Science Ace

Name: Tan Jin Neng Clarence (28)
Class : 1A3
Ace assignment in term 2

Diamond and Graphite

Introduction
Allotropy or allotropism is a behavior exhibited by some chemical elements: these elements can exist in two or more different forms, known as allotropes of that element. In each allotrope, the element's atoms are bonded together in a different manner. Allotropes are different structural modifications of an element.

Allotropes of Carbon
Carbon has eight allotropes: a) Diamond, b) Graphite, c) Lonsdaleite, d) C60 (Buckminsterfullerene or buckyball), e) C540, f) C70, g) Amorphous carbon, and h) single-walled carbon nanotube or buckytube.

Graphite and diamond are two of the most interesting allotropes. They are identical chemically—both are composed of carbon (C), but physically, they are very different in their structure and properties.
Diamond vs Graphite
Their structure and properties differ in the following ways:
1. Appearance:
Graphite is opaque and metallic- to earthy-looking, while diamonds are transparent and brilliant. Diamond is transparent over a larger range of wavelengths (from the ultraviolet into the far infrared) than is any other solid or liquid substance - nothing else even comes close.

2. Hardness:
The word “diamond” comes from the Greek word meaning unbreakable. Diamond is a perfect "10", defining the top of the hardness scale, and by absolute measures four times harder than sapphire (which is #9 on that scale). Each carbon atom in a diamond is covalently bonded to four other carbons in a tetrahedron. These tetrahedrons together form a giant 3-dimensional network of six-membered carbon rings (similar to cyclohexane), in the chair conformation, allowing for zero bond angle strain. This stable network of covalent bonds and hexagonal rings, is the reason that diamond is so incredibly strong.
On the other hand, Graphite is very soft and has a hardness of 1 to 2 on this scale. It also has a giant covalent structure but different from that of diamond. The carbon atoms are arranged in flat layers. In each layer, the carbon atoms are arranged in rings of six atoms, with each atom joined to others by strong covalent bonds. However the forces between the layers are weak and so the layers can easily slide past each other. This makes graphite soft and slippery.
Thus, Diamond is the ultimate abrasive, often used for industrial cutting and polishing tools, whereas Graphite is a very good lubricant and also used as pencil lead as the layers of atoms slide off the pencil onto the paper.



3. Thermal Conductivity
Diamond conducts heat better than anything - five times better than the second best element, Silver!

4. Electrical Conductivity
Diamond is an excellent electrical insulator as all the outer shell electrons in the atoms are used to form covalent bonds. Thus there is no free electrons and so conduction does not occur.
Graphite is the only non-metal that conducts electricity. This is due to the vast electron delocalization within the carbon layers (a phenomenon called aromaticity). These valence electrons are free to move, so are able to conduct electricity.

5. Melting Point
Both Graphite and Diamond have very high melting and boiling points because strong covalent bonds must be broken to melt and boil them.
Graphite has a higher meting point (3948 Kelvin) compared to Diamond (3820 Kelvin)

6. Crystallization
Diamond crystallizes in the Isometric system and graphite crystallizes in the hexagonal system.

7. Stability
Under the normal pressures and temperatures we experience on the Earth’s surface, Graphite is the stable form of carbon. In fact, diamonds are actually thermodynamically unstable, all diamonds at or near the surface of the Earth are currently undergoing a transformation into Graphite. This reaction, fortunately, is extremely slow for humans to notice.

8. Density & Refractive Index
Diamond is more dense than graphite with density of 3.5–3.53 g/cm3 compare to graphite’s 2.09–2.23 g/cm3
Diamond is highly refractive with an index of 2.418 (at 500 nm) whereas graphite is opaque.

9. Chemical activity
Graphite is slightly more reactive than diamond. This is because the reactants are able to penetrate between the hexagonal layers of carbon atoms in graphite. It is unaffected by ordinary solvents, dilute acids, or fused alkalis.

Conclusion
Though both are carbon, they are named differently as Diamond and Graphite probably because of the above differences. Due to their distinct differences, they are used for different purposes.
"Diamonds are a girl's best friend", it is so expensive and heavily sought after in the jewellery industry because of the "four Cs": Carat (weight), Clarity, Colour and Cut. It is also rare in that only a few survived the hazardous journey from the depths of the earth to reach the earth's surface. Because of its hardness, Diamond is used in cutting other hard solids. Rock drills and saws are studded with thousands of tiny diamonds. Diamond-tipped drills are used to cut through thousands of feet of rock to reach oil and gas deep in the Earth.
Natural graphite is mostly consumed for refractories, steelmaking, expanded graphite, brake linings, and foundry facings-lubricants. They are also used as the marking material ("lead") in common pencils, in zinc-carbon batteries, in electric motor brushes, and various specialized applications.
References
http://www.galleries.com/minerals/elements/diamond/diamond.htm
http://www.galleries.com/minerals/elements/graphite/graphite.htm
http://sciencekids.co.nz/sciencefacts/chemistry/diamond.html
http://www.enmu.edu/services/museums/miles-mineral/diamond.shtml
Chemistry Insights by Dr Rex M Heyworth ISBN: 978-981-247-824-5




Diamond Graphite