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General Chemistry I · Periodic Properties

Ionization Energy

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On this page 7 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Study tools
  7. Sources & references

In 30 seconds

Ionization energy (IE) is the energy required to remove an electron from a gaseous atom or ion. It measures how tightly an electron is held. Because effective nuclear charge rises across a period, IE generally increases left to right; because valence electrons sit in higher, more shielded shells down a group, IE decreases top to bottom. The trend has two well-known exceptions — the drops from Group 2 to 3 and from Group 15 to 16 — explained by orbital filling and electron pairing.

Why this matters

Ionization energy explains which elements form cations and how easily. Metals (low IE) readily lose electrons and form positive ions; nonmetals (high IE) hold electrons tightly. The successive-IE "jump" is the experimental fingerprint chemists use to deduce an element's group and its typical charge. IE values also feed into Born–Haber cycles and help predict chemical reactivity — the explosiveness of the alkali metals follows directly from their very low first ionization energies.

The college version

Key Ideas

  • First ionization energy (IE₁): energy to remove the first (outermost) electron from a gaseous atom: X(g) → X⁺(g) + e⁻. Always endothermic (positive).
  • Successive ionization energies (IE₂, IE₃, …): energies to remove the 2nd, 3rd, etc. electron. Each is larger than the previous one.
  • Across a period (→): IE generally increases (Zeff increases, electrons held tighter).
  • Down a group (↓): IE decreases (higher n, more shielding, weaker pull).
  • Exception 1 — Group 2 → 3: IE drops (e.g., Be 899 → B 801 kJ/mol) because the electron removed from B is in a higher-energy, more-shielded 2p orbital than Be's 2s electron.
  • Exception 2 — Group 15 → 16: IE drops (e.g., N 1402 → O 1314 kJ/mol) because O has paired 2p electrons whose repulsion makes one easier to remove, whereas N's half-filled 2p³ is relatively stable.
  • Large jump in successive IEs: occurs when removal switches from valence to core (noble-gas) electrons — this jump reveals the group number.

Equations and Variables

  • First IE: X(g) + IE₁ → X⁺(g) + e⁻, where IE₁ is the (positive) energy per mole, in kJ/mol.
  • Second IE: X⁺(g) + IE₂ → X²⁺(g) + e⁻.
  • General rule: IE₁ < IE₂ < IE₃ < … (removing each subsequent electron from an increasingly positive ion is harder).

How It Works

  1. An electron is bound by the attraction of the nucleus (measured by Zeff) reduced by its distance (shell n) and shielding.
  2. Across a period, n is constant but Zeff rises, so the valence electrons are held more tightly and IE increases.
  3. Down a group, the valence electron moves to a larger, more shielded shell, so it is held more loosely and IE decreases.
  4. The two dips arise from electron configuration: removing a lone 2p electron (B) is easier than a 2s electron (Be), and removing one of a paired set (O) is easier than one of a half-filled set (N).
  5. Once all valence electrons are removed, the next electron comes from a filled inner shell very close to the nucleus, so IE jumps enormously.

Worked Example

An element has IE₁ = 738, IE₂ = 1451, IE₃ = 7733 kJ/mol. Identify its group.

The pattern: 738 → 1451 is a modest increase (both valence electrons), then a huge jump to 7733. That jump means the third electron is being pulled from a filled, much-lower-energy shell. Two valence electrons are removed relatively easily, so the element is in Group 2 — the element is magnesium (Mg). (For comparison, Na shows the jump after only one electron: IE₁ 496 → IE₂ 4562.)

Second check — rank IE₁ for Li, Na, K, Cs: IE decreases down group 1, so Li > Na > K > Cs (Li 520, Na 496, K 419, Cs 376 kJ/mol).

Common Confusions

  • "Ionization energy is the energy released when an electron is added." — Wrong: that is electron affinity. Ionization energy removes an electron and is always endothermic.
  • "The trend increases smoothly with no exceptions." — Wrong: there are two dips (Group 2→3 and 15→16) explained by orbital configuration.
  • "Successive ionization energies can decrease." — Wrong: each successive electron is removed from a more positive ion, so IE always rises; only the size of the increase varies.
  • "Higher IE means the atom wants to lose electrons." — Wrong: high IE means the atom holds electrons tightly (typical of nonmetals), so it does not readily lose them.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Think of ionization energy as how hard you have to pull to yank a magnet off a fridge. A small, strong magnet close to the metal (a tightly held electron) is hard to pull off. Going down a group, the magnet moves farther away, so it's easier to pull. Going across a period, the fridge gets stronger, so it's harder to pull off — except for two "wiggles" where the magnet sits in an awkward, easy-to-pop spot (a lone p-electron, or a paired set that repels itself). The big jump in successive energies is like the difference between peeling off a fridge magnet (easy) and prying out a nail hammered into the wall (hard) — valence vs. core electrons. (The analogy treats electrons as single objects; the real reason is quantum shells and Zeff, but the "pull strength" picture is right.)

Key takeaways

  • IE₁ is always positive (endothermic): energy must be supplied to remove an electron.
  • IE increases left → right; decreases top → bottom.
  • Two dips: Group 2 → 3 (2s vs 2p) and Group 15 → 16 (paired vs half-filled p).
  • IE₁ < IE₂ < IE₃ < … always.
  • A large successive-IE jump signals removal of a core electron; its position reveals the group.
  • Lowest IE: Cs/Fr (lower-left); highest IE: He (and generally the noble gases).
  • Representative values (kJ/mol): Li 520, Na 496, Be 899, B 801, N 1402, O 1314, F 1681, Ne 2081.
  • IE = energy to remove an electron; IE₁ always positive.
  • Increases across a period (Zeff ↑), decreases down a group (n ↑).
  • Dips at Group 2→3 and 15→16 from orbital/electron-pairing effects.
  • Successive IEs always increase; a big jump marks the valence/core boundary.
  • The jump position identifies the element's group.
  • Low IE → metal (forms cations); high IE → nonmetal.

Keep learning

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Practice General Chemistry I

This lesson has no separate scored set. Practice draws from the subject’s question bank.

Study tools & related lessonsYou’ll learn to · Related

You’ll learn to

  • Define first and successive ionization energies and write the associated equations.
  • Explain why ionization energy increases across a period and decreases down a group.
  • Account for the two "dips" (Group 2 → 3 and Group 15 → 16) in the trend.
  • Use successive ionization energies to identify an element's group.

Sources & references

  1. OpenStax, *Chemistry 2e*, "6.5 Periodic Variations in Element Properties."
  2. Purdue University Chemistry, "Ionization Energy and Electron Affinity."
  3. LibreTexts Chemistry, "Ionization Energy."
  4. PubChem, "Sodium (Element)."

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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