Chemistry; bonding (models)

Metallic bonding

Attraction between delocalised electrons (valence electrons)This and positive metal ions:

The electrostatic attraction between these delocalised electrons and the positive metal ions explains how the atoms of metal are bonded.

•   This attraction occurs in all directions.

• As a consequence, metallic bonds are able to resist forces such as bending or hitting the metal with a hammer (malleability).

• Weak intramolecular forces

Metals have high boiling or melting points due to very strong force of attraction between the atoms and as a result, large amounts of energy are required to break this bond.

Space between atoms means delocalised electrons can roam freely in the compound, by which atoms can conduct electricity

Metals are insoluble, although some will react with water to produce a soluble product.

Ionic bonding

A chemical bond between a metal and non-metal where electrons are lost or gained , resulting in the formation of charged ions with electrostatic attraction.

• can only conduct electricty when motlen or dissolved because particles are in a fixed lattice structure

• High melting and boiling points because of a strong lattice structure held by electrostatic forces between oppositely charged particles, requiring high amounts of energy to break these bonds.

• Generally solids at room temp.

Covalent bonding

A chemical bond between a non-metal and a non-metal with sharing of valence electrons to achieve noble gas configuration or a full outer shell. No ions present.

• Weak intermolecular forces holding molecules together resulting in relatively low melting and boiling points.

• Poor conductors of electricity because of lack of free charges (electrons), unlike ionic compounds

• Typically liquid and gas or soft solids at room temp.

• Strong intramolecular forces

What are the three allotropes of carbon?

Diamond (C) graphite (C) and buckminsterfullerene (C60)

1

Mono

2

Di

3

Tri

4

Tetra

5

Penta

6

hexa

Diamond: Properties

Hard, does not conduct electricity because there are no delocalised electrons (no ions), insoluble in water (no polar molecules), high melting point, lustrous (transparent)

Diamond: chemical structure

Each carbon atom is joined to 4 other atoms by covalent bonding (tetrahedral structure).

Diamond: Uses

Cutting or drilling tools (Oil rigs)

Graphite: Properties

conducts electricity, slippery, black, shiny and opaque, high melting point

Graphite: Chemical structure

Each carbon atom is join to 3 other atoms. This forms a hexagonal structure (networked). Weak intermolecular forces and can slide over one another. Each carbon atome has a delocalised electron.

Graphite: Uses

Lubricant

Buckminsterfullerene: Properties

Cannot conduct electricity between molecules but can within itself, slippery, low melting point compared to other giant covalent structures

Buckminsterfullerene: chemical structure

60 electrons (carbon atoms) with strong covalent bonds. Molecules are spherical with hexagonal rings of carbon. Delocalised electrons stay within the structure with weak intermolecular forces. Has a hollow shape.

Buckminsterfullerene; uses

Drug delivery, storing hydrogen, fuel tank for fuel cell powered cars

ionic bonding properties

High boiling and melting point, soluble in polar solvent, a crystal lattice structure, solid form,

metallic bonding properties

malleability, ductility, good conducters of electricity, lustrousconductors

Covalent bonding properties

cannot conduct electricty, insoluble in water,

diamon: why is it hard?

Diamond is hard because it is a chemically rigid structure with a tetrahedral bond. Which signifies that one carbon atom is bonded to 4 other carbon atoms which means it takes much more energy to break the bonds.

Whys is diamond not a good conductor of electricity

diamond isnt a good conductor of electricity because it lacks the delocalised electrons moving freely within the structure hence its not a conductor of thermal/electricital conductivity.