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Xanthate sulfur as a hydrogen bond acceptor: the free xanthate anion and ligand sulfur in nickel tris ethylxanthate
Inorganica Chimica Acta 358 (2005) 633–640 www.elsevier.com/locate/ica
Xanthate sulfur as a hydrogen bond acceptor: the free xanthate anion and ligand sulfur in nickel tris ethylxanthate Marc A. Walters a,*, Justin Barad a, Anthony Sireci a, James A. Golen b, Arnold L. Rheingold b a b
Department of Chemistry, New York University, 100 Wash. Sq. E., New York, NY 10003, USA Department of Chemistry, University of California at San Diego, La Jolla, CA 92093-0358, USA Received 16 July 2004; accepted 22 September 2004 Available online 10 December 2004
Abstract Xanthates, like thiolates, form a variety of complexes with metals in which coordinating sulfur can serve as a hydrogen bond acceptor. Nickel tris xanthate complexes [Ni(xan)3], (xan = o-ethylxanthate, N-(carbamoylmethyl)ethylxanthate) have been synthesized and compared by a combination of X-ray crystallographic and spectroscopic measurements. Recent results from our studies of N–H S hydrogen bonding interactions in metal–xanthate complexes shows N–S distances to be longer than those in related thiolate complexes, indicative of weaker hydrogen bonds for the xanthates. The complex (Et4N)[N-(carbamoylmethyl)ethylxanthate)] adopts an extended conformation in both the solid state and solution and lacks either intraligand or intermolecular N–H S hydrogen bonds. The complex (CTA)[Ni(exa)3] exhibits N–H S hydrogen bonds between the amide group of the counterion and the ˚. ligand sulfur. The amide–sulfur N–H S distance is 3.567 A 2004 Elsevier B.V. All rights reserved. Keywords: Carbamoyl methyl xanthate; Nickel tris ethylxanthate; Hydrogen bond; Crystal structure
1. Introduction Xanthates have long proven useful in technology as chelators in aqueous media. Chief among their applications is the flotation of ores, and chelation therapy [1,2]. Since the early days of modeling in bioinorganic chemistry 1,1-dithiochelates, which include xanthates, have been employed as cysteine analogs [3–8]. Recently, they have been explored as alternatives to thiols in the formation of self-assembled monolayers on gold [9–11]. Although xanthates are generally employed in protic media, little is known about their participation in hydrogen bonding (H-bonding) or other dipolar interactions *
Corresponding author. Tel.: +1 212 998 8477; fax: +1 212 260 7905. E-mail address: [email protected] (M.A. Walters). 0020-1693/$ - see front matter 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.ica.2004.09.036
that might affect their character as chelators. In a continuation of our investigations of amide sulfur hydrogen bonding, we have extended our inquiry to include xanthate (xan) ligands in its free and metal complexed form. Nickel-xanthate complexes were employed in this study because of their favorable stability. In this paper, we report on the properties of xanthates relative to thiols in their capacity to form N–H S hydrogen bonds in metal complexes. The exploration of xanthates was favored by their facile synthesis, the ready availability of a large variety of alcohol precursors, and the relative stability of xanthates in air. Xanthate is unlike cysteine in its electronic structure. Therefore we will refer to data on thiolates and metal–thiolate complexes as benchmarks for the results presented here. Our aim was to prepare xanthates that are designed for intraligand H-bonding. Xanthates form amide sulfur
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N–H S hydrogen bonds in the solid state. However, the strength of the H-bond interaction is not sufficient to support the formation of intraligand N–H S hydrogen bonds. Below we discuss the formation of metal–ligand bonds and H-bonds in a set of Ni xanthate complexes and compare the xanthates with thiols in their capacity to form covalent bonds to metal ions and hydrogen bonds with amide donor groups.
2. Experimental 2.1. Synthesis of N-hydroxyethylacetamide, CH3CONHCH2CH2OH The starting alcohol was prepared by a previously established method [12]. Ethyl acetate 30 ml (0.307 mol) and ethanol amine 24.6 ml (0.410 mol, 25% excess) were combined in a 150 ml round bottom flask and refluxed for 4 h with stirring to yield a light yellow solution. The product was isolated by vacuum distillation at 108–110 C as a colorless, viscous oil; Yield: 59.9%. 2.2. Synthesis of potassium N-(carbamoylmethyl)ethylxanthate, CH 3 CONHCH 2 CH 2 OCS 2 K þ , Kcxa, (1) The xanthate, Kcxa, was synthesized by a method first described by DAmico and Bollinger [13]. The starting alcohol 1.89 g (0.018 mol) was combined with excess CS2 (15 ml). KOH 1.03 g (0.018 mol) was then added with stirring to a concentrated aqueous solution <1 ml over a period of 10 min. The reaction was allowed to stir for 24 h at room temperature. The resulting orange oil was then dried under vacuum and washed with acetone and methanol to yield a bright yellow powder. 1H NMR (DMSO): 1.8 ppm (s, 3H, acetyl), 3.3 ppm (m, 2H, methylene), 4.2 ppm (t, 2H, methylene), 8.0 ppm (s, 1H, amide); Yield: 47.5%. 2.3. Synthesis and crystallization of tetraethylammonium N-(carbamoylmethyl)ethylxanthate, Et4N, [cxa] (2) Equimolar amounts of Kcxa and Et4NBr were dissolved in acetonitrile. The resulting yellow solution was filtered to remove KBr and dried under vacuum giving a light yellow powder, Et4Ncxa. The slow diffusion of ether into a concentrated solution of the xanthate in acetonitrile yielded large, orange/yellow crystals. 2.4. Synthesis of nickel bis cxa, [Ni(cxa)2] The bis complex was prepared in a manner analogous to that reported for Ni(exa)2 [14–16]. NiCl2 Æ 6H2O 1.18 g (0.005 mol) in 10 ml water was added with stirring to an aqueous solution of Kcxa (2.16 g (0.01 mol) in 10 ml water). The resulting brown precipitate, Ni(cxa)2, was
filtered and dried under vacuum. The complex was recrystallized from DMF/ether; Yield 95%. 2.5. Synthesis of potassium nickel tris cxa, K[Ni(cxa)3] For spectroscopic and electrochemical studies the complex K[Ni(cxa)3] was formed in situ by combining [Ni(cxa)2] with Kcxa in CH3CN. The isolated complex is an oil and inconvenient to handle. The potassium nickel tris cxa complex, K[Ni(cxa)3], was prepared in a manner similar to that reported for K[Ni(exa)3] [14,15]. Ni(cxa)2 1.00 g (2.4 · 103 mol) was partially dissolved in acetonitrile. The xanthate salt Kcxa 0.524g (2.4 · 103 mol) was added with stirring to the suspension whereupon all of the [Ni(cxa)2] dissolved and a green solution resulted. The solution was filtered and reduced by rotary evaporation. The product was isolated as a green oil after extracting acetonitrile with ether and removing the ether under vacuum. 2.6. Synthesis of carbamoylmethyltrimethylammonium bromide, CTABr 1.0 g (7.24 mmol) bromoacetamide was dissolved in ethanol. An excess of trimethylamine was bubbled into the solution for about 5–7 min. The solution was left overnight in a tightly closed flask. After 24 h, a white crystalline precipitate is present in the solution. The product is filtered under nitrogen atmosphere and washed with ether; Yield 93%. 2.7. Synthesis of carbamoylmethyltrimethylammonium ethylxanthate, CTAexa (3) Potassium ethyl xanthate, Kexa, 1.01 g (6.29 mmol) and CTABr 1.24 g (6.29 mmol) [17] were dissolved in methanol. The resulting yellow solution was concentrated under vaccum. The resulting viscous material was solidified by the addition of ether and was subsequently triturated and isolated as a white powder after decanting off the ether. The powder consisted of KBr and complex 3. CH3CN, 20 ml, was added to extract the product. The solution was filtered to remove KBr and then dried under rotary evaporation to yield a yellow viscous oil that was solidified to a yellow powder by the addition of ether. The product was dried in air; Yield: 97% 1H NMR (DMSO): ethyl xanthate: 1.2 ppm (t, 3H, methyl), 4.25 (m, 2H, methylene). Carbamoylmethyltrimethylammonium chloride (CTA-Cl): 3.25 ppm (s, 9H, trimethyl), 4.1 ppm (s, 2H, a carbon), 7.8 ppm (d, 2H, amide). 2.8. Synthesis of CTA[Ni(exa)3] (4) (a) CTAexa 30 mg, (1.26 · 104 mol) was dissolved in methanol, 6.5 ml. [Ni(exa)2] 24 mg, (1.26 · 104
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mol) was dissolved in isopropanol 68 ml. The solutions were combined with stirring in a round bottomed flask and the volume was reduced to about 50 ml via rotary evaporation. The solution was chilled in refrigerator and yielded small green crystals after 24 h. The crystals were collected by vacuum filtration, washed with pentane followed by ether, and dried under a flow of air. The product was recrystallized from acetone by the addition of ether by vapor diffusion; Yield: 49%. (b) In a second method CTA(exa), 30 mg (1.26 · 104 mol), was dissolved in less than 1 ml of acetone. Complete dissolution was achieved by the brief application of gentle heat. Then, an equimolar amount of the complex [Ni(exa)2] was added as a solid. The resulting emerald green acetone solution was filtered through a plug of Kimwipe in a Pasteur pipette. The solution was transferred to a test tube which was placed in a sealed jar containing 10 ml of ether to crystallize the sample by the diffusion of ether in a refrigerator at 0 C. After about 24 h crystals were collected on a Buchner funnel, washed with hexanes, and dried in air. The product was stored in a capped vial in a freezer to retard decomposition; Yield: 60%.
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Elemental analysis was carried out by QTI. Anal. Calc. for C14H28N2O4S6Ni: C, 31.17; H, 5.23; N, 5.19. Found: C, 31.34; H, 5.15; N, 5.18%. 2.9. Cyclic voltammetry The working and counter electrodes consisted of platinum wire. Silver wire served as a quasi-reference electrode. The potential of the ferrocene/ferricinium couple, measured at room temperature was 0.93 V. Experiments were run in a 0.1 M solution of tetrabutylammonium perchlorate (TBAP) in acetonitrile. 2.10. X-ray crystallographic structure determination Crystallographic data for 2 and 4 are shown in Table 1. Data were collected with Bruker APEX equipped diffractometers and processed using the current versions of BRUKER software (Bruker AXS, Madison, WI). A noncentrosymmetric setting was chosen for 2 due to the absence of mirror-plane symmetry; refinement of the Flack parameter indicated that the reported hand is correct. Both structures were solved by direct methods and refined with anisotropic thermal parameters for all nonhydrogen atoms. Hydrogen atoms were
Table 1 Crystallographic data for Et4N(cxa) (2) and (CTA)[Ni(exa)3] (4) 2
4
Empirical formula C13H28N2O2S2 C14H28N2NiO4S6 Formula weight 308.49 539.45 Temperature (K) 218(2) 218 (2) ˚) Wavelength (A 0.71073 0.71073 Crystal system monoclinic monoclinic Space group P21 P21/n ˚) a (A 8.6621(9) 8.0348(5) ˚) b (A 9.6615(10) 16.3259(10) ˚) c (A 10.7264(11) 18.4137(11) b () 107.609(2) 93.5560(10) ˚ 3) Volume (A 855.62(15) 2410.8(3) Z 2 4 Dcalc (Mg/m3) 1.197 1.486 Absorption coefficient (mm1) 0.312 1.346 F(0 0 0) 336 1128 Crystal size (mm3) 0.30 · 0.30 · 0.30 0.40 · 0.30 · 0.20 h Range for data collection () 2.47–26.50 1.67–28.29 Index ranges 8 6 h 6 10; 12 6 k 6 11; 13 6 l 6 12 9 6 h 6 10; 21 6 k 6 21; 24 6 l 6 24 Reflections collected 5069 17 526 Independent reflections (Rint) 3025 (0.0163) 5713 (0.0231) Completeness to h 26.50, 99.4% 28.29, 95.4% Absorption correction SADABS SADABS Refinement method full-matrix least-squares on F2 full-matrix least-squares on F2 Data/restraints/parameters 3025/1/244 5713/0/244 Goodness-of-fit on F2 1.130 1.062 Final R indices [I > 2r(I)]a,b R1 = 0.0511; wR2 = 0.1383 R1 = 0.0294; wR2 = 0.0802 R1 = 0.0345; wR2 = 0.0826 R indices (all data) R1 = 0.0544; wR2 = 0.1425 ˚ 3) Largest difference peak and hole (e A 0.458 and 0.208 0.557 and 0.380 P P a R ¼ jjF o jP jF c jj= jF 0 j: P b 2 2 RðwF Þ ¼ f ½wðF o F 2c Þ2 = ½wðF 2o Þ2 g1=2 ; w ¼ 1=½s2 ðF 2o Þ þ ðaP Þ2 þ bP ; P ¼ ½2F 2c þ maxðF o ; 0Þ=3:
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treated as idealized contributions. In 2, the alkyl groups in the counterions were disordered equally over two positions.
3. Results and discussion 3.1. Structure of Et4N(cxa) (2) The ligand cxa was designed to form a seven-membered ring that is closed by the formation of an intraligand N–H S hydrogen bond. When ligand hydrogen bonding persists in solution, it allows the study of H-bonding effects on redox potentials by solution dependent methods such as cyclic voltammetry (CV). A similar strategy was employed successfully in the study of N–H S hydrogen bonding in a series of molybdenum thiolate complexes [18,19]. Earlier studies showed that intraligand N–H S hydrogen bonds can occur in both an isolated ligand and in its metal–ligand complex [20–22]. The ligand cxa however shows no propensity to form such H-bonds. In its anionic form the ligand is extended as observed in its X-ray crystal structure (Fig. 1). The C–Savg bond distance of 2 is ˚ (Tables 1a and 1b) and is essentially the same 1.680(3) A ˚ [23]. as that reported for the salt (Et4N)exa, 1.678(3) A As observed in the packing diagram of 2, close contacts are limited to C(3)–H S hydrogen bonds (Fig. 1).
Table 1a ˚ ) Et4N cxa Bond lengths (A S(1)–C(1) S(2)–C(1) O(1)–C(1) O(1)–C(2) O(1)–C(4) N(1)–C(4) N(1)–C(3) N(2)–C(12 0 ) N(2)–C(6) N(2)–C(10 0 ) N(2)–C(10) N(2)–C(8 0 )
1.686(3) 1.673(3) 1.355(3) 1.434(4) 1.230(4) 1.340(4) 1.440(4) 1.479(7) 1.488(7) 1.499(7) 1.508(5) 1.515(7)
N(2)–C(8) N(2)–C(6 0 ) N(2)–C(12) C(2)–C(3) C(4)–C(5) C(6)–C(7) C(6 0 )–C(7 0 ) C(8)–C(9) C(8 0 )–C(9 0 ) C(10)–C(11) C(10 0 )–C(11 0 ) C(12 0 )–C(13 0 ) N(1)–H(1) O(2) N(1)–H(1) O(4) N(1)–H(1) S(4)
1.537(6) 1.575(6) 1.594(8) 1.518(4) 1.497(4) 1.480(16) 1.519(10) 1.590(11) 1.460(15) 1.545(9) 1.515(11) 1.271(10) 3.201 3.006 3.567
Table 1b Bond angles () Et4N cxa
3.2. Structure of (CTA)[Ni(exa)3] (4)
C(1)–O(1)–C(2) C(4)–N(1)–C(3) C(12 0 )–N(2)–C(10 0 ) C(10 0 )–N(2)–C(8 0 ) C(6)–N(2)–C(8) C(10)–N(2)–C(8) C(12 0 )–N(2)–C(6 0 ) C(8 0 )–N(2)–C(6 0 ) C(8)–N(2)–C(6 0 ) C(6)–N(2)–C(12) C(10)–N(2)–C(12) C(8)–N(2)–C(12) O(1)–C(1)–S(2) O(1)–C(1)–S(1)
118.8(2) 122.9(3) 98.7(4) 109.8(6) 91.1(4) 109.8(6) 106.7(4) 104.9(4) 109.1(4) 106.8(5) 97.4(4) 85.4(4) 120.9(2) 113.8(2)
S(2)–C(1)–S(1) O(1)–C(2)–C(3) N(1)–C(3)–C(2) O(2)–C(4)–N(1) O(2)–C(4)–C(5) N(1)–C(4)–C(5) C(7)–C(6)–N(2) C(7 0 )–C(6 0 )–N(2) N(2)–C(8)–C(9) C(9 0 )–C(8 0 )–N(2) N(2)–C(10)–C(11) C(13)–C(12)–C(12 0 ) C(13 0 )–C(12)–N(2) N(2)–C(12 0 )–C(13)
In studies of metal–thiolate complexes we have observed N–H S bond formation between an anionic complex and its cationic counterion with donor NH groups [17,24]. The cation CTA has proved to be an ideal H-bond donating counterion for the study of amide–ligand H-bonds. In the complex (CTA)
[Ni(exa)3] (4) the CTA counterion donates hydrogen bonds to both oxygen and sulfur atoms of exa (Fig. 2). In the complex (Et4N)[Ni(exa)3] which lacks H˚ (Tabonds, the average Ni–S bond length is 2.439 A bles 1c and 1d) [12].
Fig. 1. Molecular structure packing diagram of Et4N(cxa) (2). Dotted lines denote C–H S hydrogen bonds.
125.31(17) 108.2(3) 111.9(3) 122.3(3) 122.2(3) 115.5(3) 113.1(9) 112.9(6) 112.5(6) 116.3(8) 113.5(6) 102.2(9) 114.8(6) 110.5(5)
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637
Fig. 2. Molecular structure of (CTA)[Ni(exa)3] (4). Dashed line – – – denotes N–H S hydrogen bonds.
By comparison, the average Ni–S bond length in 4 is ˚ to a value of 2.427 A ˚. decreased by 0.012 A In complex 4, three types of H-bonds are observed with N–H donation to sulfur, S(4) ether oxygen, O(2) and carbonyl oxygen, O(4), respectively. The N–H S hydrogen bond of 4 is weak. Nitrogen and sulfur are ˚ , which is at the limit of what is separated by 3.567 A generally considered to be an N–H S hydrogen bond in simple organic molecules [25]. In the N–H O(2) ˚, bond, nitrogen and oxygen are separated by 3.201 A which is a length characteristic of hydrogen bonds that are formed by ether oxygen. The N–H O(4) bond be˚ in length and is tween adjacent CTA cations is 3.006 A typical of an amide dimer [25]. In alkylthiolate complexes, H-bond donation to sulfur is expected to result in metal–ligand bond lengthening [24]. The decrease in Ni–S bond length in 4 relative to (Et4N)[Ni(exa)3] is unexpected. A significant caveat in the analysis of CTA amide hydrogen bonding is that the NH group is trifurcated as a hydrogen bond donor. We must exercise caution in the evaluation of N–H S hydrogen bonding
Table 1c ˚ ) for CTA[Ni(exa)3] Selected bond lengths (A Ni(1)–S(4) Ni(1)–S(5) Ni(1)–S(1) Ni(1)–S(3) Ni(1)–S(2) Ni(1)–S(6) S(1)–C(1) S(2)–C(1) S(3)–C(4)
2.4081(5) 2.4107(5) 2.4294(5) 2.4324(5) 2.4414(5) 2.4416(5) 1.6780(16) 1.6943(16) 1.6780(17)
S(4)–C(4) S(5)–C(7) S(6)–C(7) O(1)–C(1) O(1)–C(2) O(2)–C(4) O(2)–C(5) O(3)–C(7) O(3)–C(8)
1.6883(18) 1.6825(17) 1.6894(18) 1.327(2) 1.452(2) 1.3363(19) 1.459(2) 1.3372(19) 1.457(2)
in this system since the interaction will be affected by N–H O hydrogen bonds to a nearby xanthate and the carbonyl oxygen of a neighboring amide group. A similar investigation of related dithiocarbamates likewise showed no perturbation of Ni–S covalent bonds as a function of N–H S hydrogen bond formation [26]. Even when the H-bond acceptor is an alkane thiolate, a ˚ is observed in the average M–S small increase of 0.01 A bond length relative to M–S bonds in non-hydrogen bonded complexes [24]. 3.3. Structure of K[Ni(cxa)3] The coordination geometry of K[Ni(cxa)3], which was reproducibly isolable only as a green oil, could be
Table 1d ˚ ) for CTA[Ni(exa)3] Selected bond angles (A S(4)–Ni–(1)–S(5) S(4)–Ni(1)–S(1) S(5)–Ni(1)–S(1) S(4)–Ni(1)–S(3) S(5)–Ni(1)–S(3) S(1)–Ni(1)–S(3) S(4)–Ni(1)–S(2) S(5)–Ni(1)–S(2) S(1)–Ni(1)S(2) S(3)–Ni(1)–S(2) S(4)–Ni(1)–S(6) S(5)–Ni(1)–S(6) S(1)–Ni(1)–S(6) S(3)–Ni(1)–S(6) S(2)–Ni(1)–S(6) C(1)–S(1)–Ni(1)
95.025(18) 97.795(17) 166.038(17) 73.892(16) 91.880(18) 96.959(18) 166.514(17) 94.844(17) 73.468(15) 96.619(16) 97.887(17) 73.896(16) 98.787(17) 163.118(19) 93.666(16) 83.82(6)
C(1)–S(2)–Ni(1) C(4)–S(3)–Ni(1) C(4)–S(4)–Ni(1) C(7)–S(5)–Ni(1) C(7)–S(6)–Ni(1) C(1)–O(1)–C(2) C(4)–O(2)–C(5) C(7)–O(3)–C(8) O(1)–C(1)–S(1) O(1)–C(1)–S(2) S(1)–C(1)–S(2) O(1)–C(2)–C(3) O(2)–C(4)–S(3) O(2)–C(4)–S(4) S(3)–C(4)–S(4) O(2)–C(5)–C(6)
83.12(6) 82.78(6) 83.34(6) 83.67(6) 82.57(6) 118.96(13) 118.13(15) 117.56(13) 124.66(12) 115.84(11) 119.50(10) 107.50(15) 123.57(13) 116.83(12) 119.60(9) 107.15(19)
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inferred from its electronic absorption spectrum and the structure of the ligand salt, Et4Ncxa. The crystal structure of (Et4N)[Ni(exa)3] shows a complex that is significantly distorted from octahedral geometry, with approximate 3-fold symmetry about the central Ni2+ ion [12]. The electronic absorption spectrum of [Ni(cxa)3] in an acetonitrile solution is congruent with that of the [Ni(exa)3] complex, and is distinct from the planar complex [Ni(cxa)2] (Table 2). The bands of the nickel xanthate complexes have been assigned to ML and LM charge transfer transitions and one internal Sp ! p* transition respectively [27]. Paradoxically, the internal ligands p ! p* transition appears to be sensitive to the geometry or coordination number of the Ni(II) complexes (Table 2). From the crystal structure of the Et4Ncxa salt it is reasonable to assume that for the [Ni(cxa)3] in the solution phase the ligand is in an elongated conformation perhaps with amide to carbonyl intermolecular H-bonding. 3.4. Infrared spectra A secondary amide such as N-ethylacetamide has infrared bands at 3472 and 1720 cm1 in the gas phase that have been assigned to m(NH) and m(CO) stretching modes, respectively [29]. The mid-infrared solid state spectrum of the Kcxa salt shows corresponding bands at 3269 and 1649 cm1 characteristic of a secondary amide hydrogen bond donor (Table 3). The corresponding bands [Ni(cxa)2] occur at similarly decreased Table 2 Electronic absorption data (nm, in CH3CN solvent) on selected Ni xanthate complexes [Ni(cxa)2]
[Ni(exa)2]a
K[Ni(cxa)3]
K[Ni(exa)3]a
Assignmenta,b
196 216 252 316
218 252 316
198 218 270 316
194 218 270 312
d ! p* Sp ! p* Sr ! dr*
a
Ref. [27]. Ref. [28].
b
frequencies, 3280 and 1645 cm1, respectively, which suggests that the amides of [Ni(cxa)2] are hydrogen bonded in the solid state. In [Ni(cxa)3] the amide m(NH) band occurs at about 3404 cm1 and is likely shifted to lower frequency by interactions with the solvent acetonitrile [29]. Finally, in the complex (CTA) [Ni(exa)3] the amide group of the CTA+ counterion has four m(NH) bands in the range 3455–3166 cm1 whose provenance has not been determined. The bands assigned to the xanthate moiety are unremarkable, with the m(COC)asymm band in the region 1160–1280 cm1, and the m(CS) band in the region 1026–1070 cm1. The m(COC)asymm band is sensitive to coordination geometry and occurs in the 1270 cm1 region for the neutral planar complex [Ni(xan)2] and in the 1180 cm1 region for the anionic distorted octahedral complex [Ni(xan)3]. As discussed earlier by Mattes and Pauleickhoff, the frequency of the m(COC)asymm band is related to the resonance forms of the xanthate anion as illustrated below [30].
S-1
S-1/2 R
O
R
C
O
S-1
S-1/2
I
C
II
In resonance form II the larger bond order of oxygen results in a larger O–C force constant that is reflected in a higher frequency for m(COC)asymm. With its higher electron density, the sterically demanding sulfur of form II is expected to be more characteristic of the planar [Ni(xan)2] complex, which has an average Ni–S bond length ˚ [31]. Conversely, steric and electronic factors of 2.21 A should favor resonance form I in a Ni(II) tris xanthate complex, [Ni(xan)3], complex that has an average Ni– ˚ [12]. S bond length of 2.439(2) A
Table 3 Infrared data on selected xanthate complexes and salts Kcxaa
[Ni(cxa)2]a
3269 1649 1545 1160 1047
3280 1645 1559 1263 1050
a b c d e
KBr pellet. CH3CN solvent. Ref. [32]. Ref. [33] Ref. [34].
[Ni(exa)2]a
K[Ni(cxa)3]b
1276 1026
3405 1676 1541 1177 1069
K[Ni(exa)3]a
(CTA)[Ni(exa)3]a
Assignmentsc,d,e
1184 1041
3455, 3317, 3265, 3166 1704 1595 1181 1038
m(NH)amide m(CO)amide d (NH2)amide C–O–C CS2 out of phase
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3.5. Cyclic voltammetry
Table 5 Atomic charges on sulfur
As evidenced by X-ray crystallographic and spectroscopic data, the xanthate moiety of cxa and exa are congruent in both their structural and electronic characteristics. However they diverge in their voltammetric properties in that the redox potential of cxa is significantly anodic relative to that of exa (Table 4). The free ligand salt Kcxa undergoes an irreversible oxidation at 1.53 V vs. a Ag/Ag+ quasi-electrode as compared with an oxidation peak at 1.00 V for Kexa. Anionic xanthate is known to oxidize to a neutral radical that quickly and irreversibly dimerizes, producing the stable dixanthogen [12]. This mechanism explains the lack of a return reduction peak. The difference in oxidation potentials for the free anionic ligands is not likely due to inductive effects of the amide group in cxa but rather to differences to the free energy of dixanthogen formation from the neutral radicals. In acetonitrile the Ni2+/Ni3+ couple for the green complex [Ni(cxa)3] (formed in situ) has a value of 0.62 V with Ired/Iox = 0.23 against a silver wire quasi reference electrode at 25 C. According to Choudhury et al. [12] in work on the complex [Ni(exa)3], the anionic species is oxidized to the thermally unstable neutral Ni(III) complex, [Ni(xanthate)3], the decomposition of which accounts for Ired/Iox < 1. As with CV experiments on [Ni(exa)3], the Ired/Iox of [Ni(cxa)3] approaches 1 with decreasing temperature, which suggests a shift of equilibrium towards the tris complex. Although the redox couple at 0.62 V can be attributed to either a ligand centered redox process or a Ni(II)–Ni(III) couple, we have assigned it to the latter. Cyclic voltammetric experiments on the complexes [Zn(exa)2] and [Co(exa)3] show ligand-centered redox potentials that are more anodic than that of the free xanthate anion [14]. We therefore expect ligand centered oxidations of chelated cxa and exa to occur well above 1 V. The complex K[Ni(exa)3] has a redox couple at 0.64 V (Ired/Iox = 0.59), that is nearly identical to the redox potential of K[Ni(cxa)3] (0.62 V). Such a small difference in potential is consistent with solution IR studies
Group
Table 4 Cyclic voltammetric data on selected xanthates and Ni xanthate complexes Compound a
Kcxa Kexa Kexab K[Ni(cxa)3] K[Ni(exa)3] K[Ni(exa)3]b a b
Redox Potential (mV)/Ag at 25 C 1.53 1.00 0.23 0.62 0.64 0.11
(oxidative peak only) (oxidative peak only) (SCE) Ired/Iox = 0.23 Ired/Iox = 0.59 (SCE)
Ag wire quasi reference electrode unless otherwise indicated. Ref. [12].
Ph–S Ph–CH2–S H3C–S EtO–CS 2 ðexaÞ H3 CCðOÞNHðCH2 Þ2 OCS 2 ðcxaÞ a b
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Charge 0.844a 0.978a 1.063a 0.362b 0.465, 0.288b
Ref. [24]. This work.
of the K[Ni(cxa)3] that showed no evidence of N–H S hydrogen bonding in the [Ni(cxa)3] complex. 3.6. Computational assessment of xanthate sulfur and hydrogen bonding Xanthate ions can form hydrogen bonds to sulfur both in the metal ligated form and in the free ligand state. Sulfur is relatively is weak as an H-bond acceptor relative to the more electronegative elements nitrogen and oxygen. Thus, it is not unexpected that N–H S hydrogen bonds would be weak in xanthate. However, this type of hydrogen bond is significantly weaker than in CTA metal–thiolate complexes such as (CTA) [Co(SPh)4] that exhibits an amide–thiolate N–S average ˚ as compared with 3.6 A ˚ in complex 4 distance of 3.3 A [17]. The qualitative manifestation of hydrogen bond strength is the donor–acceptor distance. An explanation of the disparities in N–H S distances must focus on the nature of sulfur in xanthates as compared to thiolates. A qualitative assessment suggests that the charge on sulfur in xanthate is less than in thiolate. Computational evaluation supports this view (see Table 5). For computational analysis the geometries of exa and cxa were optimized using AM1. Single point energy level calculations were then carried out at the 6-31G* level. The charge on individual sulfur atoms in xanthate is approximately 0.36 e as a result of electron delocalization between the two sulfur atoms (resonance structures) and the inductive effect of oxygen. This renders xanthate a very weak H-bond acceptor in contrast to thiolate where the charge on sulfur is approximately 0.99 e. Metal complexation is expected to reduce the charge on sulfur in both cases, but with xanthate ligands the effect seriously weakens the N–H S bond strength.
4. Conclusions Xanthate as a free ion or as a ligand in a Ni complex forms relatively weak N–H S hydrogen bonds with amide donors. Ni–Sxan bonds are also relatively weak. These two phenomena are due to the relatively low charge density on xanthate sulfur as compared with thiolate sulfur. The clearest manifestation of these
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phenomena are the large N–H S and Ni–S bond distances in complex 4.
Acknowledgement Carbamoylmethyltrimethylammonium chloride was synthesized by Pamela Tadross (NYU). References [1] G.J. Jameson, Curr. Opin. Colloid Interf. Sci. 3 (1998) 351. [2] T. Stoichev, Savremenna Med. 16 (1965) 356. [3] S.J. Lippard, D. Coucouvanis, J.A. Zubieta, Inorg. Chem. 9 (1970) 2775. [4] J. Peisach, W.E. Blumberg, E.T. Lode, M.J. Coon, J. Biol. Chem. 246 (1971) 5877. [5] B. Song, J. Kim, C. Han, M. Kim, Bull. Korean Chem. Soc. 22 (2001) 413. [6] M.C. Smith, J.E. Barclay, S.C. Davies, D.L. Hughes, D.J. Evans, Dalton Trans. (2003) 4147. [7] W.E. Newton, J.L. Corbin, D.C. Bravard, J.E. Searles, J.W. McDonald, Inorg. Chem. 13 (1974) 1100. [8] W.-F. Liaw, J.-H. Lee, H.-B. Gau, C.-H. Chen, S.-J. Jung, C.-H. Hung, W.-Y. Chen, C.-H. Hu, G.-H. Lee, J. Am. Chem. Soc. 124 (2002) 1680. [9] P. Talonen, G. Sundholm, W.-H. Li, S. Floate, R.J. Nichols, Phys. Chem. Chem. Phys. 1 (1999) 3661. [10] R. Woods, G.A. Hope, G.M. Brown, Colloid Surf. A: Physicochem. Eng. Aspect 137 (1998) 339. [11] K.V. Gothelf, J. Electroanal. Chem. 494 (2000) 147. [12] S.B. Choudhury, D. Ray, A. Chakravorty, Inorg. Chem. 29 (1990) 4603. [13] J.J. DAmico, F.G. Bollinger, J. Heterocyc. Chem. 25 (1988) 1601. [14] D. Coucouvanis, J.P. Fackler, Inorg. Chem. 6 (1967) 2047.
[15] T.J.M. Mohammed, S.E. Al-Mukhtar, J. Ind. Chem. Soc. 62 (1985) 725. [16] G.W.M. Watt, J. Inorg. Nucl. Chem. 27 (1965) 898. [17] M.A. Walters, J.C. Dewan, C. Min, S. Pinto, Inorg. Chem. 30 (1991) 2656. [18] J. Huang, R.L. Ostrander, A.L. Rheingold, Y. Leung, M.A. Walters, J. Am. Chem. Soc. 116 (1994) 6769. [19] J. Huang, R.L. Ostrander, A.L. Rheingold, M.A. Walters, Inorg. Chem. 34 (1995) 1090. [20] G. Zuppiroli, C. Perchard, M.H. Baron, C. De Loze, J. Mol. Struct. 72 (1981) 131. [21] C. Perchard, G. Zuppiroli, P. Gouzerh, Y. Jeannin, F. Robert, J. Mol. Struct. 72 (1981) 119. [22] G. Zuppiroli, C. Perchard, M.H. Baron, C. De Loze, J. Mol. Struct. 69 (1980) 1. [23] B.F. Hoskins, E.R.T. Tiekink, G. Winter, Inorg. Chim. Acta 101 (1985) L11. [24] W.P. Chung, J.C. Dewan, M. Tuckerman, M.A. Walters, Inorg. Chim. Acta 291 (1999) 388. [25] W.C. Hamilton, J.A. Ibers, Hydrogen Bonding in Solids, Benjamin, New York, 1968. [26] S. Santos Jr., S. Guilardi, J.A.L.C. Resende, M.M.M. Rubinger, M.R.L. Oliveira, J. Ellena, Acta Cryst. E: Struct. Rep. Online E59 (2003) m77. [27] H. Isci, O. Dag, W.R. Mason, Inorg. Chem. 32 (1993) 3909. [28] C.K. Jorgensen, J. Inorg. Nucl. Chem. 24 (1962) 1571. [29] E.A. Cutmore, H.E. Hallam, Spectrochim. Acta, Part A: Mol. Biomol. Spectrosc. 25 (1969) 1767. [30] R. Mattes, G. Pauleickhoff, Spectrochim. Acta, Part A: Mol. Biomol. Spectrosc. 30A (1974) 379. [31] M.J. Cox, E.R.T. Tiekink, Zeitschrift fuer Kristallographie 214 (1999) 242. [32] N.B. Colthup, L.P. Powell, Spectrochim. Acta 43 (1986) 317. [33] M.R. Hunt, A.G. Kruger, L. Smith, G. Winter, Aust. J. Chem. 24 (1971) 53. [34] U. Agarwala, Lakshmi, P.B. Rao, Inorg. Chim. Acta 2 (1968) 337.
Xanthate sulfur as a hydrogen bond acceptor: the free xanthate anion and ligand sulfur in nickel tris ethylxanthate Marc A. Walters a,*, Justin Barad a, Anthony Sireci a, James A. Golen b, Arnold L. Rheingold b a b
Department of Chemistry, New York University, 100 Wash. Sq. E., New York, NY 10003, USA Department of Chemistry, University of California at San Diego, La Jolla, CA 92093-0358, USA Received 16 July 2004; accepted 22 September 2004 Available online 10 December 2004
Abstract Xanthates, like thiolates, form a variety of complexes with metals in which coordinating sulfur can serve as a hydrogen bond acceptor. Nickel tris xanthate complexes [Ni(xan)3], (xan = o-ethylxanthate, N-(carbamoylmethyl)ethylxanthate) have been synthesized and compared by a combination of X-ray crystallographic and spectroscopic measurements. Recent results from our studies of N–H S hydrogen bonding interactions in metal–xanthate complexes shows N–S distances to be longer than those in related thiolate complexes, indicative of weaker hydrogen bonds for the xanthates. The complex (Et4N)[N-(carbamoylmethyl)ethylxanthate)] adopts an extended conformation in both the solid state and solution and lacks either intraligand or intermolecular N–H S hydrogen bonds. The complex (CTA)[Ni(exa)3] exhibits N–H S hydrogen bonds between the amide group of the counterion and the ˚. ligand sulfur. The amide–sulfur N–H S distance is 3.567 A 2004 Elsevier B.V. All rights reserved. Keywords: Carbamoyl methyl xanthate; Nickel tris ethylxanthate; Hydrogen bond; Crystal structure
1. Introduction Xanthates have long proven useful in technology as chelators in aqueous media. Chief among their applications is the flotation of ores, and chelation therapy [1,2]. Since the early days of modeling in bioinorganic chemistry 1,1-dithiochelates, which include xanthates, have been employed as cysteine analogs [3–8]. Recently, they have been explored as alternatives to thiols in the formation of self-assembled monolayers on gold [9–11]. Although xanthates are generally employed in protic media, little is known about their participation in hydrogen bonding (H-bonding) or other dipolar interactions *
Corresponding author. Tel.: +1 212 998 8477; fax: +1 212 260 7905. E-mail address: [email protected] (M.A. Walters). 0020-1693/$ - see front matter 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.ica.2004.09.036
that might affect their character as chelators. In a continuation of our investigations of amide sulfur hydrogen bonding, we have extended our inquiry to include xanthate (xan) ligands in its free and metal complexed form. Nickel-xanthate complexes were employed in this study because of their favorable stability. In this paper, we report on the properties of xanthates relative to thiols in their capacity to form N–H S hydrogen bonds in metal complexes. The exploration of xanthates was favored by their facile synthesis, the ready availability of a large variety of alcohol precursors, and the relative stability of xanthates in air. Xanthate is unlike cysteine in its electronic structure. Therefore we will refer to data on thiolates and metal–thiolate complexes as benchmarks for the results presented here. Our aim was to prepare xanthates that are designed for intraligand H-bonding. Xanthates form amide sulfur
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N–H S hydrogen bonds in the solid state. However, the strength of the H-bond interaction is not sufficient to support the formation of intraligand N–H S hydrogen bonds. Below we discuss the formation of metal–ligand bonds and H-bonds in a set of Ni xanthate complexes and compare the xanthates with thiols in their capacity to form covalent bonds to metal ions and hydrogen bonds with amide donor groups.
2. Experimental 2.1. Synthesis of N-hydroxyethylacetamide, CH3CONHCH2CH2OH The starting alcohol was prepared by a previously established method [12]. Ethyl acetate 30 ml (0.307 mol) and ethanol amine 24.6 ml (0.410 mol, 25% excess) were combined in a 150 ml round bottom flask and refluxed for 4 h with stirring to yield a light yellow solution. The product was isolated by vacuum distillation at 108–110 C as a colorless, viscous oil; Yield: 59.9%. 2.2. Synthesis of potassium N-(carbamoylmethyl)ethylxanthate, CH 3 CONHCH 2 CH 2 OCS 2 K þ , Kcxa, (1) The xanthate, Kcxa, was synthesized by a method first described by DAmico and Bollinger [13]. The starting alcohol 1.89 g (0.018 mol) was combined with excess CS2 (15 ml). KOH 1.03 g (0.018 mol) was then added with stirring to a concentrated aqueous solution <1 ml over a period of 10 min. The reaction was allowed to stir for 24 h at room temperature. The resulting orange oil was then dried under vacuum and washed with acetone and methanol to yield a bright yellow powder. 1H NMR (DMSO): 1.8 ppm (s, 3H, acetyl), 3.3 ppm (m, 2H, methylene), 4.2 ppm (t, 2H, methylene), 8.0 ppm (s, 1H, amide); Yield: 47.5%. 2.3. Synthesis and crystallization of tetraethylammonium N-(carbamoylmethyl)ethylxanthate, Et4N, [cxa] (2) Equimolar amounts of Kcxa and Et4NBr were dissolved in acetonitrile. The resulting yellow solution was filtered to remove KBr and dried under vacuum giving a light yellow powder, Et4Ncxa. The slow diffusion of ether into a concentrated solution of the xanthate in acetonitrile yielded large, orange/yellow crystals. 2.4. Synthesis of nickel bis cxa, [Ni(cxa)2] The bis complex was prepared in a manner analogous to that reported for Ni(exa)2 [14–16]. NiCl2 Æ 6H2O 1.18 g (0.005 mol) in 10 ml water was added with stirring to an aqueous solution of Kcxa (2.16 g (0.01 mol) in 10 ml water). The resulting brown precipitate, Ni(cxa)2, was
filtered and dried under vacuum. The complex was recrystallized from DMF/ether; Yield 95%. 2.5. Synthesis of potassium nickel tris cxa, K[Ni(cxa)3] For spectroscopic and electrochemical studies the complex K[Ni(cxa)3] was formed in situ by combining [Ni(cxa)2] with Kcxa in CH3CN. The isolated complex is an oil and inconvenient to handle. The potassium nickel tris cxa complex, K[Ni(cxa)3], was prepared in a manner similar to that reported for K[Ni(exa)3] [14,15]. Ni(cxa)2 1.00 g (2.4 · 103 mol) was partially dissolved in acetonitrile. The xanthate salt Kcxa 0.524g (2.4 · 103 mol) was added with stirring to the suspension whereupon all of the [Ni(cxa)2] dissolved and a green solution resulted. The solution was filtered and reduced by rotary evaporation. The product was isolated as a green oil after extracting acetonitrile with ether and removing the ether under vacuum. 2.6. Synthesis of carbamoylmethyltrimethylammonium bromide, CTABr 1.0 g (7.24 mmol) bromoacetamide was dissolved in ethanol. An excess of trimethylamine was bubbled into the solution for about 5–7 min. The solution was left overnight in a tightly closed flask. After 24 h, a white crystalline precipitate is present in the solution. The product is filtered under nitrogen atmosphere and washed with ether; Yield 93%. 2.7. Synthesis of carbamoylmethyltrimethylammonium ethylxanthate, CTAexa (3) Potassium ethyl xanthate, Kexa, 1.01 g (6.29 mmol) and CTABr 1.24 g (6.29 mmol) [17] were dissolved in methanol. The resulting yellow solution was concentrated under vaccum. The resulting viscous material was solidified by the addition of ether and was subsequently triturated and isolated as a white powder after decanting off the ether. The powder consisted of KBr and complex 3. CH3CN, 20 ml, was added to extract the product. The solution was filtered to remove KBr and then dried under rotary evaporation to yield a yellow viscous oil that was solidified to a yellow powder by the addition of ether. The product was dried in air; Yield: 97% 1H NMR (DMSO): ethyl xanthate: 1.2 ppm (t, 3H, methyl), 4.25 (m, 2H, methylene). Carbamoylmethyltrimethylammonium chloride (CTA-Cl): 3.25 ppm (s, 9H, trimethyl), 4.1 ppm (s, 2H, a carbon), 7.8 ppm (d, 2H, amide). 2.8. Synthesis of CTA[Ni(exa)3] (4) (a) CTAexa 30 mg, (1.26 · 104 mol) was dissolved in methanol, 6.5 ml. [Ni(exa)2] 24 mg, (1.26 · 104
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mol) was dissolved in isopropanol 68 ml. The solutions were combined with stirring in a round bottomed flask and the volume was reduced to about 50 ml via rotary evaporation. The solution was chilled in refrigerator and yielded small green crystals after 24 h. The crystals were collected by vacuum filtration, washed with pentane followed by ether, and dried under a flow of air. The product was recrystallized from acetone by the addition of ether by vapor diffusion; Yield: 49%. (b) In a second method CTA(exa), 30 mg (1.26 · 104 mol), was dissolved in less than 1 ml of acetone. Complete dissolution was achieved by the brief application of gentle heat. Then, an equimolar amount of the complex [Ni(exa)2] was added as a solid. The resulting emerald green acetone solution was filtered through a plug of Kimwipe in a Pasteur pipette. The solution was transferred to a test tube which was placed in a sealed jar containing 10 ml of ether to crystallize the sample by the diffusion of ether in a refrigerator at 0 C. After about 24 h crystals were collected on a Buchner funnel, washed with hexanes, and dried in air. The product was stored in a capped vial in a freezer to retard decomposition; Yield: 60%.
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Elemental analysis was carried out by QTI. Anal. Calc. for C14H28N2O4S6Ni: C, 31.17; H, 5.23; N, 5.19. Found: C, 31.34; H, 5.15; N, 5.18%. 2.9. Cyclic voltammetry The working and counter electrodes consisted of platinum wire. Silver wire served as a quasi-reference electrode. The potential of the ferrocene/ferricinium couple, measured at room temperature was 0.93 V. Experiments were run in a 0.1 M solution of tetrabutylammonium perchlorate (TBAP) in acetonitrile. 2.10. X-ray crystallographic structure determination Crystallographic data for 2 and 4 are shown in Table 1. Data were collected with Bruker APEX equipped diffractometers and processed using the current versions of BRUKER software (Bruker AXS, Madison, WI). A noncentrosymmetric setting was chosen for 2 due to the absence of mirror-plane symmetry; refinement of the Flack parameter indicated that the reported hand is correct. Both structures were solved by direct methods and refined with anisotropic thermal parameters for all nonhydrogen atoms. Hydrogen atoms were
Table 1 Crystallographic data for Et4N(cxa) (2) and (CTA)[Ni(exa)3] (4) 2
4
Empirical formula C13H28N2O2S2 C14H28N2NiO4S6 Formula weight 308.49 539.45 Temperature (K) 218(2) 218 (2) ˚) Wavelength (A 0.71073 0.71073 Crystal system monoclinic monoclinic Space group P21 P21/n ˚) a (A 8.6621(9) 8.0348(5) ˚) b (A 9.6615(10) 16.3259(10) ˚) c (A 10.7264(11) 18.4137(11) b () 107.609(2) 93.5560(10) ˚ 3) Volume (A 855.62(15) 2410.8(3) Z 2 4 Dcalc (Mg/m3) 1.197 1.486 Absorption coefficient (mm1) 0.312 1.346 F(0 0 0) 336 1128 Crystal size (mm3) 0.30 · 0.30 · 0.30 0.40 · 0.30 · 0.20 h Range for data collection () 2.47–26.50 1.67–28.29 Index ranges 8 6 h 6 10; 12 6 k 6 11; 13 6 l 6 12 9 6 h 6 10; 21 6 k 6 21; 24 6 l 6 24 Reflections collected 5069 17 526 Independent reflections (Rint) 3025 (0.0163) 5713 (0.0231) Completeness to h 26.50, 99.4% 28.29, 95.4% Absorption correction SADABS SADABS Refinement method full-matrix least-squares on F2 full-matrix least-squares on F2 Data/restraints/parameters 3025/1/244 5713/0/244 Goodness-of-fit on F2 1.130 1.062 Final R indices [I > 2r(I)]a,b R1 = 0.0511; wR2 = 0.1383 R1 = 0.0294; wR2 = 0.0802 R1 = 0.0345; wR2 = 0.0826 R indices (all data) R1 = 0.0544; wR2 = 0.1425 ˚ 3) Largest difference peak and hole (e A 0.458 and 0.208 0.557 and 0.380 P P a R ¼ jjF o jP jF c jj= jF 0 j: P b 2 2 RðwF Þ ¼ f ½wðF o F 2c Þ2 = ½wðF 2o Þ2 g1=2 ; w ¼ 1=½s2 ðF 2o Þ þ ðaP Þ2 þ bP ; P ¼ ½2F 2c þ maxðF o ; 0Þ=3:
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treated as idealized contributions. In 2, the alkyl groups in the counterions were disordered equally over two positions.
3. Results and discussion 3.1. Structure of Et4N(cxa) (2) The ligand cxa was designed to form a seven-membered ring that is closed by the formation of an intraligand N–H S hydrogen bond. When ligand hydrogen bonding persists in solution, it allows the study of H-bonding effects on redox potentials by solution dependent methods such as cyclic voltammetry (CV). A similar strategy was employed successfully in the study of N–H S hydrogen bonding in a series of molybdenum thiolate complexes [18,19]. Earlier studies showed that intraligand N–H S hydrogen bonds can occur in both an isolated ligand and in its metal–ligand complex [20–22]. The ligand cxa however shows no propensity to form such H-bonds. In its anionic form the ligand is extended as observed in its X-ray crystal structure (Fig. 1). The C–Savg bond distance of 2 is ˚ (Tables 1a and 1b) and is essentially the same 1.680(3) A ˚ [23]. as that reported for the salt (Et4N)exa, 1.678(3) A As observed in the packing diagram of 2, close contacts are limited to C(3)–H S hydrogen bonds (Fig. 1).
Table 1a ˚ ) Et4N cxa Bond lengths (A S(1)–C(1) S(2)–C(1) O(1)–C(1) O(1)–C(2) O(1)–C(4) N(1)–C(4) N(1)–C(3) N(2)–C(12 0 ) N(2)–C(6) N(2)–C(10 0 ) N(2)–C(10) N(2)–C(8 0 )
1.686(3) 1.673(3) 1.355(3) 1.434(4) 1.230(4) 1.340(4) 1.440(4) 1.479(7) 1.488(7) 1.499(7) 1.508(5) 1.515(7)
N(2)–C(8) N(2)–C(6 0 ) N(2)–C(12) C(2)–C(3) C(4)–C(5) C(6)–C(7) C(6 0 )–C(7 0 ) C(8)–C(9) C(8 0 )–C(9 0 ) C(10)–C(11) C(10 0 )–C(11 0 ) C(12 0 )–C(13 0 ) N(1)–H(1) O(2) N(1)–H(1) O(4) N(1)–H(1) S(4)
1.537(6) 1.575(6) 1.594(8) 1.518(4) 1.497(4) 1.480(16) 1.519(10) 1.590(11) 1.460(15) 1.545(9) 1.515(11) 1.271(10) 3.201 3.006 3.567
Table 1b Bond angles () Et4N cxa
3.2. Structure of (CTA)[Ni(exa)3] (4)
C(1)–O(1)–C(2) C(4)–N(1)–C(3) C(12 0 )–N(2)–C(10 0 ) C(10 0 )–N(2)–C(8 0 ) C(6)–N(2)–C(8) C(10)–N(2)–C(8) C(12 0 )–N(2)–C(6 0 ) C(8 0 )–N(2)–C(6 0 ) C(8)–N(2)–C(6 0 ) C(6)–N(2)–C(12) C(10)–N(2)–C(12) C(8)–N(2)–C(12) O(1)–C(1)–S(2) O(1)–C(1)–S(1)
118.8(2) 122.9(3) 98.7(4) 109.8(6) 91.1(4) 109.8(6) 106.7(4) 104.9(4) 109.1(4) 106.8(5) 97.4(4) 85.4(4) 120.9(2) 113.8(2)
S(2)–C(1)–S(1) O(1)–C(2)–C(3) N(1)–C(3)–C(2) O(2)–C(4)–N(1) O(2)–C(4)–C(5) N(1)–C(4)–C(5) C(7)–C(6)–N(2) C(7 0 )–C(6 0 )–N(2) N(2)–C(8)–C(9) C(9 0 )–C(8 0 )–N(2) N(2)–C(10)–C(11) C(13)–C(12)–C(12 0 ) C(13 0 )–C(12)–N(2) N(2)–C(12 0 )–C(13)
In studies of metal–thiolate complexes we have observed N–H S bond formation between an anionic complex and its cationic counterion with donor NH groups [17,24]. The cation CTA has proved to be an ideal H-bond donating counterion for the study of amide–ligand H-bonds. In the complex (CTA)
[Ni(exa)3] (4) the CTA counterion donates hydrogen bonds to both oxygen and sulfur atoms of exa (Fig. 2). In the complex (Et4N)[Ni(exa)3] which lacks H˚ (Tabonds, the average Ni–S bond length is 2.439 A bles 1c and 1d) [12].
Fig. 1. Molecular structure packing diagram of Et4N(cxa) (2). Dotted lines denote C–H S hydrogen bonds.
125.31(17) 108.2(3) 111.9(3) 122.3(3) 122.2(3) 115.5(3) 113.1(9) 112.9(6) 112.5(6) 116.3(8) 113.5(6) 102.2(9) 114.8(6) 110.5(5)
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Fig. 2. Molecular structure of (CTA)[Ni(exa)3] (4). Dashed line – – – denotes N–H S hydrogen bonds.
By comparison, the average Ni–S bond length in 4 is ˚ to a value of 2.427 A ˚. decreased by 0.012 A In complex 4, three types of H-bonds are observed with N–H donation to sulfur, S(4) ether oxygen, O(2) and carbonyl oxygen, O(4), respectively. The N–H S hydrogen bond of 4 is weak. Nitrogen and sulfur are ˚ , which is at the limit of what is separated by 3.567 A generally considered to be an N–H S hydrogen bond in simple organic molecules [25]. In the N–H O(2) ˚, bond, nitrogen and oxygen are separated by 3.201 A which is a length characteristic of hydrogen bonds that are formed by ether oxygen. The N–H O(4) bond be˚ in length and is tween adjacent CTA cations is 3.006 A typical of an amide dimer [25]. In alkylthiolate complexes, H-bond donation to sulfur is expected to result in metal–ligand bond lengthening [24]. The decrease in Ni–S bond length in 4 relative to (Et4N)[Ni(exa)3] is unexpected. A significant caveat in the analysis of CTA amide hydrogen bonding is that the NH group is trifurcated as a hydrogen bond donor. We must exercise caution in the evaluation of N–H S hydrogen bonding
Table 1c ˚ ) for CTA[Ni(exa)3] Selected bond lengths (A Ni(1)–S(4) Ni(1)–S(5) Ni(1)–S(1) Ni(1)–S(3) Ni(1)–S(2) Ni(1)–S(6) S(1)–C(1) S(2)–C(1) S(3)–C(4)
2.4081(5) 2.4107(5) 2.4294(5) 2.4324(5) 2.4414(5) 2.4416(5) 1.6780(16) 1.6943(16) 1.6780(17)
S(4)–C(4) S(5)–C(7) S(6)–C(7) O(1)–C(1) O(1)–C(2) O(2)–C(4) O(2)–C(5) O(3)–C(7) O(3)–C(8)
1.6883(18) 1.6825(17) 1.6894(18) 1.327(2) 1.452(2) 1.3363(19) 1.459(2) 1.3372(19) 1.457(2)
in this system since the interaction will be affected by N–H O hydrogen bonds to a nearby xanthate and the carbonyl oxygen of a neighboring amide group. A similar investigation of related dithiocarbamates likewise showed no perturbation of Ni–S covalent bonds as a function of N–H S hydrogen bond formation [26]. Even when the H-bond acceptor is an alkane thiolate, a ˚ is observed in the average M–S small increase of 0.01 A bond length relative to M–S bonds in non-hydrogen bonded complexes [24]. 3.3. Structure of K[Ni(cxa)3] The coordination geometry of K[Ni(cxa)3], which was reproducibly isolable only as a green oil, could be
Table 1d ˚ ) for CTA[Ni(exa)3] Selected bond angles (A S(4)–Ni–(1)–S(5) S(4)–Ni(1)–S(1) S(5)–Ni(1)–S(1) S(4)–Ni(1)–S(3) S(5)–Ni(1)–S(3) S(1)–Ni(1)–S(3) S(4)–Ni(1)–S(2) S(5)–Ni(1)–S(2) S(1)–Ni(1)S(2) S(3)–Ni(1)–S(2) S(4)–Ni(1)–S(6) S(5)–Ni(1)–S(6) S(1)–Ni(1)–S(6) S(3)–Ni(1)–S(6) S(2)–Ni(1)–S(6) C(1)–S(1)–Ni(1)
95.025(18) 97.795(17) 166.038(17) 73.892(16) 91.880(18) 96.959(18) 166.514(17) 94.844(17) 73.468(15) 96.619(16) 97.887(17) 73.896(16) 98.787(17) 163.118(19) 93.666(16) 83.82(6)
C(1)–S(2)–Ni(1) C(4)–S(3)–Ni(1) C(4)–S(4)–Ni(1) C(7)–S(5)–Ni(1) C(7)–S(6)–Ni(1) C(1)–O(1)–C(2) C(4)–O(2)–C(5) C(7)–O(3)–C(8) O(1)–C(1)–S(1) O(1)–C(1)–S(2) S(1)–C(1)–S(2) O(1)–C(2)–C(3) O(2)–C(4)–S(3) O(2)–C(4)–S(4) S(3)–C(4)–S(4) O(2)–C(5)–C(6)
83.12(6) 82.78(6) 83.34(6) 83.67(6) 82.57(6) 118.96(13) 118.13(15) 117.56(13) 124.66(12) 115.84(11) 119.50(10) 107.50(15) 123.57(13) 116.83(12) 119.60(9) 107.15(19)
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inferred from its electronic absorption spectrum and the structure of the ligand salt, Et4Ncxa. The crystal structure of (Et4N)[Ni(exa)3] shows a complex that is significantly distorted from octahedral geometry, with approximate 3-fold symmetry about the central Ni2+ ion [12]. The electronic absorption spectrum of [Ni(cxa)3] in an acetonitrile solution is congruent with that of the [Ni(exa)3] complex, and is distinct from the planar complex [Ni(cxa)2] (Table 2). The bands of the nickel xanthate complexes have been assigned to ML and LM charge transfer transitions and one internal Sp ! p* transition respectively [27]. Paradoxically, the internal ligands p ! p* transition appears to be sensitive to the geometry or coordination number of the Ni(II) complexes (Table 2). From the crystal structure of the Et4Ncxa salt it is reasonable to assume that for the [Ni(cxa)3] in the solution phase the ligand is in an elongated conformation perhaps with amide to carbonyl intermolecular H-bonding. 3.4. Infrared spectra A secondary amide such as N-ethylacetamide has infrared bands at 3472 and 1720 cm1 in the gas phase that have been assigned to m(NH) and m(CO) stretching modes, respectively [29]. The mid-infrared solid state spectrum of the Kcxa salt shows corresponding bands at 3269 and 1649 cm1 characteristic of a secondary amide hydrogen bond donor (Table 3). The corresponding bands [Ni(cxa)2] occur at similarly decreased Table 2 Electronic absorption data (nm, in CH3CN solvent) on selected Ni xanthate complexes [Ni(cxa)2]
[Ni(exa)2]a
K[Ni(cxa)3]
K[Ni(exa)3]a
Assignmenta,b
196 216 252 316
218 252 316
198 218 270 316
194 218 270 312
d ! p* Sp ! p* Sr ! dr*
a
Ref. [27]. Ref. [28].
b
frequencies, 3280 and 1645 cm1, respectively, which suggests that the amides of [Ni(cxa)2] are hydrogen bonded in the solid state. In [Ni(cxa)3] the amide m(NH) band occurs at about 3404 cm1 and is likely shifted to lower frequency by interactions with the solvent acetonitrile [29]. Finally, in the complex (CTA) [Ni(exa)3] the amide group of the CTA+ counterion has four m(NH) bands in the range 3455–3166 cm1 whose provenance has not been determined. The bands assigned to the xanthate moiety are unremarkable, with the m(COC)asymm band in the region 1160–1280 cm1, and the m(CS) band in the region 1026–1070 cm1. The m(COC)asymm band is sensitive to coordination geometry and occurs in the 1270 cm1 region for the neutral planar complex [Ni(xan)2] and in the 1180 cm1 region for the anionic distorted octahedral complex [Ni(xan)3]. As discussed earlier by Mattes and Pauleickhoff, the frequency of the m(COC)asymm band is related to the resonance forms of the xanthate anion as illustrated below [30].
S-1
S-1/2 R
O
R
C
O
S-1
S-1/2
I
C
II
In resonance form II the larger bond order of oxygen results in a larger O–C force constant that is reflected in a higher frequency for m(COC)asymm. With its higher electron density, the sterically demanding sulfur of form II is expected to be more characteristic of the planar [Ni(xan)2] complex, which has an average Ni–S bond length ˚ [31]. Conversely, steric and electronic factors of 2.21 A should favor resonance form I in a Ni(II) tris xanthate complex, [Ni(xan)3], complex that has an average Ni– ˚ [12]. S bond length of 2.439(2) A
Table 3 Infrared data on selected xanthate complexes and salts Kcxaa
[Ni(cxa)2]a
3269 1649 1545 1160 1047
3280 1645 1559 1263 1050
a b c d e
KBr pellet. CH3CN solvent. Ref. [32]. Ref. [33] Ref. [34].
[Ni(exa)2]a
K[Ni(cxa)3]b
1276 1026
3405 1676 1541 1177 1069
K[Ni(exa)3]a
(CTA)[Ni(exa)3]a
Assignmentsc,d,e
1184 1041
3455, 3317, 3265, 3166 1704 1595 1181 1038
m(NH)amide m(CO)amide d (NH2)amide C–O–C CS2 out of phase
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3.5. Cyclic voltammetry
Table 5 Atomic charges on sulfur
As evidenced by X-ray crystallographic and spectroscopic data, the xanthate moiety of cxa and exa are congruent in both their structural and electronic characteristics. However they diverge in their voltammetric properties in that the redox potential of cxa is significantly anodic relative to that of exa (Table 4). The free ligand salt Kcxa undergoes an irreversible oxidation at 1.53 V vs. a Ag/Ag+ quasi-electrode as compared with an oxidation peak at 1.00 V for Kexa. Anionic xanthate is known to oxidize to a neutral radical that quickly and irreversibly dimerizes, producing the stable dixanthogen [12]. This mechanism explains the lack of a return reduction peak. The difference in oxidation potentials for the free anionic ligands is not likely due to inductive effects of the amide group in cxa but rather to differences to the free energy of dixanthogen formation from the neutral radicals. In acetonitrile the Ni2+/Ni3+ couple for the green complex [Ni(cxa)3] (formed in situ) has a value of 0.62 V with Ired/Iox = 0.23 against a silver wire quasi reference electrode at 25 C. According to Choudhury et al. [12] in work on the complex [Ni(exa)3], the anionic species is oxidized to the thermally unstable neutral Ni(III) complex, [Ni(xanthate)3], the decomposition of which accounts for Ired/Iox < 1. As with CV experiments on [Ni(exa)3], the Ired/Iox of [Ni(cxa)3] approaches 1 with decreasing temperature, which suggests a shift of equilibrium towards the tris complex. Although the redox couple at 0.62 V can be attributed to either a ligand centered redox process or a Ni(II)–Ni(III) couple, we have assigned it to the latter. Cyclic voltammetric experiments on the complexes [Zn(exa)2] and [Co(exa)3] show ligand-centered redox potentials that are more anodic than that of the free xanthate anion [14]. We therefore expect ligand centered oxidations of chelated cxa and exa to occur well above 1 V. The complex K[Ni(exa)3] has a redox couple at 0.64 V (Ired/Iox = 0.59), that is nearly identical to the redox potential of K[Ni(cxa)3] (0.62 V). Such a small difference in potential is consistent with solution IR studies
Group
Table 4 Cyclic voltammetric data on selected xanthates and Ni xanthate complexes Compound a
Kcxa Kexa Kexab K[Ni(cxa)3] K[Ni(exa)3] K[Ni(exa)3]b a b
Redox Potential (mV)/Ag at 25 C 1.53 1.00 0.23 0.62 0.64 0.11
(oxidative peak only) (oxidative peak only) (SCE) Ired/Iox = 0.23 Ired/Iox = 0.59 (SCE)
Ag wire quasi reference electrode unless otherwise indicated. Ref. [12].
Ph–S Ph–CH2–S H3C–S EtO–CS 2 ðexaÞ H3 CCðOÞNHðCH2 Þ2 OCS 2 ðcxaÞ a b
639
Charge 0.844a 0.978a 1.063a 0.362b 0.465, 0.288b
Ref. [24]. This work.
of the K[Ni(cxa)3] that showed no evidence of N–H S hydrogen bonding in the [Ni(cxa)3] complex. 3.6. Computational assessment of xanthate sulfur and hydrogen bonding Xanthate ions can form hydrogen bonds to sulfur both in the metal ligated form and in the free ligand state. Sulfur is relatively is weak as an H-bond acceptor relative to the more electronegative elements nitrogen and oxygen. Thus, it is not unexpected that N–H S hydrogen bonds would be weak in xanthate. However, this type of hydrogen bond is significantly weaker than in CTA metal–thiolate complexes such as (CTA) [Co(SPh)4] that exhibits an amide–thiolate N–S average ˚ as compared with 3.6 A ˚ in complex 4 distance of 3.3 A [17]. The qualitative manifestation of hydrogen bond strength is the donor–acceptor distance. An explanation of the disparities in N–H S distances must focus on the nature of sulfur in xanthates as compared to thiolates. A qualitative assessment suggests that the charge on sulfur in xanthate is less than in thiolate. Computational evaluation supports this view (see Table 5). For computational analysis the geometries of exa and cxa were optimized using AM1. Single point energy level calculations were then carried out at the 6-31G* level. The charge on individual sulfur atoms in xanthate is approximately 0.36 e as a result of electron delocalization between the two sulfur atoms (resonance structures) and the inductive effect of oxygen. This renders xanthate a very weak H-bond acceptor in contrast to thiolate where the charge on sulfur is approximately 0.99 e. Metal complexation is expected to reduce the charge on sulfur in both cases, but with xanthate ligands the effect seriously weakens the N–H S bond strength.
4. Conclusions Xanthate as a free ion or as a ligand in a Ni complex forms relatively weak N–H S hydrogen bonds with amide donors. Ni–Sxan bonds are also relatively weak. These two phenomena are due to the relatively low charge density on xanthate sulfur as compared with thiolate sulfur. The clearest manifestation of these
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phenomena are the large N–H S and Ni–S bond distances in complex 4.
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