How much energy does it take to remove an electron?
Drop sodium into water and it reacts violently, turning into Na⁺; neon barely reacts with anything. We say sodium “loses electrons easily”, but how easily? Chemists answer with a number they can measure: the energy needed to pull one electron off an atom.
First ionization energy
The first ionization energy第一电离能, I₁, is the minimum energy needed to remove one electron from a neutral gaseous atom in its ground state, usually given in kJ/mol: the kilojoules needed to ionize 1 mol of atoms (6.022 × 10²³ atoms; see lesson 2.2.2). For sodium, I₁ = 496 kJ/mol:
The (g) matters: ionization energy is defined for isolated gas-phase atoms, free of any neighbors.
- The lower the first ionization energy, the more easily the atom loses an electron, and the more metallic the element.
- In Example 2 of lesson 2.2.3, the Bohr model gave 2.18 × 10⁻¹⁸ J to ionize one hydrogen atom, or about 1.31 × 10³ kJ/mol, matching hydrogen’s measured first ionization energy of 1312 kJ/mol.
Trends in first ionization energy
- Across a period, first ionization energy generally increases. Zeff increases and the atoms get smaller, so the outer electrons are held more tightly. In every period, the alkali metal is lowest and the noble gas highest.
- Down a group, first ionization energy decreases. Each step down adds an energy level, so the outer electron is farther from the nucleus and easier to remove. That’s why the alkali metals get more reactive down the group (lesson 2.3.6).
- Helium has the highest first ionization energy of any element (2372 kJ/mol), and cesium, at the lower left, has the lowest (376 kJ/mol).
The dips
The overall trend rises across a period, but in four places an element is lower than the one before it:
- Be > B and Mg > Al. Boron (2s²2p¹) loses a 2p electron, which is higher in energy than a 2s electron and partly shielded by the 2s pair, so it’s easier to remove than beryllium’s 2s electron. Magnesium and aluminum work the same way (3s and 3p).
- N > O and P > S. Nitrogen’s three 2p electrons each have an orbital to themselves. In oxygen’s 2p⁴, one orbital holds a pair, and the two electrons repel each other, so one of them is easier to remove. Phosphorus and sulfur work the same way (3p³ and 3p⁴).
Chinese textbooks usually explain this another way: the full 2s² of Be and 3s² of Mg, and the half-filled 2p³ of N and 3p³ of P, are especially stable arrangements, so these elements have higher ionization energies than the next element. English books more often use the “higher-energy sublevel” and “paired-electron repulsion” explanations above. Both describe the same effect.
Successive ionization energies
After an atom loses its first electron, it can lose a second, a third, and so on. The energies needed, one after another, are the successive ionization energies, I₁, I₂, I₃, and so on.
- Successive ionization energies always increase: each ion has a larger positive charge and holds its remaining electrons more tightly.
- Once the valence electrons are gone and the next electron must come from the core, the ionization energy jumps by several times, because core electrons are closer to the nucleus and less shielded.
| Element | I₁ | I₂ | I₃ | I₄ |
|---|---|---|---|---|
| Na | 496 | 4562 | 6910 | 9543 |
| Mg | 738 | 1451 | 7733 | 10543 |
| Al | 578 | 1817 | 2745 | 11577 |
(In kJ/mol; the first value after each jump is in bold.) The electrons removed before the jump are the valence electrons: 1 for sodium, 2 for magnesium, and 3 for aluminum. That’s why these elements form Na⁺, Mg²⁺, and Al³⁺.
Click a few points on the first-ionization-energy plot and turn on the dip markers. Then switch to successive ionization energies, find each element’s jump, and finally try the mystery element:
Click a data point to see that element's first ionization energy.
- Be → B: Boron's first electron comes from the higher-energy 2p sublevel, so it's easier to remove than beryllium's 2s electron.
- N → O: Oxygen has one 2p orbital holding a pair of electrons that repel each other, so one of them is easier to remove; nitrogen's half-filled 2p is more stable.
- Mg → Al: Aluminum loses a 3p electron, which is higher in energy than magnesium's 3s electron, so it's easier to remove.
- P → S: Sulfur has a paired 3p electron that's repelled by its partner; phosphorus's half-filled 3p is more stable.
Example 1: Why is aluminum’s first ionization energy lower than magnesium’s?
Aluminum is to the right of magnesium and has a larger Zeff, yet I₁(Al) = 578 kJ/mol is lower than I₁(Mg) = 738 kJ/mol. Why?
- Write the configurations. Mg: [Ne]3s²; Al: [Ne]3s²3p¹.
- Which electron is removed? The first ionization removes the highest-energy electron: a 3s electron from magnesium, a 3p electron from aluminum.
- Compare those two electrons. 3p is higher in energy than 3s, the 3p electron is partly shielded by the 3s electrons, and it is slightly farther out on average, so aluminum’s 3p electron is easier to remove than magnesium’s 3s electron.
- The overall trend still holds. Beyond aluminum the values rise overall: Si 787, P 1012, S 1000, Cl 1251, Ar 1521 kJ/mol. The “exception” shows that ionization energy depends on which sublevel the electron comes from.
Example 2: Identifying an element from its successive ionization energies
A period 3 element X has successive ionization energies of 738, 1451, 7733, and 10543 kJ/mol. How many valence electrons does it have? What ion and what chloride does it form?
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Take ratios of neighbors. Successive ionization energies always increase; the key is where they jump:
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Find the jump. It comes between I₂ and I₃: the first 2 electrons are valence electrons, and the third must come from an inner shell much closer to the nucleus.
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Locate it. 2 valence electrons means group 2, and the group 2 element in period 3 is magnesium.
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Ion and compound. X loses its 2 valence electrons to form X²⁺, which combines 1 : 2 with Cl⁻ to give XCl₂, that is, MgCl₂.
- Na or Mg
Same period, and magnesium has the larger Zeff: Na 496, Mg 738 kJ/mol.
- Al or Mg
A dip: aluminum loses a higher-energy 3p electron. Al 578, Mg 738 kJ/mol.
- N or O
A dip: oxygen's 2p⁴ has a pair of electrons that repel each other. N 1402, O 1314 kJ/mol.
- Na or Li
Ionization energy decreases down a group, and lithium is above sodium. Li 520, Na 496 kJ/mol.
- F or Cl
Same group, and fluorine's electrons are closer to the nucleus. F 1681, Cl 1251 kJ/mol.
- B or Be
A dip: boron loses a 2p electron. B 801, Be 900 kJ/mol.
- C or N
The general trend rises across a period. C 1086, N 1402 kJ/mol.
- He or Ne
Helium's electrons are in n = 1, closest to the nucleus: He 2372, Ne 2081 kJ/mol. Helium has the highest value of any element.
- Ca or K
Same period, and calcium has the larger Zeff. Ca 590, K 419 kJ/mol.
Common mistakes
Mistake 1: “A higher first ionization energy means the atom loses electrons more easily.” It’s the reverse. Ionization energy is the energy needed to remove an electron, so the more energy it takes, the harder the electron is to remove. Metals generally have low first ionization energies.
Mistake 2: “First ionization energy rises steadily across a period.” It rises overall, but with dips at Be > B, N > O, Mg > Al, and P > S. When comparing neighbors, first check which sublevel each electron comes from.
Mistake 3: “The electron at the jump is the last valence electron.” The jump means that electron comes from the core. The valence electrons are the ones removed before the jump: magnesium’s jump is at I₃, so it has 2 valence electrons, not 3.
Quick check
Which period 2 element has the highest first ionization energy?
First ionization energy rises across a period overall, and the noble gas is highest: Ne 2081 kJ/mol, F 1681 kJ/mol.
An element has successive ionization energies of 801, 2427, 3660, and 25026 kJ/mol. Which group is it most likely in?
I₄/I₃ ≈ 6.8, so the jump comes at the fourth electron. The element has 3 valence electrons and is in group 13. It's boron.
Summary
- First ionization energy, I₁, is the minimum energy to remove one electron from a gaseous ground-state atom. The lower I₁, the more easily the atom loses an electron and the more metallic the element.
- It generally increases across a period and decreases down a group, with dips at Be > B, N > O, Mg > Al, and P > S.
- Successive ionization energies keep rising. The number of electrons removed before the big jump is the number of valence electrons, which gives a main-group metal’s usual ion charge.
On to AP (for later)
In AP U1.6, photoelectron spectroscopy (PES) measures the binding energy of the electrons in every sublevel of an atom, usually in MJ/mol (1 MJ/mol = 1000 kJ/mol). The lowest-energy peak on a PES spectrum corresponds to the first ionization energy. AP U1.7 questions often ask you to compare and explain two elements’ first ionization energies: within a period, compare Zeff; within a group, compare how far the electron is from the nucleus.