Ions are atoms with unequal numbers of protons and electrons

An ion is a charged particle formed when an atom or group of atoms loses or gains electrons. In ordinary chemical ion formation, protons are not transferred, so the nuclear charge stays the same. The electron number changes.

A cation has fewer electrons than protons, giving it an overall positive charge. Metal atoms often form cations by losing electrons. A sodium atom, for example, has 11 protons and 11 electrons, while Na⁺ has 11 protons and 10 electrons, so its charge is 1+.

An anion has more electrons than protons, giving it an overall negative charge. Non-metal atoms often form anions by gaining electrons. A sulfur atom can gain two electrons to form S²⁻, the sulfide ion.

A monatomic ion is an ion made from a single atom. The names of monatomic anions usually end in -ide: chloride, oxide, sulfide, nitride and so on. Cation names usually stay the same as the element name. If a metal can form more than one charge, Roman numerals are used in compound names, such as iron(II) and iron(III).

Predicting charge from electron configuration

An electron configuration is a notation that shows how electrons occupy shells, subshells and orbitals in an atom or ion. A useful rule at this level is that main-group atoms tend to form ions by reaching a noble gas electron configuration, as long as they can do it by losing or gaining a small number of electrons.

For cations, remove the outer-shell electrons. Sodium is 1s²2s²2p⁶3s¹, so losing the 3s electron gives Na⁺ with 1s²2s²2p⁶. Calcium is 1s²2s²2p⁶3s²3p⁶4s², so losing two 4s electrons gives Ca²⁺ with 1s²2s²2p⁶3s²3p⁶.

For anions, add electrons to complete the outer shell. Chlorine is 1s²2s²2p⁶3s²3p⁵, so gaining one electron gives Cl⁻ with 1s²2s²2p⁶3s²3p⁶. Oxygen is 1s²2s²2p⁴, so gaining two electrons gives O²⁻ with 1s²2s²2p⁶.

Isoelectronic species are atoms or ions with the same electron configuration. Na⁺ is isoelectronic with Ne; Cl⁻ and Ca²⁺ are isoelectronic with Ar. Don’t call them “the same particle”: they have different proton numbers and different charges.

The periodic table gives a quick way to predict many main-group ions:

Group 14 is awkward if you try to force the rule. Carbon would have to lose or gain four electrons to make C⁴⁺ or C⁴⁻, which is not usually favourable. Carbon more commonly shares electrons in covalent compounds.

The octet rule is a model stating that many atoms form bonds so that their valence shell contains eight electrons. It works well for many main-group ions, but it is a model, not a law. Hydrogen, for example, can lose one electron to form H⁺ or gain one electron to form H⁻, reaching the helium configuration rather than an octet.

Ion formation and redox language

Oxidation is a chemical process in which a species loses electrons. Cation formation is oxidation: Na → Na⁺ + e⁻.

Reduction is a chemical process in which a species gains electrons. Anion formation is reduction: Cl + e⁻ → Cl⁻.

That is why forming an ionic compound from its elements is a redox reaction. During the formation of sodium chloride, sodium atoms lose electrons and chlorine molecules gain electrons:

2Na(s) → 2Na⁺ + 2e⁻

Cl₂(g) + 2e⁻ → 2Cl⁻

Oxidation-state language comes later in the course, but it fits the same idea: sodium goes from 0 to +1, while chlorine goes from 0 to −1.

Transition elements can form ions with different charges

A transition element is an element that has an atom or common ion with a partially filled d subshell. Unlike most main-group metals, transition elements often form more than one cation. Iron, for example, commonly forms Fe²⁺ and Fe³⁺.

When first-row transition elements ionize, they lose the 4s electrons before the 3d electrons. So Fe, often written [Ar]4s²3d⁶, forms Fe²⁺ as [Ar]3d⁶. More electrons can then be removed because the 4s and 3d subshells are close in energy, which helps explain variable oxidation states.

Successive ionization energy data support this pattern: for transition elements, several electrons can often be removed with comparatively gradual increases before a much larger jump appears when an inner-shell electron would have to be removed.

This links back to periodic position. Main-group ion charges are closely tied to group number because atoms gain or lose electrons to reach a noble gas configuration. Transition elements are harder to predict from group number alone, because the ns and (n−1)d electrons are similar enough in energy for several cation charges to be possible.

What an ionic bond is

An ionic bond is the electrostatic attraction between oppositely charged ions in a compound. Don’t picture it as a tiny stick joining one cation to one anion. The attraction acts through the whole structure, which matters a lot when lattices come up.

A chemical bond is a strong attractive interaction that holds atoms or ions together in a substance. In the ionic model, the cation is the positively charged species, and the anion is the negatively charged species.

Ionic character and electronegativity

Electronegativity, χ, is a dimensionless measure of how strongly an atom in a covalent bond attracts the shared pair of electrons, where χ is the electronegativity value on the Pauling scale and has no unit. It generally increases across a period and up a group; fluorine has the highest value.

The electronegativity difference, Δχ = |χ₁ − χ₂|, where Δχ is the difference in electronegativity between two bonded atoms (no unit), χ₁ is the electronegativity of one atom (no unit) and χ₂ is the electronegativity of the other atom (no unit), is often used to estimate ionic character.

A larger Δχ gives a bond more ionic character. As a working guideline, bonding is mainly ionic when Δχ is greater than about 1.8. Treat that as a model, not a rule that never fails. Bonding lies on a continuum, so “metal + non-metal” is a good first guess, but it does not guarantee ionic behaviour. Aluminium chloride, AlCl₃, is the standard warning example because its properties are more covalent than you might expect from the element types alone.

Naming binary ionic compounds

A binary ionic compound is an ionic compound containing ions of two elements only. Write the cation first, then the anion. For a monatomic anion, take the end of the element name and add -ide.

Examples:

The name does not directly give the number of ions. “Magnesium fluoride” is MgF₂, not magnesium difluoride, because ionic compound names do not use covalent prefixes such as mono-, di- and tri-.

For transition metals with variable charges, the Roman numeral shows the oxidation state of the metal ion. In iron(III) oxide, iron is present as Fe³⁺, so the formula is Fe₂O₃. In copper(II) chloride, copper is Cu²⁺, so the formula is CuCl₂. The symbol charge is written as 2+ or 3+ after the number; the oxidation state is written as +2 or +3 with the sign first. Keep those formats separate.

Polyatomic ions you must know

A polyatomic ion is an ion composed of two or more covalently bonded atoms with an overall charge. When writing ionic formulas, treat a polyatomic ion as one charged unit: you may need brackets, but you must not change the atoms inside the ion.

Required polyatomic ions, with formulas and overall charges.

Polyatomic anions such as nitrate, carbonate, sulfate and phosphate are conjugate bases of acids. Their stability is connected to the acid dissociation constant, Kₐ, where Kₐ is the acid dissociation constant (mol dm⁻³ under the usual IB concentration convention). In general, a more stable conjugate base corresponds to a larger Kₐ for its conjugate acid, because acid dissociation is more favourable.

For sulfate, later work on formal charge helps explain which Lewis structure is preferred. Formal charge is a bookkeeping method for assigning electrons in a Lewis structure; the preferred sulfate representation gives the most reasonable distribution of formal charges while respecting the bonding model being used.

Deducing formulas from ions

An ionic compound is electrically neutral overall. So the total positive charge must equal the total negative charge. That’s the key idea.

A reliable method is:

1. Write the correct cation and anion, with charges.
2. Choose the smallest whole-number ratio that gives zero net charge.
3. Use brackets if more than one polyatomic ion is needed.
4. Check the charges before you write the final formula.

For magnesium chloride, Mg²⁺ needs two Cl⁻ ions to balance the charge, so the formula is MgCl₂.

For aluminium oxide, Al³⁺ and O²⁻ balance in a 2:3 ratio, because 2 × 3+ gives 6+ and 3 × 2− gives 6−. The formula is Al₂O₃.

For barium hydroxide, Ba²⁺ needs two OH⁻ ions. Since hydroxide is polyatomic and there is more than one of it, the formula is Ba(OH)₂. Do not write BaO₂H₂; that destroys the ion identity.

For ammonium phosphate, NH₄⁺ and PO₄³⁻ balance as three ammonium ions for one phosphate ion, giving (NH₄)₃PO₄.

Interconverting names and formulas

To go from formula to name, identify the ions, then name the cation followed by the anion:

To go from name to formula, don’t guess subscripts from the words. Use the ion charges:

Formation of an ionic compound from its elements is a redox reaction because a metal atom is oxidized to a cation while a non-metal atom is reduced to an anion. The ionic formula you write is the charge-balanced result of those electron-transfer processes.