Covalent Crystals
· Solid substances in which atoms are held together by covalent bond
are known as covalent crystals.
· These crystals are very stable
For example:
Bucky ball( Fullerene)
Diamond
Graphite
GRAPHITE
The planes of hexagonal rings of carbon atoms are identical. The C-C bond length is 0.141
nm but the distance between the planes is 0.335 nm.
http://www.avogadro.co.uk/structure/chemstruc/network/g-molecular.htm
http://www.chemguide.co.uk/atoms/structures/giantcov.html
It is also a crystalline form of carbon.
STRUCTURE OF GRAPHITE
In graphite each carbon atom is covalently bonded to three carbon atoms to give trigonal
geometry. Bond angle in graphite is 120oC. Each carbon atom in graphite is sp2 hybridized.
Three out of four valence electrons of each carbon atom are used in bond formation with
three other carbon atoms while the fourth electron is free to move in the structure of
graphite.
Figure
Basic trigonal units unite together to give basic hexagonal ring. In hexagonal ring C-C bond
length is 1.42Ao.In graphite these rings form flat layers. These layers are arranged in
parallel, one above the other. These layers are 3.35Ao apart and are held together by weak
van derwaals forces only. These layers can slide over one another. Thus it is very soft.
Fourth electron of each carbon atom forms delocalized -bonds which spreads uniformly
over all carbon atoms. Due to this reason graphite conducts electricity parallel to the of its
plane.
PROPERTIES
It is dark grey crystalline solid , have dull metallic luster.
It is soft and greasy.
It is used as lubricant.
Its density is 2.2 gm/cm3.
It is used in the preparation of electrodes as it conduct electricity.
It is used as 'pencil lead'.
It is used as moderator in nuclear reactors.
Due to high melting it is used to prepare crucible for making high grade steel.
http://www.citycollegiate.com/allotropyXIIb.htm
The bonding in graphite
Each carbon atom uses three of its electrons to form simple
bonds to its three close neighbours. That leaves a fourth
electron in the bonding level. These "spare" electrons in each
carbon atom become delocalised over the whole of the sheet of
atoms in one layer. They are no longer associated directly with
any particular atom or pair of atoms, but are free to wander
throughout the whole sheet.
If you are interested (beyond A'level): The bonding in
graphite is like a vastly extended version of the bonding in
benzene. Each carbon atom undergoes sp2 hybridisation,
and then the unhybridised p orbitals on each carbon atom
overlap sideways to give a massive pi system above and
below the plane of the sheet of atoms.
The important thing is that the delocalised electrons are free to
move anywhere within the sheet - each electron is no longer
fixed to a particular carbon atom. There is, however, no direct
contact between the delocalised electrons in one sheet and
those in the neighbouring sheets.
The atoms within a sheet are held together by strong covalent
bonds - stronger, in fact, than in diamond because of the
additional bonding caused by the delocalised electrons. So what
holds the sheets together?
In graphite you have the ultimate example of van der Waals
dispersion forces. As the delocalised electrons move around in
the sheet, very large temporary dipoles can be set up which will
induce opposite dipoles in the sheets above and below - and so
on throughout the whole graphite crystal.
Note: If you aren't sure about van der Waals forces follow
this link before you go on. Use the BACK button on your
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The physical properties of graphite
Graphite
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has a high melting point, similar to that of diamond. In
order to melt graphite, it isn't enough to loosen one sheet
from another. You have to break the covalent bonding
throughout the whole structure.
has a soft, slippery feel, and is used in pencils and as a
dry lubricant for things like locks. You can think of
graphite rather like a pack of cards - each card is strong,
but the cards will slide over each other, or even fall off the
pack altogether. When you use a pencil, sheets are
rubbed off and stick to the paper.
has a lower density than diamond. This is because of the
relatively large amount of space that is "wasted" between
the sheets.
is insoluble in water and organic solvents - for the same
reason that diamond is insoluble. Attractions between
solvent molecules and carbon atoms will never be strong
enough to overcome the strong covalent bonds in
graphite.
conducts electricity. The delocalised electrons are free to
move throughout the sheets. If a piece of graphite is
connected into a circuit, electrons can fall off one end of
the sheet and be replaced with new ones at the other
end.
Note: The logic of this is that a piece of graphite ought only
to conduct electricity in 2-dimensions because electrons
can only move around in the sheets - and not from one
sheet to its neighbours.
In practice, a real piece of graphite isn't a perfect crystal,
but a host of small crystals stuck together at all sorts of
angles. Electrons will be able to find a route through the
large piece of graphite in all directions by moving from one
small crystal to the next.
DIAMOND
The giant covalent structure of diamond
In diamond, each carbon shares
electrons with four other carbon atoms - forming four single bonds.
In the diagram some carbon atoms only seem to be forming two bonds (or
even one bond), but that's not really the case. We are only showing a small bit
of the whole structure.
This is a giant covalent structure - it continues on and on in three dimensions.
It is not a molecule, because the number of atoms joined up in a real diamond
is completely variable - depending on the size of the crystal
A carbon atom with covalent bonds towards the four corners of a tetrahedron...
In diamond, each C-atom is covalently bonded to four other C-atom to give a
tetrahedral unit. In diamond each C-atom is sp3-hybridized.Therefore each Catom forms four sigma bonds with
neighbouring C-atoms.
In diamond C-C-C bond angle is 109.5O.These basic tetrahedral units unite with
one another and produce a cubic unit cell.
C-C bond length in diamond is 1.54AO.
C-C bond energy is 347 kj/mole.
In diamond crystal, basic units joined to forms octahedral shape of diamond
crystal.
PROPERTIES OF DIAMOND
In diamond each C-atom utilizes its four unpaired electrons in bond formation.
These bonding electrons are localized. Due to this reason diamond is a bad
conductor of electricity.
Diamond is the hardest substance ever known.
Pure diamond is cloudless.
Its melting point is 3500OC.
Pure diamond is transparent to x-rays.
It has high refractive index i.e. 2.45.
Due to impurities it may be colored.
Its density is 3.5 gm/cm3.
http://www.worldofmolecules.com/materials/fullerene.htm
Structure
The basic C60 structure consists of 60 carbon atoms that link together to form a hollow
cage-like structure. The structure consists of 32 faces of which 20 are hexagons and 12 are
pentagons. Of these, no two pentagons share a vertex. A similar structure has been used to
make soccer balls, in particular the Telstar supplied by Adidas and used in the 1970 and
1974 World Cups.
They are NOT considered giant covalent structures and are classed as simple
molecules. They do dissolve in organic solvents giving coloured solutions (e.g.
deep red in petrol hydrocarbons, and although solid, their melting points are not
that high.
Properties
The C60 molecule is extremely stable, being able to withstand high temperatures and
pressures. The exposed surface of the structure is able react with other species while
maintaining the spherical geometry.
The hollow structure is also able to entrap other smaller species such as helium, while at the
same time not reacting with the fullerene molecule. In fact the interior of most buckyballs is
so spacious, they can encase any element from the periodic table.
Buckyballs do not bond to one another. They do however, stick together via Van der Waals
forces.
By doping fullerenes, they can be electrically insulating, conducting, semiconducting or even
superconducting.