One idea remains to complete the giant-covalent picture, and then the
chapter's strangest twist: two more forms of carbon that break every
generalisation diamond just set up.
First, close diamond's file properly. When you melt a molecular substance you
break the relatively weak intermolecular-forces — but diamond has no
intermolecular forces at all, because it has no molecules. Melting it means
breaking covalent bonds, which are very strong, and that is the whole reason
for the enormous temperatures. Generalise it: all substances with giant
covalent structures are solids with high melting and boiling points, because
a lot of energy is needed to break the strong covalent bonds running through
the entire structure. Diamond's companions on the course are graphite and
silicon dioxide (SiO₂) — sand. Two more diamond facts worth carrying: it
does not dissolve in water or any other solvent (those strong bonds would
have to break), and it conducts heat better than any other element — heat
one end and the vibrations race along the strong bonds to the other.
Diamond does not conduct electricity. Every outer-shell electron of every
carbon atom is tightly held in a covalent bond between specific atoms — none
are free to move, so there is nothing to carry a current. Keep that fact warm:
graphite, made of the very same carbon atoms, is about to break this rule,
and the contrast is one of the most examined comparisons in the course.
Graphite is also pure carbon — with the atoms arranged differently. It is
still a giant structure, but layered rather than fully three-dimensional:
picture a pack of playing cards. Within each "card" (layer), carbon atoms
are bonded by strong covalent bonds; between the cards, only weak attractions
hold the layers together. Within each layer, each carbon atom is bonded to
only three neighbours, not four.
That "only three" is the whole story of graphite's electrical fame. Each
carbon uses three of its four outer electrons for its three bonds — and the
fourth electron is free to move throughout the whole layer. Electrons
that move freely through a structure like this are called delocalised-electron, and their movement is exactly what carries a current: graphite
conducts electricity. Diamond and graphite — same element, same shared-pair
bonding, opposite conductivity — because in diamond every electron is
committed to a bond, while in graphite one electron per atom is left over and
mobile.
Graphite is soft and slippery — for the mirror-image reason. The bonds
within each layer are very strong, but the attractions between layers are
much weaker, so the layers slide over each other and flake apart with little
effort. When you write with a pencil, you are leaving a trail of graphite
layers behind on the paper (pencil "lead" is graphite mixed with clay for
hardness). Pure graphite is so slippery that powdered graphite is used as a
dry lubricant — in locks, for instance. (A refinement for the curious:
graphite's slipperiness actually depends on water molecules sticking between
the layers — graphite is not a lubricant in a vacuum.) Yet graphite still has
high melting and boiling points: to melt it you must break up the whole
layered structure, including the strong covalent bonds — weak inter-layer
attractions alone do not melt it. It is also insoluble in everything (too many
strong bonds to break) and less dense than diamond, because the gaps
between its layers are more than twice the atomic spacing within a layer — a
graphite crystal is, in a sense, full of wasted space that diamond does not
have.
And now the third form of carbon. Different fullerenes exist, whose
molecules contain different numbers of carbon atoms — which is why the name
carries a number: C₆₀ fullerene is a molecule containing exactly 60
carbon atoms joined by covalent bonds, arranged as a hollow ball. Here is
the twist: diamond and graphite have giant structures, but C₆₀ fullerene has
a simple-molecular-structure — discrete molecules, weak
intermolecular forces between them. That structural difference rebuilds every
property: C₆₀ has lower melting and boiling points than diamond and
graphite (melting it breaks only the weak intermolecular forces, not covalent
bonds); it dissolves in some solvents (again, only weak forces to break —
unlike the insoluble giants); and it does not conduct electricity. Each
of its carbons forms three bonds like graphite's, but the fourth electron can
only move around within its own C₆₀ molecule — electrons cannot jump from
molecule to molecule, so no charge can flow through the solid. One element,
three arrangements, three property profiles: structure decides everything.
Diamond, graphite and C₆₀ fullerene are allotropes — different forms
of the same element in the same physical state. The allotrope idea will
return with oxygen and ozone, and with sulfur; the carbon trio is where you
learned it.