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ChemMatters April 1990 Page 10
© Copyright 1990, American Chemical Society
Growing Diamonds
By Harold Zaugg
Alchemists of the Middle Ages sweated over their stoves and forges in
futile attempts to convert base metals like lead, copper, and zinc to
gold. Little did they know that the puny heat energy at their command
could never in a lifetime transform even one atom of lead into gold.
They would have been better advised to try turning their glowing coals
into diamonds, for both charcoal and diamond are of the same
element—carbon.
It wasn’t until the late 1700s that the carbon composition of diamond
was proven by Antoine Lavoisier and Humphry Davy, each of whom
conducted the costly experiment of burning diamonds in oxygen and
identifying the product as carbon dioxide (see The Back Burner, page 14).
Ever since, experimenters have tried to turn common forms of carbon
into diamond. They didn’t realize that nature had put formidable
barriers of temperature and pressure before them.
Carbon incognito
Carbon is found in nature in at least six crystalline forms, or allotropes.
(Allotropes are different forms of an element that have the same phase.)
The best known of these are diamond and graphite.
In diamond each carbon atom is bonded to four others, with the four
carbons at the corners of a tetrahedron around the center carbon (see
Figure 1). In contrast, in graphite each carbon atom is bonded to only
three others, the three carbons being at the corners of a triangle around
the center carbon. The result is that graphite is composed of hexagonal
rings arranged as a flat sheet (see Figure 2). In diamond, neighboring
carbon atoms are joined by single covalent bonds. The rings of graphite,
however, are aromatic, that is, the rings contain alternating single and
double bonds.
These atomic arrangements help to explain the astonishing
differences in physical properties of diamond and graphite. Because all
of the valence electrons of diamond are tied up in single bonds, there
are no mobile electrons, and diamond doesn’t conduct electricity. In
graphite the aromatic system has many mobile electrons, and these
account for its electrical and thermal conductivity, which, along with its
relative chemical inertness, makes graphite a good choice for electrodes
in electrochemical applications. Also, the numerous double bonds of
graphite cause it to absorb light of all wavelengths and explain its
opaque, black color, contrasting sharply with the colorless transparency
of pure diamond.
The tight packing and three-dimensional bonding of the atoms in
diamond make it the hardest substance known—it is defined as 10 on
the hardness scale of 0 to 10—as well as the best conductor of heat
(even better than copper and silver). In contrast, graphite is more
loosely packed because its atoms are bonded covalently only in twodimensional planes. There are no covalent bonds between the flat
sheets. Thus the sheets can slide easily over each other. For this reason
graphite is one of the slipperiest solids known, and one of the softest
(less than 1 on the 1 to 10 hardness scale shown in Figure 3). These
properties translate into a wide range of uses, from the black material in
pencils to high-temperature lubricants.
Both diamond and graphite are crystalline (their atoms are arranged
in a regular pattern), but carbon isn’t always crystalline. When
hydrocarbons are burned, the flame contains a certain amount of
unburned carbon—carbon vapor. As the combustion products leave the
flame, the carbon vapor cools too rapidly for it to be crystalline. Instead
it condenses into a powdery amorphous carbon, commonly known as
carbon black or soot. Instead of forming an orderly pattern, the carbon
atoms are scattered about in a jumbled mass. Carbon black has been
described as “poorly graphitized carbon.” It is a useful industrial
material—world production is 10 million tons per day!
Diamonds into dirt
Diamond is harder and denser than graphite but technically it is not as
permanent as graphite. Under ordinary conditions, diamond is
thermodynamically slightly less stable than graphite (by about 40
calories per gram). This means that there is a natural tendency for
diamonds to change into graphite, for gemstones to decay into black
powder! Fortunately this is not a practical problem because, to get
diamond to turn into graphite at an appreciable rate, it must be heated
to more than 1200 °C in the absence of air. This high-energy barrier of
conversion is caused by the many stable sp3 bonds of diamonds, which
must be broken before the electrons can be rearranged into the sp2
bonds of graphite.
Dirt into diamonds
As you might imagine, there is little interest in changing diamonds into
graphite, and the challenge is to go the other way. Although high
temperature alone is enough to change diamond to graphite, high
temperatures and high pressures are needed to reverse the process.
Pressures of many hundreds of thousands of pounds per square inch
are required to squeeze the graphite carbons into the tighter packing
arrangement of the diamond crystal. Many scientists believe that
natural diamonds are formed under such conditions, 100 or 200 miles
below the surface of the Earth, and are later brought by volcanic action
closer to the surface where they can be mined.
Laboratory duplication of these harsh conditions was finally
accomplished in 1954 by four General Electric scientists, Francis Budy,
Tracy Hall, Herbert Strong, and Robert Wentorf. The team of General
Electric scientists developed a small chamber that could maintain
temperatures up to 2500 °C and pressures of 1.5 million pounds per
square inch (1.0 x 107 kPa). Their first attempts not only failed to
produce diamond but also failed even to melt the graphite. The melting
temperature of graphite is above 4000 °C, far beyond the chamber’s
maximum of 2500 °C.
The researchers finally hit upon the idea of adding small amounts of
a metal that does melt below 2500 °C, in the hope that some of the
graphite might dissolve in the molten metal then crystallize out as
diamond. The strategy worked beautifully with many metals:
chromium, manganese, iron, cobalt, nickel, tantalum, ruthenium,
rhodium, palladium, iridium, and platinum. When the mixture reached
1200-2400 °C under pressures of 750,000 psi to 1.4 million psi, small
diamond crystals formed at rates as high as 0.1 mm/min. At last, a
practical synthesis of diamond had been achieved.
The crystals—usually yellow or green and less than 1 mm across—
were not gem quality, but they were as hard as natural diamond and
were useful as an industrial abrasive. Today, annual worldwide
production of synthetic diamond abrasive by the high-pressure, hightemperature (HPHT) method is measured by the ton, far outstripping
natural diamonds in this application. Synthetic HPHT diamonds as
large as 8 mm across have been made, but the long heating and
pressing time they require—seven days—makes them more costly than
natural diamonds of the same size. Therefore, the large diamonds used
in jewelry are still natural diamonds.
Vapor into diamonds
We have all seen frost form on a windowpane during a cold night. The
water vapor in the air has directly formed crystals of ice. Similarly,
carbon vapor can condense into amorphous carbon, and this is what
happens when soot forms on the glass chimney of a kerosene lamp.
Scientists have long been attracted to the idea of condensing carbon
vapor into crystalline carbon—growing diamonds from vapor. In the
mid-1950s W. G. Eversole at the Union Carbide Corporation passed
methane gas, CH4, over a hot tungsten filament and then onto some
seed diamond. He hoped that the carbon vapor, formed by stripping
hydrogen atoms from the methane, would condense as a new layer of
diamond over the original diamond substrate. Unfortunately, more than
99% came out as graphite. He did learn, however, that the graphite
could be removed selectively by reducing it with hydrogen at high
temperature and pressure, leaving behind a very thin layer of new
diamond.
H2 + heat from filament ⇒ 2H
C(graphite) + 4H ⇒ CH4
C(diamond) + H ⇒ no reaction
He eventually achieved a 60% increase in the weight of the diamond
seed, but only after 85 cycles of crystal growth and graphite removal.
Meanwhile, in Moscow, Boris Derjaguin and his co-workers had
been performing similar experiments since about 1955. While
conducting basic studies of the way diamond grew on diamond seeds,
Derjaguin began to wonder whether atomic hydrogen was responsible
for selectively removing graphite from the material deposited on the
seed diamond.
In the late 1970s Derjaguin tried the crucial experiment of using only
a small amount of methane (about 1%) in a large excess of hydrogen
(see Figure 4). He was gratified to discover that essentially No graphite
was deposited but that a microcrystalline film of diamond formed on
the seeds. Graphite removal was not needed, and the diamond film
could be grown continuously. Thanks to Derjaguin’s insight, growing
diamond films became a practical reality. Pure diamond could be grown
at the rate of 0.01 mm/h, forming ice-like films over objects of any
shape. Derjaguin’s lab made pure diamond films that were a full
millimeter thick.
Today there is much active research in growing diamonds from
vapor, and new methods have been devised. Many techniques besides
the use of a hot filament can generate a gas of carbon and hydrogen
atoms. These include use of microwaves, laser beams, electric discharge,
and ultraviolet light. Also, experimenters have found that methane is
not the only suitable source of carbon: Many organic compounds and
even solid carbon nicely suffice. And diamond is not the only surface
on which the films will grow: Silicon, quartz, and other materials can be
used, although the material must be able to withstand the high
temperatures of the process. Recently a Japanese experimenter reported
that a beam of hydrogen ions can be used to grow diamond films at
room temperature. This holds the potential for diamond coating on
heat-sensitive substrates.
Diamond coated
Uses for diamond films are easy to imagine. Because diamond is
transparent to infrared, ultraviolet, and microwave radiation, it can be
used to make rugged lenses and windows for lasers, telescopes, and
other scientific instruments. Knives and razor blades coated with
diamond films would never need sharpening.
Of great importance would be the use of diamond chips in the
integrated circuits of future computers. Diamond-coated computer
chips would be extremely efficient and sturdy, allowing them to operate
under severe conditions of heat and stress. Their unmatched ability to
conduct heat would permit closer spacing of integrated circuits than is
now possible with chips made of silicon. Future supercomputers could
be substantially reduced in size and weight.
The future of diamonds has never been brighter. It’s not hard to
imagine, some years from now, an engineer sitting at the console of a
diamond-chip-powered computer and using a light pen, with a
diamond optical window, to draw on the computer’s diamond-coated,
scratch-proof screen, while she designs a new, long-lasting lawn mower
blade that has a diamond-hard edge. Diamonds are this engineer’s best
friend.
SIDE BAR
Hot stuff
Andrew Good, a high school student—and inventor—from Adelaide,
Australia, recently developed a new process for growing diamond
films. Andrew used an oxyacetylene plasma (a high-temperature flame
similar to a welding torch) into which he injected a stream of hot
methane gas (the source of carbon) that was surrounded by a stream of
cool hydrogen gas. He focused the plasma electrostatically. The process
grew a hard diamond coating on stainless steel at the rate of 0.5 µm per
minute. Andrew won Australia’s Beyond 2000 Award in Science and
Technology, which included a trip to the United States as prize.
CAPTIONS
Marilyn Monroe singing “Diamonds are a Girl’s Best Friend”, in the film Gentlemen Prefer
Blondes.
Figure 1. Diamond is the hardest substance known because of the bonding pattern of its
carbon atoms. Each atom has four covalent bonds, which join it to four neighboring atoms.
The bonds are oriented at angles of 109° to each other, as though they were pointing toward
the corners of a tetrahedron (sp3-hybridized geometry).
Figure 2. Graphite, the black material in lead pencils, has little physical strength because its
atoms are bonded to each other in only two dimensions (sp2-hybridized geometry). The
resulting flat sheets readily slide over each other. The rings of graphite are aromatic and,
decades ago, were believed to be made up of alternating single and double bonds, as shown
here (for ease of illustration). It is now believed that some of the bonding electrons are evenly
distributed above and below the graphite plane and that these electrons can move freely
throughout the graphite sheet. This explains why graphite is the only nonmetal found in
nature that conducts electricity.
Figure 3. The limits of the Mohs hardness scale were originally defined by assigning talc a
value of 1 and diamond a value of 10. On this scale, diamond’s crystalline allotrope, graphite,
has a hardness between 0.5 and 1.
Figure 4. Russian scientist Boris Derjaguin was the first to grow diamond from vapor. When
he passed a gaseous mixture of 1% CH4, 99% H2 over a hot wire and onto a seed crystal, a
film of diamond grew on the crystal. Diamond films, which look like ice coatings, have been
grown as thick as a millimeter.
BIOGRAPHY
Harold Zaugg, now retired from Abbott Laboratories where he was a
senior research fellow, is a free-lance science writer. He extends a thank
you to Rustum Roy, Pennsylvania State University, for assisting with
this article.
REFERENCES
Bovenkerk, H. P.; Bundy, F. P.; Hall, H. T.; Strong, H. M.; Wentorf, R. H. Nature 1976, 184, 1094.
DeVries, R. C. “Synthesis of Diamond Under Metastable Conditions,” General Electric
Corporation Technical Information Series (Class I), March 1987.
Messier, R.; Badzian, A. R.; Badzian, T.; Spear, K. E.; Backmann, R.; Roy, R. “From Diamondlike Carbon to Diamond Coatings,” unpublished preview.
Robinson, A. L. Science 1986, 234, 1074.
Sneed and Brasted, Comprehensive Inorganic Chemistry, Vol. VII, Chapter 1, “Carbon,” Van
Nostrand: Princeton, NJ.