chemistry sac 2

carbon materials

it is the building block of life because of its ability to form many complex but stable molecules.

Some forms of pure carbon that exists include: diamond, graphite, graphene and amorphous carbon.

they are allotropes of carbon

• all made up pure carbon but have very different properties

allotropes

different ways of structuring molecules of the same element

diamond

a form of carbon where every carbon atom has four single covalent bonds to other carbon atoms arranged into a 3D covalent network lattice. This structure is very strong.

properties of diamond

• high melting point

• hard

• brittle

• no electrical conductivity

• high thermal conductivity

• insoluble

high melting point

a number of strong covalent bonds have to be broken in order to break the lattice. therefore it is difficult to disrupt the structure of diamond. This results in a high melting point

hard

the rigid lattice structure of diamond makes diamond hard

brittle ( hard but able to break easily)

the rigid lattice structure does not allow diamond to be bent. It will just break.

no electrical conductivity ( flow of electricity through the movement of charged particles)

each carbon takes part in four covalent bonds. This means that there are no free electrons available to conduct electricity and therefore, there is no movement of charged particles to conduct electricity.

high thermal conductivity (ability to transfer heat)

the strong covalent bonds between the carbon atoms allow diamond to conduct heat

insoluble ( unable to be dissolved/ no aqueous state)

the strength of the covalent bonds in the structure of diamond cannot be overcome by the intermolecular bonds of a solvent like water

applications

• cutting tools

  the hardness of diamond resists wear and enhances durability

• thermal conductor in electrical components

  diamond’s strong covalent bonding enables thermal conductivity

• optical components

  diamond lasers can be produced due to diamond’s ability to transmit heat and light very effectively

• abrasive

  diamond is able to induce friction on other objects without experiencing wear itself due to its hardness

graphite

where carbon atoms are covalently bonded to 3 other carbon atoms, but only in 2D layers. There are also delocalised electrons moving around the different layers

this structure is called a covalent layer lattice

dispersion forces between graphite

properties of graphite

• high melting point

• high thermal conductivity

• soft, slippery feeling

• lower density than diamond

• insoluble

high melting point - important

a number of strong covalent bonds have to be broken in order to disrupt the structure of graphite. this results in a high melting point

high thermal conductivity

when heat is applied, the delocalised electrons are able to move quickly and bump into other electrons in the lattice structure, thereby passing on the thermal energy from an area of high temp to low temp

soft, slippery feeling

the layers of graphite readily slide other each other due to the lack of covalent bonds

lower density than diamond

this is due to the greater distance between the layers of graphite, which are held together by intermolecular forces of dispersion which are far weaker than the covalent bonding in the diamond lattice

insoluble

the strength of the covalent bonds in the structure of graphite cannot be overcome by the intermolecular bonds of a solvent like water

applications of graphite

• carbon brushes in electrical motors

  graphite is able to conduct electricity, and therefore can transfer current from a stationary wire to moving parts in electrical motors

• electrode in batteries

  graphite is an inert (unreactive) and electrically conductive material.

• industrial lubricant

  layers of graphite are able to slide over each other, which, when used as a lubricant, reduces friction in machinery

graphene

• single layer of graphite.

• known as carbon nanomaterial

• similar but slightly different properties to graphite because it is only a single layer.

• delocalised electrons move incredibly fast when voltage is applied, making it a very good conductor of electricity.

properties of Graphene

• thin and light

• high thermal conductivity

• high melting point

• electrical conductivity

• flexible

thin and light

graphene consists of a single-layered network of carbon atoms.

high thermal conductivity

similar to graphite, the movement of delocalised electrons increases as heat energy is applied. This results in electrons moving quicker, bumping into other electrons and thereby conducting thermal energy

high melting point

the strong covalent bonds holding the lattice structure together require a large amount of energy to break

electrical conductivity

Graphene possesses a layer of delocalised electrons. The movement of these electrons when electrical energy is applied allows graphene to conduct electricity

flexible

since graphene is very thin, it is easy to bend. Additionally because of the strength of the covalent bonds within graphene, it is difficult to break graphene through bending. Therefore, graphene is very flexible.

applications of graphene

• solar cells

  graphene’s flexibility could produce a breakthrough in the solar industry, with solar cells being able to be applied to any surface, including curved ones

• filtration

  graphene membranes have been shown to be an effective filter

• drug delivery in medicine

  graphene- based carriers of drugs for cancer therapy have been shown to target cancer cells more effectively.

amorphous carbon

• can be formed from the burning wood and other plant matter when air ( oxygen gas) is not plentiful

• known as carbon black (soot)

• usually does not have a consistent structure and can be used as a fuel

diamond vs graphite lattice