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Textbooks often say that sodium “gives” its outer electron to chlorine so that both achieve a full outer shell. It’s a useful picture, but it hides the real reason ions form, and it’s actually misleading: pulling an electron off a sodium atom costs energy, and chlorine gaining that electron doesn’t pay it all back. Ionic compounds form because of what happens after the ions are made. This article follows the energy through each step, and explains how and where ions form in the real world.
For the basics of what ions are, see what is an ion?.
Making a positive ion: ionisation energy
To remove an electron from an atom in the gas phase, you must supply energy. The first ionisation energy is the energy needed to remove one electron from each atom in a mole of gaseous atoms:
Na(g) → Na⁺(g) + e⁻ ΔH = +496 kJ/mol
Ionisation energies are always positive (energy is taken in), because the electron is attracted to the nucleus and must be pulled away.
- They’re lowest for group 1 metals, whose single outer electron is far from the nucleus and well shielded.
- They increase across a period as nuclear charge increases.
- Successive ionisation energies increase, and there’s a huge jump when an electron must be removed from a full inner shell. Sodium’s second ionisation energy is 4,562 kJ/mol, over nine times the first, which is why sodium forms Na⁺ and never Na²⁺.
See ionization energy trend and effective nuclear charge.
Making a negative ion: electron affinity
When a gaseous atom gains an electron, energy is usually released. The first electron affinity is the energy change when each atom in a mole of gaseous atoms gains one electron:
Cl(g) + e⁻ → Cl⁻(g) ΔH = −349 kJ/mol
Non-metals near the top right of the periodic table (the halogens especially) release the most energy, because the incoming electron is attracted strongly by a high effective nuclear charge. See electron affinity.
But adding a second electron is a different story:
O⁻(g) + e⁻ → O²⁻(g) ΔH ≈ +800 kJ/mol
The second electron is being forced onto an ion that’s already negative, so it’s repelled. Energy must be supplied. Oxide ions (O²⁻) exist in compounds only because other energy terms make up for this.
The puzzle: isolated ions cost energy
Put the two steps together for sodium and chlorine:
| Step | Energy (kJ/mol) |
|---|---|
| Na(g) → Na⁺(g) + e⁻ | +496 |
| Cl(g) + e⁻ → Cl⁻(g) | −349 |
| Net | +147 |
Making a pair of separate gaseous ions costs 147 kJ/mol. And that’s before counting the energy needed to turn solid sodium into gaseous atoms (+107 kJ/mol) and to break chlorine molecules into atoms (+122 kJ/mol for half a mole of Cl₂).
So why does sodium burn so vigorously in chlorine, releasing a lot of energy?
The payoff: lattice energy
When gaseous Na⁺ and Cl⁻ ions come together to form a solid crystal, each ion is surrounded by ions of opposite charge. The strong electrostatic attraction releases a large amount of energy, the lattice energy:
Na⁺(g) + Cl⁻(g) → NaCl(s) ΔH ≈ −787 kJ/mol
This is far larger than the cost of making the ions. Adding everything up (a Born–Haber cycle):
| Step | ΔH (kJ/mol) |
|---|---|
| Na(s) → Na(g) (atomisation) | +107 |
| ½Cl₂(g) → Cl(g) (atomisation) | +122 |
| Na(g) → Na⁺(g) + e⁻ (first ionisation energy) | +496 |
| Cl(g) + e⁻ → Cl⁻(g) (first electron affinity) | −349 |
| Na⁺(g) + Cl⁻(g) → NaCl(s) (lattice formation) | −787 |
| Na(s) + ½Cl₂(g) → NaCl(s) | −411 |
The overall reaction is strongly exothermic (−411 kJ/mol, the standard enthalpy of formation of sodium chloride). The ions form because the attraction between them in the lattice releases more than enough energy to pay for making them. See exothermic vs endothermic.
Lattice energy depends on charge and size
Lattice energy is larger for:
- higher charges: MgO (2+ and 2−) has a lattice energy of about −3,800 kJ/mol, nearly five times that of NaCl
- smaller ions: ions pack closer together, so attraction is stronger
That’s why magnesium forms Mg²⁺ and oxygen forms O²⁻ in MgO, even though forming O²⁻ requires energy: the enormous lattice energy of a 2+/2− compound more than compensates. See ionic radius.
Ions in solution: hydration
In water, ions don’t need a crystal lattice to be stabilised. Water molecules are polar: their slightly negative oxygen atoms surround cations, and their slightly positive hydrogen atoms surround anions. This hydration releases energy, the hydration enthalpy. When a salt dissolves, the energy released by hydrating the ions is compared with the lattice energy needed to separate them; the balance helps decide whether the salt dissolves. See polar vs nonpolar molecules and solubility rules.
Where ions form in the real world
- In chemical reactions between metals and non-metals, forming ionic compounds. See ionic vs covalent bonds.
- In solution, when acids ionise (HCl + H₂O → H₃O⁺ + Cl⁻) or salts dissolve and dissociate.
- In flames and plasmas, where high temperatures ionise atoms, used in ICP-MS and seen in lightning.
- In batteries and electrolysis, where metals lose electrons at one electrode and ions gain electrons at another. See oxidation and reduction.
- In the upper atmosphere, where ultraviolet and X-rays from the Sun ionise gases, forming the ionosphere, which reflects radio waves.
- In mass spectrometers, where molecules are deliberately ionised so they can be steered by fields. See mass spectrometry.
- In radiation detectors, where alpha, beta and gamma radiation ionise gas atoms, producing a measurable current.
Predicting which ions form
Putting the energy picture together gives some reliable rules of thumb:
- Group 1, 2 and 13 metals form ions with the charge of their group (+1, +2, +3), because removing further electrons would mean breaking into an inner shell, which costs far more energy than any lattice could repay.
- Group 15, 16 and 17 non-metals form ions of charge −3, −2 and −1, reaching the next noble gas configuration; beyond that, extra electrons would enter a new shell and be very weakly held.
- Transition metals often form several ions (Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺), because successive ionisation energies of their d electrons rise more gradually.
See how to predict the charge of an ion.
Common misconceptions
- “Sodium wants to lose an electron.” Removing the electron costs 496 kJ/mol; it only happens because of the energy released later.
- “Chlorine gaining an electron releases enough energy to pay for sodium losing one.” It doesn’t (−349 vs +496); lattice energy makes the difference.
- “A full outer shell is always the lowest-energy state.” O²⁻ has a full shell, but forming it from O⁻ costs energy.
- “Ions exist as separate pairs in a crystal.” In NaCl, each Na⁺ is surrounded by six Cl⁻ ions, and vice versa; there are no individual “NaCl molecules”.
Key takeaways
- Forming cations requires ionisation energy (always positive); forming most anions releases electron affinity energy, but second electron affinities are positive.
- Making isolated gaseous ions usually costs energy overall.
- Lattice energy, released when oppositely charged ions attract in a crystal, pays the bill; for NaCl the overall formation is −411 kJ/mol.
- Higher charges and smaller ions give larger lattice energies, which is why ions such as Mg²⁺ and O²⁻ form.
- In solution, hydration of ions by water plays the stabilising role that the lattice plays in solids.
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