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The dot-and-cross diagram for sodium chloride is probably the most drawn picture in school chemistry. One electron hops from a sodium atom to a chlorine atom, square brackets go round both ions, and a small charge sits in each corner. It is a useful diagram. It is also, if it is the only picture students ever see, a quiet source of several of the most stubborn misconceptions in the subject.
Students who learn ionic bonding only through that picture tend to conclude three things: that ions form because atoms “want” a full outer shell, that salt is made of NaCl pairs, and that the bond exists between the one sodium and the one chlorine that swapped the electron. Science-education researchers, notably Keith Taber, have documented these ideas for decades. None of them is true, and all of them cause trouble later with lattice energy, solubility and electrolysis.
This guide describes a two-lesson sequence for students of about 14 to 16 that keeps the dot-and-cross diagram but puts it in its proper place: as a record of electron transfer, not a picture of the bond.
Learning goals
Students should be able to:
- Describe how metal atoms lose electrons and non-metal atoms gain them to form ions, and draw dot-and-cross diagrams for simple cases.
- Define ionic bonding as the electrostatic attraction between oppositely charged ions, acting in all directions.
- Describe the structure of a solid ionic compound as a giant lattice in which each ion is surrounded by ions of opposite charge.
- Explain that the formula NaCl gives the ratio of ions, not a molecule.
- Use the lattice model to explain high melting points, brittleness and the conditions under which ionic compounds conduct.
Prior knowledge
- Electron arrangements for the first twenty elements.
- Metals on the left of the periodic table, non-metals on the right.
- The basic idea that opposite charges attract and like charges repel.
A short retrieval quiz on electron arrangements at the start of lesson 1 is worth the time. The interactive electron configuration tool is handy for students who want to check.
Lesson 1: from transfer to attraction (60 minutes)
Starter: what holds salt together? (5 minutes)
Hand round a few crystals of table salt and a hand lens. Ask: “Salt is made of sodium and chlorine. What is holding the particles together in this crystal?” Collect answers without comment. Many will mention “sharing” or “swapping electrons”. Keep the list for the plenary.
Direct teaching: making ions (15 minutes)
Draw the dot-and-cross diagrams for NaCl, MgO and CaCl₂ in the usual way (see dot-and-cross diagrams for conventions). Then add the step most textbooks skip: what happens next? The diagram shows the electron moving. It does not yet show a bond.
Ask where the bond is. Guide students to the answer that the bond is the attraction between a positive and a negative ion, and that nothing about that attraction is limited to the partner that supplied or received the electron.
Activity: the neighbour count (20 minutes)
Give each group a tray of two colours of modelling clay balls, one colour slightly larger than the other (larger for chloride, smaller for sodium), and ask them to build a flat layer in which no two balls of the same colour touch. They will produce a checkerboard. Then build a second layer on top, with each ball sitting above one of the opposite colour, and a third.
Now pick one small ball near the centre and count the large balls touching it. The answer is six: four in its own layer, one above and one below. Ask: “Which of these six is its ‘partner’?” There isn’t one. Each sodium ion is attracted equally to all six neighbours, and each chloride ion to six sodium ions. This is the moment the “one Na, one Cl” idea usually collapses.
If you use cocktail sticks to join the balls, discuss the limitation straight away: the sticks show where ions sit, not bonds between pairs. Some teachers prefer no sticks at all, pressing the clay balls gently together, precisely to avoid that impression. Wash hands afterwards; no other safety steps are needed.
Direct teaching: formula as a ratio (10 minutes)
Count the balls in a finished model. However big it is, the numbers of each colour are equal (or very nearly, at the edges). That is what NaCl means: one sodium ion for every chloride ion across the whole crystal, not a molecule of two atoms. Contrast with MgCl₂: one Mg²⁺ for every two Cl⁻, because that is what makes the charges cancel. Ionic lattice structure is a good follow-up reading.
Plenary (10 minutes)
Return to the starter list. Students mark each suggestion as correct, partly correct or incorrect, and rewrite the incorrect ones.
Lesson 2: why it happens, and what it explains (60 minutes)
Starter: does sodium “want” to lose an electron? (10 minutes)
Pose the question as a vote. Most students will say yes, “to get a full outer shell”. Then share a single fact: removing an electron from a gaseous sodium atom costs energy, about 496 kJ/mol. Chlorine gaining an electron releases about 349 kJ/mol, which is less. So transferring one electron between isolated atoms is uphill overall.
Ask: “If it costs energy, why does sodium chloride form at all?” Let them struggle for a minute.
Direct teaching: the energy story (15 minutes)
The missing piece is the lattice. When the new ions are pulled together into a crystal, a very large amount of energy is released, around 787 kJ/mol for NaCl. That release pays for the ion-making step with plenty to spare, and it is the real driving force. The full outer shell is a description of the ions that result, not a reason for them forming.
Keep this qualitative for most groups: an energy bar with “cost to make ions” going up a little and “energy released when ions attract” going down a lot. Stronger students can meet the numbers, and later the Born–Haber cycle.
A comparison that lands well: magnesium oxide needs even more energy to make its 2+ and 2− ions, yet it forms readily, because a lattice of doubly charged ions releases far more energy still. The article on NaCl vs MgO is written for exactly this discussion.
Activity: explain the properties (25 minutes)
Run a carousel of four stations. Each station has a short task and a question to answer using the lattice model.
- Melting point. A card with the melting point of sodium chloride (801 °C) and of sugar (which melts or decomposes at a far lower temperature). Why the difference? (Many strong attractions in every direction must be overcome in NaCl.)
- Brittleness. Students press a crystal of rock salt with the back of a spoon on a tile, then slide one layer of their clay model sideways by one ball. What ends up next to what? (Like charges meet and repel, so the crystal splits.)
- Conductivity. If you have a low-voltage conductivity tester, test dry salt crystals, then salt dissolved in water. Otherwise use a card with results. Why does the solution conduct but the solid not? (Ions are free to move only when the lattice is broken up.)
- Dissolving. A diagram of water molecules surrounding ions. What do the water molecules have to compete with? (The attraction between the ions in the lattice.)
All four stations use table salt and water only. The properties of ionic compounds page summarises the answers for revision.
Exit ticket (10 minutes)
Two questions: “Draw the arrangement of ions around one Na⁺ in a sodium chloride crystal” and “Explain in one sentence why NaCl is not a molecule.”
Misconceptions to expect
| Misconception | Why it’s wrong | What to do |
|---|---|---|
| “Atoms form ions because they want a full shell.” | Atoms have no wants; making Na⁺ from Na costs energy | Share the ionisation energy and the lattice energy; describe the full shell as a result, not a cause |
| “Sodium chloride is made of NaCl molecules.” | The solid is a continuous lattice with no separate units | Count ions in a model; the formula is a ratio |
| “The ionic bond is between one Na⁺ and one Cl⁻.” | Each ion attracts all its neighbours, six in NaCl | The neighbour-count activity |
| “An ion can only form as many bonds as electrons it gave or took.” | Attraction is not tied to the electron’s history | Ask how Na⁺ “knows” which Cl⁻ took its electron |
| “Ionic bonding only happens between atoms that swapped electrons.” | Ionic compounds can form from ions made in other ways, e.g. precipitation from solution | Show that solutions of two salts can form a new ionic solid |
| “Solid salt conducts because it contains ions.” | Ions in a solid lattice cannot move | Conductivity station |
| “Ionic bonds are weaker than covalent bonds because salt dissolves.” | Dissolving involves attraction to water, not simply breaking weak bonds | Compare the high melting point with the ease of dissolving |
Differentiation
Support
- Provide pre-drawn outlines for dot-and-cross diagrams so students focus on the electrons, not the circles.
- Use only NaCl in the modelling activity, then extend to MgO with the same model.
- Offer sentence starters: “The bond is the attraction between…”, “Each ion is surrounded by…”.
Stretch
- Predict the formula of an ionic compound from the ion charges, then explain why its ratio differs from NaCl’s (for example, CaF₂).
- Explain why MgO has a much higher melting point than NaCl using both charge and ionic size.
- Discuss the limits of the model: most “ionic” compounds have some covalent character, and the electron is never cleanly transferred in every case.
Questions to ask
- “Where exactly is the bond in this diagram?”
- “How many chloride ions is this sodium ion attracted to?”
- “If you split a crystal of salt in half, do you get two NaCl molecules?”
- “Does making ions from atoms release energy or take it in? So what makes up the difference?”
- “What would have to be free to move for this solid to conduct?”
Question bank
Recall
- Define ionic bonding. (The electrostatic attraction between oppositely charged ions.)
- How many ions of opposite charge surround each ion in sodium chloride? (Six.)
Application 3. Draw a dot-and-cross diagram for magnesium oxide, showing outer shells only. 4. Write the formula of the compound formed between calcium and chlorine and explain the ratio. (CaCl₂: Ca²⁺ needs two Cl⁻ to balance its charge.)
Explanation 5. A student says “salt is made of NaCl molecules”. Explain why this is incorrect. (Solid NaCl is a giant lattice; each ion is attracted to six neighbours; the formula gives the ratio of ions.) 6. Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride does. (Ions are fixed in the solid; in the melt they can move and carry charge.)
Challenge 7. Removing an electron from a sodium atom takes more energy than chlorine releases when it gains one. Explain why sodium chloride still forms. (The lattice releases a large amount of energy when the ions come together, more than making up for the cost.)
Summary
- Keep dot-and-cross diagrams, but present them as a record of electron transfer, not as a picture of the bond.
- Define the bond as attraction between ions in all directions, and make the six-neighbour arrangement physical with a model.
- Teach the formula of an ionic compound as a ratio, and say explicitly that there are no molecules.
- Replace “atoms want a full shell” with a simple energy story in which the lattice pays for making the ions.
- Use the lattice model to explain melting point, brittleness, conductivity and dissolving, so the structure earns its place.
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