Ionization Energy and Electron Affinity — Quick Reference

Ionization Energy

  • Definition: Ionization energy is the minimum energy required to remove an electron from a gaseous atom or ion in its ground state. The first ionization energy is the energy to remove the first electron from a neutral atom.
    • Process for the first ionization: ext{A(g) → A$^{+}$(g) + e$^{-}$} with energy I1I_1.
    • For the second ionization: ext{A$^{+}$(g) → A$^{2+}$(g) + e$^{-}$} with energy I2I_2.
  • Examples:
    • Sodium: the first ionization energy corresponds to the process ext{Na(g) → Na$^{+}$(g) + e$^{-}$}.
    • The general trend is that successive ionization energies increase because electrons are being removed from particles with increasing positive charge.
  • Variation in successive ionization energies
    • After removing outer-shell electrons, a much larger jump occurs when an inner-shell electron is removed (core electron).
    • Example (silicon): removal energies increase from 786 to 4356 kJ/mol for the four outer electrons, but removal of a 2p electron (inner shell) is about 16,091 kJ/mol.
    • Reason: inner-shell electrons experience a much higher effective nuclear charge (Z_eff) and are closer to the nucleus.
  • Periodic trends (first ionization energy, I):
    • Across a period (left to right): I generally increases.
    • Alkali metals have the lowest I in each period; noble gases have the highest.
    • Down a group: I generally decreases.
    • Noble gases: He > Ne > Ar > Kr > Xe in order of IE.
    • s- and p-block elements show a larger range of I values than transition metals; transition metals increase more slowly across a period; f-block metals show small variation.
    • Smaller atoms tend to have higher ionization energies because the outer electron is closer to the positively charged nucleus and is less shielded.
    • Factors affecting IE: a) increasing effective nuclear charge (Z_eff) and decreasing atomic radius across a period increase IE; b) increasing radius down a group decreases IE.
  • Irregularities within a period
    • Be → B: IE drops slightly because B’s outer electron enters the 2p subshell (higher energy than 2s).
    • N → O: IE drops slightly due to electron-electron repulsion when pairing occurs in the 2p orbitals (Hund’s rule; half-filled vs filled subshell considerations).
  • Electron configurations of ions
    • When removing electrons to form cations, electrons are removed first from the orbitals with the largest principal quantum number n.
    • Example: Li → Li$^+$ removes the 2s electron: ext{Li(1s$^2$ 2s$^1$) → Li$^+$ (1s$^2$) + e$^-$}.
    • Fe: ext{Fe([Ar] 4s$^2$ 3d$^6$) → Fe$^{2+}$([Ar] 3d$^6$) + 2e$^-$}.
    • If a further electron is removed (Fe$^{3+}$), it comes from a 3d orbital (n = 3) since 4s and 3d shells with n = 4 are empty.
    • Electrons added to form anions are added to the empty or partially filled orbital having the lowest value of n.
    • Example: F + e$^-$ → F$^-$; added to the remaining vacancy in the 2p subshell.
    • Common ion configurations illustrate the ideas: Ca → Ca$^{2+}$ isoelectronic with Ar: [Ar]; Co → Co$^{3+}$ (removing 4s electrons first, then 3d): [Ar] 3d$^6$; S → S$^{2-}$: [Ne] 3s$^2$ 3p$^6$ = [Ar]; Sn examples show changes in ns and nd ordering during ionization.
  • Electron affinity (EA)
    • Definition: Electron affinity is the energy change when an electron is added to a gaseous atom.
    • Process: ext{A(g) + e$^{-}$ → A$^{-}$(g)} with energy change EAEA.
    • Cl example: ext{Cl(g) + e$^{-}$ → Cl$^{-}$(g)}, EA=−349extkJ/molEA = -349 ext{ kJ/mol} (exothermic; energy released).
    • Sign conventions
    • Thermodynamic convention (used here): negative EA indicates exothermic addition of an electron.
    • Historical convention: EA would be +349 kJ/mol for Cl under a different sign convention.
    • Trends and interpretation
    • Atoms readily gain electrons to form anions if the process is energetically favorable (large, negative EA).
    • Halogens generally have large (favorable) electron affinities; noble gases have very small or positive EA values (unfavorable to add an electron to an already full shell).
  • Distinctions: IE vs EA
    • IE measures energy required to remove an electron; EA measures energy change when an electron is added.
    • Both relate to an atom’s tendency to gain or lose electrons and influence chemical bonding and reactivity.

Periodic Trends in First Ionization Energies (summary)

  • Across a period: I₁ generally increases from left to right; alkali metals lowest, noble gases highest.
  • Down a group: I₁ generally decreases with increasing atomic size.
  • Irregularities: small shifts due to subshell occupancy and electron-electron repulsion (e.g., Be→B, N→O patterns).
  • Across the table, smaller atoms tend to have higher I due to greater attraction between nucleus and outer electrons; radius and Z_eff drive the trends.

Electron Configurations of Ions (quick rules)

  • Cations: remove electrons from orbitals with the largest n first; if multiple orbitals share the same n, remove from the one with higher energy first (and, in transition metals, 4s electrons are removed before 3d electrons).
    • Example: Ca → Ca$^{2+}$: [Ar]
    • Example: Co → Co$^{3+}$: [Ar] 3d$^6$
  • Anions: add electrons to the empty or partially filled orbital with the lowest n.
    • Example: F → F$^{-}$: completes the 2p shell to [Ne] 2s$^2$ 2p$^6$
  • Common transitions show how electron configurations change when electrons are removed or added (e.g., Sn: removal from 5s before 5p; resulting ion configurations can differ from the neutral atom's order).

Electron Affinity (EA) essentials

  • Definition: Energy change when an electron is added to a gaseous atom.
    • Process: ext{A(g) + e$^{-}$ → A$^{-}$(g)}; EA is the energy change for this process.
  • Sign conventions
    • In this text: exothermic gain of electron → EA < 0 (negative value).
    • Historical: EA would be reported as a positive value for the amount of energy released.
  • Typical values and trends
    • Halogens have large, favorable electron affinities (high tendency to gain an electron).
    • Noble gases have small or unfavorable EA values due to filled subshell stability.
  • Compare IE and EA
    • IE: energy to remove an electron (positive input required).
    • EA: energy released or absorbed when adding an electron (often negative, indicating energy released).

Quick practice and reference points

  • Example ordering (from lowest to highest first ionization energy): K < Na < P < Ar < Ne (consistent with across a period and group trends).
  • First ionization energy irregularities are explained by subshell occupancy and electron-electron repulsion in partially filled subshells.
  • For ions, remember the removal/addition rules: remove from highest n when forming cations; add to lowest available n when forming anions.
  • Typical first ionization energies (kJ/mol) for reference: H ≈ 1312, He ≈ 2372, Li ≈ 520, Be ≈ 900, B ≈ 800, C ≈ 1086, N ≈ 1402, O ≈ 1314, F ≈ 1681, Ne ≈ 2081, Na ≈ 496, Mg ≈ 738, Al ≈ 578, Si ≈ 786, P ≈ 1012, S ≈ 1000, Cl ≈ 1251, Ar ≈ 1521.