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Diamond is extremely strong owing to its crystal structure, known as diamond cubic, in which each carbon atom has four neighbors covalently bonded to it. Bulk cubic boron nitride (c-BN) is nearly as hard as diamond. Diamond reacts with some materials, such as steel, and c-BN wears less when cutting or abrading such material. [4]
These electrons are free to move, so are able to conduct electricity. However, the electricity is only conducted along the plane of the layers. In diamond, all four outer electrons of each carbon atom are 'localized' between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current.
Liquids made of compounds with only covalent bonds cannot conduct electricity. Certain organic ionic liquids , by contrast, can conduct an electric current. While pure water is not an electrical conductor, even a small portion of ionic impurities, such as salt , can rapidly transform it into a conductor.
However, a substance such as diamond, which conducts heat quite well, is not an electrical conductor. This is not a consequence of delocalization being absent in diamond, but simply that carbon is not electron deficient. Electron deficiency is important in distinguishing metallic from more conventional covalent bonding.
Unless otherwise noted, this article describes the stable form of an element at standard temperature and pressure (STP). [b]While arsenic (here sealed in a container to prevent tarnishing) has a shiny appearance and is a reasonable conductor of heat and electricity, it is soft and brittle and its chemistry is predominately nonmetallic.
Diamond is the hardest material on the qualitative Mohs scale. To conduct the quantitative Vickers hardness test, samples of materials are struck with a pyramid of standardized dimensions using a known force – a diamond crystal is used for the pyramid to permit a wide range of materials to be tested. From the size of the resulting indentation ...
In diamond all four outer electrons of each carbon atom are 'localized' between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current. In graphite, each carbon atom uses only 3 of its 4 outer energy level electrons in covalently bonding to three other carbon atoms in a plane ...
These definitions are equivalent to stating that metals conduct electricity at absolute zero, as suggested by Nevill Francis Mott, [2]: 257 and the equivalent definition at other temperatures is also commonly used as in textbooks such as Chemistry of the Non-Metals by Ralf Steudel [3] and work on metal–insulator transitions. [4] [5]