C5: Covalent Bonding
Covalent bonding, simple molecules and giant covalent structures
C5 Covalent Bonding
📋 Key Definitions
Covalent bond: A shared pair of electrons between two non-metal atoms. Each atom contributes one electron to the shared pair. Both atoms achieve a full outer shell.
Simple molecular substance: A substance made up of separate molecules, each containing a fixed number of atoms joined by covalent bonds (e.g. H₂O, CO₂).
Giant covalent structure: A huge network of atoms all joined by covalent bonds in a giant lattice (e.g. diamond, graphite).
Intermolecular forces: Weak forces of attraction between molecules (NOT between atoms within a molecule). These are much weaker than covalent bonds.
Delocalised electron: An electron that is not attached to a particular atom and is free to move through the structure.
🔗 How Covalent Bonding Works
Covalent bonding occurs between non-metals. Instead of transferring electrons (as in ionic bonding), the atoms share pairs of electrons so each atom can achieve a full outer shell.
Key principle: Each atom contributes one electron to each shared pair. A single covalent bond is one shared pair of electrons. A double bond is two shared pairs (e.g. O=O in O₂). A triple bond is three shared pairs (e.g. N≡N in N₂).
🔵 Dot and Cross Diagrams for Molecules
Dot and cross diagrams for covalent bonding show shared pairs of electrons as overlapping circles, with one atom's electrons as dots and the other as crosses.
Hydrogen (H₂)
Each H atom has 1 electron. They share a pair so each has 2 in its outer shell.
H● ×H → H—H (one shared pair, single bond)
Chlorine (Cl₂)
Each Cl atom has 7 outer electrons. They share one pair so each has 8 in its outer shell.
Cl● ×Cl → Cl—Cl (one shared pair, single bond)
Water (H₂O)
Oxygen (6 outer electrons) shares one pair with each of two hydrogen atoms (1 electron each).
H—O—H (oxygen shares 2 pairs, forming 2 single bonds)
Water has a bent shape because oxygen has 2 lone pairs that repel the bonding pairs.
Methane (CH₄)
Carbon (4 outer electrons) shares one pair with each of four hydrogen atoms.
H—C—H (carbon forms 4 single bonds, tetrahedral shape)
Ammonia (NH₃)
Nitrogen (5 outer electrons) shares one pair with each of three hydrogen atoms. Nitrogen keeps one lone pair.
H—N—H (nitrogen forms 3 single bonds + 1 lone pair, pyramidal shape)
Carbon Dioxide (CO₂)
Carbon (4 outer electrons) shares two pairs (double bond) with each of two oxygen atoms.
O=C=O (carbon forms 2 double bonds, linear shape)
🧪 Simple Molecular Substances
Most covalent substances are simple molecular substances. Examples include H₂, Cl₂, H₂O, CH₄, NH₃, CO₂, and many organic compounds.
Properties of simple molecular substances: Low melting and boiling points because the intermolecular forces between molecules are weak (NOT because covalent bonds are weak). They do NOT conduct electricity because there are no free electrons or ions.
| Property | Explanation |
|---|---|
| Low melting and boiling points | Weak intermolecular forces between molecules require little energy to overcome |
| Mostly liquids or gases at room temperature | Due to low melting/boiling points |
| Do NOT conduct electricity | No delocalised electrons or free ions to carry charge |
| Often insoluble in water | Many covalent molecules do not interact with water (some exceptions like sugar) |
Crucial distinction: When a simple molecular substance melts or boils, the covalent bonds WITHIN the molecules do NOT break. Only the weak intermolecular forces BETWEEN the molecules break. This is why little energy is needed.
Worked Example - Explaining Low Boiling Points
Why does water boil at only 100°C when the O—H covalent bonds are very strong?
When water boils, the covalent bonds within each H₂O molecule do not break. Only the weak intermolecular forces between separate H₂O molecules are overcome. These intermolecular forces are weak, so relatively little energy is needed, giving water a relatively low boiling point.
💎 Giant Covalent Structures
A few covalent substances form giant covalent structures (macromolecules) where every atom is bonded to neighbouring atoms by strong covalent bonds in a huge network.
Diamond
Diamond: Each carbon atom forms 4 covalent bonds with other carbon atoms in a rigid tetrahedral structure. This makes diamond very hard and gives it a very high melting point.
- Very hard - each atom bonded to 4 others in a rigid 3D structure
- Very high melting point (about 3550°C) - strong covalent bonds throughout
- Does NOT conduct electricity - no delocalised electrons (all electrons are used in bonding)
Graphite
Graphite: Each carbon atom forms 3 covalent bonds with other carbon atoms in layers. One electron per atom is delocalised. The layers can slide over each other.
- Soft and slippery - layers can slide because only weak intermolecular forces hold the layers together
- Used as a lubricant - layers slide easily
- Very high melting point - strong covalent bonds within each layer
- Conducts electricity - delocalised electrons can move between layers
Graphene
Graphene: A single layer of graphite - one atom thick sheet of carbon atoms in a hexagonal lattice. Extremely strong, lightweight, flexible, and an excellent conductor of electricity and heat.
Fullerenes
Fullerenes are molecules of carbon atoms shaped like hollow balls, tubes or cages.
Examples of Fullerenes
- Buckminsterfullerene (C₆₀): A hollow sphere of 60 carbon atoms arranged in pentagons and hexagons, like a football. Used in drug delivery and lubricants.
- Carbon nanotubes: Cylindrical fullerenes - very long, thin tubes of carbon atoms. Extremely strong, lightweight, and conduct electricity. Used in electronics, materials and medicine.
| Property | Diamond | Graphite | Graphene |
|---|---|---|---|
| Carbon bonds per atom | 4 | 3 | 3 |
| Arrangement | Rigid 3D tetrahedral | Layers of hexagons | Single layer of hexagons |
| Hardness | Very hard | Soft / slippery | Very strong but flexible |
| Electrical conductivity | Does not conduct | Conducts (delocalised electrons) | Excellent conductor |
| Melting point | Very high | Very high | Very high |
| Delocalised electrons | None | Yes (1 per C atom) | Yes (1 per C atom) |
🧵 Polymers
Polymers are very large molecules (macromolecules) made from many small molecules (monomers) joined together by covalent bonds. Examples include polythene, PVC and nylon.
Polymers: Long chain molecules where monomers are linked by covalent bonds. The chains are held together by intermolecular forces. Polymers have higher melting points than simple molecular substances because the intermolecular forces between the long chains are stronger than those between small molecules.
❓ Practice Questions
Q1: Foundation Describe how a covalent bond forms. How is it different from an ionic bond?
Q2: Foundation Draw a dot and cross diagram for a molecule of ammonia (NH₃).
Q3: Foundation Explain why simple molecular substances have low melting and boiling points. Why is it important to mention intermolecular forces?
Q4: Higher Explain why graphite conducts electricity but diamond does not, even though both are made of carbon atoms.
Q5: Higher Compare the structure and properties of diamond, graphite and graphene. Include at least three comparisons.
Q6: Foundation What are fullerenes? Describe the structure of buckminsterfullerene (C₆₀) and give one use of carbon nanotubes.
✅ Answers
- A covalent bond forms when two non-metal atoms share a pair of electrons, with each atom contributing one electron to the shared pair. This is different from an ionic bond, where electrons are transferred from a metal to a non-metal forming charged ions that are attracted to each other.
- In NH₃, nitrogen (5 outer electrons) shares one pair with each of three hydrogen atoms (1 outer electron each). Nitrogen contributes one electron to each shared pair, and each hydrogen contributes one. Nitrogen keeps one lone pair of electrons. The diagram shows N in the centre with three H atoms, each H—N bond shown as a shared pair (one dot, one cross).
- Simple molecular substances have low melting/boiling points because the intermolecular forces between molecules are weak and require little energy to overcome. It is crucial to mention intermolecular forces because the covalent bonds WITHIN the molecules are actually very strong - they do NOT break when the substance melts or boils. Only the weak forces BETWEEN molecules are overcome.
- Graphite conducts electricity because each carbon atom forms only 3 covalent bonds, leaving one electron per atom delocalised (free to move throughout the structure). These delocalised electrons can carry charge. Diamond does not conduct electricity because each carbon atom forms 4 covalent bonds - all 4 outer electrons are used in bonding, so there are no delocalised electrons to carry charge.
- Three comparisons: (1) Bonding: diamond has 4 bonds per C atom (tetrahedral), graphite and graphene have 3 bonds per C atom (hexagonal layers). (2) Conductivity: diamond does not conduct (no delocalised electrons), graphite and graphene do conduct (one delocalised electron per C atom). (3) Hardness: diamond is very hard (rigid 3D structure), graphite is soft (layers slide), graphene is very strong and flexible (single layer). (4) Melting point: all three have very high melting points because of the strong covalent bonds.
- Fullerenes are molecules of carbon atoms shaped like hollow balls, tubes or cages. Buckminsterfullerene (C₆₀) is a hollow sphere of 60 carbon atoms arranged in pentagons and hexagons (like a football). Carbon nanotubes are cylindrical fullerenes that are very strong and conduct electricity - used in electronics, materials science and medicine.
🎯 Exam Tips
- The most common mistake: saying "covalent bonds are weak" when explaining low melting points. It is the intermolecular forces that are weak, NOT the covalent bonds
- Always specify "intermolecular forces between molecules" when explaining properties of simple molecular substances
- For diamond vs graphite questions, always mention the number of bonds per carbon atom AND delocalised electrons
- Graphite conducts because of delocalised electrons - do NOT say "free electrons" as this is less precise
- Graphene is a single layer of graphite - remember this for the exam
- When drawing dot and cross diagrams, make sure shared pairs are clearly shown in the overlap between atoms
🔢 Maths Skills
Mathematical Skills
Counting shared electrons: In a covalent bond, each atom contributes one electron per shared pair. A single bond = 1 shared pair (2 electrons). A double bond = 2 shared pairs (4 electrons). A triple bond = 3 shared pairs (6 electrons).
Example: In CO₂, carbon shares 2 pairs with each oxygen (O=C=O), so carbon contributes 4 electrons total and each oxygen contributes 2 electrons.
⚠️ Common Misconceptions
Watch Out!
Covalent bonds are weak. Wrong: covalent bonds are weak Correct: covalent bonds within molecules are very strong; it is the intermolecular forces between molecules that are weak
Diamond and graphite are different elements. Wrong: diamond and graphite are different elements Correct: both are allotropes of carbon — they are the same element with different structures
✍️ 6-Mark Question
Extended Answer
6 marks: Compare simple molecular and giant covalent structures.
Simple molecular structures (e.g. H₂O, CO₂) consist of separate molecules with strong covalent bonds within each molecule but weak intermolecular forces between molecules. This gives them low melting and boiling points because little energy is needed to overcome the weak intermolecular forces. They do not conduct electricity because there are no free electrons or ions. Giant covalent structures (e.g. diamond, graphite) have a huge network of atoms all joined by strong covalent bonds throughout. This gives them very high melting points because a vast amount of energy is needed to break all the covalent bonds. Diamond does not conduct (no delocalised electrons) but graphite does (one delocalised electron per carbon atom).
Mark scheme: 1 mark for describing simple molecular structure; 1 mark for low mp/bp explained by intermolecular forces; 1 mark for no conductivity; 1 mark for describing giant covalent structure; 1 mark for very high mp explained by covalent bonds; 1 mark for conductivity comparison.
📊 AO3: Analyse & Evaluate
Analysis and Evaluation
Substance A melts at 80°C and does not conduct in any state. Substance B melts at 3550°C and does not conduct. Substance C melts at 3600°C and conducts electricity.
Question: Classify each substance as simple molecular, giant covalent (diamond-like) or giant covalent (graphite-like). Justify your answers.
Answer: A = simple molecular (low mp, no conductivity). B = giant covalent diamond-like (very high mp, no delocalised electrons so no conductivity). C = giant covalent graphite-like (very high mp, conducts because of delocalised electrons).
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