The purpose of this report is to examine the reasons why Au is found in nature in the pure metal form and not as an oxide. This analysis will examine the topic by considering the morphology of Au and the thermodynamics of the Au/O reduction-oxidation reactions.
A review of the literature indicated a scarcity of research on the reactivity of Au with oxygen.
Gold oxides are unstable compounds and are usually formed by indirect methods. Gold(I) oxide, Au2O, is a covalent compound and is formed by heating AuOH to 200oC. Gold(I) oxide easily reduces to the gold metal. Gold (I) hydroxide, AuOH, is prepared from a Au(I) solution by the addition of potassium hydroxide solution. AuOH dissolves in excess alkali to form aurates, i.e. KAu(OH)2. Gold(III) oxide, Au2O3, is formed by heating Au(OH)3 at 100oC in the presence of a dehydrating agent. Au2O3 also is easily reduced to the metal and dissolves in excess of alkali hydroxide, forming an aurate with the ion [Au(OH)4]-. Gold(III) hydroxide, Au(OH)3, is precipitated by the addition of potassium hydroxide solution in equivalent amount, to a solution of chloroauric acid. Au(OH)3 is insoluble in water, gives many of the reactions of Au2O3, and may be a hydrous form of that compound. Gold(II) oxide is believed to consist of gold(I) and gold(III).[5,9,10]
Structurally, Au2O3, the most stable Au oxide, exists as a square planar compound, exhibiting distorted octahedral geometry coordination on the Au3+ ion and a tetrahedral coordination on the O2- ion. The distorted geometry occurs as a result of d8 orbital splotting, the Jahn-Teller effect. Additionally, radius ratio rules support a distorted octahedral geometry by predicting an octahedral coordination for Au3+ (r = 0.85 A) and a tetrahedral coordination for O2- (r = 1.32 A). [3,5,7] Despite its observed geometry, Au2O3 and other AuxOy compounds exist in an amorphous state. Attempts were made to crystallize Au2O3 by heating the amorphous powder in air but the oxide decomposed into metallic Au. [4]
Buehrer and Roseveare [8] investigated the free energy of formation of auric oxide, Au2O3. They used direct measurements of the potential of the auric oxide electrode against a hydrogen electrode at 25oC. The following reported results combined the free energy of the reaction taking place with the free energy of formation of one mole of liquid water from its elements:
Subtracting (2) from (1) to obtain:
This high positive free energy of formation of the auric oxide indicates a very unstable compound. Buehrer and Roseveare end their discussion with a warning that the auric oxide involved in the above equations is very probably not Au2O3, but a hydrated form of the oxide.
In addition to free energy calculations, the stability of Au/O compounds in the presence of either gasous O2 or water can be considered through high pressure oxidation reactions under increasing temperature and predominance area (Pourbaix) diagrams.
A study by Muller et al. [4] on crystalline Au oxides included heating Au(OH)3, AuI, AuO(OH), and Au2O3 at high oxygen pressures. None of the starting compounds, except for AuO(OH), yielded crystalline oxide before dissociating to the metal. Muller states that AuO(OH) heating resulted in a "very poorly crystallized" product. Results from runs using Au2O3 as the starting material are shown in Fig. 2. The Au2O3 dissociated under pressures of approximately 250 to 2500 atm. and temperatures of 125oC to 225oC into AuOx and AuOy compounds. Under same pressures with increasing temperature, the Au2O3 dissociated into the Au metal.
These results do not appear to support the premise that the Au oxide compounds have greater stability than the Au metal. Two points concerning this diagram should be noted: 1) the data does not capture reactions occurring below O2 pressures of 100 atm and 2) the number of data points below 1000 atm O2 pressure does not yield a good confidence level. One hypothesis that could be made for lower pressure / lower temperature reactions is that the positive slope of the near verticle phase boundary for the stability of the Au metal would decrease to the point that the area of stability of the Au metal would be dominant. By regraphing the data on a ln PO2 vs T plot, a linear relationship would probably be seen with the Au metal area of stability located at the pressure and temperature ranges existing in nature.
To illustrate the stability of gold and its oxides at equilibrium in the presence of water at 25oC, the Pourbaix diagram in Figure 1 plots the potentials of Au/O compounds as a function of pH. [5,6,9] This data indicates that the Au metal would be more stable than both the Au oxide compounds and ionized Au atoms under the conditions for stability of H2O. Since the Au metal has no coexistence boundaries with the Au oxides within the H2O boundaries (signified by the dotted lines a and b), the oxides would not be produced directly from the metal by oxidation in the presence of water. Above a pH of 9, the limiting boundaries of relative predominance of dissolved Au+, H2AuO3-, and HAuO3- enter the area of H2O stability.
2. Cox, P.A. (1992). Transition Metal Oxides, Clarendon Press, Oxford.
3. West, Anthony. (1984). Solid State Chemistry and its Applications, Wiley, New York.
4. Muller, O. et al., (1969). J. Inorg. Nucl. Chem, 31, 2966-2970.
5. Douglas, B. (1983). Concepts and Models of Inorganic Chemistry, 2nd ed., Wiley, New York.
6. Pourbaix, Marcel. (1966). Atlas of Electrochemical Equilibria, Pergamon Press, New York.
7. Weast. (1992). Handbook of Chemistry and Physics, CRC Press.
8. Buehrer, T. F. and Roseveare, W. E. (1927), J. Amer Chem Soc., 1989.
9. Puddephatt, Richard. (1978). The Chemistry of Gold, Elsiver, New York.
Figure 2. PO2 -T conditions used to synthesize gold oxide phases using Au2O3 as the starting material.