Is Diamond A Network Solid

Covalent Network Solids

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Covalent Network Solids are giant covalent substances like diamond, graphite and silicon dioxide (silicon(Four) oxide). This page relates the structures of covalent network solids to the physical properties of the substances.

Diamond

Carbon has an electronic arrangement of 2,4. In diamond, each carbon shares electrons with four other carbon atoms - forming iv single bonds.

In the diagram some carbon atoms simply seem to exist forming two bonds (or even ane bond), but that's not really the example. We are only showing a small bit of the whole construction. This is a behemothic covalent structure - information technology continues on and on in iii 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.

How to draw the structure of diamond

Don't try to exist as well clever past trying to draw too much of the construction! Learn to depict the diagram given in a higher place. Do it in the following stages:

Exercise until yous can practice a reasonable free-hand sketch in well-nigh 30 seconds.

Physical Backdrop of Diamond

has a very high melting point (almost 4000°C). Very potent carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs.
is very hard. This is again due to the need to intermission very potent covalent bonds operating in 3-dimensions.
doesn't conduct electricity. All the electrons are held tightly between the atoms, and aren't gratuitous to motility.
is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently spring carbon atoms.

Graphite

Graphite has a layer structure which is quite hard to draw assuredly in iii dimensions. The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced.

Notice that you cannot actually draw the side view of the layers to the same scale as the atoms in the layer without one or other office of the diagram being either very spread out or very squashed. In that instance, it is important to give some idea of the distances involved. The distance betwixt the layers is about two.5 times the distance betwixt the atoms within each layer. The layers, of course, extend over huge numbers of atoms - not just the few shown in a higher place.

You lot might contend that carbon has to form 4 bonds because of its four unpaired electrons, whereas in this diagram it only seems to be forming 3 bonds to the neighboring carbons. This diagram is something of a simplification, and shows the arrangement of atoms rather than the bonding.

The Bonding in Graphite

Each carbon atom uses three of its electrons to class elementary bonds to its iii shut neighbors. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon cantlet get delocalized over the whole of the sheet of atoms in 1 layer. They are no longer associated direct with any particular cantlet or pair of atoms, but are free to wander throughout the whole canvass. The important thing is that the delocalized electrons are gratuitous to move anywhere inside the sheet - each electron is no longer fixed to a particular carbon atom. At that place is, however, no direct contact between the delocalized electrons in one sheet and those in the neighboring sheets. The atoms inside a canvass are held together by strong covalent bonds - stronger, in fact, than in diamond considering of the additional bonding caused by the delocalized electrons.

So what holds the sheets together? In graphite you have the ultimate instance of van der Waals dispersion forces. As the delocalized electrons movement around in the sheet, very large temporary dipoles can be gear up up which will induce reverse dipoles in the sheets above and below - and then on throughout the whole graphite crystal.

Graphite has a loftier melting betoken, similar to that of diamond. In order to melt graphite, it isn't plenty to loosen one sheet from another. You accept to break the covalent bonding throughout the whole structure. Information technology has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. Yous tin think of graphite rather similar a pack of cards - each card is stiff, simply the cards volition slide over each other, or even autumn off the pack birthday. When you lot utilise a pencil, sheets are rubbed off and stick to the newspaper. Graphite has a lower density than diamond. This is considering of the relatively large amount of space that is "wasted" betwixt the sheets.

Graphite is insoluble in water and organic solvents - for the same reason that diamond is insoluble. Attractions between solvent molecules and carbon atoms will never exist stiff enough to overcome the strong covalent bonds in graphite. conducts electricity. The delocalized electrons are gratis 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.

Silicon dioxide: SiO₂

Silicon dioxide is also known as silica or silicon(IV) oxide has three different crystal forms. The easiest one to remember and depict is based on the diamond structure. Crystalline silicon has the same structure equally diamond. To plow it into silicon dioxide, all you need to practice is to change the silicon structure by including some oxygen atoms.

Notice that each silicon atom is bridged to its neighbors past an oxygen atom. Don't forget that this is just a tiny part of a giant construction extending on all three dimensions.

Silicon Dioxide has a high melting betoken - varying depending on what the item structure is (recollect that the structure given is only one of 3 possible structures), merely around 1700°C. Very strong silicon-oxygen covalent bonds take to be cleaved throughout the structure before melting occurs. Morevoer, it difficult due to the need to intermission the very strong covalent bonds. Silicon Dioxide does not acquit electricity since there aren't whatever delocalized electrons with all the electrons are held tightly between the atoms, and are not gratis to motility.Silicon Dioxide is insoluble in h2o and organic solvents. There are no possible attractions which could occur betwixt solvent molecules and the silicon or oxygen atoms which could overcome the covalent bonds in the giant construction.