Ions, Cations and Anions

The naming is the first obstacle and it is pure accident. Cation looks as though it should be the ion that has taken something on; it is in fact the one that has lost electrons and carries a positive charge. Nothing in the word signals that.

The reason is historical. Michael Faraday, advised by the polymath William Whewell, built the vocabulary of electrochemistry in 1834 around the direction ions travel in an electric field. A cation moves toward the cathode; an anion moves toward the anode. The Greek roots mean "down" and "up", and they describe migration, not charge. Once you know that, two mnemonics stop being arbitrary: the t in cation can be read as a plus sign, and anion contains the phrase "a negative ion".

An ion is not a broken atom

The second obstacle is a picture rather than a word. Students often read ionisation as damage — the atom has lost a piece and is now unstable, waiting to be repaired.

The opposite is true. A sodium ion is a more comfortable arrangement than a sodium atom in almost every chemical environment it will ever meet. Sodium's outermost electron is loosely held and shielded from the nucleus by the ten electrons beneath it; removing it costs relatively little energy and leaves behind a compact, closed-shell structure. Chloride is similarly settled. Table salt is not two damaged atoms clinging together, it is two species that have each reached a lower-energy arrangement and are held together electrostatically.

The energetic accounting behind that — why the transfer pays for itself — belongs to ionic bonding. What matters here is the mental picture: ions are stable, ordinary, and vastly more common in the world than neutral atoms. Most of the mass of the ocean's dissolved solids, most of the material of a bone, and the signalling in every nerve you own, are ions.

Reading the charge off the group number

For main-group elements the charge is predictable from position, because it is whatever gets the atom to the nearest noble-gas electron count.

• Group 1 loses one electron: Li⁺, Na⁺, K⁺.
• Group 2 loses two: Mg²⁺, Ca²⁺, Ba²⁺.
• Group 13 loses three: Al³⁺.
• Group 15 gains three: N³⁻, P³⁻.
• Group 16 gains two: O²⁻, S²⁻.
• Group 17 gains one: F⁻, Cl⁻, Br⁻, I⁻.
• Group 18 does neither, having arrived already.

Monatomic anions are renamed with an -ide ending — chlorine becomes chloride, oxygen becomes oxide, nitrogen becomes nitride. Cations simply keep the element's name, so Na⁺ is "sodium ion" with nothing added.

Group 14 is the gap in the pattern, and textbooks paper over it. Carbon and silicon are not going to lose four electrons or gain four; the energy required is prohibitive either way, and a bare C⁴⁻ or C⁴⁺ has no ordinary chemical existence. Group 14 elements share electrons instead, which is why carbon chemistry is covalent from top to bottom. Reading "+4 or −4" off the group number for carbon will produce formulas that do not exist.

The heavy elements that keep two electrons back

Below period 4, group 13 and 14 metals stop following the pattern in a systematic way that is worth learning as a class rather than as a list of oddities.

Thallium sits under aluminum, so the naive prediction is Tl³⁺. In practice Tl⁺ is the dominant form of thallium in its compounds. Lead, in group 14, is overwhelmingly Pb²⁺ rather than Pb⁴⁺; bismuth favours Bi³⁺ over Bi⁵⁺; tin forms both Sn²⁺ and Sn⁴⁺ with the lower state increasingly favoured on descending the group.

The pattern is the inert pair effect: the outermost s² pair in these heavy atoms is held far more tightly than its position suggests, partly because the 6s orbital is contracted and stabilised by relativistic effects at high nuclear charge. The two s electrons stay put, and the element ionises to a charge two lower than its group implies.

Transition metals need their own treatment, because most of them form several ions. Iron gives Fe²⁺ and Fe³⁺, copper gives Cu⁺ and Cu²⁺, and the compound's name must therefore state which — iron(III) chloride, copper(I) oxide. An older system, still met in medicine and industry, used -ous for the lower charge and -ic for the higher: ferrous and ferric, cuprous and cupric. It is unambiguous only if you already know which two charges exist, which is why Roman numerals replaced it. Three transition metals are reliably fixed and never take a numeral in practice: silver is Ag⁺, zinc is Zn²⁺, and scandium is Sc³⁺.

Sodium shrinks, chloride swells

Forming an ion changes an atom's size dramatically, and in opposite directions.

A sodium atom has a radius of about 186 pm. Na⁺ measures about 102 pm — barely half. Losing that single 3s electron does not shave a sliver off the outside; it removes the entire outermost shell, so the ion's radius is set by the shell beneath. Eleven protons then pull on ten electrons instead of eleven, tightening what remains.

A chlorine atom's covalent radius is about 99 pm; the chloride ion measures about 181 pm. The added electron enters the shell already occupied, so the nuclear charge is unchanged while the mutual repulsion between electrons increases, and the cloud expands.

The consequence is a reversal that catches people out. Neutral sodium is much larger than neutral chlorine. Sodium ion is much smaller than chloride. Any question about relative size must specify whether it means atoms or ions, and the ranking can invert completely between the two.

Ten electrons, seven different sizes

A set of species with the same electron count is called isoelectronic, and it isolates the effect of nuclear charge with everything else held constant.

N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺ and Al³⁺ all have exactly ten electrons in exactly the same arrangement. Their radii, in picometres, run: N³⁻ about 146, O²⁻ 140, F⁻ 133, Na⁺ 102, Mg²⁺ 72, Al³⁺ about 54.

The order is strictly by proton count. Seven protons hold ten electrons loosely; thirteen protons hold the same ten electrons in a cloud less than half the width. The rule for ranking any isoelectronic set is therefore short and reliable: more protons, smaller ion, and it does not matter at all whether the species is a cation, an anion or a noble gas.

The oxygen families: -ite, -ate, hypo- and per-

Most ions in real chemistry are polyatomic — a group of covalently bonded atoms carrying an overall charge, which travels and reacts as a single object. Sulfate is SO₄²⁻, nitrate is NO₃⁻, carbonate is CO₃²⁻, phosphate is PO₄³⁻, hydroxide is OH⁻, and ammonium, NH₄⁺, is the common polyatomic cation.

Their names look arbitrary and are not. Within a family of oxyanions of the same central element, the endings and prefixes count oxygen atoms while the charge stays fixed:

• hypo-…-ite — two fewer oxygens than the -ate form: ClO⁻, hypochlorite
• -ite — one fewer: ClO₂⁻, chlorite
• -ate — the reference form: ClO₃⁻, chlorate
• per-…-ate — one more: ClO₄⁻, perchlorate

All four carry a 1− charge. The same scheme runs through sulfite and sulfate, nitrite and nitrate, phosphite and phosphate. Learning the -ate member of each family and the two rules gets you the rest, which is a much smaller memory load than the twenty separate items it looks like.

What none of these names tells you is the charge itself — you have to know that sulfate is 2− and nitrate is 1−. Those, unavoidably, are memorised. Combining them with cations into neutral formulas is the next step, and it is set out with the lattices they build on the ionic bonding page.