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Educerie · IB Diploma · Chemistry
Structure 2 Models of bonding and structure · S2.3 The metallic model
What you must be able to do
| You must be able to | Level | What it looks like in the exam |
|---|---|---|
| Describe the metallic bond as the attraction between a lattice of cations and delocalized electrons | SL, HL | "Describe metallic bonding" (2 marks) |
| Explain the electrical conductivity, thermal conductivity and malleability of metals | SL, HL | "Explain why metals are malleable" (2 marks) |
| Relate the characteristic properties of metals to their uses | SL, HL | "Suggest why aluminium is used for overhead power cables" (2 marks) |
| Explain how the charge and radius of the metal ion set the strength of the metallic bond | SL, HL | "Explain why magnesium has a higher melting point than sodium" (3 marks) |
| Explain trends in melting points of s-block and p-block metals | SL, HL | Paper 1B: describe and explain a trend in a table or graph (2–3 marks) |
| Explain the high melting points and electrical conductivity of transition elements | HL only | "Explain why iron has a much higher melting point than potassium" (3 marks) |
Before you start
You need the idea from S2.1 that metal atoms lose their valence electrons to form cations, and that the attraction between opposite charges grows with charge and falls with distance. Figure 3 compares metals with ionic crystals, so the ionic lattice from S2.1 should be fresh.
1The idea in one paragraph
A metal atom holds its few valence electrons loosely. In a lump of metal those electrons leave their own atoms and spread through the whole structure, and what is left behind is a regular lattice of positive ions. The negative electrons, shared by every ion at once, hold the positive ions together: that attraction is the metallic bond. Because the electrons are free to move, metals conduct electricity and heat. Because the bond does not care which ion sits next to which, layers of ions can slide without the bond breaking, so metals bend instead of shattering. And because the strength of the attraction depends on the charge on the ions, how many electrons they give up and how closely they pack, the melting points of metals follow patterns you can predict.
2The metallic bond
A metallic bond is the electrostatic attraction between a lattice of cations and delocalized electrons.
Figure 1 draws the model. Each atom has given up its valence electrons and become a cation; the cations pack into a regular lattice. The electrons they released are delocalized: they are not tied to any one atom or any one bond, but move through the whole lattice. They are often called a "sea" of electrons.
Three things in that picture earn marks.
It is an electrostatic attraction, between the positive cations and the negative electron sea, just as the ionic bond is an attraction between opposite charges. What differs is that the negative charge is not in fixed anions but spread through the whole structure.
It is non-directional. Every cation is attracted to the electron sea around it in all directions. There is no bond between particular pairs of atoms, as there is in a covalent molecule.
The electrons are delocalized, not free atoms. The lattice is made of cations, not atoms. Writing "a lattice of metal atoms in a sea of electrons" loses the mark, because an atom that still had its valence electrons would have nothing to contribute to the sea.
3Electrical and thermal conductivity
Electrical conductivity. An electric current is a flow of charge. In a metal the charge carriers are the delocalized electrons. With no potential difference across a wire, they move at random in all directions and there is no overall flow. Connect a cell, and the electrons drift towards the positive terminal while new electrons enter from the negative end: a current flows. Figure 2 shows both situations.
The cations do not move. This is the contrast with a molten ionic compound in S2.1, where ions carry the current and the substance is changed chemically at the electrodes. A metal conducts in the solid and in the liquid state, and is not changed by conducting.
Thermal conductivity. When one end of a metal bar is heated, the delocalized electrons there gain kinetic energy. Because they move freely through the lattice, they carry that energy quickly to cooler parts. The closely packed cations also pass on vibrations to their neighbours. Both mechanisms operate, but the mobile electrons are the reason metals conduct heat so much better than non-metals such as wood or plastic, where only the vibration route exists.
4Malleability and ductility
A malleable material can be hammered or pressed into a new shape without breaking; a ductile one can be drawn out into a wire. Metals are both, and Figure 3 shows why by setting a metal beside an ionic crystal.
When a force pushes one layer of cations over another in a metal, the layer moves to a new position. But the new position is exactly like the old one: cations in a sea of delocalized electrons. The attraction acts in all directions and does not depend on which cation is beside which, so the bonding is not broken, only rearranged. The metal changes shape and stays in one piece.
In an ionic crystal the same small shift brings ions of the same charge side by side. They repel, and the crystal splits along that plane. The difference is non-directional bonding: metals have it without conditions, ionic lattices only as long as every ion stays surrounded by opposite charges.
5From properties to uses
The guide asks you to relate a metal's properties to what it is used for. The answer always has the same shape: name the property, explain it from the bonding if asked, and say why the use needs it.
| Metal | Use | The properties that matter |
|---|---|---|
| Copper | electrical wiring | excellent electrical conductivity; ductile, so drawn into wire |
| Aluminium | overhead power cables | good electrical conductivity with low density, so long spans are light; resists corrosion |
| Aluminium, copper, stainless steel | saucepans | high thermal conductivity; malleable, so pressed into shape; high melting point |
| Gold | electrical contacts in electronics | very good conductor that does not corrode |
| Iron, as steel | bridges, building frames | strong and can be shaped; cheap (the alloy is in S2.4) |
| Tungsten | old-style lamp filaments | the highest melting point of any metal |
Lustre, the shine of a freshly cut metal surface, also comes from the delocalized electrons, which absorb and re-emit light. Jewellery uses it. It is not one of the three properties the guide asks you to explain, so it will not carry marks on its own.
A good use answer weighs more than one property. Copper conducts better than aluminium, but aluminium is chosen for overhead cables because its low density keeps the weight on the pylons down. Say that, and you have shown the examiner you know a use is a trade-off.
6How strong is a metallic bond?
The metallic bond is an attraction between positive ions and negative electrons, so the same two factors that set the strength of an ionic lattice set this one too.
A metallic bond is stronger when the cation has a higher charge and a smaller radius.
Charge. A metal that gives up more valence electrons forms a cation with a higher charge and puts more electrons into the sea. Sodium gives one electron per atom, magnesium two, aluminium three. More charge on the cation and a denser sea of electrons mean a stronger attraction. The guide calls this "a simple treatment in terms of charge of cations and electron density", and that is all you need.
Radius. A smaller cation lets the delocalized electrons come closer to its nucleus, so the attraction is stronger.
The melting point is the handiest measure of bond strength. To melt a metal, the cations must gain enough energy to move past each other, which means overcoming (in part) the attraction to the electron sea. Stronger bonding, higher melting point.
Down a group, melting points fall. Figure 4 shows group 1. Every element forms a 1+ ion and gives one electron to the sea, so charge is constant. But each ion down the group has an extra shell and is larger, from Li⁺ at 76 pm to Cs⁺ at 167 pm. The delocalized electrons are further from the nuclei, the attraction is weaker, and the melting point falls from 181 °C for lithium to 28 °C for caesium, which would melt in your hand.
Across a period, melting points rise from group 1 to group 13. Figure 5 puts sodium, magnesium and aluminium side by side. Going across, the cation's charge rises from 1+ to 2+ to 3+, each atom contributes one, two, then three electrons to the sea, and the ions get smaller, 102 pm to 72 pm to 54 pm. All three changes strengthen the bond, and the melting point rises from 98 °C to 650 °C to 660 °C.
Look at the last two numbers, though. The simple model predicts a large step from magnesium to aluminium, and the real step is 10 °C. The model gets the order right, which is what the exam asks for, but not the size, because melting also depends on how the particular lattice is packed and how much of the bonding survives in the liquid. That gap is a fair example of what a model is: a picture good enough to predict a direction, and honest about where it stops. Beyond group 13 the period 3 elements are not metals at all, so the trend does not continue; Structure 3.1 picks up that story.
Write the explanation with three links in the chain. For magnesium against sodium: Mg²⁺ has a higher charge than Na⁺ and is smaller, and each magnesium atom delocalizes two electrons rather than one; so the electrostatic attraction between the cations and the delocalized electrons is stronger; so more energy is needed to overcome it, and the melting point is higher.
7HLTransition elements: delocalized d-electrons
SL students can skip to section 8.
Transition elements, in the d-block, have much higher melting points than the s-block metals of the same period, and they are good electrical conductors. Figure 6 shows the jump across period 4.
The reason is that the 3d and 4s sub-levels of a transition element are very close in energy, so a transition metal atom can release electrons from both into the sea, not just its 4s electrons. Potassium delocalizes one electron per atom and calcium two; titanium, iron and their neighbours delocalize more. More delocalized electrons, attracted to cations that are also fairly small and highly charged, make a much stronger metallic bond: iron melts at 1538 °C against 842 °C for calcium and 63 °C for potassium. The same large number of mobile electrons makes the transition metals good electrical conductors.
The rise across the d-block is not smooth, and the guide asks why the trend is "less evident" there. Across the s- and p-block metals each step adds one more delocalized electron. Across the d-block the number of electrons that are actually delocalized does not keep rising with the number of d electrons, and the lattice structure changes from element to element, so melting points level off, mostly between about 1250 °C and 1900 °C, and fall away at copper. At zinc, with a full, stable 3d¹⁰ sub-level, the d electrons are no longer delocalized, and the melting point drops back to 420 °C.
The chemical properties of the transition elements, their coloured compounds, variable oxidation states and catalytic activity, belong to Reactivity 3.4, not here.
8Where marks are lost
"A lattice of metal atoms." The lattice is of cations. The atoms have given up their valence electrons to the sea.
"Free electrons" with nothing more. Say the electrons are delocalized and move through the lattice when a potential difference is applied. "Free" alone is usually accepted, but "free ions" is wrong: ions carry the current in molten ionic compounds, never in metals.
Explaining malleability by "weak bonds". Metallic bonds are strong. Metals are malleable because the bonding is non-directional, so the layers can slide without the bonding being broken.
Using radius alone across a period. Across period 3 the charge on the cation and the number of delocalized electrons rise; that is the main point. Radius is the supporting point.
Calling the metallic bond "intermolecular". Metals contain no molecules. Melting a metal overcomes metallic bonding, not intermolecular forces.
Giving a use with no property, or a property with no use. "Copper is used in wires because it is a metal" scores nothing. Name the property that the use needs.
HL: saying transition metals have "more electrons" without saying which. The point is that both 3d and 4s electrons can be delocalized.
9Draw it right
- Draw the cations as circles with a charge, in a regular arrangement. Label them "cations" or "positive ions", not "atoms".
- Draw the delocalized electrons between the cations, spread through the whole structure, and label them "delocalized electrons".
- For a comparison of two metals, show the difference that matters: a higher charge on the cation, more electrons in the sea, a smaller cation.
- For malleability, draw before and after a layer has moved, and show that the arrangement after is still cations in an electron sea.
- For conduction, show the electrons moving towards the positive terminal; the cations stay put.
10Try it
Marks in brackets. Answers and marker's notes are at the end.
Q1. Which statement describes metallic bonding? 1 mark
A. Attraction between positive and negative ions in a lattice
B. Attraction between a lattice of cations and delocalized electrons
C. A shared pair of electrons between two metal nuclei
D. Attraction between metal atoms and free protons
Q2. Explain why metals conduct electricity and why they are malleable. 4 marks
Q3. The table shows the melting points of three group 2 metals and the radii of their ions. 3 marks
| Metal | Radius of the 2+ ion / pm | Melting point / °C |
|---|---|---|
| calcium | 100 | 842 |
| strontium | 118 | 777 |
| barium | 135 | 727 |
(a) Describe the relationship between ionic radius and melting point in the table. 1 mark
(b) Explain the relationship. 2 marks
Q4. Explain why magnesium has a higher melting point than sodium. 3 marks
Q5. Copper is a better electrical conductor than aluminium, yet overhead power cables are usually made mainly of aluminium. Suggest two reasons. 2 marks
Q6 (HL). Explain why iron has a much higher melting point than potassium, and why iron is a good electrical conductor. 3 marks
11In one breath
A metal is a regular lattice of cations held together by their electrostatic attraction to a sea of delocalized electrons, the valence electrons every atom gave up; the bond acts in all directions. The delocalized electrons move when a potential difference is applied, so metals conduct electricity, solid or liquid, without changing; they also carry kinetic energy through the lattice, so metals conduct heat. The bonding is non-directional, so layers of cations can slide into new positions that are just as well bonded, which makes metals malleable and ductile, while an ionic crystal shatters when like charges meet. Uses follow from these properties: copper for wiring, aluminium for light overhead cables, metal pans for heat. The bond is stronger with a higher cation charge, more delocalized electrons and a smaller cation, so melting points fall down group 1 as the ions grow and rise from sodium to magnesium to aluminium as the charge grows. HL: transition elements delocalize both 3d and 4s electrons, so their metallic bonding is much stronger, their melting points much higher and their conductivity good.
Answers
Q1. B. A describes ionic bonding and C covalent bonding; D has the charges on the wrong particles. B only.
Q2. A metal contains a lattice of cations surrounded by delocalized electrons. When a potential difference is applied, the delocalized electrons move through the lattice towards the positive terminal, carrying charge, so the metal conducts. The metallic bonding is non-directional: the attraction between the cations and the electron sea does not depend on which cation is next to which. So when a force is applied, layers of cations can slide over each other and the bonding is not broken, and the metal changes shape instead of breaking. 1 for delocalized electrons, 1 for the electrons moving and carrying charge, 1 for non-directional bonding or attraction in all directions, 1 for layers sliding without the bonding breaking. "Free ions" scores 0 for the conduction marks.
Q3. (a) As the radius of the ion increases, the melting point decreases. (b) All three form 2+ ions and delocalize two electrons per atom, so the difference is size. A larger cation means the delocalized electrons are further from the nucleus, so the electrostatic attraction between the cations and the delocalized electrons is weaker, and less energy is needed to overcome it. 1 for (a); in (b) 1 for the same charge with larger ions, 1 for weaker attraction between cations and delocalized electrons needing less energy. Mention of intermolecular forces scores 0 for (b).
Q4. Magnesium forms Mg²⁺ ions, sodium forms Na⁺, so magnesium's cations have a higher charge and each atom contributes two delocalized electrons rather than one. Mg²⁺ is also smaller than Na⁺. The electrostatic attraction between the cations and the delocalized electrons is therefore stronger in magnesium, and more energy is needed to overcome it, so the melting point is higher. 1 for higher charge or more delocalized electrons, 1 for a smaller ion, 1 for stronger attraction needing more energy. Stated without the link to the attraction between cations and electrons, capped at 2.
Q5. Aluminium has a much lower density than copper, so a long cable is much lighter and the pylons can be further apart. Aluminium is also cheaper, and it resists corrosion because of a protective oxide layer. It is still a good conductor. 1 for low density with the consequence, 1 for a second valid reason such as cost or corrosion resistance. "Aluminium is a better conductor" scores 0, since the question says it is not.
Q6 (HL). In a transition element such as iron, the 3d and 4s electrons are close in energy, so both 3d and 4s electrons are delocalized, whereas potassium delocalizes only its one 4s electron. Iron therefore has many more delocalized electrons and smaller, more highly charged cations, so the metallic bonding is much stronger and more energy is needed to melt it. The large number of delocalized electrons that can move through the lattice also makes iron a good electrical conductor. 1 for both 3d and 4s electrons delocalized, 1 for stronger attraction and more energy to melt, 1 for many mobile delocalized electrons explaining conductivity.
Educerie · written from the published IB Diploma Programme Chemistry guide, first assessment 2025, section S2.3 The metallic model. Original text, examples and questions. Diagrams drawn by Educerie. Last reviewed 25 September 2026.