C6: Covalent Bonding
Shared electrons, simple molecules, giant covalent structures, polymers, dot and cross diagrams and intermolecular forces
📋 Key Definitions
Covalent bond: A bond formed when two atoms share a pair of electrons, so that both atoms achieve a full outer shell. Covalent bonding occurs between non-metal atoms.
Molecule: A group of two or more atoms held together by covalent bonds. Molecules can be made of the same element (e.g. H₂, O₂) or different elements (e.g. H₂O, CO₂).
Intermolecular forces: Weak forces of attraction between molecules. These are much weaker than covalent bonds and are broken when simple molecular substances melt or boil.
Simple molecular substance: A substance made up of separate molecules, each containing a fixed number of atoms joined by covalent bonds. Examples include H₂O, CO₂ and CH₄.
Giant covalent structure: A structure in which all the atoms are bonded to each other by strong covalent bonds in a giant three-dimensional lattice. Examples include diamond and graphite.
✏️ Dot and Cross Diagrams for Covalent Bonding
In covalent bonding, atoms share pairs of electrons. Each shared pair counts towards the outer shell of both atoms. Dot and cross diagrams show the shared pair using one dot and one cross to show each electron comes from a different atom.
Worked Example - Hydrogen (H₂)
Each hydrogen atom has 1 electron in its outer shell and needs 2 for a full shell (first shell only holds 2).
The two H atoms share their single electrons, forming one covalent bond (a single bond).
H· + ·H → H⨉H (shared pair shown as one dot and one cross between the two H nuclei)
The shared pair gives each H atom a full outer shell of 2 electrons.
Worked Example - Water (H₂O)
Oxygen (2,6) needs 2 electrons to complete its outer shell. Each hydrogen (1) needs 1 electron.
Oxygen shares 2 of its electrons, one with each H atom, forming two single covalent bonds.
H⨉O⨉H (oxygen shares one pair with each hydrogen)
O has 4 non-bonding electrons (2 lone pairs) plus 4 shared electrons = 8 in outer shell. Each H has 2 shared electrons = full outer shell.
Worked Example - Methane (CH₄)
Carbon (2,4) needs 4 electrons to complete its outer shell. Each hydrogen (1) needs 1 electron.
Carbon shares one electron with each of 4 hydrogen atoms, forming four single covalent bonds.
H⨉C⨉H with H atoms above, below and to both sides of C
C has 8 shared electrons in its outer shell. Each H has 2 shared electrons. All atoms have full outer shells.
Worked Example - Carbon Dioxide (CO₂)
Carbon (2,4) needs 4 electrons. Each oxygen (2,6) needs 2 electrons.
Carbon shares two electrons with each oxygen, forming two double covalent bonds (C=O).
O⨉⨉C⨉⨉O (two shared pairs between C and each O = double bond)
C has 8 shared electrons. Each O has 4 shared electrons + 4 non-bonding (2 lone pairs) = 8 in outer shell.
🔬 Simple Molecular Substances
Simple molecular substances contain small molecules with strong covalent bonds within the molecules but weak intermolecular forces between the molecules.
Important distinction: When a simple molecular substance melts or boils, only the weak intermolecular forces between molecules are overcome. The strong covalent bonds within the molecules are NOT broken.
| Property | Simple molecular substance | Explanation |
|---|---|---|
| Low melting and boiling points | Yes | Only weak intermolecular forces between molecules need to be overcome; these require little energy |
| Do not conduct electricity | Generally yes | No free ions or electrons to carry charge; molecules are overall neutral |
| Usually gases or liquids at room temperature | Yes | Weak intermolecular forces mean little energy is needed to separate molecules |
| Often insoluble in water | Often | Many are non-polar and do not interact with water molecules |
When explaining low melting points of simple molecules, you must mention that intermolecular forces between molecules are weak and require little energy to overcome. Do NOT say covalent bonds are weak - they are very strong!
💎 Giant Covalent Structures
Giant covalent structures (also called macromolecular structures) contain many atoms bonded together in a giant lattice of strong covalent bonds. These structures have very different properties from simple molecular substances.
Diamond: Each carbon atom forms four covalent bonds with other carbon atoms in a rigid tetrahedral structure. This makes diamond extremely hard. All the covalent bonds must be broken to melt diamond, giving it a very high melting point. Diamond does not conduct electricity because there are no free electrons or ions.
Graphite: Each carbon atom forms three covalent bonds with other carbon atoms in layers of hexagonal rings. One electron from each carbon is delocalised and can move freely between the layers. This makes graphite a good conductor of electricity. The layers are held together by weak intermolecular forces, so they can slide over each other easily - this is why graphite is soft and slippery.
Graphene: A single layer of graphite - one atom thick sheet of carbon atoms in a hexagonal arrangement. It is extremely strong, lightweight, flexible and an excellent conductor of electricity and heat.
Fullerenes: Molecules of carbon shaped like hollow tubes or spheres. Buckminsterfullerene (C₆₀) has 60 carbon atoms arranged in a hollow sphere. Fullerenes can be used for drug delivery, lubricants and catalysts.
Carbon nanotubes: Cylindrical fullerenes with very high length-to-diameter ratios. They are very strong, good conductors of electricity and can be used in electronics and strengthening materials.
| Property | Diamond | Graphite |
|---|---|---|
| Bonding | 4 covalent bonds per C atom | 3 covalent bonds per C atom + 1 delocalised electron |
| Structure | Rigid 3D tetrahedral lattice | Layered hexagonal rings |
| Hardness | Very hard | Soft and slippery |
| Melting point | Very high | Very high |
| Electrical conductivity | Does not conduct | Conducts electricity |
| Reason for conductivity | No free electrons | Delocalised electrons can move between layers |
| Reason for hardness/softness | Strong covalent bonds in rigid 3D structure | Weak intermolecular forces between layers allow them to slide |
| Uses | Cutting tools, jewellery | Pencils, lubricants, electrodes |
When comparing diamond and graphite, make sure you link the structure to the property. For example: "Graphite conducts electricity because each carbon atom forms only 3 covalent bonds, leaving one delocalised electron that can move between the layers and carry charge."
🧬 Polymers
Polymers are very large molecules (macromolecules) made up of many small repeating units called monomers joined together by covalent bonds.
Polymer properties: Polymers are solid at room temperature because the intermolecular forces between polymer chains are stronger than between small simple molecules (due to the large size of the molecules). However, these forces are still weaker than covalent or ionic bonds, so polymers generally have lower melting points than ionic or giant covalent substances.
Common Polymers
- Poly(ethene) - made from ethene monomers, used in plastic bags
- Poly(propene) - made from propene monomers, used in food containers
- PVC (poly(chloroethene)) - used in window frames and pipes
- PTFE (poly(tetrafluoroethene)) - non-stick coating for pans
📊 Simple Molecular vs Giant Covalent
| Property | Simple Molecular | Giant Covalent |
|---|---|---|
| Examples | H₂O, CO₂, CH₄, Cl₂ | Diamond, graphite, silicon dioxide |
| Melting/boiling point | Low | Very high |
| What is overcome when melting | Weak intermolecular forces between molecules | Strong covalent bonds throughout the structure |
| Electrical conductivity | Does not conduct (no free charges) | Most do not conduct (diamond, SiO₂); graphite does conduct |
| State at room temperature | Usually gas or liquid, some low-melting solids | Solid |
| Strength | Soft or brittle | Very hard (diamond) or soft layers (graphite) |
The key difference: simple molecular substances have weak forces between molecules (low MP), while giant covalent structures have strong covalent bonds throughout the entire structure (very high MP). Always specify which type of force is being overcome.
❓ Practice Questions
Q1: Foundation Explain why carbon dioxide is a gas at room temperature while silicon dioxide (SiO₂) is a solid with a very high melting point.
CO₂ is a simple molecular substance. Weak intermolecular forces between the CO₂ molecules require little energy to overcome, so it is a gas at room temperature. SiO₂ is a giant covalent structure. Strong covalent bonds throughout the entire structure require a very large amount of energy to overcome, giving it a very high melting point.
Q2: Foundation Draw a dot and cross diagram for a molecule of hydrogen chloride (HCl). Explain how the covalent bond forms.
Hydrogen (1 electron) and chlorine (2,8,7) each need one more electron for a full outer shell. They share one pair of electrons - one from H (shown as a dot) and one from Cl (shown as a cross). H achieves 2 electrons in its outer shell and Cl achieves 8. The diagram shows H and Cl with one shared pair of electrons between them, Cl also has 3 lone pairs of electrons.
Q3: Higher Explain why graphite can conduct electricity but diamond cannot. Refer to the bonding in each substance.
In graphite, each carbon atom forms 3 covalent bonds with other carbon atoms, leaving one electron per carbon atom delocalised. These delocalised electrons are free to move between the layers and can carry charge, making graphite a conductor. In diamond, each carbon atom forms 4 covalent bonds. All four outer electrons are used in bonding, so there are no free or delocalised electrons to carry charge, so diamond does not conduct electricity.
Q4: Higher Describe the structure and bonding in graphene and explain two properties that result from its structure.
Graphene is a single layer of carbon atoms arranged in hexagonal rings. Each carbon atom forms 3 covalent bonds, with one delocalised electron per atom. Properties: (1) It is extremely strong because the covalent bonds in the 2D sheet are very strong. (2) It is an excellent conductor of electricity because the delocalised electrons are free to move through the layer and carry charge. (3) It is very lightweight and flexible because it is only one atom thick.
🎯 Exam Tips
Never confuse intermolecular forces with covalent bonds. Intermolecular forces are weak and exist between molecules. Covalent bonds are strong and exist within molecules. Breaking intermolecular forces = melting/boiling simple molecules. Breaking covalent bonds = decomposing the substance.
For questions about why a simple molecular substance has a low boiling point, the correct sequence is: (1) it is a simple molecular substance, (2) weak intermolecular forces between molecules, (3) little energy needed to overcome these forces. Do NOT say "weak covalent bonds" - the covalent bonds are strong!
When asked to draw dot and cross diagrams: only show outer shell electrons, use different symbols for electrons from different atoms, and make sure each atom ends up with a full outer shell (2 for H, 8 for others).
🧮 Maths Skills
Counting Shared Electrons in Covalent Bonds
Single bond = 1 shared pair (2 electrons): In H₂, each H contributes 1 electron → 1 shared pair = 2 electrons total. In H₂O, O shares 2 pairs (one with each H) → 4 shared electrons from O + 1 from each H = 4 shared electrons in bonds.
Double bond = 2 shared pairs (4 electrons): In CO₂, C shares 2 pairs with each O. Total shared electrons = 8 (4 per double bond). In O₂, each O shares 2 pairs → O=O is a double bond.
Drawing Dot and Cross Diagrams
Rules for dot and cross diagrams: (1) Only show outer shell electrons. (2) Use dots for one atom's electrons and crosses for the other. (3) Each shared pair has one dot and one cross. (4) Count total outer electrons to check: H₂O has 6 (from O) + 1 + 1 (from two H) = 8 electrons shown. (5) Each atom must end up with a full outer shell (2 for H, 8 for others).
Counting electrons in structures: In CH₄, C has 4 outer electrons, 4 H atoms each contribute 1 → total = 8 outer electrons. C forms 4 single bonds (8 shared electrons). Each H has 2 shared electrons. All atoms have full outer shells.
Interpreting data tables: When given melting point data, identify the order of magnitude difference (e.g. CO₂ sublimes at −78°C vs SiO₂ melts at 1713°C — a difference of roughly 1800°C) and link this to the type of bonding and forces involved. Simple covalent = low MP; giant covalent = very high MP.
❌ Common Misconceptions
Misconceptions About Covalent Bonding and Structure
Wrong: Covalent bonds are weak — that is why simple molecules have low boiling points
Correct: Intramolecular covalent bonds (within molecules) are very strong. It is the weak intermolecular forces between molecules that cause low boiling points. When a simple molecular substance boils, only the weak intermolecular forces are overcome — the covalent bonds within the molecules are NOT broken.
Wrong: Diamond and graphite are different elements
Correct: Both diamond and graphite are allotropes of carbon — they are made of only carbon atoms but have different structures. Diamond has each C bonded to 4 others in a rigid 3D lattice. Graphite has each C bonded to 3 others in layers with delocalised electrons between the layers. Same element, different properties due to different structures.
Wrong: Molecules and lattices are the same thing
Correct: A molecule is a fixed group of atoms bonded by covalent bonds (e.g. H₂O has exactly 3 atoms). A giant covalent lattice has no fixed "molecule" — atoms are bonded in a network extending throughout the entire structure (e.g. diamond, SiO₂).
✍️ 6-Mark Extended Question
Question
Compare the structure and properties of diamond and graphite. Explain the differences in terms of bonding.
Similarities: Both diamond and graphite are giant covalent structures made entirely of carbon atoms. Both have very high melting points because large amounts of energy are needed to break the strong covalent bonds throughout the structure [1 mark].
Diamond — structure and properties: Each carbon atom forms 4 covalent bonds in a rigid tetrahedral 3D lattice. This makes diamond extremely hard because all bonds are equally strong in all directions. Diamond does not conduct electricity because all 4 outer electrons per carbon are used in bonding, leaving no delocalised electrons to carry charge [2 marks].
Graphite — structure and properties: Each carbon atom forms 3 covalent bonds in layers of hexagonal rings. One electron per carbon atom is delocalised and can move between the layers, making graphite a good conductor of electricity. The layers are held together by weak intermolecular forces, allowing them to slide over each other easily — this makes graphite soft and slippery, useful as a lubricant and in pencils [2 marks].
Key difference explained: The different properties arise because diamond uses all 4 outer electrons in bonding (hard, non-conducting) while graphite uses only 3, leaving one delocalised electron (soft, conducting) [1 mark].
Mark scheme: 1 mark for similarity (giant covalent, high MP); 2 marks for diamond (4 bonds, hard, no free electrons); 2 marks for graphite (3 bonds + delocalised electron, conducts, layers slide); 1 mark for linking bonding difference to property difference.🔍 AO3: Analyse and Evaluate
Identifying Structure Type from Property Data
Scenario: A student is given data for three substances: Substance A has a melting point of −188°C and does not conduct electricity in any state. Substance B has a melting point of 1614°C and does not conduct electricity. Substance C has a melting point of 801°C, does not conduct when solid but conducts when molten. Classify each substance's bonding type and justify your answer.
Analysis: Substance A is a simple molecular substance. The very low melting point (−188°C) indicates that only weak intermolecular forces need to be overcome, and the lack of conductivity means there are no free charged particles — consistent with small neutral molecules. Example: F₂ or HCl.
Substance B is a giant covalent structure. The very high melting point (1614°C) indicates that strong covalent bonds throughout the structure must be broken. The lack of conductivity means no delocalised electrons — consistent with a giant covalent lattice like SiO₂ (not graphite, which conducts). Example: silicon dioxide.
Substance C is an ionic compound. The high melting point (801°C) indicates strong electrostatic forces between ions. Does not conduct when solid (ions fixed in lattice) but conducts when molten (ions free to move). Example: NaCl.
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