A bond with nobody to give and nobody to take
Every bond so far has had a clear story. In sodium chloride, sodium gives an electron and chlorine takes it. Opposite ions attract. In a chlorine molecule, two atoms each put one electron into a shared pair. The pair holds them together.
Now look at a block of pure sodium. Every atom is the same. No atom wants to take an electron from its neighbour, because sodium atoms do not attract extra electrons strongly. And a sodium atom has only one outer electron, so it cannot share pairs with all the neighbours around it. Neither story works. Yet sodium is a solid, and some metals are among the hardest, toughest solids we use.
The lectures on intermolecular forces ended with a promise: metallic bonding is a different kind of attraction altogether, not between separate molecules, but throughout a whole piece of metal. This lecture explains what that attraction is, and why it is strong in some metals and weak in others.
Electrons that belong to the whole metal
Start with what metal atoms are like. Most metal atoms have only one, two or three electrons in their outer shell. The lectures on ionisation energy showed that these outer electrons are held quite loosely. Little energy is needed to remove them.
Now pack many metal atoms together in a solid. Each atom has neighbours on every side, close and touching. Its loose outer electron is no closer to its own nucleus than to the nuclei around it. So that electron is no longer tied to one atom. It can move from the region around one atom to the next, and the next, through the whole piece of metal.
Electrons that are free to move through the whole structure are called delocalised electrons. "Localised" means fixed in one place; "delocalised" means spread out, belonging to no single atom.
When an atom's outer electrons become delocalised, what is left behind is a positive ion. In magnesium, each atom gives two electrons to the shared pool, so the ions are Mg²⁺. These ions sit in a regular, repeating pattern, a giant metallic lattice, and the delocalised electrons move through the spaces between them. Chemists often call this the "sea of electrons". It is a picture, not a claim that electrons behave like water. It means only that the electrons are spread through the whole structure and free to move.
What holds this together? Every positive ion is attracted to the negative electrons around it. Every electron is attracted to the positive ions around it. No ion is paired with a particular electron. The attraction acts everywhere at once.
Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and the delocalised electrons around them.
That is it. That is the whole idea. The glue in a metal is the attraction of positive ions to a shared pool of electrons. Ionic bonding has two kinds of ion. Covalent bonding has shared pairs between two atoms. Metallic bonding has one kind of ion and electrons that belong to all of them.
A piece of magnesium is made of positive Mg²⁺ ions. Is the piece of metal positively charged?
Show the answer
No. Every Mg²⁺ ion was made by an atom giving two electrons to the pool, and those electrons are still in the metal. In the diagram, 15 ions carry a charge of 15 × 2 = +30, and there are 30 electrons, a charge of −30. The total is zero. A piece of metal is electrically neutral. The electrons have not left the metal. They have only stopped belonging to one atom.
Two features of this bond matter for the next lecture. First, the electrons can move, so they can carry an electric current. Second, the bond has no particular direction. A covalent bond points from one atom to one other atom. A metallic bond is the pull of the whole pool on each ion, the same in every direction. That is why layers of ions can slide past each other without the metal shattering. The next lecture uses both features to explain the properties of metals.
Before going on, check how metallic bonding compares with the ionic bonding you met earlier.
Put each statement in a box: true only of metallic bonding, true only of ionic bonding, or true of both. Every card explains itself when you place it.
Ten statements, sorted.
| Statement | True of |
|---|---|
| Positive ions are arranged in a regular lattice. | Both |
| It is held together by electrostatic attraction between opposite charges. | Both |
| The attraction acts in all directions, not along particular bonds. | Both |
| The bonding is stronger when the positive ions have a higher charge. | Both |
| Negative ions sit in the lattice. | Ionic only |
| It forms when electrons transfer from one element to another. | Ionic only |
| It conducts electricity only when melted or dissolved. | Ionic only |
| Electrons are free to move through the whole solid. | Metallic only |
| It conducts electricity as a solid. | Metallic only |
| It can hold atoms of a single element together. | Metallic only |
The rule: both are giant structures held by attraction between opposite charges. The difference is what carries the negative charge: fixed negative ions in an ionic lattice, a mobile pool of electrons in a metal.
What makes a metallic bond stronger?
Metallic bonds are not all equally strong. Mercury is a liquid at room temperature, and tungsten melts only at 3422 °C. The model explains why, with the same question you asked about ionic lattices: how strongly do the positive ions and the electrons pull on each other?
Three things make the pull stronger.
- A higher charge on the ions. A 2+ ion pulls on electrons more strongly than a 1+ ion.
- More delocalised electrons per atom. If each atom gives two electrons instead of one, the pool is twice as crowded with negative charge, so each ion is surrounded by more of it.
- Smaller ions. In a small ion, the nucleus is closer to the electrons in the pool, so the attraction is stronger.
The first two always go together: an atom that gives two electrons becomes a 2+ ion. Look at both in a worked example. Melting and boiling points are taken from one table (see the note under the data).
Sodium melts at 97.8 °C and boils at 883 °C. Magnesium, next to it in Period 3, melts at 650 °C and boils at 1090 °C. Explain the difference.
Charge and electrons. Sodium is in Group 1. Each atom gives one electron to the pool, leaving Na⁺. Magnesium is in Group 2. Each atom gives two electrons, leaving Mg²⁺. So magnesium has ions with twice the charge and a pool with twice as many electrons per ion.
Size. Both are in Period 3, so their atoms have the same number of shells. Magnesium has one more proton, which pulls the electrons in more tightly. And once its two outer electrons are in the pool, the Mg²⁺ ion that remains is smaller than Na⁺.
Conclusion. Higher charge, more delocalised electrons and smaller ions all make the attraction stronger in magnesium. More energy is needed to separate its ions, so its melting and boiling points are higher.
Notice the shape of the answer. It names the particles (ions and delocalised electrons), the attraction between them, and why that attraction is stronger. Then it links the strength to the energy needed. Examiners look for exactly that chain.
Down a group: bigger is weaker
Now a prediction. It tests a habit you may have picked up recently.
Lithium and caesium are both in Group 1, so each atom gives one electron to the pool. A lithium atom has 3 electrons in total, and a caesium atom has 55. Which metal has the higher melting point?
Melting points down Group 1 fall steadily, from lithium at 180.5 °C to caesium at 28.5 °C.
This is an important contrast. Between molecules, a bigger molecule has more electrons that can form temporary dipoles, so the London forces get stronger. In a metal, only the outer electrons join the pool, and their number does not change down a group. What changes is the size. Each step down adds a shell, so the ion gets larger and the pool is further from the nucleus. The attraction weakens, and the melting point falls.
Use the data to describe and explain the trend.
| Metal | Melting point (°C) | Boiling point (°C) |
|---|---|---|
| Lithium | 180.5 | 1342 |
| Sodium | 97.8 | 883 |
| Potassium | 63.5 | 759 |
| Rubidium | 39.3 | 688 |
| Caesium | 28.5 | 671 |
Trend. Both melting and boiling points fall down the group. The biggest drop is between lithium and sodium. The drops then get smaller.
Explanation. Every atom gives one electron to the pool, so every ion is 1+. Down the group, each atom has one more shell, so the ion is larger. The pool of electrons is further from the nucleus, and the attraction between the ions and the pool is weaker. Less energy is needed to separate the ions.
Check. The explanation predicts a fall, and every value falls. It does not predict the exact sizes of the drops. That needs more detail than this model has.
Data note: melting points are the recommended values from the Wikipedia data page "Melting points of the elements", and boiling points are from the WebElements row of "Boiling points of the elements". Other tables can differ by a few degrees.
Iron and mercury: two ends of the scale
Iron melts at 1538 °C and boils at 2861 °C. Mercury melts at −38.83 °C and boils at 356.73 °C. What does the model say?
Iron is a transition metal, in the d-block. As well as its two outer 4s electrons, some of its 3d electrons also take part in the delocalised pool. More electrons in the pool, held by small, highly charged ions, give strong metallic bonding and a high melting point. (How strong iron is as a material, and why steel is stronger still, depends on how the layers of ions slide. That belongs to the next lecture.)
Mercury is the opposite case. Its outer 6s electrons are held unusually tightly, so they are reluctant to join a shared pool. The bonding between mercury atoms is weak, and room temperature gives enough energy to let the atoms move past one another. The full reason involves physics beyond this level. In 2013, F. Calvo and colleagues used computer simulations to test it. Without the effects of Einstein's theory of relativity, the simulated melting point was about 355 K. With relativity included, it was about 250 K, close to the measured 234 K (−39 °C). In a heavy atom such as mercury, relativity makes the electrons near the nucleus move very fast, and the effect pulls the 6s electrons in closer.
The model is the same at both ends of the scale. What changes is how strongly each metal's ions hold the pool of electrons.
When a metal melts, are its metallic bonds broken?
Show the answer
Not completely. Liquid mercury and molten iron still conduct electricity, so the delocalised electrons are still there, and the ions are still attracted to them. Melting lets the ions move past one another while staying in the pool. Only boiling pulls the atoms fully apart into a gas. This is why boiling points often follow the strength of metallic bonding more closely than melting points do: melting also depends on how the lattice is arranged.
Where the model came from
The idea of electrons moving freely inside a metal is surprisingly old. J. J. Thomson discovered the electron in 1897. Only three years later, in 1900, the German physicist Paul Drude proposed a model of metals. In it, the electrons move freely between fixed positive ions, like the molecules of a gas in a box. His model explained why metals obey Ohm's law, and (partly by a lucky cancelling of errors) why good conductors of electricity are also good conductors of heat.
It also got something badly wrong. If the electrons behaved like a gas, heating a metal should take far more energy than it really does, about a hundred times more for the electrons' share. The puzzle was solved in 1927, when Arnold Sommerfeld rebuilt the model using the new quantum mechanics. The electrons are still delocalised and free to move, but they obey quantum rules, and only a few of them can take up extra energy. Today the electrons in a metal are described as waves spread through the whole crystal. The simple picture you have learned, positive ions in a pool of mobile electrons, is the part of Drude's idea that survived. It is enough to explain everything at this level.
Three classic mistakes
A metal has no molecules. It is a giant structure held by metallic bonding throughout. Write about the attraction between positive ions and delocalised electrons, never about intermolecular forces.
That describes ionic bonding. A metal contains only positive ions. The negative charge is carried by delocalised electrons. The examiner's words are "positive ions" and "delocalised electrons".
That rule is for London forces between molecules. In metals, count only the electrons each atom gives to the pool, and look at the size of the ion. Caesium has 55 electrons, but it gives only one to the pool, and its large ion holds that pool weakly.
Work each one on paper first. Open a hint only when you are stuck.
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Calcium is in Group 2. In a small piece of calcium containing 1000 atoms, what ions are present, and how many delocalised electrons are there? What is the total charge of the piece?
Hint 1
A Group 2 atom has two outer electrons.
Hint 2
Each atom gives both of its outer electrons to the pool.
Full answer
1000 Ca²⁺ ions and 1000 × 2 = 2000 delocalised electrons. The ions carry +2000 and the electrons −2000, so the total charge is zero. Check: the electrons came from the atoms and stayed in the metal, so the metal must be neutral.
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Potassium (Group 1) and calcium (Group 2) are next to each other in Period 4. Predict which has the higher melting and boiling points, and explain. Then check your prediction: potassium melts at 63.5 °C and boils at 759 °C; calcium melts at 842 °C and boils at 1484 °C.
Hint 1
Compare the charge on the ions and the number of delocalised electrons per atom.
Hint 2
Both atoms have four shells, but calcium has one more proton.
Full answer
Calcium. Each calcium atom gives two electrons to the pool, forming Ca²⁺, while potassium gives one, forming K⁺. Calcium's ions have twice the charge, the pool holds twice as many electrons per ion, and Ca²⁺ is smaller than K⁺ because of the extra proton. So the attraction between ions and delocalised electrons is stronger in calcium, and more energy is needed to separate the ions. The data agree: calcium melts 778.5 °C higher and boils 725 °C higher.
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Caesium melts at 28.5 °C. A sealed glass tube of caesium held in a warm hand (about 37 °C) slowly turns liquid. Lithium in the same tube does not. Explain both, using the model.
Hint 1
Compare each melting point with 37 °C.
Hint 2
Then explain why caesium's metallic bonding is so weak compared with lithium's.
Full answer
37 °C is above caesium's melting point (28.5 °C) but far below lithium's (180.5 °C). Both metals give one electron per atom to the pool, so both have 1+ ions. But Cs⁺ has five shells and Li⁺ has only one, so in caesium the pool is much further from the nucleus. The attraction between ions and delocalised electrons is much weaker, so a hand's warmth gives enough energy for the ions to move past one another. (The tube is sealed because caesium reacts violently with air and water.)
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A student says: "Aluminium forms Al³⁺ ions, so its metallic bonding must be much stronger than magnesium's. Its melting point should be far higher." The data: magnesium melts at 650 °C and boils at 1090 °C; aluminium melts at 660 °C and boils at 2519 °C. Use the data to judge the claim.
Hint 1
Compare the two differences: 660 − 650 and 2519 − 1090.
Hint 2
Think about what melting needs and what boiling needs. Look back at the second Stop-and-think.
Full answer
The melting points differ by only 10 °C, but the boiling points differ by 1429 °C. The boiling points support the student's idea: with three electrons per atom in the pool and small Al³⁺ ions, aluminium's bonding is much stronger, so pulling its atoms fully apart takes much more energy. The melting points do not support it, because melting only lets the ions slide past each other while the pool remains. How easily that happens depends on how the lattice is arranged, not only on the strength of the bonding. So the claim is half right: aluminium's metallic bonding is stronger, but "its melting point should be far higher" does not follow. Boiling points are the better guide to bond strength.
You have now met three ways that atoms hold together, and all three are electrical. In an ionic lattice, positive and negative ions attract. In a covalent bond, two nuclei attract a shared pair of electrons between them. In a metal, a lattice of positive ions attracts a pool of electrons that belongs to all of them. Each is the same force, the attraction between opposite charges, arranged in a different way.
The arrangement decides the behaviour. Because a metal's electrons are free to move, and because its bonding has no direction, metals conduct, shine, bend and can be drawn into wires. The next lecture follows each of those properties back to the pool of electrons.
- In a metal, the outer electrons of each atom become delocalised: they belong to the whole structure, not to one atom.
- Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and the delocalised electrons around them.
- A piece of metal is electrically neutral, because the delocalised electrons balance the charge on the ions.
- Metallic bonding is stronger when the ions have a higher charge, give more electrons to the pool, and are smaller.
- Down Group 1 the ions get larger, so the bonding weakens and melting points fall, the opposite of the trend for London forces.
- Metallic bonding has no particular direction, which, with the mobile electrons, explains the properties of metals.