Ionic bonding: why ions form, how they pack, and why magnesium oxide melts at 2 850 °C
By the end of this lesson you will be able to
- Explain why metal atoms lose electrons and non-metal atoms gain them, and deduce the charge on an ion from the electron configuration or the group.
- Draw dot-and-cross diagrams for ionic compounds such as NaCl, MgO, CaF₂ and Na₂O, with brackets and charges.
- Write the formula of any ionic compound from its ions, including polyatomic ions and transition-metal ions named with Roman numerals.
- Describe the giant ionic lattice and explain the ionic bond as electrostatic attraction acting in every direction.
- Explain and predict differences in melting point from ion charge and ion size, and explain brittleness and conductivity with the lattice model.
You met this before
- An atom has protons in a tiny nucleus and the same number of electrons in shells round it. For the first 20 elements the shells fill 2, then 8, then 8: sodium is 2,8,1, chlorine 2,8,7, calcium 2,8,8,2.
- For groups 1 and 2 the number of outer (valence) electrons equals the group number; for groups 13 to 17 it is the group number minus 10. The noble gases (group 18) have full outer shells and hardly react.
- An ion is an atom (or group of atoms) that has lost or gained electrons, so the numbers of protons and electrons no longer match. The nucleus never changes in a chemical reaction.
- Opposite charges attract and like charges repel; the force is stronger when the charges are bigger and when they are closer.
- Sodium chloride is a giant lattice: it melts at 801 °C, and it conducts electricity when molten or dissolved, but not as a solid (Year 3).
Write the electron configurations of magnesium (Z = 12) and oxygen (Z = 8), and predict the charge on the ion each forms.
Reveal
Mg: 2,8,2. Losing its two outer electrons leaves 2,8 — the configuration of neon — with 12 protons and 10 electrons, so the ion is Mg²⁺.
O: 2,6. Gaining two electrons gives 2,8, also like neon, with 8 protons and 10 electrons, so the ion is O²⁻. Both ions have the same configuration; only the nuclear charge differs.
Revise it in full: Year 3 · Atomic structure and bonding · Lesson 2 — Electronic configuration: the rule of 2, 8, 8 · Year 3 · Atomic structure and bonding · Lesson 3 — Ionic bonding: metals give, non-metals take, and the lattice holds
The puzzle
Sodium chloride and magnesium oxide look almost the same: white, crystalline, made of a metal and a non-metal, each with a simple 1 : 1 formula and the same arrangement of ions. Yet table salt melts in a kitchen blowtorch at 801 °C, while magnesium oxide survives to about 2 850 °C.
That is why the inside of a steel-making furnace is lined with magnesia bricks. Same structure, same kind of bond: where do the extra two thousand degrees come from?
Predict which of these two changes would raise the melting point of an ionic solid more, and why: (i) doubling the charge on both ions; (ii) making both ions a little smaller.
Compare your answer
Both raise it; doubling both charges does more. The attraction between two ions depends on the product of their charges, so 2+ with 2− attracts four times as strongly as 1+ with 1− at the same distance.
Smaller ions sit closer together, which also strengthens the attraction, but by less. Magnesium oxide has both advantages over sodium chloride: charges 2+ and 2−, and smaller ions. Section 1.5 turns this into numbers.
The science — structured notes
1.1 Why atoms form ions
Atoms bond because the bonded arrangement has less energy — it is more stable — than the separate atoms. For many elements the stable arrangement is a full outer shell like a noble gas: the “octet” of eight outer electrons (two for the first shell).
A metal atom has only one, two or three outer electrons, held weakly because they are far from the nucleus and shielded by the inner shells. It reaches a noble-gas configuration by losing them, forming a positive ion (cation).
A non-metal atom has five, six or seven outer electrons and a strong pull from its nucleus. It reaches a noble-gas configuration by gaining electrons, forming a negative ion (anion). When a metal meets a non-metal, electrons are transferred from the metal to the non-metal.
The charge on the ion is simply protons minus electrons. A sodium ion has 11 protons and 10 electrons, so its charge is +1. A chloride ion has 17 protons and 18 electrons, so its charge is −1. Notice the name changes for the anion: chlorine atom, chloride ion.
| Element | Group | Atom | Electrons lost or gained | Ion | Ion configuration (like) |
|---|---|---|---|---|---|
| Lithium | 1 | 2,1 | loses 1 | Li⁺ | 2 (helium) |
| Sodium | 1 | 2,8,1 | loses 1 | Na⁺ | 2,8 (neon) |
| Magnesium | 2 | 2,8,2 | loses 2 | Mg²⁺ | 2,8 (neon) |
| Aluminium | 13 | 2,8,3 | loses 3 | Al³⁺ | 2,8 (neon) |
| Calcium | 2 | 2,8,8,2 | loses 2 | Ca²⁺ | 2,8,8 (argon) |
| Nitrogen | 15 | 2,5 | gains 3 | N³⁻ (nitride) | 2,8 (neon) |
| Oxygen | 16 | 2,6 | gains 2 | O²⁻ (oxide) | 2,8 (neon) |
| Sulfur | 16 | 2,8,6 | gains 2 | S²⁻ (sulfide) | 2,8,8 (argon) |
| Fluorine | 17 | 2,7 | gains 1 | F⁻ (fluoride) | 2,8 (neon) |
| Chlorine | 17 | 2,8,7 | gains 1 | Cl⁻ (chloride) | 2,8,8 (argon) |
Where the stability really comes from. Taking an electron from a sodium atom costs energy (496 kJ per mole of atoms), and a chlorine atom gives back only 349 kJ mol⁻¹ when it takes one. On their own, the two ions are less stable than the atoms.
The pay-off is the attraction: when a mole of Na⁺ and Cl⁻ ions come together into a lattice, about 787 kJ is released. The electrostatic attraction between the ions, not the octet itself, is what makes sodium chloride stable. The octet tells you which ions form.
Deduce the ion formed by each of these atoms, with its charge and electron configuration: potassium (atomic number 19), sulfur (16), aluminium (13) and nitrogen (7).
Reveal answer
K: 2,8,8,1 → K⁺, configuration 2,8,8.
S: 2,8,6 → S²⁻ (sulfide), 2,8,8.
Al: 2,8,3 → Al³⁺, 2,8.
N: 2,5 → N³⁻ (nitride), 2,8. In each case count protons minus electrons to check: Al³⁺ has 13 protons and 10 electrons.
1.2 Dot-and-cross diagrams for ionic compounds
A dot-and-cross diagram shows where every outer electron ends up. Electrons that came from one atom are drawn as crosses (×), from the other as dots (•) — all electrons are identical, the symbols only keep the bookkeeping honest. Four rules earn the marks:
- Draw each ion inside square brackets with its charge at the top right (2+, not +2).
- The anion’s outer shell shows eight electrons, including the crosses it received.
- The cation is drawn with an empty outer shell, or with its complete inner shell; either is accepted if the charge is right.
- Show the right number of each ion: CaF₂ needs one [Ca]²⁺ and two [F]⁻, or a “2” in front of [F]⁻.
Draw the dot-and-cross diagram of calcium fluoride and deduce its formula.
Configurations. Ca 2,8,8,2 has two outer electrons; F 2,7 needs one.
Transfer. One calcium atom gives one electron to each of two fluorine atoms.
Ions. [Ca]²⁺ and two [F]⁻, each fluoride with seven dots and one cross in its outer shell.
Formula. Charges must balance: (+2) + 2 × (−1) = 0, so the formula is CaF₂.
A student drew sodium oxide as: [Na]⁺ [O]²⁻, with the oxide ion showing six dots and one cross. Find two errors and describe the correct diagram.
Reveal answer
Error 1: only one sodium ion is shown. Oxygen needs two electrons, and each sodium atom has only one to give, so there must be two [Na]⁺ ions (formula Na₂O).
Error 2: the oxide ion must show eight outer electrons — six dots from oxygen and two crosses, one from each sodium atom.
Correct: 2[Na]⁺ [O]²⁻, with the oxide ion’s outer shell holding 6 dots and 2 crosses and the charge 2− at the top right of its bracket.
1.3 Formulae: balancing charges, polyatomic ions and Roman numerals
An ionic compound has no overall charge, so the formula is the simplest whole-number ratio of ions whose charges add to zero. Two further kinds of ion appear constantly in examinations.
Polyatomic ions are groups of covalently bonded atoms carrying an overall charge. They stay together as one unit, so when you need more than one of them you put the group in brackets: Ca(OH)₂, not CaOH₂.
| Positive polyatomic ion | Negative polyatomic ions | ||
|---|---|---|---|
| ammonium NH₄⁺ | hydroxide OH⁻ | nitrate NO₃⁻ | hydrogencarbonate HCO₃⁻ |
| carbonate CO₃²⁻ | sulfate SO₄²⁻ | phosphate PO₄³⁻ |
Transition metals often form more than one ion. The name gives the charge as a Roman numeral: iron(II) is Fe²⁺, iron(III) is Fe³⁺, copper(I) is Cu⁺ and copper(II) is Cu²⁺. Zinc (Zn²⁺) and silver (Ag⁺) have one common ion each and need no numeral.
Write the formula of aluminium sulfate.
Ions. Al³⁺ and SO₄²⁻.
Lowest common multiple of the charges. 6.
Two Al³⁺ give +6; three SO₄²⁻ give −6.
Formula. Al₂(SO₄)₃ — brackets round the sulfate because there are three of the whole group.
Bone and tooth enamel are built mainly from a calcium phosphate mineral. Write the formula of calcium phosphate, and calculate the percentage by mass of calcium in it (relative atomic masses: Ca 40.1, P 31.0, O 16.0).
Ions. Ca²⁺ and PO₄³⁻; the lowest common multiple of 2 and 3 is 6.
Three Ca²⁺ give +6; two PO₄³⁻ give −6: Ca₃(PO₄)₂.
Formula mass. 3 × 40.1 + 2 × (31.0 + 4 × 16.0)
Formula mass = 120.3 + 190.0
Formula mass = 310.3
Percentage of calcium. 120.3 ÷ 310.3 × 100
Percentage of calcium = 38.8 %
The real bone mineral, hydroxyapatite, also contains hydroxide ions, Ca₅(PO₄)₃OH; in fluoride toothpaste some OH⁻ is swapped for F⁻, giving a harder, less acid-soluble fluorapatite.
Write the formulae of: ammonium nitrate; iron(III) oxide; magnesium hydroxide; copper(II) nitrate; ammonium phosphate.
Reveal answer
NH₄NO₃ (charges +1 and −1).
Fe₂O₃ (2 × +3 balances 3 × −2).
Mg(OH)₂ (brackets essential).
Cu(NO₃)₂.
(NH₄)₃PO₄ (three +1 ions balance one −3 ion).
1.4 The giant ionic lattice
Ions do not pair off. Each ion attracts every oppositely charged ion around it and repels every ion of the same charge. The arrangement with the lowest energy is a regular three-dimensional giant lattice, in which every ion is surrounded by ions of the opposite charge.
In sodium chloride every Na⁺ has six Cl⁻ neighbours and every Cl⁻ has six Na⁺ neighbours. A grain of salt 1 mm across contains roughly 4 × 10¹⁹ ions.
This is what the IB means by “the ionic bond is due to electrostatic attraction between oppositely charged ions”. The bond is not a stick between one Na⁺ and one Cl⁻: it is the net attraction of the whole lattice, acting in all directions.
The formula NaCl is therefore a ratio of ions (1 : 1), not a molecule. There is no such thing as “a molecule of sodium chloride” in the solid.
“In solid sodium chloride each sodium ion is bonded to one chloride ion, which is why the formula is NaCl.”
Reveal answer
False. Each Na⁺ is attracted equally to six neighbouring Cl⁻ ions (and more weakly to more distant ones). The formula records only the 1 : 1 ratio of the ions in the lattice.
X-ray diffraction, in 1913, was the evidence that settled this: no sodium ion is closer to one particular chloride ion than to the other five.
1.5 How strong? Charge and size decide
To melt an ionic solid the ions must be given enough energy to break free of their fixed positions, overcoming the attractions of the lattice. The stronger the attraction, the higher the melting point. For two ions, the force of attraction depends on:
F ∝ q₁ × q₂d²
where q₁ and q₂ are the charges on the ions and d is the distance between their centres (roughly the sum of the ionic radii). Two rules follow. Larger charges mean stronger attraction, so oxides of 2+ metals melt far higher than halides of 1+ metals. Smaller ions sit closer, so the attraction is stronger: NaF melts higher than NaI.
Use the ionic radii Na⁺ 102 pm, Cl⁻ 181 pm, Mg²⁺ 72 pm and O²⁻ 140 pm to estimate how many times stronger the attraction between one Mg²⁺ and one O²⁻ is than between one Na⁺ and one Cl⁻.
Distances. d(NaCl) = 102 + 181 = 283 pm
d(MgO) = 72 + 140 = 212 pm
Charge factor. (2 × 2) ÷ (1 × 1) = 4
Distance factor. (283 ÷ 212)²
Distance factor = 1.335² = 1.78
Ratio of forces. 4 × 1.78 = 7.1
Charge contributes most of the difference; size adds most of the rest. That is where magnesium oxide’s extra two thousand degrees come from.
Rank potassium chloride (K⁺ 138 pm), calcium oxide (Ca²⁺ 100 pm) and sodium chloride in order of increasing melting point, and justify with charges and sizes. Then check against the data: KCl 770 °C, NaCl 801 °C, CaO about 2 600 °C.
Reveal answer
KCl < NaCl < CaO. KCl and NaCl both have charges 1+ and 1−, but K⁺ is larger than Na⁺, so the ions in KCl are further apart (319 pm against 283 pm) and attract less strongly.
CaO has charges 2+ and 2− — four times the charge product — and its ions are close (240 pm), so its lattice is much the strongest. The data agree.
1.6 What the lattice model explains
| Property | Observation | Explanation from the model |
|---|---|---|
| Melting point | High: NaCl 801 °C, MgO about 2 850 °C | Many strong electrostatic attractions in every direction must be overcome |
| Hardness and brittleness | Hard, but shatter along flat faces when struck | A shift of one ion brings like charges face to face; they repel and the crystal splits |
| Conductivity of the solid | Does not conduct | The ions are charged but held in fixed positions; nothing charged can move |
| Conductivity when molten or dissolved | Conducts; is decomposed at the electrodes (electrolysis) | The ions are free to move and carry charge towards the oppositely charged electrodes |
| Solubility | Many dissolve in water, few in oil or hexane | Polar water molecules attract the ions strongly enough to pull them out of the lattice (Lesson 4) |
Solid calcium chloride does not conduct electricity. The same compound conducts when molten or when dissolved in water. Explain both observations.
Reveal answer
Calcium chloride is ionic: a giant lattice of Ca²⁺ and Cl⁻ ions held by strong electrostatic attraction. A current needs charged particles that are free to move.
In the solid the ions are charged but locked in fixed positions, so no charge can flow. When the solid melts, or dissolves, the lattice breaks up and the ions are free to move: Ca²⁺ towards the negative electrode and Cl⁻ towards the positive electrode, carrying the current.
Aluminium oxide has ions Al³⁺ (54 pm) and O²⁻ (140 pm), so its charge product is 3 × 2 = 6, larger than magnesium oxide’s 4. Yet Al₂O₃ melts at 2 072 °C, below MgO. Calculate the force factor q₁q₂ ÷ d² for both relative to NaCl, and suggest why the simple model fails here.
Reveal answer
NaCl: 1 ÷ 283² — call this 1.00.
MgO: 4 × (283 ÷ 212)² = 7.1
Al₂O₃: 6 × (283 ÷ 194)²
Al₂O₃: 6 × 2.13 = 12.8
On this model Al₂O₃ should melt far above MgO, and it does not. The model treats ions as hard charged spheres.
A very small, highly charged Al³⁺ ion pulls the electron cloud of the neighbouring O²⁻ towards itself, so electrons are partly shared: the bonding has some covalent character, and the structure is less symmetrical than rock salt. Real bonding is a continuum between ionic and covalent (Lesson 3); the pure ionic model works best for large, singly or doubly charged ions.
Try it — are the ions really what carries the current?
The lab dissolves a weighed mass of salt in 100 cm³ of distilled water and passes a current between two graphite electrodes 3 cm apart, connected to a 6 V supply and a milliammeter.
You can change the solute (sodium chloride, sugar or copper sulfate) and test the dry solid instead of a solution. The milliammeter reads to 0.1 mA.
Typical mean currents from the lab were: sodium chloride 2 g → 46.2 mA, 4 g → 88.4 mA, 6 g → 126.8 mA, 8 g → 161.4 mA, 10 g → 192.1 mA (all per 100 cm³). Copper sulfate at 4 g gave 19.5 mA and at 8 g gave 33.5 mA. Sugar at 10 g gave 0.1 mA and dry salt 0.0 mA.
(a) Test whether the current is directly proportional to the concentration of sodium chloride.
(b) Suggest one reason why a gram of copper sulfate gives a smaller current than a gram of sodium chloride.
(c) Explain the sugar and dry-salt results.
Reveal answer
(a) Current ÷ concentration: 23.1, 22.1, 21.1, 20.2 and 19.2 mA per g per 100 cm³ — falling steadily, so the current is not exactly proportional. Doubling 4 g to 8 g multiplies the current by 161.4 ÷ 88.4 = 1.83, not 2.
The graph is nearly straight at low concentration and bends below the line at high concentration, as crowded ions slow one another.
(b) A gram of CuSO₄ (formula mass 159.6) contains fewer formula units — fewer ions — than a gram of NaCl (formula mass 58.5), about 2.7 times fewer.
(c) Sugar is a covalent molecular compound: it dissolves as neutral molecules, so there are no ions to carry charge and the reading is that of the water itself. Dry salt contains ions, but they are fixed in the lattice and cannot move.
Common mistakes to avoid
- “The ions share electrons” — ionic bonding is transfer followed by attraction; sharing is covalent (Lesson 2).
- Leaving out the attraction — “sodium loses an electron and chlorine gains it” is only half; add “the oppositely charged ions are held by electrostatic attraction”.
- Wrong charge notation — write Mg²⁺ and O²⁻; in diagrams put the charge outside the bracket, top right.
- Missing brackets on polyatomic ions — Ca(OH)₂ and Al₂(SO₄)₃, never CaOH₂ or Al₂SO₄₃.
- Calling NaCl a molecule — it is a giant lattice; the formula is a ratio.
- “Electrons carry the current in salt solution” — ions do.
- Confusing chlorine and chloride — chlorine is the element (Cl₂ or a Cl atom); chloride is the ion Cl⁻.
Study smart: how this is examined
Study routine. Today: learn the polyatomic ions and write ten formulae from scratch; draw NaCl, MgO, CaF₂ and Na₂O as dot-and-cross diagrams. In two days: without notes, write the four-point explanation of ionic bonding for lithium oxide, then rank NaF, NaCl, MgO and CaO by melting point with a reason.
In a week: run the lab as a full investigation, do the Think task, then the quiz.
How we found out — ions nobody believed in
In 1834 Michael Faraday, studying electrolysis at the Royal Institution in London, needed words for the things that carried current through a solution. With the help of the Cambridge scholar William Whewell he introduced “ion” (Greek for “goer”), “anion”, “cation”, “electrode” and “electrolyte”. Faraday thought the current itself tore molecules apart into ions.
In 1884 the young Swedish chemist Svante Arrhenius argued in his doctoral thesis that salts split into ions as soon as they dissolve, with no current needed. His examiners gave the thesis almost the lowest passing grade.
The idea that a solution of salt could contain free, electrically charged sodium seemed absurd — sodium metal explodes in water. Measurements of conductivity and of freezing points steadily supported him, and Arrhenius received the Nobel Prize in Chemistry in 1903.
In 1913 William Henry Bragg and his son William Lawrence Bragg fired X-rays through crystals of sodium chloride and worked out the positions of the atoms from the pattern of reflections. There were no NaCl pairs: each sodium was surrounded by six chlorines at equal distances. They shared the 1915 Nobel Prize in Physics.
Some chemists refused to accept it. In 1927 Henry Armstrong wrote to Nature that the idea was “repugnant to common sense”. The evidence won.
The explanation of why ions form came in 1916, when Walther Kossel in Germany proposed that atoms transfer electrons to reach the configuration of the nearest noble gas — in the same year that Gilbert N. Lewis in California proposed the shared pair (Lesson 2).
Where it is used now
- Salt and soda ash from the Makgadikgadi. At Sua Pan in Botswana, Botash pumps brine from beneath the pan, evaporates it in solar ponds and crystallises sodium chloride and sodium carbonate — two ionic compounds sold across southern Africa for glass, detergents, water treatment and food.
- Furnace linings. Magnesium oxide bricks line steel furnaces and cement kilns because their lattice survives temperatures that melt the steel inside.
- Oral rehydration salts. The WHO formula dissolves sodium chloride, potassium chloride, trisodium citrate and glucose in a litre of water; the ions and glucose together restore what diarrhoea removes, and the mixture saves many thousands of children’s lives each year.
- Fertilisers. Ammonium nitrate, potassium chloride (potash) and ammonium phosphates are ionic, dissolve in soil water and deliver N, K and P as ions that roots can absorb.
- Teeth. Sodium fluoride in toothpaste supplies F⁻ ions that replace OH⁻ in enamel, making it harder to dissolve in the acids bacteria produce.
Science now — and the scientist you need to be
What is happening in 2026
Batteries are ionic chemistry: in every lithium-ion cell, Li⁺ ions shuttle through an electrolyte between two electrodes. Two changes are under way. Sodium-ion batteries, which use Na⁺ from salt instead of scarcer lithium, began to enter mass production in China in 2024–2025, first for grid storage and small cars.
Many carmakers are also developing solid-state batteries, whose electrolyte is an ionic solid through which small ions can hop, and several have announced plans to use them later this decade.
Finding a solid that lets ions move quickly while staying stable is a search through millions of possible lattices. In 2023 Google DeepMind’s GNoME model predicted about 2.2 million new crystal structures, of which some 380 000 were judged stable. In early 2024 Microsoft and the Pacific Northwest National Laboratory used AI screening of about 32 million candidates to find a new sodium-and-lithium solid electrolyte, then made and tested it.
The chemist of the next decade will use such tools daily — and will still need to judge, from charges and sizes, whether a predicted lattice makes sense.
The mindshift
What you need to be able to do
Go from the periodic table to the ion, from the ions to the formula and the dot-and-cross diagram, and from the charges and sizes to a justified prediction of melting point — then use the lattice to explain why the compound is brittle and when it conducts.
See an example
Toothpaste: sodium fluoride NaF (Na⁺, F⁻). Sports drink: sodium chloride NaCl, potassium citrate, magnesium carbonate MgCO₃ (Mg²⁺, CO₃²⁻). Antacid: magnesium hydroxide Mg(OH)₂ and calcium carbonate CaCO₃.
A correct MgCl₂ diagram shows one [Mg]²⁺ and two [Cl]⁻, each chloride with seven dots and one cross, charges at the top right. A common AI error is to show a single chloride with two crosses.
Key terms
| Ion | An atom or group of atoms with an overall charge, because the number of electrons differs from the number of protons. |
| Cation | A positive ion, formed when a metal atom loses its valence electrons. |
| Anion | A negative ion, formed when a non-metal atom gains electrons. |
| Ionic bond | The electrostatic attraction between oppositely charged ions. |
| Giant ionic lattice | A regular three-dimensional arrangement of very many ions in which each ion is surrounded by ions of opposite charge. |
| Octet rule | The tendency of atoms to gain, lose or share electrons until they have eight electrons in their outer shell. |
| Polyatomic ion | A charged group of covalently bonded atoms that behaves as one unit, such as SO₄²⁻ or NH₄⁺. |
| Dot-and-cross diagram | A diagram showing the outer electrons of each atom or ion, with the origin of each electron shown by a dot or a cross. |
| Brittle | Shattering, rather than bending, when struck — in ionic solids because a shift brings like charges together. |
| Electrolyte | A molten or dissolved ionic compound that conducts electricity by the movement of ions. |
Think like a scientist
Criterion A Aluminium oxide (alumina, Al₂O₃) is the mineral in sapphires and rubies and the ore from which aluminium metal is made. It melts at 2 072 °C. In the Hall–Héroult process it is dissolved in molten cryolite (Na₃AlF₆) at about 960 °C and electrolysed; aluminium forms at the negative electrode.
(a) State the electron configurations of an aluminium atom and an oxygen atom and of the ions each forms. [2]
(b) Draw a dot-and-cross diagram for aluminium oxide, showing outer shells only, and deduce its formula. [3]
(c) Explain why aluminium oxide has a very high melting point. [3]
(d) Explain why the alumina must be molten or dissolved, not solid, for electrolysis to work, and suggest why the industry dissolves it in cryolite rather than melting it. [3]
(e) A student predicts, using charges and ionic radii, that Al₂O₃ should melt at a higher temperature than MgO (about 2 850 °C). Evaluate this prediction against the data. [3]
Model answer
(a) Al 2,8,3 → Al³⁺ 2,8; O 2,6 → O²⁻ 2,8 [1 for the atoms, 1 for the ions].
(b) Two aluminium atoms give three electrons each (six in total) to three oxygen atoms, which take two each.
Diagram: 2[Al]³⁺ and 3[O]²⁻; each oxide ion has an outer shell of 6 dots and 2 crosses, each ion in brackets with its charge top right [2]. Formula Al₂O₃, since 2 × (+3) + 3 × (−2) = 0 [1].
(c) Aluminium oxide is a giant ionic lattice of Al³⁺ and O²⁻ ions [1]. Every ion is held by strong electrostatic attraction to many oppositely charged neighbours, in all directions [1].
The charges are large (3+ and 2−) and the ions small, so the attractions are especially strong and a great deal of energy is needed to free the ions from their positions [1].
(d) Electrolysis needs the ions to move to the electrodes; in the solid they are fixed in the lattice, but when molten or dissolved they are free to move [1]. Al³⁺ ions then travel to the negative electrode, gain electrons and form aluminium [1].
Melting alumina needs over 2 072 °C, a huge energy cost and very hard on the equipment; dissolved in cryolite the mixture is liquid at about 960 °C, saving energy and money [1].
(e) On the simple model the force factor for Al₂O₃ (charges 3 and 2, distance 54 + 140 = 194 pm) is about 12.8 times NaCl’s, against 7.1 for MgO, so the prediction follows logically from the model [1].
But the data contradict it: 2 072 °C is about 780 °C below MgO [1]. The model assumes hard spheres with complete charge transfer; the very small, highly charged Al³⁺ distorts the oxide ion’s electrons, giving partly covalent bonding and a less symmetrical lattice, so the model overestimates [1].
Probe deeper — exam-style questions
Attempt every question before opening the answer. Marks in brackets show what a full answer is worth.
State the electron configuration of a sulfur atom and of the ion it forms in sodium sulfide, and give the formula of sodium sulfide.
Reveal answer
S 2,8,6; S²⁻ 2,8,8 [1].
Na₂S [1].
Draw a dot-and-cross diagram for calcium chloride, showing outer electrons only.
Reveal answer
One [Ca]²⁺ with an empty outer shell (or its full inner shell), charge 2+ top right [1].
Two [Cl]⁻ (or 2[Cl]⁻), each with 7 dots and 1 cross [1].
Brackets and charges correct on every ion [1].
Write the formulae of ammonium sulfate and iron(III) hydroxide.
Reveal answer
(NH₄)₂SO₄ [1].
Fe(OH)₃ [1]. Brackets required in both.
Explain how lithium and oxygen bond to form lithium oxide.
Reveal answer
Each lithium atom (2,1) loses one electron, forming Li⁺ [1].
Each oxygen atom (2,6) gains two electrons, forming O²⁻, so two lithium atoms are needed per oxygen atom [1].
An ionic bond forms [1], which is the electrostatic attraction between the oppositely charged Li⁺ and O²⁻ ions, in a giant lattice [1].
A crystal of sodium chloride shatters when struck with a hammer. A piece of copper bends. Explain why the sodium chloride shatters.
Reveal answer
In the lattice each ion is surrounded by oppositely charged ions [1].
A blow shifts one layer of ions along by one ion, so that ions of the same charge come next to each other [1].
The like charges repel strongly, so the layers are pushed apart and the crystal splits along a plane [1].
The melting points of sodium fluoride and sodium iodide are 993 °C and 661 °C. Explain the difference.
Reveal answer
Both are giant ionic lattices with ions of charge 1+ and 1− [1].
The fluoride ion (133 pm) is much smaller than the iodide ion (220 pm), so in NaF the ions are closer together [1].
The electrostatic attraction between closer ions is stronger, so more energy is needed to overcome it and the melting point is higher [1].
7. Three white solids are tested. W: melts at about 2 850 °C and does not dissolve in water. X: melts at 801 °C, dissolves, and the solution conducts. Y: melts (and decomposes) at about 186 °C, dissolves, and the solution does not conduct. Identify which could be sodium chloride, sucrose and magnesium oxide, and suggest one further test that would distinguish W from a giant covalent solid such as silicon dioxide.
Reveal answer
X is sodium chloride: ionic lattice, 801 °C, soluble, and its ions conduct in solution [1].
Y is sucrose: low melting point (weak forces between molecules) and a non-conducting solution of neutral molecules [1].
W is magnesium oxide: a very high melting point from 2+ and 2− ions close together, and almost insoluble because its lattice is so strong [1].
Test: melt W and test the conductivity of the liquid. A molten ionic compound conducts because its ions move; molten silicon dioxide, which has no ions, does not [1].
8. Lithium fluoride has ions Li⁺ (76 pm) and F⁻ (133 pm); sodium fluoride has Na⁺ (102 pm) and F⁻. Using F ∝ q₁q₂ ÷ d², predict which should melt higher and calculate the ratio of the two forces. The measured melting points are LiF 845 °C and NaF 993 °C. Comment.
Reveal answer
d(LiF) = 76 + 133 = 209 pm
d(NaF) = 102 + 133 = 235 pm [1]
Ratio of forces = (235 ÷ 209)²
Ratio of forces = 1.124² = 1.26 — LiF predicted to melt higher [1].
The data show the opposite: NaF melts 148 °C higher [1].
The simple model predicts correctly for large changes of charge (NaCl against MgO) and for a series with one ion varied (NaF to NaI), but here it fails.
The attraction in solid LiF really is stronger (its lattice energy is larger than that of NaF), yet melting point also depends on how the ions behave in the liquid, and the very small Li⁺ ion distorts (polarises) the electron cloud of the fluoride ion. A model is a tool with limits; the trend holds only when other things are similar [1].
Summary
- Metals lose their valence electrons to form cations and non-metals gain electrons to form anions; the electron configuration (or group) decides the charge: 1+, 2+, 3+; 3−, 2−, 1−.
- The ionic bond is the electrostatic attraction between oppositely charged ions, acting in every direction in a giant lattice; the formula is the ratio of ions.
- Dot-and-cross diagrams: brackets, charge top right, eight in the anion’s outer shell, the right number of each ion.
- Formulae balance the charges; polyatomic ions go in brackets when there is more than one; Roman numerals give transition-metal charges.
- Larger charges and smaller ions give stronger attraction and a higher melting point; ionic solids are brittle, and conduct only when molten or dissolved, because then the ions can move.
Next time → Covalent bonding: why two non-metals share instead of transfer, how to draw a Lewis structure for any small molecule in five steps, and why a triple bond makes nitrogen gas almost impossible to break apart.