The shape of the lesson
By the end of the lesson, learners can
- Ions in a sea: Choose sodium, magnesium and then aluminium in the simulator.
- Properties from the sea: Every classic metal property traces back to electrons that are free to move and bonding that has no fixed direction.
- Metals and ionic solids compared: Both are giant lattices held by electrostatic attraction, so students often mix them up.
- Definition: attraction between positive ions and delocalised electrons.
- Stronger with more electrons per atom, higher charge, smaller ions.
- Conducts: delocalised electrons move.
- Malleable: layers slide, bonding is non-directional.
One picture every idea returns to
Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons. The outer electrons of each atom are released into a “sea” that belongs to the whole giant lattice. The bonding gets stronger as more electrons per atom are delocalised, as the ionic charge rises and as the ions get smaller, so melting points rise from sodium to magnesium to aluminium. The mobile electrons carry current and heat, and because the bonding is non-directional the layers of ions can slide without the lattice breaking.
Forty-five minutes, phase by phase
| Time | Phase | What happens in the room | Grouping |
|---|---|---|---|
| 0 to 5 min | Hook: predict | Open with the question in the title, A sea of electrons.. Metal atoms give up their outer electrons, but not to anyone in particular. Put the figure on the board: it is the one picture every idea in the lesson returns to. | Whole class, then pairs |
| 5 to 13 min | Teach and check: Ions in a sea | Choose sodium, magnesium and then aluminium in the simulator. Project tab 4 of the lab (Chemical bonding lab, metallic bonding) and let a student drive it while the class predicts. | Whole class, then pairs |
| 13 to 21 min | Teach and check: Properties from the sea | Every classic metal property traces back to electrons that are free to move and bonding that has no fixed direction. | Pairs |
| 21 to 30 min | Teach and check: Metals and ionic solids compared | Both are giant lattices held by electrostatic attraction, so students often mix them up. Project tab 8 of the lab (Chemical bonding lab, structures and properties) and let a student drive it while the class predicts. | Mix-Pair-Share |
| 30 to 36 min | Check | The four Stage 1 questions from the student lesson, answered on mini whiteboards or by show of fingers. Re-teach the one most get wrong. | Individual, then whole class |
| 36 to 43 min | Plenary | The exam-style questions from Stage 2, one at a time: thirty seconds alone, then a random call. The feedback lines in the key give the examiner's wording. | Individual, random call |
| 43 to 45 min | Exit and homework | Set the worksheet below (6 questions, 14 marks). Learners self-assess against the objectives before they leave. | Individual |
Protect the parts that carry the learning
The sequence carries about fifty minutes of material, so the lab is driven by one student while the class predicts rather than explored freely. Times are a guide.
Protected: the check at 30 minutes. It is the only individual read on whether the key idea landed, and it is not cut.
Cut first if the build overruns: the last check in Metals and ionic solids compared, never the check or the plenary.
In a 60 minute block: run worksheet questions 1 to 3 in class as a Rally Coach set and model both the hardest worked example and its check live.
Mix-Pair-Share
Students mix around the room until you call pair; they pair with the nearest classmate, discuss the prompt, then mix again. Three prompts of rising difficulty, three different partners.
The full structure, with its variations and the other twelve, is on the cooperative-learning page.
What to head off, and how
| Trap learners fall into | Teaching move that pre-empts it |
|---|---|
| Say positive ions in a metal lattice, not atoms, and describe the electrons as delocalised. | Answers that describe “metal atoms sharing electrons” or “ions attracting each other” do not score: the bond is ion to electron sea. |
Metallic bonding: worksheet
6 questions, 14 marks. Written in the style of the papers; no past paper question is reproduced.
Describe the structure and bonding in a metal.
Explain why magnesium has a much higher melting point than sodium.
Explain why the melting point of potassium (63 °C) is lower than that of sodium (98 °C).
Explain why aluminium is used for overhead power cables, referring to its bonding.
Explain why a metal can be bent without breaking, whereas an ionic solid shatters.
A student writes: “In a metal, the metal atoms share their electrons in covalent bonds.” Identify two errors in this statement.
Answer key and marking notes
- A giant lattice of positive metal ions surrounded by delocalised electrons.
- Metallic bonding is the electrostatic attraction between the positive ions and the delocalised electrons.
- Magnesium forms Mg2+ ions and contributes two delocalised electrons per atom; sodium forms Na+ and contributes one.
- Mg2+ is also smaller than Na+.
- The attraction between the ions and the delocalised electrons is stronger in magnesium, so more energy is needed to melt it.
- Both form 1+ ions and give one delocalised electron per atom.
- K+ is larger, so the attraction between the ions and the delocalised electrons is weaker.
- Aluminium has delocalised electrons that can move through the lattice, so it is a good electrical conductor.
- (It is also low density and can be drawn into wires because its layers of ions slide: it is ductile.)
- In a metal the layers of positive ions slide over each other, and the delocalised electrons still surround them, so the metallic bonding is maintained.
- In an ionic solid a shift brings ions of the same charge together.
- They repel, and the crystal breaks.
- The lattice is made of positive ions, not neutral atoms.
- The electrons are delocalised through the whole structure, not shared in localised covalent bonds between pairs of atoms.
8 multiple-choice questions on this lesson
From the topic question bank. Assign them in the app for a marked score in your gradebook, or run the topic as a class game.
Which statement best describes metallic bonding?
- Athe electrostatic attraction between positive metal ions and a sea of delocalised electrons
- Bthe electrostatic attraction between positive and negative ions in a lattice
- Cpairs of electrons shared between each metal atom and its neighbours
- Dthe attraction of each metal nucleus for its own outer-shell electrons only
The melting points of sodium, magnesium and aluminium are 98 °C, 650 °C and 660 °C. Which statement explains why the melting point of magnesium is higher than that of sodium?
- Aeach magnesium atom provides two delocalised electrons and forms a smaller 2+ ion, so the metallic bonding is stronger
- Bmagnesium atoms are larger than sodium atoms, so there is more contact between them
- Cmagnesium has more electron shells than sodium, so there are more delocalised electrons
- Dmagnesium ions have a lower charge than sodium ions, so they repel each other less
Why do solid metals conduct electricity?
- Athe delocalised electrons can move through the structure when a potential difference is applied
- Bthe positive metal ions are free to move through the lattice
- Cthe metal ions and electrons both move towards the positive terminal
- Dvibrating positive ions pass the charge from one ion to the next
Why are metals malleable (can be hammered into shape without breaking)?
- Alayers of positive ions can slide over each other while the delocalised electrons keep attracting them, so the bonding is not broken
- Bmetallic bonds are weak, so the atoms separate easily when hit
- Cthe electrons are fixed between particular pairs of ions, so the ions stay in place
- Dwhen the layers move, ions of the same charge come next to each other and repel
Which Group 1 metal has the highest melting point?
- Alithium
- Bsodium
- Cpotassium
- Drubidium
The chart shows the melting points of the first four elements in Period 3. Why is the melting point of silicon so much higher than that of aluminium?
- Asilicon has a giant covalent structure, and many strong covalent bonds must be broken to melt it
- Bsilicon provides four delocalised electrons per atom, so its metallic bonding is the strongest
- Csilicon forms Si4+ ions, which have the highest charge in the period
- Dsilicon molecules are larger, so they have stronger intermolecular forces
How many delocalised electrons are present in 2.43 g of magnesium? (Ar: Mg 24.3; L = 6.02 × 1023 mol−1)
- A1.20 × 1023
- B6.02 × 1022
- C3.01 × 1022
- D7.22 × 1023
Element Y conducts electricity as a solid and as a liquid. Its oxide does not conduct as a solid but does conduct when molten. What do these observations show?
- AY has metallic bonding with delocalised electrons in both states, and its oxide is ionic with ions that can only move in the liquid
- BY and its oxide both contain ions that are free to move in every state
- CY is ionic, and its oxide has a giant covalent structure
- DY has delocalised electrons only when molten, and its oxide releases electrons when it melts
Quiz key
Each metal atom releases its outer electrons, which become delocalised through the whole structure, leaving a lattice of positive ions; the bonding is the attraction between these ions and the delocalised electrons. There are no negative ions in a pure metal, so the second option describes ionic bonding. Localised shared pairs between neighbouring atoms would be covalent bonding. If each atom kept its own outer electrons, there would be no bonding between the atoms and no free electrons to conduct electricity.
Magnesium contributes two outer electrons per atom, so the sea is denser, and the Mg2+ ions are smaller and more highly charged than Na+. The attraction between the ions and the delocalised electrons is therefore stronger and more energy is needed to melt the metal. Magnesium atoms are smaller than sodium atoms, not larger, and bigger size would weaken the bonding. Both elements are in Period 3 and have three shells. Mg2+ has a higher charge than Na+, not a lower one.
The delocalised electrons are not tied to any particular ion, so they drift through the lattice towards the positive terminal, carrying the current. The positive ions stay in fixed positions in the solid; they only vibrate. Positive ions would move towards the negative terminal, not the positive one, and in any case they are fixed. Lattice vibrations carry heat, not electric charge.
The attraction between the ions and the sea of electrons does not depend on which particular ions are neighbours, so layers can slide into new positions and the metal changes shape without the structure falling apart. Metallic bonding is generally strong, as the high melting points of most metals show. Electrons fixed between pairs of ions would be localised covalent bonds, which would make the solid brittle. Like charges meeting and repelling is the reason ionic solids are brittle, not a property of metals.
Each Group 1 metal provides one delocalised electron per atom and forms a 1+ ion, so the difference is ion size. Li+ is the smallest, so the delocalised electrons are closest to the ion centres and the attraction is strongest (Li 181 °C, Na 98 °C, K 63 °C, Rb 39 °C). Sodium, potassium and rubidium have progressively larger ions and weaker metallic bonding. Choosing rubidium assumes that heavier atoms must have higher melting points, but mass does not decide bond strength.
Na, Mg and Al are metals, and their melting points rise as more electrons are delocalised. Silicon is not a metal: each atom forms four covalent bonds in a giant three-dimensional lattice (like diamond), and melting requires breaking many strong covalent bonds, so its melting point (1410 °C) is far higher. Silicon does not have a sea of delocalised electrons, which is why it is only a semiconductor. It does not form Si4+ ions. It is not made of simple molecules, so intermolecular forces are not what is overcome.
n(Mg) = 2.43 ÷ 24.3 = 0.100 mol. Each Mg atom gives two electrons to the sea, so there are 0.200 mol of delocalised electrons = 0.200 × 6.02 × 1023 = 1.20 × 1023. 6.02 × 1022 is the number of atoms, assuming only one delocalised electron per atom. 3.01 × 1022 divides the amount of magnesium by 2 instead of multiplying by 2. 7.22 × 1023 counts all 12 electrons in each atom, but the inner electrons are not delocalised.
A solid that conducts, and still conducts when molten, has delocalised electrons in both states, which is metallic bonding. A compound that conducts only when molten has ions that are locked in a lattice in the solid and become mobile in the liquid, which is ionic bonding. If the oxide had mobile ions in every state it would also conduct as a solid. An ionic element is impossible, since an element cannot contain two kinds of oppositely charged ions, and a giant covalent oxide would not conduct even when molten. Solid metals already conduct, so the electrons are delocalised before melting, and molten ionic compounds conduct by moving ions, not released electrons.
What to project
The student lesson embeds these tabs. Open them full screen on the projector; every slider has a drag handle on the stage.
Slides for this lesson
A projectable deck for every pack is on the way. Until then, present the student lesson full screen: its sections, figure and lab tabs follow the same order as the plan above.