IBDP Giant Covalent, Diamond , Graphite, Fullerene, Graphene, Allotropes of Carbon, Silica SiO2

Описание: Most covalent substances exist as discrete molecules with a nite number of atoms.
However, there are some that have a very different structure, a crystalline lattice
in which the atoms are linked together by covalent bonds. Effectively, the crystal
is a single molecule with a regular repeating pattern of covalent bonds, so has no
finite size. It is referred to as a giant molecular or network covalent structure or
macromolecular structure. As we might expect, such crystalline structures have
very different properties from other smaller covalent molecules. A few examples are
considered here.
Allotropes are different forms of an element in the same physical state, such as
oxygen (O2) and ozone (O3) which both exist as gases. Different bonding within these
structures gives rise to distinct forms with different properties.

Did you know that diamond is a form of the element carbon, the same
element contained in your pencil leads? Diamond is the hardest known
natural substance. Artificial diamonds can be made by heating pure carbon
to very high temperatures under enormous pressures. 'Industrial diamonds'
made like this are embedded in the drills used by oil companies. They have
to drill through layers of rock to get to the crude oil deep underground.
Many covalently bonded substances are made up of individual
molecules. However, some substances, such as diamond, form very
different structures. These do not have relatively small numbers of atoms
arranged in simple molecules. They form huge networks of atoms held
together by strong covalent bonds in giant covalent structures.

Bonding in graphite
Carbon is not always found as diamonds. Another form is graphite, the
form of carbon used in pencil lead. In graphite, carbon atoms are only
bonded to three other carbon atoms. They form hexagons, which are
arranged in giant layers. There are no covalent bonds between the layers,
only weak intermolecular forces, so the layers can slide over each other
quite easily. It is a bit like the effect of cards sliding off a pack of playing
cards. This makes graphite a soft material that feels slippery to the touch.
As the carbon atoms in graphite's layers are arranged in hexagons, each
carbon atom forms three strong covalent bonds (Figure 4). Carbon atoms
have four electrons in their outer shell available for bonding. This leaves
one spare outer electron on each carbon atom in graphite.
These mobile electrons can move freely along the layers of carbon atoms.
The mobile electrons found in graphite are called delocalised electrons.
They no longer belong to any one particular carbon atom. They behave
rather like the electrons in a metallic structure (which you will look at in
detail in Topic C3.9).

These delocalised electrons allow graphite to conduct electricity.
The electrons will drift away from the negative terminal of a battery
and towards its positive terminal when put into an electrical circuit.
Diamond - and most other covalently-bonded substances - cannot conduct electricity.
This is because their atoms have no free electrons,as
all their outer shell electrons are involved in covalent bonding.


FULLERENES are hollow-shaped molecules of carbon . The structure of fullerenes is based on hexagonal rings of carbon atoms, as in graphite. However, they
may also have rings of five (pentagonal) or seven (heptagonal) carbon atoms.

GRAPHENE
If you could separate a single sheet of carbon atoms from graphite, you would get a layer of inter-locking hexagonal rings of carbon atoms . It would be just one atom thick. Scientists at Manchester University managed to do this in 2004, basically by using a
piece of sticky-tape. They stuck the tape across a piece of graphite, pulled
it off, and looked at the tape under a powerful electron microscope. They
had managed to isolate a 2D material - the thinnest ever made.