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Our Aim

Much of the current work is aimed at the development and exploitation of solvents for the preparation of small crystals of new nitrides, especially multinary nitrides containing transition metals (TM) and/or rare earths (REs). Small crystals are generally needed to determine the crystal structure and we get the opportunity to learn how to perform all the measurements and solve the structure ourselves.

We continue to look at sodium based fluxes but, in order to synthesize TM containing nitrides we alloy this with such elements as zinc, gallium and indium. We have already had some interesting results from this area, including the compound shown below, AE2[Cr2N6] (AE = Ca or Sr). This highlights that different compounds can be made from the same set of elements, depending upon the synthetic route. The sodium-gallium flux allowed us to use lower temperatures than had been used before in our synthesis of Ca3CrN3 and hence we make a different compound. A pure sodium flux continues to be explored in an effort to synthesize large (cm scale) single crystals of GaN, a compound with tremendous promise in all areas of electronics. We are also looking at Li3N as a flux for transition metal nitrides, something that our collaborators in Germany have been doing for some time now.

In addition we are also interested in luminescent nitrides.

Our Challenges

The main challenges of nitride chemistry reside in the seemingly contradictory refractory behavior and the propensity for decomposition by loss of N2. That is, nitrides typically have high melting points (Tm) but at the same time have high nitrogen equilibrium pressures at Tm. This is a direct consequence of the large triple bond energy of dinitrogen (941 kJ/mole). This energy is very large compared to dioxygen (498 kJ/mole) or disulphide (352 kJ/mole); in fact, only CO has a stronger bond at 1072 kJ/mole. Thus, decomposition occurs at much lower temperatures for nitrides than for oxides,

Consider: xM + y/2 N2 = MxNy

as always ΔGf = ΔHf - TΔSf. Since N2 is a gas, ΔSf is large and negative. The essential point for nitrides is that the magnitude of ΔHf is also small (but negative) due to the high energy cost of breaking the dinitrogen bond in the reaction. Thus the free energy of formation, ΔGf, for nitrides becomes positive at much lower temperatures than for oxides (oxides have a similar ΔSf, but more negative ΔHf). In most cases, the melting points (Tm) of nitrides are sufficiently high that at one atmosphere pressure ΔGf at Tm is already quite positive. Of course, the entropy of the gas is lowered by increasing the pressure, thus increasing the decomposition temperature. For most nitrides the equilibrium N2 pressure at melting (for ΔGf = 0 at Tm) is hundreds or even many thousands of atmospheres. There are several known exceptions: Li3N melts at 850 °C and Li3BN2 melts congruently at about 915 °C, both under 1 atm of N2.

Loss of N2 from nitrides on heating may proceed in steps, rather than decomposing directly to the elemental species; e.g. MxNy = MxNy +z/2 N2. For example, Ta3N2 under 1 atm. of N2 will lose N2 to form TaN at about 850 °C. Ternary (or quaternary) nitrides may also decompose into binary (or binary/ternary) phases on nitrogen loss. Consequently, to obtain the largest number of nitrides, especially nitrides at the highest nitrogen content, the synthesis temperature must be kept as low as possible.

However, low synthesis temperatures pose a kinetic challenge. In many reactions the activation barrier for N2 bond breaking on the surface of the reactants is very high (again consistent with the large dinitrogen bond energy) and high temperatures (1200 °C and above) are required for reasonable reaction rates. In a number of cases, such high temperatures are above the decompostion temperatures of the target phases. In those cases the nitrogen must be introduced to the reaction in more reactive forms than N2 or the decomposition of N2 must be catalyzed. Ammonia is a suitable reagent to introduce reactive nitrogen (but depending on the conditions, H may also be introduced in the product); in other cases the presence of certain elements allows the catalytic decomposition of N2. For example, clean Li metal will react with N2 even at room temperature! Alkaline earths also reaact with N2 at relatively low temperatures, starting at about 500 °C. Yet reasonable reaction rates also depend on modest solid state diffusion rates. However, due to the refractory nature of many nitrides, the diffusion constant for N is the solid is usually very low - typically lower than that for O in oxides at the same temperature. Unless there is a high cation diffusion rate in the product nitride, the overall reaction rate with N2 will be very slow. As examples of possible extremes, Li diffuses very rapidly in Li3N allowing the continued reaction of N2 with Li at room temperature, but the NbN coating on Nb metal only grows to about 1 micron thick after 24 hours at 1400 °C in 1 atm. of N2. Even when the reactants are liquids or dissolved in fluxes, a thin "crust" of solid nitride may form on the surface greatly impeding further reaction, e.g. on molten Al even at 1400 °C.