Diamond is not an average substance. We are all familiar with the characteristics of jewelry-grade diamonds that make them among the most coveted gems: their extreme hardness (diamonds are the hardest substance known), resistance to tarnishing, exquisite clarity, clean facetting, and lack of color (unless they contain small amounts of metal impurities which can impart slight tints and may increase their value even more). Diamonds demonstrate distinctive properties of which you may not be aware: e.g., diamonds are excellent conductors of heat but poor conductors of electricity. The heat conductivity of diamond is so extraordinary that modern tests to determine whether a gem is a genuine diamond or a fake is based on measurement of the gem's heat conductivity.
Diamonds play another role in modern society: an industrial workhorse. The unusual properties of diamond makes it technologically very important. Because of its extreme hardness diamonds are excellent for surgical cutting tools and coatings on cutting tools for drilling, mining, and industrial productions. A rapidly emerging technology centers on thin films of diamond. Diamond films have been marketed on tweeters in stereo speakers and as scratch-proof coatings on watches. Diamond-coated windows for infrared-scanning systems and light-filtering masks are also on the market. Diamonds offer promise for new electronics materials. Diamonds containing small amounts of other elements (generally called dopants) dramatically change the electrical properties of the diamonds. Doped diamonds are semiconductors, much like the electronics based on silicon and compounds such as GaAs, in contrast to undoped diamonds which are non-conducting insulators. Diamond-based chips are desireable because they can operate at 1000 °C, handle very high power, and operate in highly radioactive environments. Figure 4.1 shows headlines from some of these current areas of interest in diamond films. Despite the tremendous promises of diamond film technology, applications of diamond films are not widespread yet.
Figure 4.1 Newspaper headlines
Why aren't all eyeglasses coated with diamond films to keep them from being scratched? Why aren't diamond-based electronics chips used to regulate the engines in all automobiles? Even Paul Simon's suggestion of "Diamonds on the Soles of Her Shoes" suggests a possible wear-limiting application. Put simply, large-area, high quality diamond films are very hard but hard to make. In 1990, diamond was named "molecule of the year" by Science magazine in recognition of the progress that has been made toward practical synthesis of diamond films. Although diamond is not a molecule and the synthesis is still not ideal, the world market for synthetic diamonds was already between five hundred million and one billion dollars in 1990. As diamond film methodology improves, scientists predict that the value of the market for diamond films will be four to five billion dollars by the year 2000.
On the basis of their macroscopic properties, it seems impropable that diamonds and graphite are made of exactly the same atoms. This begs the question, how do we know that diamonds are indeed carbon? Surprisingly, this has been known for over 200 years. In 1772 (before the signing of the Declaration of Independence) a talented French chemist, Antoine Lavoisier, burned diamond (ca. 150 milligrams) in a sealed container. Diamonds do not burn readily, so Lavoisier used the high heat provided by a large lens placed in the sunlight. By a simple analysis of the gases produced by burning, Lavoisier determined that the gas we now know as CO2 was produced. Furthermore, the same amount of this gas was produced when 150 milligrams of amorphous carbon in the form of charcoal was burned. Thus, Lavoisier concluded that charcoal and diamond are the same. However Lavoisier was reluctant to publicize such a radical and counterintuitive result. It took more than 20 years before Lavoisier's conclusions were confirmed by Tenant, an English chemist. Let us examine Lavoisier's conclusion in the context of the times. Lavoisier's discovery predated Dalton's formulation that all matter is made of elements by 29 years and Mendeleev's formulation of the periodic table by almost 100 years. Our modern understanding of chemistry allows us to formulate Lavoisier's experiments in a compact way, the chemical equation.
These equations indicate that burning either diamond or graphite causes the bonds of the reactants (the carbon material and dioxygen) to be rearranged to form the product (CO2) and energy in the form of heat. In general, a chemical reaction involves changes in the bonding between atoms in going from reactants to products. Let's contrast these combustion equations with the analogous equation for burning buckyballs.
Note how the buckyball equation has a subscript of 60 on the carbon, indicating that this form of carbon is molecular. As a result, 60 dioxygen molecules are required to convert the sixty atoms of carbon in each buckyball into 60 molecules of CO2. In a proper, balanced chemical equation the number and types of atoms on the reactant (left) side of the equation equal those on the product (right) side. Because heat is released in this chemical reaction, heat is written as a product. Reactions that generate heat are called exothermic.
Ever since the late eighteenth century experimenters have tried to synthesize diamond. Most of the earliest efforts (and indeed most of the industrial processes today) focused on converting graphite into diamond at high pressure. Under normal conditions, graphite is the most stable form of carbon because the bonds in graphite are stonger. In part this stability can be attributed to resonance. Having the double bond character spread evenly throughout the entire structure adds extra stability to graphite. When several different forms of matter with the same composition exist, the most stable structure is called the thermodynamically favored form. Under the usual conditions of atmospheric pressure and room temperature, the more stable or thermodynamically favored form of carbon is graphite. However, at high pressure, diamond becomes more stable than graphite. We can understand why the preferred arrangement of atoms may change with pressure. On average the carbon atoms in diamond are closer together than the carbon atoms in graphite. As a result diamonds are more dense (that is, more mass in a given volume) than graphite. When a piece of graphite is subjected to high pressure, the external forces compress the graphite or push the carbon atoms closer together. Under these conditions it is more favorable for the carbon atoms to rearrange themselves into the diamond structure. If graphite is squished (very high pressure is applied) then the graphite will be transformed into diamond.
One common way of depicting the relative stability of different materials is by using an energy diagram. It is generally true that more stable materials have lower potential energies. Figure 4.2 shows an energy diagram for the conversion between diamond and graphite. Compounds that are more stable are placed lower on the diagram (i.e., at lower potential energy) while the less stable compounds (higher potential energy compounds) are higher on the diagram. The x-axis in the energy diagram is the reaction coordinate and indicates the progress of the reaction. Thus, moving from left to right in Figure 4.2 shows the conversion of graphite to diamond. Because the density changes in coverting from graphite to diamond, one could quantitate the reaction coordinate for this transformation with the density.
Figure 4.2 Energy diagram for conversion of graphite to diamond and diamond burning
4.4 How do stability and energy affect diamond formation?
Because graphite is lower in energy than diamond, the conversion requires addition of energy to the system. Such reactions are called endothermic reactions. Contrast this situation with the burning of carbon allotropes; combustion reactions are exothermic. To a good approximation, we can rationalize the endothermicity of the graphite to diamond conversion by comparing the relative bond strengths. On average the C-C bonds of graphite are stronger than the bonds of diamond. If we were to turn our perspective around and look at the conversion of diamond to graphite, we would say the reaction is exothermic. That is, it is energetically more favorable for diamonds to simply turn into graphite.
If carbon atoms would rather be arranged as the graphite structure rather than diamond, why don't these priceless gems discolor and crumble into low value graphite powder? Examine Figure 4.2. In order to trace from diamond and graphite one moves from right to left on the diagram. Notice that a high energy barrier (like a hill) must be passed over. The top of this barrier is called the transition state. The transition state represents the highest energy structure involved in a reaction. It is intrinsically unstable and can't be isolated. What might the transition state for the diamond to graphite conversion look like? We might imagine that it is a structure in which the some of the C-C bonds are stretched and the carbon atoms are beginning to arrange into the planes of graphite. At the transition state these planes will be buckled due to stretched bonds between the ultimately non-bonded planes of carbon atoms of graphite. As a result the stabilization afforded by four single C-C bonds at each C atom has begun to be lost, but the the resonance stabilization of graphite has not been fully realized. As a result the C-C bonds in the transition state on average are weaker than the bonds of either diamond or graphite. Therefore, the potential energy of the transition state is greater than that of reactant and product. The energy difference between the starting material (graphite) and the transition state is the activation energy, denoted by the symbol Eact. Note that the value of the activation energy depends on the direction of the reaction; the activation energy for the diamond to graphite transformation is lower than that for the graphite to diamond transformation. If the activation energy for a reaction is high, the reaction will occur slowly. The rate of reaction is determined by how frequently the atoms collect enough energy to "climb" over the hill. Therefore we expect that the higher the barrier the slower the reaction. Also we expect that increasing the temperature, or kinetic energy of the atoms, will increase the reaction rate because the atoms will collect the energy needed to "climb" the barrier more frequently.
Note that the rate of reaction is not determined by the reaction thermodynamics. An exothermic reaction can be very slow if the activation energy is high and an endothermic reaction can be fast if the activation energy is low. Because the activation energy in the reaction converting diamond to grapthite is high, much energy in the form of high temperature is needed to convert the starting materials (diamond) to the transition state so that the product (graphite) can be formed. The potential energy change of the reaction as a whole is the difference in energy between the starting material (diamond) and the product (graphite). The potential energy change of this reaction is -2,900 Joules for every 12.01 g of carbon. A Joule is a unit of energy that corresponds to the energy needed to lift 2 kg of mass by 10 cm in the presence of the earth's gravitational field. By convention, the plus sign indicates that the reaction is endothermic (or uphill) and a negative sign for the potential energy change indicates that the reaction is exothermic. Although diamonds convert to graphite exothermically, the reaction is slow due to the high activation energy; diamonds stored at ambient temperatures do not crumble into black graphite.
Diamonds are formed deep inside the earth's interior where crushing pressure and blistering heat work together to create the diamond lattice. Although scientists have long puzzled about how diamonds are transported to the earth's surface, a recent discovery of diamonds in Canada may have shed some light on this issue (Figure 4.3). Scientists now believe that narrow volcanic pipes running down into the earth's interior allowed diamonds to be transported via violent eruptions to the earth's surface. The eruptions were so fast and so violent that the diamonds were coughed straight to the surface. Although the pressure release as the diamonds rose to the surface conceivably could have allowed for transformation to graphite, the explosions were believed to be so fast that the diamonds reached cool temperatures at the surface quickly. Due to the rapid cooling, the back-transformation of the diamonds to more stable graphite was too slow to occuren route . Slow kinetics for the diamond-to-graphite transformation allowed the less stable allotrope of carbon to make it to the earth's surface.
Figure 4.3 Newspaper article
Synthetic diamonds can be formed in much the same way as natural diamonds; graphite is heated to temperatures exceeding 1500 C at about 60,000 atmospheres of pressure (that is a pressure that is 60,000 times greater than the pressure exerted by our atmosphere). Even at these temperatures and pressures, diamond formation is not easy. Addition of small amounts of the metallic elements, iron or nickel, speeds up the reaction. Why are metals added? The role of the metallic additives is to reduce the amount of energy needed to form diamond from graphite. Figure 4.4 shows an energy diagram for the conversion of graphite to diamond with and without metals. The metal in the reaction with the lower activation energy acts as a catalyst in the reaction. A catalyst is a substance that affects the rate of a reaction without being consumed by the overall reaction process. Chemists say that catalysts do not affect the thermodynamics of the reaction, rather they affect the kinetics. In diamond formation, the metal catalysts are trapped in the diamond lattice as it forms. This is why many synthetic diamonds are colored while naturally occuring diamonds which lack impurities are colorless.
Figure 4.4 Catalyst Energy Diagram
It has recently been shown that, if buckminsterfullerene is used instead of graphite, lower pressure (20+5 Gpa or about 20,000 atmospheres) and ambient temperature is enough to cause the transformation to diamond. This synthesis is mild enough that catalyts are not necessary. The synthesis of diamonds from buckyballs differs their synthesis from graphite in two important ways. First, the buckyball conversion to diamond (and to graphite) is exothermic. The C-C bonds of buckyball are weaker than those of either diamond or graphite. Second, the activation energy for the buckyball conversion is much lower than that for graphite. Both features combine to allow a low temperature pathway.
The extensive network of carbon atoms makes diamond a very hard material. The only material that can cut diamond is diamond. Despite its hardness, if diamond is hit at just the right angle with a chisel, it will cleave to form two stones with perfectly flat faces. Until better tools were available, jewelers relied upon this ability to break diamonds with the sharp raps of a hammered chisel to make diamonds of desired sizes and shapes.
Why do diamonds break when force is applied at certain angles? If enough pressure is added to one plane of carbon atoms in a diamond, then that plane of atoms will begin to move (Figure 4.5). Essentially, the mechanical force will cause some of the C-C bonds made by atoms in this plane to stretch and some to compress from their normal values. As the atoms are further displaced the interactions eventually become so undesireable that the diamond splits apart along the planes where the bond stretching and compression is most undesirable. The key to obtaining a clean break is to orient the chisel such that force is applied along the direction of a plane of carbon atoms.
Figure 4.5 Diamond Cleavage.
More Info: A description of Jewelry process with Figure 5 Jewelry processes
For many years, only high pressure synthesis of diamond from graphite was considered possible. However, it has recently been discovered that diamond films can be made at low pressures if one adopts a different strategy. Rather than starting with graphite, the reactant is methane (CH4), a simple molecule containing carbon and hydrogen atoms. Note that methane satisfies the bonding rules of carbon by having four bonds to carbon and one to each hydrogen. In this reaction, methane is decomposed into diamond rather than graphite even though graphite is more stable. How can this occur? In a reaction, formation of the thermodynamically favored product (graphite) can be slower than formation of a less stable product. When more than one product can be formed, the product that forms faster is called the kinetic product. Such a situation occurs for the methane decomposition because the activation energy for forming diamond films from methane is lower than the activation energy for making graphite. Such a situation is not uncommon. Indeed, when an ordinary candle burns, an amorphous, sooty carbon form is synthesized. This soot is higher in energy than graphite, but once formed it is very slow to convert further to graphite. The decomposition of methane to give diamond films is similar. Figure 4.6 depicts the reaction energy profiles for a situation in which the kinetic product is not the thermodynamic product.
Figure 4.6 Energy diagram displaying kinetic vs thermodynamic control.
Synthesis of diamond films at low temperature must be controlled so that graphite formation does not compete with diamond formation. Although the reasons are not fully understood, the most successful way to keep graphite from forming is to form the diamond under a constant stream of hydrogen. CVD, or chemical vapor deposition, is the most common low temperature method for synthesizing diamond films. Figure 4.7 shows a picture of the deposition vacuum chamber for CVD growth of diamond films. Inside the chamber is a tungsten filament source of intense heat. This heat decomposes a carbon source such as methane as it is carried through the chamber on a stream of hydrogen. Carbon fragments are deposited on target surfaces at 600 to 900 °C that are also in the chamber. It appears that the role of hydrogen atoms may be to etch away any graphite that is formed so that only diamond crystals can grow. A diamond film made of countless tiny diamond crystals builds up slowly. The target surface can be silicon, graphite, diamond or tiny diamond chips, or even C70. Microwave discharges and oxyacetylene torches have been used in place of the tungsten filament heat source, and the methane gas carbon source has recently been replaced by buckyballs to give high quality diamond films.
Figure 4.7 Deposition Chamber and Schematic Diagram
Due to limitations of the CVD method, scientists are searching for lower temperature and pressure methods for the synthesis of flawless diamond films. The CVD process requires high temperatures and pressures and produces a diamond-like film where many small diamond crystals are aligned together rather than one large single crystal of diamond. As a result many materials cannot be coated with diamond films because they will melt or decompose in the coating process. A potential alternate process involves passing carbon soot slowly through a laser beam. This process has been shown to make diamond films. An intriguing new process takes a polymeric carbon source called poly(phenylcarbyne) and heats it in a furnace to produce a diamond-like film (Figure 4.8).
Figure 4.8 Newspaper article.
The first single crystal film was grown in 1991 by using an accelerator to ram carbon ions into a copper surface at a density of about one billion ions per square centimeter. Then the carbon-laden copper was blasted with a series of powerful laser bursts which rapidly liquified the topmost copper layer. Once the laser pulses stopped, the copper solidified and the carbon atoms aligned on its surface to make a diamond film. The film was too small for any electronic applications, but at 100 square microns in area it was a big advance.
More info: endothermic/exothermic.
Diamond films are manmade replicas of the diamonds made by nature. Like natural diamonds, they have an extensive lattice or orderly array of carbon atoms bonded together in a tetrahedral arrangement. Each carbon has four bonds to neighboring carbons. (Figure 4.9) Diamond films, however, are not as ordered as natural diamonds. If they are made at high pressure, they include metal impurities that create slight imperfections in the crystal lattice. If diamond films are made by CVD, they are actually many tiny crystals of diamond that have grown against each other. These diamond films can be grown to be at least twelve inches long, and their thickness can vary from just a few nanometers to many microns or even larger. Which technique is used to grow the diamond film affects the size of the diamond more than anything else. Figure 4.10 shows several examples of diamond films.
Figure 4.9 Diamond lattice
Figure 4.10 Diamond films
Even though diamond films are not as ordered as natural diamonds, they mostly retain the diamond lattice of natural diamonds. It is not surprising, then, that they also retain the properties of natural diamonds. As in natural diamonds, the extensive network of bonded tetrahedral carbon atoms make diamond films very strong and difficult to break or compress. Also, the diamond lattice of diamond films spaces atoms at an ideal distance to serve as connections in electrical circuitry if metal impurities have been included during diamond growth.
Although many big advances have been made, several problems still keep diamond films from being useful as electrical chips and other devices. One major obstacle is expense. CVD diamonds cost well over a hundred dollars a carat to produce. (High pressure diamonds sell for slightly over one dollar a carat, but they contain too many flaws for electronic circuitry and many other applications.) Before CVD diamond films find wide practical use, better adhesion to the substrate must be achieved and lower temperature processes must be developed.
Recent work with diamond films has led to synthesis of similar compounds composed of carbon and nitrogen atoms. In 1989, computer calculations suggested that a compound composed of carbon and nitrogen with the formula C3N4 in a diamond-like structure (Figure 4.11) would be even harder than diamond if it could be made. It wasn't until 1993 that the synthesis of this compound was achieved, and still only very small amounts of C3N4 have been made. Because only small quantities of this material, commonly called carbon nitride, are available, it remains to be proven that the substance is indeed harder than diamond.
Figure 4.11 The structure of Carbon Nitride (Chem and Eng. News)
In order to make a compound harder than diamond, the compound should have a denser packing of tightly bound atoms than diamond. This suggests that the material must have bonds that shorter and stronger than those in diamond. The idea that a carbon-nitrogen mixture might allow for very short bonds and very hard materials springs from the fact that silicon-nitrogen mixtures of the same type are known to be very strong and very hard. They actually rival diamond in hardness. Recall that elements in the same row of the periodic table have similar properties; because silicon is in the same row on the periodic table as carbon, the properties of these elements are similar. Because silicon is lower in the periodic table than carbon, it is larger. The larger atoms of silicon cannot fit together as closely as carbon atoms can, so their bonds must be longer. Carbon atoms bonded to nitrogen can have shorter bonds than silicon atoms bonded to nitrogen simply because they are smaller. The shorter the bond, the more energy is required to break the bond and the harder the material.
If materials such as C3N4 truly are harder than diamond, then the future of diamond films will also include films of this material. Whether diamond films or related films will be most applicable remains to be seen. In any event, diamond films are an exciting new area with potential that has only begun to be realized, and use of diamond films is expected to skyrocket in the future.
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