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Chemistry

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Ionic bonding

Learning objectives · Understand how ions are formed by electron loss or gain; Know the charges of these ions: metals in Groups 1, 2 and 3; non-metals in Groups 5, 6 and 7; Ag⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Zn²⁺; hydrogen (H⁺), hydroxide (OH⁻), ammonium (NH₄⁺), carbonate (CO₃²⁻), nitrate (NO₃⁻), sulfate (SO₄²⁻); Write formulae for compounds formed between the ions listed above; Draw dot-and-cross diagrams to show the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3, 5, 6 and 7 (only outer electrons need be shown); Understand ionic bonding in terms of electrostatic attractions; Understand why compounds with giant ionic lattices have high melting and boiling points; Know that ionic compounds do not conduct electricity when solid, but do conduct electricity when molten and in aqueous solution

Ionic bonding — the chapter map

SB p.75

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.

Forming ions — sodium chloride

SB p.75

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−.

Key point

Two anchors for the mechanism. (1) The formation reaction is a simplification — in reality you don't react sodium atoms with chlorine atoms but with chlorine molecules, Cl₂; the atom-level story is the model, not the laboratory recipe. (2) You can check the 1+ charge from the Periodic Table: sodium has atomic number 11, so 11 protons give an 11+ nucleus; the atom's 11 electrons cancel it to neutral, but the ion has only 10 electrons — 11+ against 10− leaves an overall 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).

Key point

Two definitions of record. Electrostatic attraction simply means that positively charged particles attract negatively charged particles. And ionic bonding is often shown using dot-and-cross diagrams, where the transferred electrons are drawn as dots for one atom and crosses for the other. There is absolutely no difference between the electrons in reality — dots and crosses only show which atom each electron came from (two colours of dots, or two colours of crosses, would serve equally well).

Exam hint

One classic trap lives here: there are one or two exceptions — ionic compounds that do not contain a metal, for example those containing the ammonium ion (such as NH₄Cl, (NH₄)₂SO₄). So "contains a metal" is a good recognition rule but never a definition. In "state whether the bonding in X is ionic" questions, ammonium salts are ionic despite having no metal — the examiners have used this exact trap.

Common misconception

"Chlorine turns negative because it gains a proton." Protons never move in chemistry — they are the fixed identity of the nucleus. Gain or loss of electrons is the only charge-changing move: Cl has 17 protons and 17 electrons; Cl⁻ has 17 protons and 18 electrons. The same discipline explains the sodium side: Na⁺ is still sodium (11 protons) — it has simply lost one of its 11 electrons.

Quick check

Question-bank join: spec 1.37 items — describing ion formation by electron loss or gain for a named metal/non-metal pair, and explaining the charge on a named ion from its proton and electron counts. Spec 1.41 items — stating that ionic bonding is electrostatic attraction between oppositely charged ions. Self-test: draw the dot-and-cross diagram for NaCl from the two electronic configurations, then label both ion charges.

Key terms in this section

ion: — A charged particle formed when an atom, or a group of atoms, loses or gains electrons. Ions can carry a positive or a negative charge.
cation: — A positive ion, formed when an atom loses electrons — for example Na⁺. All metal ions are cations.
anion: — A negative ion, formed when an atom gains electrons — for example Cl⁻.
electrostatic attraction: — The attraction between positively and negatively charged particles — the force that holds oppositely charged ions together throughout an ionic lattice.
dot-and-cross diagram: — A diagram of ionic bonding showing outer-shell electrons only, with one atom's electrons as dots and the other's as crosses, and the charges on the ions shown.

Keep goingFormation of Ions Ionic compounds: Bonds, Structure & Properties 1-3-atomic-structure 1-4-the-periodic-table 1-6-ionic-bonding

Ionic bonding in magnesium oxide

SB p.76

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⁻.

Key point

A word worth knowing: isoelectronic means have the same number of electrons. When sodium forms its ion, electronic configuration [2, 8], it becomes isoelectronic with neon, [2, 8]. The same idea runs through this whole lesson: Mg²⁺ is isoelectronic with neon, and O²⁻ is isoelectronic with neon too — three different particles, one electron count.

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.

Exam hint

Dot-and-cross mark schemes for 1.40 are nearly always the same three lines: outer shells only (spec words: only outer electrons need be shown); correct electron counts in each ion's outer shell; charges on both ions. For MgO: magnesium bracket [2, 8]²⁺ with empty outer shell, oxygen bracket [2, 8]²⁻ with eight outer electrons — two of them crosses. Draw big, clear square brackets; tiny cramped brackets are where marks go missing.

Common misconception

"Mg²⁺ means the magnesium atom gained two electrons, and O²⁻ lost two." Read the charges against the process: losing electrons leaves more protons than electrons, so loss makes positive (Mg lost 2 → 2+); gaining electrons makes the balance negative (O gained 2 → 2−). The superscript records the balance after the transfer — never the direction of the move itself.

Quick check

Question-bank join: spec 1.40 items — drawing the MgO dot-and-cross diagram (outer electrons only, charges shown); spec 1.41 items — identifying the electrostatic attraction between Mg²⁺ and O²⁻ as the ionic bond. Self-check before moving on: can you say why the charges are 2+ and 2− from the group numbers alone (Group 2 → 2+; Group 6 → 2−)?

Key terms in this section

dot-and-cross diagram: — A diagram of ionic bonding showing outer-shell electrons only, with one atom's electrons as dots and the other's as crosses, and the charges on the ions shown.
isoelectronic: — Having the same number of electrons: the sodium ion [2, 8] is isoelectronic with the neon atom [2, 8].

Keep goingIonic Bonds: Dot & Cross Diagrams Ionic compounds: Bonds, Structure & Properties 1-6-ionic-bonding 1-8-metallic-bonding 2-1-group-1-alkali-metals

Noble gas electronic configurations — and the ions that break the pattern

SB p.77

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.

Key point

Spec 1.38's charge table in one view: Groups 1/2/3 lose 1/2/3 electrons → 1+/2+/3+; Groups 5/6/7 gain 3/2/1 → 3−/2−/1−. The group number is the outer-electron count; the charge is the cheapest route to a full outer shell.

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).

Exam hint

"Noble gas configuration" is a description, not a cause you must argue at this level: for the first 20 elements the stable arrangement explains the group law, and that is the full depth the spec wants. The exam trap is over-applying it — a question that offers Fe³⁺ or Zn²⁺ is testing whether you know the named-ion charges (1.38) rather than the group rule. Learn both lists and never mix them: group charges for Groups 1/2/3/5/6/7, learned charges for the six named metals plus H⁺ and the four complex ions.

Common misconception

"Every ion must be isoelectronic with a noble gas — so Fe³⁺ can't exist." It exists, it's common, and it is not isoelectronic with anything noble — that is exactly why it gets flagged. The noble-gas rule is a trend for the first 20 elements, not a law of the universe. Transition-metal and heavy-metal ions (Fe, Cu, Zn, Ag, Pb) have their own charges that are learned, not derived.

Quick check

Question-bank join: spec 1.37 items — predicting whether an element loses or gains electrons, and how many, from its position; spec 1.38 items — stating the charges of Group 1/2/3 metals and Group 5/6/7 non-metals, and recalling the named-ion list. Self-test: cover this section and derive each result — Rb (Group 1, [2, 8, 18, 8, 1]) → which noble gas? I (Group 7) → which?

Key terms in this section

isoelectronic: — Having the same number of electrons: the sodium ion [2, 8] is isoelectronic with the neon atom [2, 8].

Keep goingFormation of Ions Common Ions 1-1-states-of-matter 1-3-atomic-structure 1-4-the-periodic-table

Other examples of ionic bonding — lithium fluoride and calcium chloride

SB p.78

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.

Key point

The general law hiding in these two examples: the ratio in the formula is set by how many electrons must move and how many each partner can accept or lose. Li gives 1, F takes 1 → LiF. Ca gives 2, Cl takes 1 each → two Cl per Ca → CaCl₂. Section 6 turns this counting law into the general formula shortcut.

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.

Exam hint

For a "draw the dot-and-cross diagram for calcium chloride" question (1.40), the mark scheme wants: calcium with empty outer shell labelled 2+, two chlorine outer shells each labelled 1− and each showing eight outer electrons (one of them a cross from calcium). The classic error is drawing one chlorine receiving two electrons — a chlorine outer shell holds only eight, and the 1− charge says it took one. Two transfer arrows to two separate ions is the correct picture.

Common misconception

"Calcium shares its two electrons, one with each chlorine." Sharing is covalent bonding (Chapter 8); here the electrons transfer and stay — each chlorine walks away with both of its new electrons permanently, and the calcium is left 2+. The tell is in the products: charged ions (Ca²⁺, 2 Cl⁻), which is why the compound conducts when molten and needs the electrostatic-lattice description of the sections that follow.

Quick check

Question-bank join: spec 1.40 items — dot-and-cross diagrams for LiF and CaCl₂ (Groups 1/2 × Group 7, outer electrons only); spec 1.41 items — identifying the electrostatic attraction between the named ions as the bond. If you can draw CaCl₂ correctly, every Group-2 × Group-7 combination in past papers is the same skill with different symbols.

Key terms in this section

electrostatic attraction: — The attraction between positively and negatively charged particles — the force that holds oppositely charged ions together throughout an ionic lattice.

Keep goingIonic Bonds: Dot & Cross Diagrams Ionic compounds: Bonds, Structure & Properties 1-6-ionic-bonding 1-8-metallic-bonding 2-1-group-1-alkali-metals

Formulae for ionic compounds

SB p.78

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:

Group in the Periodic Table	Charge on ion	Example
1	1+	Na⁺
2	2+	Mg²⁺
3	3+	Al³⁺
5	3−	N³⁻
6	2−	O²⁻
7	1−	Br⁻

Key point

Two reading rules for that table. Any element in Group 2 has 2 outer electrons and loses them all → 2+; any element in Group 6 has 6 outer electrons and room for 2 more → 2−; similar arguments fix the other groups shown. And the reminder that stands behind every row: all metals form positive ions.

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.

Key point

The metal-free exception in its full picture: ammonium chloride (NH₄Cl) is ionic — there is ionic bonding between the NH₄⁺ and Cl⁻ ions — yet inside the NH₄⁺ ion the N–H links are covalent (Chapter 8). An ionic compound can therefore contain covalent bonds in its polyatomic ion; the two descriptions apply at different scales.

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):

Charge	Substance	Ion
positive	zinc	Zn²⁺
	silver	Ag⁺
	hydrogen	H⁺
	ammonium	NH₄⁺
negative	nitrate	NO₃⁻
	hydroxide	OH⁻
	carbonate	CO₃²⁻
	sulfate	SO₄²⁻

Key point

Two closers for the method. (1) These learn-list ions are the difficult ones — learn both the formula and the charge for each; you will meet others later, but this eight is the working set. (2) The brackets law: when a formula needs more than one of a multi-atom complex ion, write brackets — barium nitrate is Ba(NO₃)₂; without the brackets, "BaNO32" would read as 1 barium, 1 nitrogen and 32 oxygens. Anywhere else, brackets are unnecessary.

Exam hint

One habit removes the last excuse: you will always have a copy of the Periodic Table, even in an exam — so the group charges are always available. What the table cannot tell you: lead forms Pb²⁺ (Group 4 elements form few ionic compounds and the situation is more complicated) — Pb²⁺ rides on the named-ion list. Exam routine: group charges from the Periodic Table; named charges from memory; Roman numerals from the name. And one warning that earns its own headline: not looking carefully at word endings is one of the most common mistakes students make when they start to write formulae — the -ide/-ate endings get their own section next.

Common misconception

"Ionic formulae are 100% memorisation." Only the eight learn-list charges are memory; everything else is derivation. Balance the charges and count: a 2+ ion meeting a 1− ion needs two of the 1− ion; a 3+ meeting a 2− needs three and two respectively (6+ balances 6−). Learners who memorise formulae one at a time drown by Unit 2; learners who balance charges write unfamiliar formulae — aluminium sulfate, sodium carbonate — on demand.

Quick check

Question-bank join: spec 1.38 items — charges of the learn-list ions and group charges; spec 1.39 items — formulae from ion pairs (both-sides drilling: Na⁺/O²⁻; Ca²⁺/Cl⁻; NH₄⁺/SO₄²⁻). The two tables are the recall targets; drill until writing Na₂SO₄ and Al₂O₃ takes under ten seconds.

Keep goingCommon Ions Formulae for Ionic Compounds 1-1-states-of-matter 1-5-chemical-formulae-equations-calculations 1-6-ionic-bonding

To find the formula for sodium oxide

SB p.80

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.

Exam hint

Write the ion charges ABOVE the symbols before scaling: Na⁺ over Na, O²⁻ over O. The intermediate "2 Na⁺ per 1 O²⁻" line is what the mark scheme's "balancing charges" credit rewards — and it catches the slide back to "NaO" because the 2− makes you look for two positives, not one. Subscripts belong to the atom that needs multiplying: Na₂O, never NaO₂ (two oxygens would give 4− against 2+).

Common misconception

"Na⁺ and O²⁻ — the + and − cancel 1-for-1, so NaO." The cancellation is between total charge magnitudes, not between ions. One Na⁺ carries one unit of positive charge; one O²⁻ carries TWO units of negative. Neutral happens at two positives per one negative — 2+ against 2−. The charge number (the superscript) is how many units each ion brings; the formula counts how many ions bring it.

Quick check

Question-bank join: spec 1.38 items — recalling Na⁺ (Group 1) and O²⁻ (Group 6); spec 1.39 items — writing Na₂O from the balanced-charge argument. Extension joins from the same skill: K₂O, Li₂O — then the reverse-reading check: why is the peroxide Na₂O₂ not the plain oxide? (Its ion is not O²⁻ — beyond the spec, but a good test that the endings law from section 7 stays in scope.)

Keep goingCommon Ions Formulae for Ionic Compounds 1-1-states-of-matter 1-5-chemical-formulae-equations-calculations 1-6-ionic-bonding

To find the formula for barium nitrate

SB p.80

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.

Key point

One question learners are often shy to ask — why don't formulae show the charges? They can: sodium chloride is sometimes written Na⁺Cl⁻ to make a point about the ions present. But for most purposes the charges are left out: in an ionic compound the charges are there, whether you write them or not. The formula records the neutral ratio; the charge labels live on the ions, and on dot-and-cross diagrams.

Exam hint

Brackets law in one line: brackets appear only when a complex ion is multiplied. Ba(NO₃)₂ yes; Na₂SO₄ yes (sulfate doubled); NaNO₃ no (one nitrate — brackets would be harmless but non-standard); NaCl no. When a mark scheme says "correct use of brackets", it is checking exactly this rule — and the (NO₃)₂ must be followed by the subscript outside the bracket, never inside.

Common misconception

"Ba(NO₃)₂ and BaNO₃₂ are the same formula written differently." They name different substances: the bracketed form is 1 Ba, 2 N, 6 O; the unbracketed run-together reads 1 Ba, 1 N, 32 O — physically nothing. The 2 in a formula always multiplies what it follows: following a bracket it doubles the whole ion; following an element symbol it multiplies only that element. That subtle reading is why the bracket exists.

Quick check

Question-bank join: spec 1.38 items — recalling NO₃⁻ and Ba²⁺; spec 1.39 items — writing bracketed formulae for Group-2 × nitrate/hydroxide/ carbonate pairs (Ba(NO₃)₂, Ca(OH)₂, Na₂CO₃, K₂SO₄). The oxygen-count check (6 oxygens in Ba(NO₃)₂) is a fast self-audit the examiners also use in "explain why the formula is wrong" items.

Key terms in this section

complex ion: — An ion made of more than one atom — for example NO₃⁻, CO₃²⁻, SO₄²⁻, OH⁻ and NH₄⁺. Brackets are written when a formula needs more than one of them.

Keep goingCommon Ions Formulae for Ionic Compounds 1-1-states-of-matter 1-5-chemical-formulae-equations-calculations 1-6-ionic-bonding

To find the formula for iron(III) sulfate

SB p.81

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−.

Key point

The crossover shortcut: to work out complicated formulae, cross over the numbers in the charges — the 3 from Fe³⁺ becomes the sulfate count; the 2 from SO₄²⁻ becomes the iron count. Fe³⁺ + SO₄²⁻ → Fe₂(SO₄)₃ in one move. The crossover is the balanced-charge arithmetic in disguise; the totals check (6+/6−) is always available to prove it.

Maths skills

The crossover is a digit-swap shortcut with a built-in audit. Swap the magnitudes — 3 lands on the sulfate count, 2 on the iron count — then verify by totals: 2 × 3+ = 6+ and 3 × 2− = 6−. The audit is the safety net the swap lacks: applied blindly to Ca²⁺ + O²⁻ it yields Ca₂O₂, and the totals check (4+ = 4−) says "consistent but not lowest" — simplify to CaO. Swap, then always re-add the charges; if the sums are not equal, the formula is wrong, whatever the swap said.

Exam hint

The crossover shortcut is fast but it is presentational — the marks sit in the balancing statement. Write "6+ and 6−" (or "2 × 3+ balances 3 × 2−") before the formula: that sentence earns the marks even if the final formula slips, and it catches the classic slip of crossing over to Fe₃(SO₄)₂ (which gives 9+/4− — instantly unbalanced). In "explain why Fe₂(SO₄)₃ is the correct formula" items, the totals line IS the answer.

Common misconception

"3+ and 2− balance each other directly, so the formula is FeSO₄ with the charges noted." FeSO₄ is iron(II) sulfate — a different salt made of Fe²⁺. Charges balance as totals across the whole formula unit, not ion-by-ion: 3+ against 2− needs the lowest common multiple (6), which takes 2 irons and 3 sulfates. The Roman numeral in the name is the charge of the metal ion — and iron(II) versus iron(III) formulas are the standard exam pair for testing exactly this.

Quick check

Question-bank join: spec 1.38 items — Fe³⁺ from the name iron(III), SO₄²⁻ from the learn-list; spec 1.39 items — writing Fe₂(SO₄)₃ and siblings (Al₂(SO₄)₃, 3+/3− pairs, CuCO₃ vs CuSO₄). Drill the totals check: every formula you write, verify the charge sums to zero before moving on.

Keep goingCommon Ions Formulae for Ionic Compounds 1-1-states-of-matter 1-5-chemical-formulae-equations-calculations 1-6-ionic-bonding

Calcium chloride provides another example

SB p.81

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.)

Exam hint

Crossover protocol for exams: (1) write both charges; (2) if the two charge magnitudes are equal, stop — the formula is one of each (CaO, BaO, NaF); (3) otherwise cross over and simplify nothing (the crossed numbers are already the lowest pair for the exam's compounds); (4) brackets only around multi-atom ions (Fe₂(SO₄)₃, not Fe₂SO₄₃). The Ca₂O₂ trap is a real "spot-the-error" item format — now you know both halves of the answer: the ratio 1:1, and why.

Common misconception

"Ca₂O₂ simplifies to CaO, so writing Ca₂O₂ shows the ions better and should score." Formulae of ionic compounds are written in their lowest whole-number ratio — that is the convention every mark scheme applies, because the compound is a giant lattice with no discrete Ca₂O₂ unit to point at. The ion charges are shown on diagrams and in explanations, not inside the compound's formula.

Quick check

Question-bank join: spec 1.39 items — formulae from ion pairs where charges are equal vs unequal (CaO vs CaCl₂; MgO vs MgCl₂; BaO vs Ba(NO₃)₂). Self-test: write the six formulae for the pairs {Mg²⁺, Ba²⁺} × {O²⁻, Cl⁻, OH⁻} — three need brackets, three don't.

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Giant ionic structures

SB p.81

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.

Exam hint

This next rule is bold-urgent: never talk about molecules of an ionic compound. In the exam it is marked wrong and can cost you all the marks for a question. Say instead: ions, lattice, giant structure, formula unit. The habit matters in every later answer — "NaCl molecules exist in solution", "one molecule of MgO" — all zero-mark phrasings. In electrolysis and structure questions (Units 1 and 2), the words "ions free to move" only make sense once "no molecules" is reflex.

Common misconception

"A crystal of NaCl is a huge molecule made of lots of NaCl units stuck together." There are no NaCl units to stick together — that is precisely what "giant" rules out. The structure is one continuous array of Na⁺ and Cl⁻ ions with no natural boundary or repeat unit you could lift out and call a molecule; any fragment you pick is arbitrary. The chemical formula NaCl records the 1:1 ratio, not a particle that exists on its own.

Cross-curricular links

The lattice is mathematics made solid: a regular repeating array is exactly what a tessellation or wallpaper pattern is in geometry — one unit cell (here, the 6-around-6 sodium-chloride arrangement you will meet in the next section) translated again and again to fill space with no gaps. Crystallography, the science built on that idea, won Nobel prizes by applying X-ray diffraction to precisely the lattices this lesson describes — the same repeating-array mathematics also underlies tiling problems and the way metals are forged. When chemistry reaches for a lattice fence to explain the word, it is pointing at one of chemistry's deepest exports to mathematics.

Quick check

Question-bank join: spec 1.42 items — describing the giant ionic lattice (regular arrangement of oppositely charged ions, held by electrostatic attraction) and explaining why the word "molecule" never applies. This description is the setup for the melting-point argument in section 14 — the two are nearly always examined together.

Key terms in this section

lattice: — A regular, repeating array of particles; an ionic lattice packs positive and negative ions together in a repeating three-dimensional pattern.
giant ionic lattice: — A structure in which vast numbers of oppositely charged ions are held by electrostatic attraction extending in all directions; there are no individual molecules.
electrostatic attraction: — The attraction between positively and negatively charged particles — the force that holds oppositely charged ions together throughout an ionic lattice.

Keep goingIonic compounds: Bonds, Structure & Properties 1-6-ionic-bonding 1-8-metallic-bonding 1-9-electrolysis 2-1-group-1-alkali-metals

The structure of magnesium oxide

SB p.82

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.)

Key point

The 6-around-6 number and the "no molecules" law are the two facts examiners probe in structure questions: describe the coordination (each ion surrounded by six of the opposite charge) and name the force (electrostatic attraction between oppositely charged ions, extending throughout the giant lattice). The same-structure-different-strength comparison of MgO vs NaCl is the classic "explain why magnesium oxide has a higher melting point than sodium chloride" answer at this level: higher charges → stronger attraction → more energy to break the lattice.

Exam hint

When explaining melting-point differences (spec 1.42), anchor every step to the lattice: charge → attraction strength → energy to separate → melting point. "MgO has stronger ionic bonds because the ions are charged 2+/2−" scores; "MgO's bonds are just stronger" scores nothing — the examiner wants the property of the ions (charge) linked to the property of the lattice (attraction) to the observable (melting point). Keep the exploded-view detail in reserve for "describe the structure" items: 6 of each, repeating, no molecules.

Common misconception

"The lines in the exploded lattice diagram are the ionic bonds — six bonds per sodium ion." The lines are not bonds, they show touching/arrangement. Each ion's attraction is not six private links but an electrostatic pull on all nearby oppositely charged ions, extending throughout the giant lattice. Drawing conventions ≠ physical bonds — the same discipline as dot-and-cross dots (section 2): notation is bookkeeping, the physics is the attraction.

Quick check

Question-bank join: spec 1.41 items — naming the force holding the lattice (electrostatic attraction between oppositely charged ions); spec 1.42 items — explaining MgO's higher melting point against NaCl via the 2+/2− versus 1+/1− charge comparison. Sketch the exploded NaCl/MgO lattice from memory with the 6-around-6 counts labelled.

Key terms in this section

giant ionic lattice: — A structure in which vast numbers of oppositely charged ions are held by electrostatic attraction extending in all directions; there are no individual molecules.
electrostatic attraction: — The attraction between positively and negatively charged particles — the force that holds oppositely charged ions together throughout an ionic lattice.

Keep goingIonic compounds: Bonds, Structure & Properties 1-6-ionic-bonding 1-8-metallic-bonding 1-9-electrolysis 2-1-group-1-alkali-metals

The physical properties of ionic substances

SB p.82

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.

Key point

The vocabulary note worth keeping: organic solvents include ethanol (alcohol) and hydrocarbons such as hexane and the mixture in petrol. The property pair to recall: soluble in water; insoluble in organic solvents.

Exam hint

The word molten just means the salt has been melted — it is a liquid. It appears in every conductivity and electrolysis question: "molten sodium chloride" = melted NaCl, ions free to move; "aqueous" = dissolved in water, same conclusion. Both words are route words, not chemistry content — know them cold so a vocabulary wobble never costs a physics-of-ions mark. (Spec 1.42's melting-point explanation gets its marks from the lattice-energy chain of section 13, not from the word "strong".)

Common misconception

"Ionic compounds have high melting points because an ionic bond is stronger than a covalent bond." The comparison is not bond-versus-bond — it is structure-versus-structure. Melting means separating the particles of the lattice, and in a giant lattice every ion is attracting many oppositely charged neighbours: you must overcome a vast network of attractions, not one pair. State it structurally — strong electrostatic attraction throughout the giant lattice, much energy to break — and the mark scheme's levels are all covered.

Quick check

Question-bank join: spec 1.42 items — explaining high melting/boiling points via the giant lattice; predicting property differences (brittle; water-soluble vs organic-insoluble) from the ionic structure. Quick audit of the four properties: for each, can you name the lattice feature that causes it without looking back?

Key terms in this section

molten: — Melted — the solid has been heated until it is a liquid.

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The electrical conductivity of ionic substances — and the lesson close

SB p.83

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.

Summary

In one page: formation — metals lose electrons, non-metals gain them; the charged particles are ions (cation +, anion −). Diagrams — dot-and-cross, outer shells only, charges labelled. Charges — Groups 1/2/3 lose 1/2/3 electrons (1+/2+/3+); Groups 5/6/7 gain 3/2/1 (3−/2−/1−); plus the learned list Zn²⁺, Ag⁺, H⁺, NH₄⁺, NO₃⁻, OH⁻, CO₃²⁻, SO₄²⁻ (not all ions are noble-gas configurations — Fe³⁺ and friends). Formulae — equal pluses and minuses; brackets only when a complex ion is multiplied; -ide = two elements, -ate = oxygen present; crossover with the same-charges trap (CaO, not Ca₂O₂). Structure — giant ionic lattice, no molecules, 6-around-6. Properties — high melting/boiling points (strong electrostatic attraction throughout the lattice), crystalline, brittle, water-soluble/organic-insoluble, and conducting only when molten or aqueous because only then are the ions free to move.

Can you…?

Explain how a sodium atom and a chlorine atom become Na⁺ and Cl⁻, and why the compound is neutral overall? (1.37)
State the charges of ions for Groups 1, 2, 3 and 5, 6, 7, and recall the eight learn-list ions? (1.38)
Write formulae on demand — Na₂O, CaCl₂, Ba(NO₃)₂, Fe₂(SO₄)₃, CuSO₄, (NH₄)₂SO₄ — with brackets in the right places? (1.39)
Draw dot-and-cross diagrams for NaCl, MgO and CaCl₂, outer shells only, charges shown? (1.40)
Say what ionic bonding is, in one phrase, without naming electrons as the carrier? (1.41)
Explain high melting points from the giant lattice, and why MgO beats NaCl? (1.42)
State when ionic compounds conduct and why — using "ions free to move", never "electrons"? (1.43)

Exam hint

The exam shape for this topic is stable and worth knowing in advance: electronic configurations of the ions in named reactions, then dot-and-cross diagrams for Group-1/2 × Group-7 combinations you have met here as families (K/F, Ca/Br, Mg/I are the classics). Nearly every ionic-bonding question is the same three moves in the same order: configurations → diagram → formula. Drill the moves, not the examples.

Common misconception

"Solid ionic compounds don't conduct because their electrons can't move." Swap the carrier word and the answer is correct: the ions are the charge carriers, and in the solid they are locked in the lattice. Writing "electrons" in a 1.43 answer costs the mark even though the sentence sounds physical — the spec words are ions fixed in position (solid) and ions free to move (molten/aqueous). The forbidden word is the tell that the learner has confused this with metallic conduction (Chapter 9).

Quick check

Question-bank join: spec 1.43 items — the three-state conductivity rule and its explanation in the mark scheme's own words; cross-links to the electrolysis material of u1:c10 where the mobile ions finally get put to work. Final self-audit: can you answer all seven can-you prompts from the summary alone, without reopening a single section?

Key terms in this section

molten: — Melted — the solid has been heated until it is a liquid.
aqueous solution: — A substance dissolved in water; its state symbol is (aq).

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Ionic bonding

Ionic bonding — the chapter map

Forming ions — sodium chloride

Ionic bonding in magnesium oxide

Noble gas electronic configurations — and the ions that break the pattern

Other examples of ionic bonding — lithium fluoride and calcium chloride

Formulae for ionic compounds

Confusing endings — -ide and -ate

To find the formula for sodium oxide

To find the formula for barium nitrate

To find the formula for iron(III) sulfate

Calcium chloride provides another example

Giant ionic structures

The structure of magnesium oxide

The physical properties of ionic substances

The electrical conductivity of ionic substances — and the lesson close