Learn to draw the diagram given above. Do it in the following stages:. Practice until you can do a reasonable free-hand sketch in about 30 seconds. Graphite has a layer structure which is quite difficult to draw convincingly in three dimensions. The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced.
Notice that you cannot really draw the side view of the layers to the same scale as the atoms in the layer without one or other part of the diagram being either very spread out or very squashed. In that case, it is important to give some idea of the distances involved. The distance between the layers is about 2. The layers, of course, extend over huge numbers of atoms - not just the few shown above.
You might argue that carbon has to form 4 bonds because of its 4 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. 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? An edible sea snail found all around European coasts.
Diamonds are cut with specialized tools that make use of diamond tipped phosphor bronze or diamond dusted steel blades. Such tools are used to exploit the structural weakness of the diamond by grooving and striking along specific tetrahedral planes. According to the Mohs Hardness Scale, tooth enamel earns a 5.
For reference, diamonds are the strongest substance on earth, ranking 10 on the Mohs scale. There is no such thing as perfect toughness. Any gem will break, not just chip, if it is hit hard enough.
Diamonds are very tough, but remember that if a cutter can purposely cleave split a diamond by giving it a sharp blow in the right direction, you can achieve the same thing if you hit it hard enough accidentally. The structure of boron nitride in its wurtzite configuration is stronger than diamonds. Boron nitride can also be used to construct nanotubes, aerogels, and a wide variety of other fascinating applications.
It is a myth; you cannot die if you try to lick a diamond. Diamonds do not emit toxically or release toxic substances. However, a person may die if he swallows a diamond because a diamond is tough and has sharp edges, and can cut some part of the intestine in the stomach.
Diamond does not conduct electricity because it has no charged particles that are free to move. Graphite does conduct electricity because it has delocalised electrons which move between the layers. Diamond and graphite Diamond and graphite are different forms of the element carbon. Diamond Structure and bonding Diamond is a giant covalent structure in which: each carbon atom is joined to four other carbon atoms by strong covalent bonds the carbon atoms form a regular tetrahedral network structure there are no free electrons Carbon atoms in diamond form a tetrahedral arrangement Properties and uses The rigid network of carbon atoms, held together by strong covalent bonds, makes diamond very hard.
Graphite Structure and bonding Graphite has a giant covalent structure in which: each carbon atom forms three covalent bonds with other carbon atoms the carbon atoms form layers of hexagonal rings there are no covalent bonds between the layers there is one non-bonded - or delocalised - electron from each atom Dotted lines represent the weak forces between the layers in graphite Properties and uses Graphite has delocalised electrons, just like metals.
These atoms have two types of interactions with one another. In the first, each carbon atom is bonded to three other carbon atoms and arranged at the corners of a network of regular hexagons with a degree C-C-C bond angle. These planar arrangements extend in two dimensions to form a horizontal, hexagonal "chicken-wire" array. In addition, these planar arrays are held together by weaker forces known as stacking interactions. The distance between two layers is longer 3.
This three-dimensional structure accounts for the physical properties of graphite. Unlike diamond, graphite can be used as a lubricant or in pencils because the layers cleave readily. It is soft and slippery, and its hardness is less than one on the Mohs scale. Graphite also has a lower density 2. The planar structure of graphite allows electrons to move easily within the planes. This permits graphite to conduct electricity and heat as well as absorb light and, unlike diamond, appear black in color.
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