The chemistry of coinage metals is rich and surprising, and considerable differences are found between copper, silver, and gold.1 Among the compounds of Group 11 metals, those of divalent silver are in many ways unique and owe their properties to the remarkable electron-withdrawing power of oxidizing silver(II).2–4 As a consequence, many simple binary and ternary connections of divalent silver have not yet been prepared. Particularly intriguing is that a sulfate of silver(II) has not yet been synthesized, given that anhydrous CuSO4 has been meticulously characterized in the past,5 and the related AuSO4 has been known for a decade.6 This inconsistency is further emphasized by the fact that few other oxo derivatives of silver (II) are known, including fluorosulfate, triflate,7 and oxoargentate(III), Ag[AgO2]2.8 Indeed, it has been recently shown that inorganic sulfates resist the presence of AgII from AgF2 up to fairly high temperatures of about 300 °C.9 Herein we present two distinct synthetic pathways to AgSO4, which is a genuine sulfate of divalent silver and neither a mixed-valence AgI/AgIII compound nor the peroxodisulfate of AgI.10
Black AgSO4 can be prepared by metathesis in anhydrous HF solvent:((1 a))
or by displacement of a weaker acid by a stronger one:((1 b))
The product of Reaction (1 a) is contaminated by substantial amounts of amorphous KSbF6 (see the Supporting Information); therefore, all results reported herein are for the superior-quality product of obtained from Reaction (1 b).
AgSO4 crystallizes in the triclinic P space group (Table 1 and Figure 1; for Rietveld refinement see the Supporting Information). The unit cell of silver(II) sulfate is the second smallest amongst the transition metal sulfates after CdSO4. The closest Ag⋅⋅⋅Ag separation is as large as 4.6835(2) Å, thus preventing formation of the short metal–oxygen–metal bridges seen in homologous CuSO4. The structure of AgSO4 is also remarkably different from that of AuSO4, which contains short AuAu bonds. All oxygen atoms of the sulfate anion in AgSO4 are utilized for bonding, in a similar fashion to CdSO4 or PdSO4, serving as OSO linkers between silver(II) cations, forming a three-dimensional network.10 There are two non-equivalent silver atoms in the structure that are surrounded by oxide anions in a square-planar fashion, with AgO separations falling in a narrow range between 2.094(10) Å and 2.198(5) Å. Similarities in the coordination spheres of both silver centers suggests the AgIISO4 formulation in contrast to the binary oxide “AgO”, which is in fact the mixed-valence species AgIAgIIIO2.11 The increased Lewis acidity of AgSO4 (AgO⋅SO3) appears to prevent disproportionation12 in a similar fashion to some gold(II) derivatives.13
|a, b, c [Å]:||4.6923(1)||4.7535(1)||8.0125(2)|
|α, β, γ [°]||103.403(1)||76.478(1)||118.078(1)|
|Rp=6.12 %||Rwp=6.26 %|
The sulfate anion has four distinct SO bonds that vary between 1.427(12) Å and 1.503(8) Å. This feature, superimposed on mutual interactions between similar frequency oscillators inside the unit cell (Z=2), leads to splitting of the broad sulfate band (950–1200 cm−1, SO stretching vibrations) and of the bands corresponding to the O-S-O bending vibrations (500–700 cm−1) in the MIR spectrum (Figure 2; for wavenumbers of the bands in the IR and Raman spectra of AgSO4, see the Supporting Information).
A characteristic increase of absorbance above 1420 cm−1 (broad feature leveling off at about 7500 cm−1; see Figure 2, inset) corresponds to a low-energy electronic excitation across the band gap of circa 0.18 eV. The charge-transfer nature of this excitation (O2−→Ag2+; AgII is a potent oxidizer) is discussed below and in the Supporting Information. The low-energy electronic absorption explains deep black color of the compound,14 but it also heralds its thermodynamic instability. Indeed, as deduced from thermogravimetric and differential scanning calorimetry (DSC) profiles (Figure 3 and Supporting Information), and from simultaneous evolved-gas analysis (consisting of IR and MS; Supporting Information), AgSO4 decomposes exothermally above 120 °C (Q=−26.5 kJ mol−1) with evolution of molecular oxygen. This observation suggests that AgSO4 is thermodynamically unstable at ambient conditions owing to both enthalpy and entropy contributions to the Gibbs free energy.10, 15
The mass loss during thermal decomposition of AgSO4 is 3.8 wt %, which corresponds to the theoretical value of 3.9 wt % associated with the chemical reaction:((2))
The crystalline decomposition product was confirmed to be the previously ill-characterized compound silver(I) pyrosulfate, Ag2S2O7, by XRD and IR spectroscopy (Supporting Information). Upon further heating, this compound decomposes endothermally in two steps in the 250–370 °C temperature range with loss of 18.3 wt % (theoretical: 19.6 wt %16) in the form of SO3:((3))
The paramagnetic nature of divalent silver (electronic configuration 4d9) in AgIISO4 raises questions about the strength of spin interactions. The temperature dependence of the magnetic susceptibility (Figure 4) shows that the magnetic behavior of AgSO4 above 40 K can be fit well with a Bonner–Fisher model17 of strongly antiferromagnetically coupled one-dimensional chains with isotropic interactions between nearest neighbors (H=−2 JΣSiSj). The obtained exchange integral is J=−217 K per pair of interacting Ag2+ cations (i.e., 9.5 meV per formula unit), whilst a value g=2.087 was obtained from the ESR data.18 The antiferromagnetic coupling, which is similar to that observed for various derivatives of Cu2+ or V4+,19 persists even at temperatures around 100 °C at the onset of thermal decomposition. This observation is consistent with ESR spectra of AgSO4 (Figure 4, inset), which show presence of a very broad (1.0–2.6 kG) antiferromagnetic signal centered around 3250 G (g=2.087) in the entire 2.4–292 K temperature range.20 (For the temperature dependence of parameters of the ESR signal, see the Supporting Information).
The exchange integral |J| between the nearest neighbors for AgSO4 is over one order of magnitude higher than those previously observed for other transition metal sulfates. For example, sulfates of nickel(II), iron(II), and cobalt(II) order antiferromagnetically at 37 K, 21 K, and 15.5 K, respectively.21 As each of these compounds contains between two or four unpaired electrons per transition metal cation (compared to one for AgSO4), superexchange for AgSO4 must thus be viewed as unusually strong. Even more striking is the contrast between AgSO4 and CuSO4: Anhydrous CuSO4 orders antiferromagnetically below 36 K5 despite the fact that magnetic superexchange in this compound takes place over direct oxo bridges, and not much longer covalent OSO linkers as in AgSO4.
To gain insight into the nature of magnetic interactions and electronic band structure of AgSO4, we performed spin-polarized DFT calculations using the GGA and GGA+U methods. Pathways for magnetic interactions are very complex; we therefore describe only the simplest model (Figure 5) with a magnetic unit cell identical to crystal unit cell. The artificially enforced antiferromagnetic state converged to a metallic solution within the GGA framework, as is typical for transition metal open-shell systems. However, the GGA+U calculations confirm an antiferromagnetic ground state of AgSO4, with ferromagnetic and metallic solutions lying higher in energy by 3.1 meV and 20.3 meV per formula unit, respectively. The derived magnetic superexchange constant is 6.2 meV per one silver atom in the one-dimensional coupling model, and it is thus 1.5-fold underestimated relative to the experimental value. This discrepancy could certainly be decreased by allowing variation of semiempirical Mott–Hubbard U parameters for Ag(d), O(p), and S(p) electrons. The superexchange pathway involves the OSO bridge, but is largely over a direct O⋅⋅⋅O polarization, thus omitting the sulfur atom (Figure 5).
The calculated magnetic moments on the silver atoms (−0.39 μB and +0.44 μB) are much smaller than those calculated for the copper atoms in CuSO4 (±0.80 μB), because spin density is considerably smeared over oxygen atoms for the former compound, but much less for the latter (see the Supporting Information).22 Indeed, strongly oxidizing silver(II) easily introduces holes to oxygen-based bands, resulting in the lack of thermodynamic stability of AgSO4 and evolution of O2 upon its heating.
The band structure of AgSO4 shows that the uppermost two spin-majority and two spin-minority states are considerably separate from each other in the vicinity of the Fermi level. This result can be traced to the presence of two non-equivalent silver cations in the magnetic unit cell (Supporting Information). In consequence, the band gap between occupied (O-predominated) and unoccupied (Ag-predominated) α bands is much smaller than the corresponding gap for β states. The calculated band gap at the Fermi level of 0.18 eV agrees strikingly well with the experimental estimate, rendering AgSO4 a narrow-band-gap magnetic semiconductor; it also explains the spin-allowed optical charge-transfer α→α absorption responsible for its black color. The band gap of AgSO4 is one order of magnitude smaller than that of 2.34 eV calculated for CuSO4 with the LSDA+U method (Supporting Information). Attempts are now ongoing in our laboratory to form a pressure-induced or chemical-doping-induced metallization of the AgSO4 antiferromagnet.