry; elements with low electronegativity, like caesium, are electron-generous.
Electron Affinity: Trends and Exceptions
Electron affinity (\(EA\)) is defined as the energy change when a gaseous atom gains an electron to form a gaseous anion. By convention, a negative value indicates an exothermic process (energy released), which is the case for most non-metals. The trend across a period generally follows the increase in \(Z_{\text{eff}}\). As we move from left to right, the nucleus exerts a stronger pull on the added electron, making the process more exothermic. However, there are significant exceptions due to electron-electron repulsion in small orbitals.
Consider the halogens. Fluorine has a less negative electron affinity than chlorine. This is because the fluorine atom is very small, and the added electron enters the compact \(2p\) subshell. The repulsion between the new electron and the existing \(2p\) electrons is significant, reducing the net energy release. In chlorine, the added electron enters the larger \(3p\) subshell, where repulsion is less severe, resulting in a more exothermic process. Similarly, noble gases have positive electron affinities because adding an electron requires placing it in a new, higher-energy shell, which is energetically unfavourable.
Electronegativity: The Pauling Scale
Electronegativity is a dimensionless number that represents the tendency of an atom to attract a bonding pair of electrons. The most common scale is the Pauling scale, where fluorine is assigned the highest value of 4.0. The scale is derived from bond dissociation energies of various molecules. The trend in electronegativity mirrors the trend in ionisation energy and electron affinity: it increases from left to right across a period and decreases from top to bottom down a group.
This trend is driven by the same factor: effective nuclear charge. As \(Z_{\text{eff}}\) increases across a period, the nucleus holds onto its own electrons more tightly and attracts shared electrons more strongly. Down a group, the increase in principal quantum number \(n\) increases the distance between the nucleus and the valence electrons, and the shielding effect reduces the nuclear pull, leading to lower electronegativity. For example, in Group 17, fluorine (4.0) is more electronegative than chlorine (3.0), which is more electronegative than bromine (2.8).
Connecting to Periodic Trends
The properties of electron affinity and electronegativity are not independent; they are both manifestations of the effective nuclear charge experienced by valence electrons. The result from the previous lecture, \(Z_{\text{eff}}\), provides the unifying explanation. When \(Z_{\text{eff}}\) is high, the atom has a strong attraction for electrons, leading to high ionisation energy, high (negative) electron affinity, and high electronegativity. When \(Z_{\text{eff}}\) is low, these values are lower. This consistency allows chemists to predict the behaviour of elements in compounds. For instance, a large difference in electronegativity between two atoms (\(\Delta \chi > 1.7\)) typically indicates an ionic bond, while a small difference (\(\Delta \chi < 0.5\)) indicates a non-polar covalent bond.
The result
Periodic trends
Define the effective nuclear charge felt by an electron as \(Z_{\text{eff}}=Z-S\), where \(S\) (the shielding constant) is computed from Slater's rules: for an electron in an \(ns\) or \(np\) orbital, sum \(0.35\) for every other electron in the same \(n s,np\) group, \(0.85\) for every electron in the \(n{-}1\) shell, and \(1.00\) for every electron in shells \(n{-}2\) and below. Then, moving left to right across a period, \(Z_{\text{eff}}\) for the valence electron increases (atomic radius decreases, first ionisation energy increases); moving down a group, \(Z_{\text{eff}}\) stays roughly constant while the valence principal quantum number \(n\) increases (atomic radius increases, first ionisation energy decreases).
A common misconception is that electron affinity always increases across a period. This is incorrect: in beryllium the extra electron must enter the higher-energy 2p subshell, and in nitrogen it must pair up inside the half-filled 2p subshell, leading to less negative (or even positive) electron affinities compared to their neighbours. Always consider the electronic configuration when predicting trends. C
Worked example
Problem. Calculate the electronegativity difference between sodium (\(\chi = 0.9\)) and chlorine (\(\chi = 3.0\)) and determine the bond type based on the standard threshold of 1.7 for ionic character.
- Identify the electronegativity values: \(\chi_{\text{Na}} = 0.9\) and \(\chi_{\text{Cl}} = 3.0\).
- Calculate the difference: \(\Delta \chi = |3.0 - 0.9| = 2.1\).
- Compare with the threshold: \(2.1 > 1.7\).
- Conclusion: The bond is predominantly ionic.
Units check. Electronegativity is dimensionless, so the difference is also dimensionless.
Three quick tests
- Which element has a more negative electron affinity, fluorine or chlorine?
Answer
Chlorine, due to less electron-electron repulsion in the larger 3p orbital. - How does electronegativity change down Group 1?
Answer
It decreases, as atomic radius increases and shielding increases. - What is the electronegativity difference between carbon (\(\chi = 2.5\)) and hydrogen (\(\chi = 2.1\))?
Answer
0.4, indicating a non-polar covalent bond.
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Every lecture and every result in this unit, with prerequisites and threads.
Lecture notes drafted with a local model and checked line by line by the editor; every number in the worked example was recomputed.