The shape of the lesson
By the end of the lesson, learners can
- Diamond and graphite: Switch between the two carbon structures in the simulator.
- Silicon(IV) oxide: Extended.
- Metallic bonding: Extended.
- Diamond: 4 bonds per C, no free electrons: hard, non-conductor.
- Graphite: 3 bonds per C, weak forces between layers, delocalised electrons.
- SiO2 (Extended): 4 O per Si, 2 Si per O, like diamond.
- Metals (Extended): positive ions in a sea of delocalised electrons.
One picture every idea returns to
In a giant covalent structure a huge number of atoms are joined by strong covalent bonds in a continuous network. In diamond each carbon atom is bonded to four others, so it is very hard, has a very high melting point and does not conduct. In graphite each carbon atom is bonded to three others in flat layers of hexagons; there are weak forces between the layers and one delocalised electron per atom, so graphite is soft, slippery and conducts. (Extended) Metallic bonding is the electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons.
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, Bonded all the way through.. Diamond and graphite are both pure carbon, yet one cuts glass and the other slides off onto paper. 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: Diamond and graphite | Switch between the two carbon structures in the simulator. Project tab 5 of the lab (States of matter lab, carbon allotropes) and let a student drive it while the class predicts. | Whole class, then pairs |
| 13 to 21 min | Teach and check: Silicon(IV) oxide | Extended. | Pairs |
| 21 to 30 min | Teach and check: Metallic bonding | Extended. Project tab 4 of the lab (Chemical bonding lab, metallic) and let a student drive it while the class predicts. | Cooperative Concept Mapping |
| 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, 16 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 Metallic bonding, 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.
Cooperative Concept Mapping
Teams receive the key terms on cards and arrange them on a large sheet, drawing and labelling the links. Compare maps between teams before showing the model map.
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 |
|---|---|
| Graphite is soft because the forces between its layers are weak, not because its covalent bonds are weak. | The bonds within each layer are very strong, which is why graphite still has a very high melting point. 'Graphite has weak covalent bonds' is a classic lost mark. |
| Metals conduct because delocalised electrons move, not ions. | In a solid metal the positive ions stay in the lattice. Mixing this up with ionic compounds, which conduct only when their ions are free to move, is a common error. |
Giant covalent structures and metallic bonding: worksheet
6 questions, 16 marks. Written in the style of the papers; no past paper question is reproduced.
Describe the structure and bonding in diamond.
Use the structure of graphite to explain why it is used as a lubricant and as an electrode.
Explain why diamond does not conduct electricity but graphite does.
(Extended) In silicon(IV) oxide each silicon atom is bonded to four oxygen atoms and each oxygen atom to two silicon atoms. Explain how this gives the formula SiO2, and predict two physical properties.
(Extended) Describe metallic bonding.
(Extended) Copper is used for electrical wiring. Use its metallic bonding to explain why copper conducts electricity and why it can be drawn into wires.
Answer key and marking notes
- A giant covalent structure of carbon atoms.
- Each carbon atom is covalently bonded to four other carbon atoms (tetrahedrally).
- Graphite has layers of hexagons in which each carbon atom is covalently bonded to three others.
- The forces between the layers are weak, so the layers slide over each other: a lubricant.
- Each carbon atom has one delocalised electron,
- which can move through the structure and carry charge, so graphite conducts: an electrode.
- In diamond all four outer electrons of each carbon are used in covalent bonds, so there are no free electrons.
- In graphite each carbon uses only three electrons in bonding, leaving one delocalised electron per atom that can move.
- Each Si is surrounded by 4 O, but each O is shared between 2 Si, so there are 4 ÷ 2 = 2 O atoms for every Si atom: SiO2.
- Very high melting point (many strong covalent bonds to break).
- Hard, and does not conduct electricity (no free electrons or ions).
- A lattice of positive metal ions in a sea of delocalised electrons.
- Metallic bonding is the electrostatic attraction between the positive ions and the delocalised electrons.
- Copper contains delocalised electrons that are free to move through the structure and carry charge.
- The layers of positive ions can slide over each other,
- while the delocalised electrons keep attracting them, so the metallic bonding is not broken and the metal is ductile.
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.
A substance is extremely hard, has a very high melting point, and does not conduct electricity when solid or when molten. Which substance could it be?
- Adiamond
- Bgraphite
- Csodium chloride
- Dcopper
The diagram shows part of the structure of graphite: layers of carbon atoms arranged in rings of six. What do the red dashed lines between the layers represent?
- Aweak forces between the layers
- Bstrong covalent bonds between the layers
- Cionic bonds between carbon ions
- Dpaths of the delocalised electrons
Diamond and graphite are both made only of carbon atoms. Why is diamond hard but graphite soft?
- AIn diamond every atom is held by strong covalent bonds in all three dimensions; in graphite the layers are held together only by weak forces
- BThe covalent bonds in diamond are much stronger than the covalent bonds in graphite
- CDiamond contains ionic bonds but graphite contains covalent bonds
- DGraphite's delocalised electrons weaken the covalent bonds within each layer
Silicon(IV) oxide, SiO2, melts at about 1700 °C, but carbon dioxide, CO2, is a gas at room temperature. Which explains the difference?
- ASiO2 is a giant covalent structure, so melting breaks many strong covalent bonds; CO2 is made of simple molecules held by weak intermolecular forces
- BThe bonds in SiO2 are ionic, while the bonds in CO2 are covalent
- CCO2 molecules contain no covalent bonds
- DSiO2 molecules are heavier than CO2 molecules, so their intermolecular forces are stronger
Which pair of properties do diamond and silicon(IV) oxide share, and for what reason?
- ABoth are hard with very high melting points, because both are giant covalent structures with strong covalent bonds in three dimensions
- BBoth conduct electricity, because both contain delocalised electrons
- CBoth are soft and slippery, because both have layers that slide over each other
- DBoth have low melting points, because both are made of small molecules
Why do most metals have high melting points?
- AA lot of energy is needed to overcome the strong electrostatic attraction between the positive ions and the delocalised electrons
- BA lot of energy is needed to break the strong covalent bonds between the metal atoms
- CA lot of energy is needed to overcome the attraction between positive and negative metal ions
- DThe weak forces between metal molecules are hard to overcome
Solid copper conducts electricity, but solid copper(II) chloride does not. Why?
- ACopper has delocalised electrons that are free to move; in solid copper(II) chloride the ions are held in fixed positions
- BCopper ions move through solid copper, but copper(II) chloride contains no ions
- CSolid copper(II) chloride contains no charged particles at all
- DCopper atoms are joined by covalent bonds that carry the current
A substance melts at 1085 °C, conducts electricity both as a solid and as a liquid, and can be hammered into thin sheets. What is its structure and bonding?
- Ametallic: a lattice of positive ions in a sea of delocalised electrons
- Bionic: a giant lattice of positive and negative ions
- Cgiant covalent: a network of atoms joined by covalent bonds
- Dsimple molecular: small molecules held by weak intermolecular forces
Quiz key
Diamond is a giant covalent structure: many strong covalent bonds make it very hard with a very high melting point, and every outer electron is used in bonding, so it never conducts. Graphite conducts because of its delocalised electrons and is soft. Sodium chloride conducts when molten, because its ions can then move. Copper conducts when solid.
Within each layer, every carbon atom forms strong covalent bonds to three others (the solid lines). Between the layers there are only weak forces, shown by the dashed lines, so the layers can slide over each other: graphite is soft and slippery. If the layers were held by strong covalent bonds, graphite would be hard like diamond. Graphite contains no ions, and the delocalised electrons move along the layers rather than being fixed lines between them.
Diamond has a 3D network of strong covalent bonds, each carbon bonded to four others, so there is no easy direction for it to break or deform. Graphite has strong bonds only within its layers; the layers are held by weak forces and slide easily, making it soft. The bonds within graphite's layers are just as strong as those in diamond (in fact slightly stronger), neither substance is ionic, and the delocalised electrons do not weaken the layers.
In silicon(IV) oxide each Si is bonded to four O atoms and each O to two Si atoms in a continuous giant network, so melting means breaking huge numbers of strong covalent bonds. CO2 consists of separate small molecules; only the weak forces between them are overcome when it changes state. Both compounds are covalent (non-metal with non-metal). SiO2 has no separate molecules, so comparing molecular masses does not apply.
In both substances every atom is joined to its neighbours by strong covalent bonds throughout a 3D network (each C to 4 C in diamond; each Si to 4 O and each O to 2 Si in silicon(IV) oxide). This makes both hard with very high melting points. Neither has delocalised electrons, so neither conducts. Sliding layers describe graphite, and low melting points describe simple molecular substances.
In a metal, positive ions sit in a lattice surrounded by a sea of delocalised electrons. The electrostatic attraction between the ions and the electrons is strong and acts throughout the structure, so a lot of energy is needed to melt the metal. Metals are not held together by shared pairs (covalent), they contain no negative ions, and they are not made of molecules.
In copper the delocalised electrons can move through the lattice of positive ions and carry charge, so the solid conducts. Copper(II) chloride is ionic: it does contain charged particles (Cu2+ and Cl− ions), but in the solid they are fixed in the lattice and cannot move, so it only conducts when molten or dissolved. In a metal it is the electrons, not the positive ions, that move, and metallic bonding is not covalent.
Conducting as a solid needs mobile charge carriers in the solid, and being malleable means layers can slide without the bonding breaking. Both point to metallic bonding: delocalised electrons carry the current and keep attracting the positive ions as the layers slide (this substance is copper). Ionic solids do not conduct and are brittle; giant covalent substances (apart from graphite) do not conduct and are not malleable; simple molecular substances have low melting points.
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.