Concept 05 / 05Classification of Elements and Periodicity
Electron Gain Enthalpy and Electronegativity
Definition and sign convention of electron gain enthalpy, trends across periods and down groups, why chlorine has a more negative value than fluorine, positive values for noble gases and second electron gain, electronegativity as a bonded-atom property (Pauling scale), its trends and uses for bond polarity and oxide character.
Electron gain enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom: . It is negative (exothermic) when the atom readily accepts an electron. Electron affinity is often quoted as the same quantity with the opposite sign (energy released).
Trends.
Across a period, generally becomes more negative, because increases and atoms are smaller; the halogens have the most negative values, one electron short of a noble-gas configuration.
Down a group it generally becomes less negative, as the added electron is farther from the nucleus.
Noble gases have positive values ( , kJ/mol): the added electron must go into a new, higher shell. Group elements and nitrogen also have near-zero or positive values, because of stable and half-filled configurations.
The fluorine anomaly. of Cl ( kJ/mol) is more negative than that of F (). Fluorine is so small that its subshell is compact and crowded; the incoming electron experiences strong repulsion from the electrons already there. In the larger subshell of chlorine this repulsion is less. The same effect makes S () more negative than O (). So the most negative electron gain enthalpy in the periodic table belongs to chlorine.
Second electron gain enthalpy is always positive: adding an electron to a negative ion requires overcoming repulsion. is strongly endothermic (about kJ/mol). Oxides nevertheless form because the large lattice energy of compounds more than compensates.
Electronegativity is the tendency of an atom in a molecule to attract the shared pair of electrons towards itself. Unlike ionisation and electron gain enthalpies, it is not a measurable property of an isolated atom; it is a relative scale. On the Pauling scale (based on bond energies), fluorine is highest (about ), followed by O (), then N and Cl (), C () and H (); caesium and francium are lowest (about ). The Mulliken scale averages ionisation energy and electron affinity; the Allred-Rochow scale uses .
Trends: electronegativity increases across a period and decreases down a group, because it follows and size. Fluorine, not chlorine, is the most electronegative element; the size effect that lowers F's electron gain enthalpy does not apply in the same way when the electron is shared in a bond.
Uses of electronegativity.
Bond polarity: the larger the difference between two atoms, the more polar the bond; a difference above about usually indicates predominantly ionic bonding.
Nature of oxides and hydroxides: metals of low electronegativity give basic oxides (), non-metals of high electronegativity give acidic oxides (), and elements in between give amphoteric oxides (, ).
Metallic and non-metallic character: low electronegativity correlates with metallic character.
It also depends on hybridisation (more s character, higher electronegativity: sp carbon > sp > sp) and oxidation state (higher oxidation state, higher electronegativity).
Key Takeaways:
- negative means energy released on gaining an electron.
- Most negative: Cl (), not F (), because of repulsion in the small subshell.
- Noble gases, group 2, N: positive or near-zero .
- Second electron gain enthalpy is always positive.
- Electronegativity: property of a bonded atom; F highest (Pauling ).
- large: ionic; oxides become more acidic as electronegativity rises.
- Pitfall: stating that F has the most negative electron gain enthalpy.
Brain-Teaser 1: Arrange F, Cl, Br, I in order of (a) electronegativity, (b) negative electron gain enthalpy.
Solution: (a) F > Cl > Br > I. (b) Cl > F > Br > I (values ). The two orders differ only in the position of F, which is exactly the anomaly caused by fluorine's small size.
Brain-Teaser 2: Why is the electron gain enthalpy of nitrogen close to zero while that of carbon is negative ( kJ/mol)?
Solution: Carbon () accepts an electron into an empty orbital to form , a stable half-filled set, so energy is released. Nitrogen () is already half-filled; the added electron must pair up in an occupied orbital, facing extra repulsion and destroying the stable arrangement, so almost no energy is released. The same configuration argument that explains ionisation exceptions explains electron gain exceptions.
Brain-Teaser 3: Predict the acidic or basic nature of the oxides of the third-period elements , MgO, , , , , .
Solution: As electronegativity increases across the period, oxides change from strongly basic (), basic (MgO), amphoteric (), weakly acidic (), acidic (, ) to strongly acidic (, the anhydride of ). An electropositive element releases (base), while an electronegative one holds oxygen covalently and its oxide reacts with water to give an oxoacid.
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