Covalent Bonding (and Metallic) Flashcards
How are covalent bonds formed:
When two metal atoms share a pair of electrons to fill the outer shell
Covalent bonding definition:
SEFABSPEAN
Strong electrostatic force of attraction between a shared pair of electrons and nuclei
Why can Fullerene not conduct electricity
Although in every molecule every carbon is only covalently bonded to 3 others, having delocalized electrons, these electrons cannot jump between different molecules
Properties of Graphite
Graphite can conduct electricity because each carbon atom has a strong covalent bond to 3 other carbons, with one delocalized electron. All the electrons form a sea of delocalized electrons
Graphite is soft and slippery because the layers have weak IMFs, allowing them to slide over each other
Why do structures with simple molecular structures have low melting or boiling points
If question does not state:
1. State that it has a simple molecular structure
2. There are weak intermolecular forces
3. Which require little energy to break
Why does the melting and boiling points of substances with simple molecular structures increase, in general, with increasing relative molecular mass
Larger molecules have more forces of attraction needed to overcome
Why are substances with giant covalent structures solids with high melting and boiling points (e.g. Diamond)
- Diamond has a giant covalent structure
- With many strong covalent bonds (4 per carbon)
- That require a lot of energy to break
Why can’t covalent compounds usually conduct electricity
Covalent compounds do not have any free-moving, charged particles
They are often used as insulators
Metallic Bonding definition:
SEFABPISODE
Strong electrostatic force of attraction between positive ions and a sea of delocalized electrons
Physical properties of metals
Good conductors - have delocalized electrons which are free to move
Malleable - The layers of ions can slide (easily) over each other
