Electron affinity Knowledge-Based Study

Study notes

1. Electron affinity (EA) is a fundamental periodic property that measures the energy change associated with adding an electron to a gaseous atom.

   First Electron Affinity

   The first electron affinity is formally defined as the enthalpy change when one mole of electrons is added to one mole of isolated gaseous atoms to form one mole of uninegative gaseous ions:

   $\text{X}(g) + e^- \rightarrow \text{X}^-(g)$

   For most neutral atoms, adding an electron results in an electrostatic attraction between the incoming electron and the nucleus. Therefore, the first electron affinity is typically an exothermic process ($\Delta H < 0$). In standard thermodynamic conventions, a negative enthalpy change corresponds to energy release, although chemistry texts often define electron affinity as the quantity of energy released, quoting positive values.

   Like atomic radius and ionization energy, electron affinity follows distinct patterns across the periodic table governed by effective nuclear charge and atomic size.

   Across a Period (Left to Right)

   Moving from left to right across a period, electron affinity generally becomes more exothermic (more negative $\Delta H$). This trend arises because:

   • The effective nuclear charge ($Z_{\text{eff}}$) increases while valence electrons enter the same principal energy level.
   • The atomic radius decreases, bringing the valence shell closer to the positively charged nucleus.
   • The attraction between the incoming electron and the nucleus becomes stronger, releasing more energy upon capture.

   Down a Group (Top to Bottom)

   Moving down a group, electron affinity generally becomes less exothermic (less negative $\Delta H$). Key factors include:

   • Additional electron shells increase the principal quantum number ($n$).
   • The atomic radius increases and inner electron shells provide substantial shielding.
   • The incoming electron is situated farther from the nucleus, weakening the electrostatic attraction.

   Deviations from the expected periodic trends occur due to electron-electron repulsion and specific subshell stabilities.

   Group 2 and Group 18 Elements

   Alkaline earth metals (Group 2) have filled $s$ subshells ($ns^2$), meaning an added electron must enter a higher-energy $p$ subshell, resulting in very low or positive $\Delta H$. Noble gases (Group 18) possess fully stable octets ($ns^2 np^6$); adding an electron requires opening an entirely new principal shell, making the process decidedly endothermic.

   Group 15 Elements

   Nitrogen and other Group 15 elements have a half-filled $p$ subshell ($np^3$). Adding an electron forces pairing in an already occupied $p$ orbital, producing interelectronic repulsion that reduces the overall energy released compared to preceding Group 14 elements.

   The Fluorine versus Chlorine Anomaly

   Although halogen affinity generally decreases down Group 17, chlorine has a more exothermic electron affinity than fluorine:

   • Fluorine's valence electrons occupy the extremely compact $2p$ subshell.
   • High electron density in this small volume causes significant electron-electron repulsion, which partially offsets the attraction to the nucleus.
   • Chlorine accepts the electron into a larger, more diffuse $3p$ subshell where interelectronic repulsion is considerably smaller.

   Certain nonmetals, such as oxygen and sulfur, can form polynegative ions (e.g., $\text{O}^{2-}$, $\text{S}^{2-}$). This requires the addition of a second electron.

   Second Electron Affinity

   The second electron affinity represents the addition of an electron to a singly charged anion:

   $\text{X}^-(g) + e^- \rightarrow \text{X}^{2-}(g)$

   Unlike the first electron affinity, the second electron affinity is always endothermic ($\Delta H > 0$) for all elements. Energy must be supplied to overcome the strong electrostatic repulsion between the incoming electron and the negatively charged ion.

   Stable ionic compounds containing oxide ($\text{O}^{2-}$) or sulfide ($\text{S}^{2-}$) ions exist exclusively in the solid state, where the substantial endothermic cost of the second electron affinity is overcome by the high lattice energy of the resulting ionic crystal lattice.