Carbon and Its Allotropes

 

Susan Evans

 

Chemistry 405

 

April 30, 2001

            Carbon, element six on the periodic table and having the symbol C, is the fundamental building block of all living organisms.  Thus, all carbon-containing compounds are known as organic compounds.  Excluding hydrogen, carbon is able to form more compounds than any other element.  “[Carbon] is the major component of coal, petroleum, asphalt, limestone, and most materials made by plants and animals.”(“Carbon”, Encarta Online, 2001)  Most natural and synthetic fibers contain carbon.  Building materials such as wood and plastics, heat and energy sources such as natural gas and coal, pesticides, medicines, and many foods are composed in small or large part of carbon.  “Carbon atoms from part or all of the backbone for the major molecules of all living things on Earth, including sugars, proteins, fats, and deoxyribonucleic acids (DNA).”(“Carbon”, Encarta Online, 2001)  In industry carbon is used to make steel from iron, purify metals, and add strength to rubber. 

“In 1961 the international unions of physicists and chemists agreed to use the mass of the isotope carbon-12 as the basis for atomic weight.”(“Carbon”, Encarta 2000)  Carbon-12, carbon-13, and carbon-14 are the three isotopes of carbon that can be found in nature.  Carbon-12 is the most abundant isotope, “accounting for about 98.89% of all carbon.”(“Carbon”, Encarta Online, 2001)  Although less plentiful, carbon-13 and carbon-14 have important uses.  The nucleus of carbon-13 is magnetic, which allows scientists to identify carbon-13 nuclei using nuclear magnetic resonance.  This technique enables them to study the structure of molecules containing this isotope of carbon.  Carbon-14 is radioactive with a half-life of 5730 years.  This property of carbon-14 lends itself to a technique called carbon dating, which is used to determine the age of fossils, among other things.  While an organism is alive, it contains a ratio of one atom of carbon-14 for every 1012 atoms of carbon-12.  When an organism dies, no additional carbon-14 is taken in.  Carbon-14 decreases with time, so by measuring the carbon-12 - carbon-14 ratio, the approximate date of death can be determined.

              Four allotropes of carbon have been determined.  “Allotropes differ in the way the atoms bond with each other and array themselves into a structure.  Because of their different structures, allotropes have different physical and chemical properties.”(“Carbon”, Encarta Online, 2001)  Carbon exists as diamond, graphite, amorphous carbon, and the recently discovered buckminsterfullerene.  Each allotrope of carbon will be discussed in this paper, with particular emphasis on diamond and graphite. 

            The structure and bonding in graphite lends itself to the properties that make it a good lubricant, among other uses, which will be visited later.  “In graphite, the atoms form planar, or flat, layers.  Each layer is made up of rings containing six carbon atoms.  Each atom has three sigma bonds and belongs to three neighboring rings.  The fourth electron of each atom becomes part of an extensive pi bond system.”(“Carbon”, Encarta Online, 2001)  A structural diagram of graphite is shown  below, in Figure 1.

Figure 1
 

 


           

The free movement of pi bonds throughout the molecule enable graphite to conduct electricity.  “Bonds between atoms within a layer of graphite are strong, but the forces between the layers are weak.”(“Carbon”, Encarta Online, 2001)  Because of this inter-layer weakness, the layers are able to move past each other giving rise to one of graphite’s physical properties, softness.  It is precisely this property that makes graphite a good lubricant, alone or mixed with grease, oil, or water.  The layers in graphite can be easily removed.  To see this, all one must do is write with a pencil.  The lead in a pencil is not actually lead, but a graphite-clay mixture. 

            Graphite has many uses other than those previously mentioned.  “Graphite is used as electrodes in electrochemical industries where corrosive gases are given off, and for electric furnaces that reach extremely high temperatures.”(“Graphite”, Encarta Online, 2001)  Graphite can be used in these high-temperature furnaces because is it a poor conductor of heat.  Graphite is also used in high-temperature crucibles, some paints, and as a moderator in atomic reactors.  Because of the many uses of graphite, a method has been created for making the allotrope in the laboratory.  “Graphite is made artificially by baking a mixture of petroleum coke and coal tar pitch at 9500C (17400F) for 11 to 13 weeks, then transferring the baked product to electric graphitizing furnaces and heating it to about 28000C (56000F) for 4 to 5 weeks.”(“Graphite”, Encarta Online, 2001)

            Graphite is the most stable allotrope of carbon.  The other most widely-used allotrope of carbon, diamond, is continuously undergoing a reaction into graphite.  Fortunately, the process is very slow.  A faster method of converting graphite to diamond is used by diamond makers.  “Extremely high pressure (more than 100,000 times the atmospheric pressure at sea level) and temperature (about 30000C or 50000F) [are applied to graphite].  High temperatures break the strong bonds in graphite so that the atoms can rearrange themselves into a diamond lattice.  About 90% of the diamonds used in tools in the United States are made this way.”(“Carbon”, Encarta Online, 2001)  Diamonds of gem quality have been created, but the best diamonds are found in nature. 

            The structure of diamond is quite different from that of graphite, giving rise to their many differences.  “In diamond, each carbon atom bonds tetrahedrally to four other carbon atoms to form a three-dimensional lattice.  The shared electron pairs are held tightly in sigma bonds between adjacent atoms.”(“Carbon”, Encarta Online, 2001)  The diamond structure is shown in  Figure 1.  Diamond is rather inert and reacts with few chemicals this is due to its compact crystal lattice.  “Diamond doesn’t react with chemical reagents like acids and alkalis.  However, diamond reacts with sodium carbonate at high temperatures to form carbon monoxide.”(“Crystalline Allotropes of Carbon”, 2000) 

Diamond is the hardest substance, “four times harder than the next hardest natural mineral, corundum.”(“The Mineral Diamond”, 1996)  It defines the upper limit of the Mohs hardness scale, having a value of 10.  Sought-after diamonds are colorless, but some rare diamonds can be green, blue or red in color.  “The color in a diamond is caused by the presence of minor elements.”(“Diamond”, Encarta Online, 2001)  Unlike graphite, diamonds are good conductors of heat and poor electrical conductors.  “[In fact], diamond conducts heat better than anything-five times better than the second best element, silver.”(“The Mineral Diamond”, 1996)  Because of its hardness, diamond is often used as an abrasive.  “Diamond coatings can also be synthetically produced by heating carbon dioxide over a metal surface with a series of lasers.  These diamond coatings have the potential to greatly extend the lifetime of precision dies, drills, and saw blades.”(Encarta, 2000)

            An interesting physical property of diamond is its melting point.  “Diamond has the highest melting point (3820 K) [of any substance].”(“The Mineral Diamond”, 1996)  Another fascinating property of diamond is its lattice density.  “The atoms of diamond are packed closer together than are the atoms of any other substance.”(“The Mineral Diamond”, 1996)  A table contrasting the physical characteristics of graphite and diamond follows.

 

Physical Characteristics

Diamond

Graphite

Color

variable-pale yellows, browns, grays, and also white, blue, black, reddish, greenish and colorless

Black silver

Luster

adamantine to waxy

metallic to dull

Transparency

crystals are transparent to translucent in rough crystals

crystals are opaque

Crystal System

Isometric

hexagonal

Crystal Habits

cubes and octahedrons

massive lamellar veins and earthy masses

Hardness(Mohs)

10

1-2

Specific Gravity

3.5 (above average)

2.2 (well below average)

Cleavage

perfect in 4 directions forming octahedrons

perfect in 1 direction

Fracture

Conchoidal

Flaky

Streak

White

Black gray to brownish gray

Best Field Indicator

Extreme Hardness

Softness, luster, density and streak

(The data contained in this table was taken from the following site:  http://mineral.galleries.com/Minerals/elements/. )

           

 

A phase diagram for carbon, showing graphite and diamond follows.

P

R

E

S

S

U

R

E
 
 

 

 

 

 

 

 

 

 

 


Graphite
 
           

 

TEMPERATURE
 
 

 

 

 


Graphite and diamond also have different thermodynamic properties.  Since graphite is the stable form of carbon, it has a heat of formation of zero.  To contrast, diamond has a heat of formation of  +1.895 kJ/mol.  Other thermodynamic data is represented in the table below, which includes carbon in the gas phase to serve as a comparison.  Notice that the entropy of graphite is larger than that of diamond; this is due again to diamond’s compact crystal lattice.  In other words, it is much more ordered than the graphite structure.  Carbon is found in nature in solid form, this accounts for the largely positive enthalpy and Gibbs values.  A lot of energy is needed to vaporize carbon.

 

 

Standard Enthalpy of Formation(kJ/mol)

Standard Gibbs Energy of Formation(kJ/mol)

Standard Molar Entropy(J/Kmol)

Graphite

0

0

5.740

Diamond

+1.895

+2.900

2.377

Gaseous Carbon

+716.68

+671.26

158.10

(The data found in this table was taken from Elements of Physical Chemistry, Peter Atkins.)

 

“Amorphous carbon is actually made up of tiny crystal-like bits of graphite with varying amounts of other elements, which are considered impurities.”(“Carbon”, Encarta Online, 2001)  Some examples of amorphous carbons are charcoal, soot, and a coal-derived fuel called coke.  The higher the carbon content in coal, the more energy is released in combustion.  “The coal industry divides coal up into various grades depending on the amount of carbon in the coal and the amount of impurities.  The highest grade, anthracite, contains about 90% carbon.  Lower grades include bituminous coal, which is 76% to 90% carbon, subbituminous coal, with 60% to 80% carbon and lignite, with 55% to 73% carbon.”(“Carbon”, Encarta Online, 2001)  Coal was formed from buried fossils that were exposed to high pressure and temperature over a long time span. 

            “In 1985 chemists created a new allotrope of carbon by heating graphite to extremely high temperatures.”(“Carbon”, Encarta Online, 2001)  This new allotrope, C60, was given the name buckminsterfullerene after an architect-engineer, R. Buckminster Fuller who designed geodesic domes.  A structural diagram of buckminsterfullerene is shown in Figure 1. 

“Because the C60 sphere is hollow, other atoms can be trapped within it.  When a graphite sheet soaked in LaCl3 solution is subjected to vaporization-condensation experiments, a substance with formula LaC60 is formed.”(“Buckyballs”)  Experiments using other metal salts in which laser pulses were used to shrink the ball of carbon atoms yielded such compounds as CsC48 and KC44.  “Other experiments have produced new materials with C60.  For example, C60 doped with potassium is a superconductor below 18 K.”(“Buckyballs”)  Three-dimensional polymers and tubes of carbon called nanotubes have also been made using this new allotrope of carbon.  As for the future of C60, “derivatives of buckminsterfullerene have been found to be biologically active and have been used to attack cancer.”(“Buckminsterfullerene”, Encarta Online, 2001)


Bibliography

1.                  “Carbon”.  Encarta Online Deluxe Encyclopedia.  Available WWW: http://encarta.msn.com/find/Concise.asp?z=1&pg=2&ti=761577017.

 

2.                  “Graphite”.  Encarta Online Deluxe Encyclopedia.  Available WWW: http://encarta.msn.com/find/Concise.asp?z=1&pg=2&ti=761552936.

 

3.                  “Crystalline Allotropes of Carbon”.  Available WWW: http://www.schoolcircle.com/studycircle/chemistry/ch_sr_b_icse_carbon2_pageone.htm

 

4.                  “The Mineral Diamond”.  Available WWW:  http://mineral.galleries.com/minerals/elements/diamond/diamond.htm.

 

5.                  “Diamond”.  Encarta Online Deluxe Encyclopedia.  Available WWW: http://encarta.msn.com/find/Concise.asp?z=1&pg=2&ti=761557986.

 

6.                  “Buckyballs”.  Available WWW:  http://scifun.chem.wisc.edu/chemweek/buckball/buckball.html

 

7.                  Atkins, Peter (1993). The Elements of Physical Chemistry (3rd ed.). New York:  W.H. Freeman and Company.