why do network solids have high melting points

Each carbon atom uses three of its electrons to form simple bonds to its three close neighbors. That leaves a fourth electron in the bonding level. These spare electrons in each carbon atom become delocalized 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. The important thing is that the delocalized 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 delocalized electrons in one sheet and those in the neighboring 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 delocalized electrons. So what holds the sheets together? In graphite you have the ultimate example of. As the delocalized 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. 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.


It 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 delocalized 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. Hardness: Very hard, due to the strong covalent bonds throughout the lattice (deformation can be easier, however, in directions that do not require the breaking of any covalent bonds, as with flexing or sliding of sheets in graphite or mica). Melting point: High, since melting means breaking covalent bonds (rather than merely overcoming weaker intermolecular forces).


Solid-phase : Variable, depending on the nature of the bonding: network solids in which all electrons are used for (e. g. diamond, quartz) are poor conductors, as there are no delocalized electrons. However, network solids with delocalized (e. g. graphite) or can exhibit metal-like conductivity. Liquid-phase electrical conductivity: Low, as the macromolecule consists of neutral atoms, meaning that melting does not free up any new charge carriers (as it would for an ionic compound). Solubility: Generally insoluble in any solvent due to the difficulty of solvating such a very large molecule.

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