Ionic bonding

Chapter 7 takes the compound idea from Chapter 2 and finally explains what "chemically combined" means for the most important family of compounds on the course. Start from the everyday proof: sodium is a dangerously reactive metal stored under oil, chlorine is a poisonous, reactive gas — yet the compound they form, sodium chloride, is safe to eat in small quantities. The properties of a compound are nothing like the properties of its elements, and this lesson explains why: in ionic compounds the atoms have transferred electrons and become charged ions.

All of the objectives of record for this lesson are core specification statements — examined in both Paper 1 (1C) and Paper 2 (2C), and shared with Science Double Award:

• how ions are formed by electron loss or gain (spec 1.37);
• the charges on ions of Group 1/2/3 metals, Group 5/6/7 non-metals, and the named ions Ag⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Zn²⁺, H⁺, OH⁻, NH₄⁺, CO₃²⁻, NO₃⁻, SO₄²⁻ (spec 1.38);
• writing formulae for compounds between those ions (spec 1.39);
• dot-and-cross diagrams for electron transfer, outer electrons only (spec 1.40);
• ionic bonding as electrostatic attraction (spec 1.41);
• why giant ionic lattices have high melting and boiling points (spec 1.42);
• conductivity: no conduction when solid, conduction when molten or in aqueous solution (spec 1.43).

The lesson runs in three movements: forming ions and showing the transfer with dot-and-cross diagrams, then the formula shortcut — equal pluses and minuses, with the learn-list of ion charges — then the giant-lattice structure and the properties that follow from it. Master all three and every ionic-bonding question on the course becomes routine.

Sodium chloride is the flagship example: an ionic compound formed when a metal combines with a non-metal. You can usually recognise ionic compounds because they contain a metal — magnesium oxide, calcium fluoride and zinc bromide all follow the metal-plus-non-metal pattern. Watch for the ammonium exceptions flagged below, so the rule never gets over-learned.

Here is the formation story for NaCl, one step at a time. A sodium atom has electronic configuration [2, 8, 1] — one outer electron. A chlorine atom has [2, 8, 7] — room for exactly one more. Chlorine attracts electrons more strongly than sodium does, so one electron is transferred from the sodium's outer shell to the chlorine's outer shell. Losing an electron leaves sodium with one more proton than electron: charge 1+. Gaining one leaves chlorine with one more electron than proton: charge 1−.

The charged particles formed are called ions — charged particles made when atoms (or groups of atoms) lose or gain electrons. A positive ion is a cation (for example Na⁺); a negative ion is an anion (for example Cl⁻). When an ionic compound forms, electrons are transferred from a metal atom to a non-metal atom, and the compound is held together by strong electrostatic-attraction — the attraction between the positive and negative ions just formed (spec 1.41's whole idea in one line).

Magnesium oxide (MgO) is the second worked example, and it adds two new ideas at once: a 2+/2− ion pair, and the charge labels that make a dot-and-cross-diagram complete. Magnesium ([2, 8, 2]) has two outer electrons; oxygen ([2, 6]) has room for exactly two. Two electrons are transferred from a magnesium atom to an oxygen atom, forming Mg²⁺ and O²⁻.

Draw the diagram yourself: magnesium's two outer electrons become crosses inside oxygen's outer shell (or dots-and-crosses in any colour convention you like — a mark scheme cares about the electron count and the charges, not the symbol you used). Both ions go inside big square brackets with their charges outside: [2, 8]²⁺ and [2, 8]²⁻. The single most common lost mark on these diagrams is a correct electron layout with no charges on the ions — a complete answer always labels Mg²⁺ and O²⁻.

Notice what the transfer did not do: it did not make a molecule. One Mg²⁺ paired with one O²⁻ is the formula unit the diagram shows — the actual compound (section 12) is a giant lattice of very many Mg²⁺ and O²⁻ ions. The diagram is the bookkeeping for one transfer; the compound is the sum of enormous numbers of those transfers' results.

Look back at the ions formed so far — Na⁺, Cl⁻, Mg²⁺, O²⁻ — and a pattern emerges: every one of them has a noble gas electronic configuration, [2, 8] (neon) or [2, 8, 8] (argon). That is the deep "why" behind ion formation for the first 20 elements: atoms lose or gain electrons so that they achieve a noble gas electronic configuration — the already-stable arrangement of Group 0.

The group law follows. Elements in Groups 1, 2 and 3 lose their outer shell electrons to form 1+, 2+ and 3+ ions; elements in Groups 5, 6 and 7 gain electrons to form 3−, 2− and 1− ions. Two worked corners: aluminium ([2, 8, 3]) loses three electrons → Al³⁺; nitrogen ([2, 5]) gains three → the nitride ion N³⁻. And elements in Groups 1, 2, 6 and 7 always form ions isoelectronic with the nearest noble gas: rubidium (Group 1) loses its one outer electron to become isoelectronic with krypton; iodine (Group 7) gains one electron to become isoelectronic with xenon.

Then the caveat that keeps 1.38 honest: a lot of common ions do not have noble gas structures. Fe²⁺, Fe³⁺, Cu²⁺, Zn²⁺, Ag⁺ and Pb²⁺ are all ions you will meet across the International GCSE course, and not one of them has a noble gas structure (the Fe²⁺ ion is, as the classic caption puts it, definitely not a noble gas structure!). You are not required to write their electronic configurations — only their charges, learned as data (section 6's learn-list).

Two more examples pair a 1+/1− case with a 2+/1− case, so you see the same transfer logic producing different formulas.

Lithium fluoride, LiF. The lithium atom has 1 outer electron that is easily lost; fluorine has space to receive exactly one. One electron is transferred from lithium to fluorine, giving Li⁺ and F⁻. Lithium fluoride is held together by the strong electrostatic-attraction between positive Li⁺ ions and negative F⁻ ions. One of each — the simplest possible ratio.

Calcium chloride, CaCl₂. The calcium atom [2, 8, 8, 2] has two outer electrons, but each chlorine atom [2, 8, 7] has room to take only one. You therefore need two chlorines for every calcium: the two electrons are transferred from one calcium atom to two separate chlorine atoms, one each. The formula is CaCl₂ — and because the Ca²⁺ ion is doubly charged, the electrostatic attractions holding the ions together are very strong.

Note also what the diagrams do not show: the CaCl₂ diagram draws one calcium and two chlorines because that is the ratio, not because a CaCl₂ unit floats around as a thing. Like every ionic compound, calcium chloride is a giant lattice (section 12) — the dot-and-cross diagram is the accounting for one calcium's two electrons.

There are far too many ionic compounds on the course to learn formulae one by one — you need a simple, shortcut method. The method rests on one non-negotiable law.

The law: equal numbers of pluses and minuses. Ions are formed when atoms or groups of atoms lose or gain electrons, and every compound is electrically neutral — so an ionic compound must contain the right number of each ion for the total positive charge to cancel the total negative charge exactly. To apply it you need the charges, and there are two ways to get them.

Where you can work the charge out. The group law from section 4, in one table:

Where the name tells you the charge. For metals that form more than one ion, the name carries the charge in Roman numerals: lead(II) oxide contains Pb²⁺; iron(III) chloride contains Fe³⁺; copper(II) sulfate contains Cu²⁺. The number in brackets is the number of charges.

The learn-list. Some charges cannot be worked out at all — memorise this table (both the formula and the charge of every ion on it):

Two name-endings decide what a compound contains, and mixing them up is one of the most common formula errors on the course.

-ide: nothing complicated. An -ide name means the compound contains exactly the elements named and nothing else. Sodium chloride is just sodium and chlorine combined; copper(II) sulfide contains copper and sulfur only — Cu²⁺ and S²⁻ ions, formula CuS.

-ate: oxygen is there too. An -ate name means oxygen (and possibly other things) is present as well. Copper(II) sulfate contains copper, sulfur and oxygen — Cu²⁺ and SO₄²⁻, formula CuSO₄.

The method now runs on worked examples, and the first one is deliberately the smallest case where a plain "1:1 cancels" reflex fails.

Sodium oxide. Sodium is in Group 1, so the ion is Na⁺. Oxygen is in Group 6, so the ion is O²⁻. To equalise the charges you need two sodium ions (2 × 1+ = 2+) for one oxide ion (2−): the 2+ from two sodiums cancels the 2− of one oxide. The formula is therefore Na₂O.

The worked example is short because the method is short: name the ions, count the charges, scale to balance, write the formula with the counts as subscripts. Forty seconds in an exam, one mark banked.

The second worked example adds the one piece of syntax the method needs: brackets around a multi-atom ion when more than one of it appears.

Barium nitrate. Barium is in Group 2, so the ion is Ba²⁺. Nitrate ions are NO₃⁻ — a complex-ion (an ion made of more than one atom) whose charge you remember, not derive (the learn-list, section 6). To equalise charges you need two nitrate ions for each barium ion (2+ against 2 × 1−). The formula is Ba(NO₃)₂.

The brackets are load-bearing, not decoration: they say "everything inside is doubled". Written without them the formula reads BaNO₃₂ — which would read as 1 barium, 1 nitrogen and 32 oxygens! With brackets: 1 barium, 2 nitrogens, 6 oxygens.

The third worked example is the hard case the method was built for: a 3+ metal against a 2− complex ion.

Iron(III) sulfate. Iron(III) tells you the metal ion is Fe³⁺ (the name carries the charge — section 6's Roman-numeral rule). Sulfate ions are SO₄²⁻. Balancing: you need two iron(III) ions for every three sulfate ions — 2 × 3+ = 6+ and 3 × 2− = 6−, equal and opposite. The formula is Fe₂(SO₄)₃.

Both syntax laws fire at once: the subscript 2 on iron (multiplied because 2 ions are needed) and the brackets with subscript 3 on sulfate (a complex ion multiplied). Total audit: 2 Fe, 4 × 3 = 12 O, charges 6+ = 6−.

The formulae block closes by running the crossover method back onto calcium chloride — and by showing where the shortcut must NOT be trusted.

CaCl₂ by crossover. Ca²⁺ and Cl⁻: cross the 2 to the chlorine and… nothing to the calcium, because there is no extra number in front of the charge on Cl⁻ — we do not tend to write in a 1. The charge "1−" carries an unwritten 1, so the crossover gives Ca₁Cl₂, written simply CaCl₂.

The trap: same charges. Try the crossover on calcium oxide and it hands you Ca₂O₂ — 2+ crossed to oxygen, 2− crossed to calcium. The correction: the formula of calcium oxide is CaO and not Ca₂O₂. When the charges on the two ions are the same, there is already 1 of each in the neutral ratio — nothing needs swapping, and the formula is the plain pair: CaO, MgO, NaCl, LiF. (Ca₂O₂ isn't wrong as an arithmetic fact — 2+ per 2− — it is simply never written: formulae use the lowest whole-number ratio.)

Everything so far described the bookkeeping of electron transfer. This section zooms out to the actual substance: what a crystal of sodium chloride physically is.

All ionic compounds form crystals consisting of lattices of positive and negative ions packed together in a regular way. A lattice is a regular array of particles — and the everyday picture is exactly what the word suggests: a lattice fence, a repeating pattern that extends as far as you care to look. The ionic lattice is held together by the strong electrostatic-attraction between the positively and negatively charged ions, and in a giant-ionic-lattice that attraction runs throughout the whole structure, in all directions.

Take one crystal and dissect it. In diagrams the ions of sodium chloride are usually drawn in an "exploded" view: each sodium ion is surrounded by 6 chloride ions, and each chloride ion is surrounded by 6 sodium ions — and that 6-around-6 pattern repeats throughout the structure over vast numbers of ions. The lines in the exploded diagram are not bonds; they only show which ions touch.

This is the giant-ionic-lattice in its full sense: giant is not "big" — it means a structure with no individual molecules. All the sodium ions attract all the chloride ions; you cannot pick out a sodium chloride molecule, because none exists. The bonding extends throughout the structure in every direction, with no limit to the particle count — all we know is that the numbers of sodium and chloride ions must be equal.

Magnesium oxide has exactly the same structure — Mg²⁺ in place of Na⁺, O²⁻ in place of Cl⁻, the same 6-around-6 lattice. The difference is force: the magnesium oxide lattice is held together by stronger electrostatic-attraction, because 2+ ions are attracting 2− ions, where sodium chloride's attractions run only between 1+ and 1− ions. (Going further: Mg²⁺ is also smaller than Na⁺ and O²⁻ smaller than Cl⁻, which raises the attraction further — but the charge effect is the more important one.)

Four properties, one cause: the giant lattice.

High melting and boiling points. Ionic compounds have high melting points and boiling points because of the strong electrostatic forces of attraction holding the lattice together: a lot of energy must be supplied to break the attractions between the many oppositely charged ions in the giant structure. (Spec 1.42's whole demand.)

Crystalline. Ionic compounds tend to be crystalline — the visible echo of the regular lattice. Some crystals are too small to see: magnesium oxide is always a white powder because its individual crystals are microscopic.

Brittle. Ionic crystals tend to be brittle: any small distortion of the crystal brings ions of the same charge alongside each other — like charges repel — and the crystal splits itself apart along that plane.

Solubility pattern. Ionic substances tend to be soluble in water and insoluble in organic solvents (ethanol, hexane and the petrol hydrocarbons are the everyday examples of the organic-solvent family you will meet properly in Unit 4). Going further: water is polar — its shared electrons sit nearer the oxygen, leaving oxygen slightly negative and the hydrogens slightly positive — so water molecules form strong enough attractions to ions to pull the lattice apart. Hexane is non-polar and cannot. Not every ionic salt obeys: magnesium oxide is insoluble in water — its 2+/2− attractions are too strong for the water molecules to overcome.

The last spec point is a two-state rule with one forbidden word.

When solid: ionic compounds do not conduct electricity — the ions are fixed in position in the lattice and are not free to move around. When molten (melted) or in aqueous-solution (dissolved in water): they do conduct — because the ions then become free to move and can carry the charge through the liquid.

One language rule carries the whole answer: the charge carriers are ions. Do not say "electrons" — the examiners require the exact words: the ions are free to move. That single substitution is the difference between the mark and no mark.