IB MYP 4-5 Chemistry -Simple covalent molecules and giant structures- Study Notes - New Syllabus
IB MYP 4-5 Chemistry -Simple covalent molecules and giant structures- Study Notes
Key Concepts
Simple Covalent Molecules and Giant Structures
Covalent substances are formed when non-metal atoms share electrons to achieve full outer electron shells. These can exist as either simple covalent molecules (with few atoms) or giant covalent structures (with huge networks of atoms bonded together).
Simple Covalent Molecules
A simple covalent molecule consists of a few atoms held together by strong covalent bonds within the molecule, but with weak intermolecular forces between molecules.
Examples: Hydrogen (\( \mathrm{H_2} \)), Oxygen (\( \mathrm{O_2} \)), Water (\( \mathrm{H_2O} \)), Carbon dioxide (\( \mathrm{CO_2} \)), Ammonia (\( \mathrm{NH_3} \)), Methane (\( \mathrm{CH_4} \)).
Characteristics:
- Bonding: Strong covalent bonds inside the molecule, weak forces between molecules (van der Waals forces).
- Melting and boiling points: Low, because little energy is needed to overcome intermolecular forces.
- Electrical conductivity: Do not conduct electricity — no free ions or electrons.
- Solubility: Often soluble in non-polar solvents (like oil, ethanol).
- Physical state: Usually gases or liquids at room temperature.
Reason for Low Melting/Boiling Points:
The strong covalent bonds are inside each molecule, but when melting or boiling, only weak intermolecular forces are broken — not the covalent bonds themselves.
Exceptions:
- \( \mathrm{H_2O} \) has a relatively high boiling point for a small molecule due to hydrogen bonding.
- \( \mathrm{NH_3} \) and \( \mathrm{HF} \) also have higher boiling points because of hydrogen bonds.
Giant Covalent Structures
Giant covalent (macromolecular) structures consist of billions of atoms bonded together by strong covalent bonds in a continuous 3D network.
Key Idea: All atoms are connected by covalent bonds throughout the entire structure — there are no individual molecules.
Examples:
- Diamond — a form of carbon.
- Graphite — another form of carbon.
- Silicon dioxide (\( \mathrm{SiO_2} \)) — sand/quartz.
Diamond (Carbon)
- Each carbon atom forms four covalent bonds (tetrahedral structure).
- All bonds are identical and strong → giant 3D network.
- Properties:
- Extremely hard (used in cutting tools).
- High melting and boiling points (many covalent bonds to break).
- Does not conduct electricity (no free electrons).
- Transparent, crystalline solid.
Structure summary: \( \mathrm{C-C\ tetrahedral\ bonding} \)
Graphite (Carbon)
- Each carbon atom forms three covalent bonds → hexagonal layers.
- One electron per atom remains delocalized → conducts electricity.
- Layers are held by weak forces → slide easily.
- Properties:
- Soft and slippery (used as a lubricant and in pencils).
- Good conductor of electricity (delocalized electrons move between layers).
- High melting point (strong covalent bonds within layers).
Structure summary: \( \mathrm{C-C\ hexagonal\ layers\ with\ delocalized\ e^-} \)
Silicon Dioxide (\( \mathrm{SiO_2} \)) — Quartz
- Each silicon atom bonds to four oxygen atoms.
- Each oxygen atom bonds to two silicon atoms → continuous lattice.
- Properties:
- Very high melting point (strong Si–O covalent bonds).
- Hard, brittle solid.
- Does not conduct electricity (no free electrons).
Structure summary: \( \mathrm{SiO_2\ tetrahedral\ lattice} \)
Comparison: Simple vs. Giant Covalent Structures
| Property | Simple Covalent Molecules | Giant Covalent Structures |
|---|---|---|
| Number of Atoms | Few atoms per molecule | Millions of atoms bonded in a network |
| Bonding Type | Strong covalent bonds within molecules; weak forces between | Strong covalent bonds throughout |
| Melting/Boiling Points | Low | Very high |
| Electrical Conductivity | Poor (no ions/electrons) | None (except graphite) |
| Solubility | Often soluble in non-polar solvents | Insoluble in most solvents |
Example:
Explain why water (\( \mathrm{H_2O} \)) has a higher boiling point than oxygen (\( \mathrm{O_2} \)).
▶️ Answer / Explanation
Step 1: Both are simple covalent molecules, but water forms hydrogen bonds.
Step 2: Hydrogen bonds are stronger than normal intermolecular forces.
Final Answer: Water has a higher boiling point because more energy is needed to break hydrogen bonds between molecules.
Example :
Explain why graphite conducts electricity but diamond does not, although both are forms of carbon.
▶️ Answer / Explanation
Step 1: In graphite, each carbon forms 3 bonds — one electron is delocalized and free to move.
Step 2: In diamond, all 4 electrons are used in covalent bonds — no free electrons.
Final Answer: Graphite conducts electricity due to delocalized electrons; diamond does not because all electrons are fixed in bonds.
Example :
Silicon dioxide (\( \mathrm{SiO_2} \)) and carbon dioxide (\( \mathrm{CO_2} \)) have similar chemical compositions, yet their physical states differ. Explain why.
▶️ Answer / Explanation
Step 1: \( \mathrm{CO_2} \) is a simple covalent molecule with weak forces between molecules → gas at room temperature.
Step 2: \( \mathrm{SiO_2} \) forms a giant covalent lattice — every Si atom bonded to 4 O atoms.
Step 3: Many covalent bonds must be broken to melt \( \mathrm{SiO_2} \).
Final Answer: \( \mathrm{CO_2} \) is a gas because it has weak intermolecular forces; \( \mathrm{SiO_2} \) is a solid because it has strong covalent bonds throughout its structure.