Ag(I), Cu(II), Co(III) and Hg(II) complexes and metal-assisted products derived from 4-methyl-piperidine-carbodithioate: syntheses, structures, thermal analyses, redox behaviour and fluorescence properties

Abstract

Four new complexes [Ag₂(4-mpipdtc)₂(PPh₃)₂] (1), [Cu(4-mpipdtc)₂] (2), [Co(4-mpipdtc)₃]·CHCl₃ (3) and [PhHg(4-mpipdtc)] (4) and two new products, bis(4-methyl piperidinethiocarbonyl) disulfide {(4-mpipdtc)₂} (5) and (4-methyl-piperidin-1-yl) carbothioylsulfanyl-methyl (4-methyl)-piperidine-1-carbodithioate {CH₂(4-mpipdtc)₂} (6) have been obtained in this study. The syntheses of compounds 5 and 6 were assisted by Mn(II) and Ag(I) ions, respectively, and were obtained from potassium 4-methyl-piperidine-carbodithioate {K⁺(4-mpipdtc)⁻} {where, 4-mpipdtc⁻ = 4-methyl piperidine carbodithioate}. All new compounds have been characterized by elemental analyses, IR, NMR, magnetic susceptibility and single X-ray crystallography techniques. These compounds are stabilized by intermolecular C–H⋯S, S⋯S, C–H⋯π and C–H⋯N interactions. The most interesting feature in complex 4 is that the ligand bound phenylmercury cation is stabilized via intermolecular as well as intramolecular Hg⋯S secondary interactions. Compounds 3 and 5 are highly fluorescent in a solution when compared to the free ligand and emit violet/violet-blue light at 372 and 413 nm upon excitation at 328 and 300 nm, respectively. The course of the thermal degradation of complexes 1–4 have been investigated by TG-DTA, which indicates that metal sulphide is formed as the final residue. The results obtained from the electronic structure calculations at the density functional theory level corroborate our experimental findings obtained from the IR data. Frontier molecular orbital analysis reveals that complexes 1 and 3 are softer and more reactive than complexes 2 and 4. Cyclic voltammetry shows that complex 2 exhibits a reversible Cu(II)/Cu(I) redox process at 0.515 V, whereas the ligand and complexes 1, 3 and 4 show irreversible redox behaviour.

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1. Introduction

Dithiocarbamates are versatile ligands capable of forming a variety of complexes with most of the transition elements in various oxidation states. Metal complexes of dithiocarbamato ligands have been attracted a lot of attention in the field of coordination and organometallic chemistry due to their applications in molecular electrical conductivity, optical, magnetic properties and as a single source precursor for the preparation of metal sulphide nanoparticles. They are also applicable in industrial areas such as rubber vulcanization accelerators, flotation agents in metallurgy and petroleum additives, and have importance in biological processes.^1–10 Ag(I) dithiocarbamato complexes have extensively been used in the formation of Ag₂S nanoparticles and nanoscale metals.¹¹ Cobalt(III) dithiocarbamato complexes have been used as catalysts in organic transformations such as the enamination reaction.¹² Furthermore, fluorescent nanowires of cobalt(III) dithiocarbamato complexes have been synthesized and used in optoelectronic devices.¹³ Organomercury(II) dithiocarbamato complexes exhibit a great affinity for sulfur donor sites present in protein frameworks.^14–16 Piperidine dithiocarbamato ligands generally form four membered S,S-chelate rings and bridge two metal centres due to the dominant contribution of the resonance form III (Scheme 1).¹⁷ The thiol form may undergo aerial oxidation to form a disulphide or coupling of two thiol groups via a metal-catalyzed reaction. The most common route for creating disulphide bonds is via the oxidation of thiol groups, which provides stability to protein dimers, polymers and complexes in which the sulphide bonds helps in protein folding.¹⁸ Metal ion-assisted cyclization of thiosemicarbazides into their corresponding oxadiazole/thiadiazole and disulphide bond formation have been reported.^19,20 The synthesis of bis(dialkylaminethiocarbonyl) disulfides has been reported but their structural characterization has not been performed. Moreover, these syntheses involve sodium nitrate and hydrochloric acid, which resulted in poor yields of the products.²¹ Disulfide bond formation of methylene bis-dithiocarbamates has been reported either in the absence^21–23 or the presence of metal salts.²⁴ Cyclic and open-chain amines react with carbon disulfide and methylene halides in the presence of an ionic liquid ([pmIm]Br) at room temperature to produce the corresponding dithiocarbamates.²³ These symmetrical bis-dithiocarbamates are used further as valuable intermediates in the synthesis of several biologically active molecules.²⁵ Some complexes of the methyl-piperidine carbodithioates that have been reported are studied on the basis of their electronic and infrared spectra but no data pertaining to single X-ray crystallography is available.²⁶ In view of the interesting bonding modes and various properties as discussed above, herein we report on the syntheses, photoluminescence, thermal behaviour and structural characterization of Ag(I), Cu(II), Co(III) and Hg(II) complexes of the 4-methyl piperidine dithiocarbamato ligand. Additionally, the formation of the disulfide bridged product 5 obtained using Mn²⁺ and another product 6 containing the S–CH₂–S group, catalysed by Ag⁺ ions are also reported.

Scheme 1 The resonance structures of 4-methyl-piperidine-carbodithioate.

2. Results and discussion

Compounds 1–6 were synthesized in a mixture of methanol–chloroform via a simple and very convenient method at room temperature. These compounds have been characterised by various physiochemical techniques. For the structural studies, crystals of 1–6 were grown at room temperature using a slow solvent evaporation technique. A methanol solution of the ligand, 4-methyl-piperidine-carbodithioate reacted with AgNO₃ to form a white precipitate, which dissolved upon the addition of a chloroform solution of PPh₃ yielding [Ag₂(4-mpipdtc)₂(PPh₃)₂] (1). A methanol-chloroform solution of the ligand reacted with Cu(OAc)₂·H₂O/CoCl₂/PhHgOAc to form clear solutions yielding the complexes [Cu(4-mpipdtc)₂] (2), [Co(4-mpipdtc)₃]·CHCl₃ (3) and [PhHg(4-mpipdtc)] (4), respectively. Scheme 2 depicts the formation of complexes 1–4 in which complex 1 contains triphenylphosphine as a co-ligand. {(4-mpipdtc)₂} (5) and {CH₂(4-mpipdtc)₂} (6) were synthesized via the reaction of the 4-methyl-piperidine-carbodithioate with Mn(OAc)₂·4H₂O and AgNO₃ in methanol and methanol–dichloromethane mixture, respectively, at room temperature with simple stirring of the reagents. Scheme 3 depicts the formation of compounds 5 and 6. We have synthesized bis(dialkylaminethiocarbonyl) disulfide from the potassium salt of dithiocarbamates assisted by Mn(II) acetate at room temperature, this method is different from the reported synthesis.²¹ Both the syntheses are simple one-pot and multi-component reactions completed in two steps. The present route is relatively safer involving less hazardous chemicals that can be easily handled. Furthermore, the formation of symmetrical methylene bis(dithiocarbamates) reported in the literature have been the by-products of the reaction between transition metal halides and sodium or potassium salts of dithiocarbamates when methylene chloride was used as the solvent. Generally, methylene chloride can react with strong nucleophiles such as the sodium salt of thiols to give the corresponding symmetrical bis-dithio compounds, however, it requires long reaction times and gives low product yields. In the present reaction, methylene chloride acts as an electrophile and the sodium/potassium salt of thiols/dithiocarbamates acts as strong nucleophiles. Although the reaction proceeded equally well with CH₂Cl₂, CH₂Br₂ and CH₂I₂ for the synthesis of compound 6, CH₂Cl₂ was preferred due to its widespread availability and low cost. Our work reports the silver(I) nitrate assisted synthesis of symmetrical bis-dithiocarbamates in the presence of dichloromethane with product yields in the range of 80–90%. Complex 1 is soluble in chloroform and DCM and complex 2 is soluble in chloroform and DMSO. Compounds 2, 3, 5 and 6 are soluble in methanol, acetonitrile and DMSO. Compounds 1, 2, 3, 4, 5 and 6 melt at 196, 213, 309, 248, 102 and 140 °C, respectively.

Scheme 2 Syntheses of complexes 1–4.

Scheme 3 Syntheses of the Mn(II)- and Ag(I)-assisted products 5 and 6.

2.1 Magnetic moments and electronic spectra

Complexes [Ag₂(4-mpipdtc)₂(PPh₃)₂] (1), [Co(4-mpipdtc)₃]·CHCl₃ (3) and [PhHg(4-mpipdtc)] (4) are diamagnetic indicating the presence of low spin Ag(I), Co(III) and Hg(II) centres. Complexes 1, 3 and 4 showed absorptions in the region of 263–328 nm due to intra-ligand/charge transfer transitions Fig. 1(a). The complex, [Cu(4-mpipdtc)₂] (2) showed a magnetic moment of 1.75 BM, which indicates the presence of an unpaired electron. The presence of a broad band around 439 nm was assigned to the envelope of the ²B_1g → ²A_1g, ²B_2g and ²E_g transitions and suggests a square planar geometry for the complex (ESI Fig. 1†). Another high energy band observed at 272 nm was assigned to π → π intra-ligand/charge transfer transitions.²⁷

Fig. 1 (a) The absorption spectra of the compounds at 2 × 10⁻⁵ M and (b) the emission spectra of the compounds at 2 × 10⁻⁵ M excitation wavelength was 300 nm for compounds 1–6, and the ligand and also at 328 nm for compound 3 in graph 3(b).

2.2 IR spectra

The IR spectrum of the potassium derivative of the 4-methyl-piperidine-carbodithioate ligand, {K⁺(4-mpipdtc)⁻}, showed absorptions due to the stretches attributed to C–H (2951 cm⁻¹), C–N (1585 cm⁻¹) and C=S (1015 cm⁻¹) bonds, respectively. The IR spectra of [Ag₂(4-mpipdtc)₂(PPh₃)₂] (1), [Cu(4-mpipdtc)₂] (2), [Co(4-mpipdtc)₃]·CHCl₃ (3), [PhHg(4-mpipdtc)] (4), {(4-mpipdtc)₂} (5) and {CH₂(4-mpipdtc)₂} (6) showed aliphatic ν(C–H) bands at 2989, 2948, 2949, 2947, 2950, 2924 cm⁻¹, respectively. The aromatic ν(C–H) band for 1 and 4 appeared at 3058 and 3062 nm, respectively. The IR spectra of compounds 1–6 revealed ν(C–N) bands at 1469, 1504, 1493, 1484, 1481 and 1434 cm⁻¹, respectively. The ν(C=S) band was observed at 962, 965, 966 and 968 cm⁻¹ for complexes 1–4, respectively. A new band for complexes 1–4, attributable to ν(M − S) appeared at 320 (calc.), 429, 428 and 430 cm⁻¹, respectively. The ν(C=S) band suffered a negative shift of 53, 50, 49 and 47 cm⁻¹ in complexes 1, 2, 3, and 4, respectively, when compared to the free ligand indicating the involvement of the dithio sulfur atoms in bonding. Compounds 5 and 6 showed two bands at 959, 816 and 998, 714 cm⁻¹, respectively indicating the presence of C–S and C=S bonds (Table 1) (ESI Fig. 2–8†).

Table 1 The important IR bands (cm⁻¹) for the {K⁺(4-mpipdtc)⁻} ligand and compounds 1–6

Compound	ν(C–H)	ν(C–N)	ν(C=S/C–S)	ν(M − S)
K^+(4-mpipdtc)^−	2951	1585	1015	—
[Ag2(4-mpipdtc)2(PPh3)2] (1)	3058, 2989	1469	962	320
[Cu(4-mpipdtc)2] (2)	2948	1504	965	429
[Co(4-mpipdtc)3]·CHCl3 (3)	2949	1493	966	428
[PhHg(4-mpipdtc)] (4)	3062, 2947	1484	968	430
(4-mpipdtc)2 (5)	2950	1481	959/816	—
CH2(4-mpipdtc)2 (6)	2924	1434	998/714	—

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2.3 NMR studies

The ¹H NMR spectra of {K⁺(4-mpipdtc)⁻}, 1, 3, 4, 5 and 6 showed a triplet at δ 3.76, 3.07, 2.89, 3.10, 3.32 and 2.95 ppm, respectively due to the axial proton C₁H. The signal for the equatorial protons C₁H appeared as a doublet at δ 5.67, 5.30, 4.66, 4.95, 5.00 and 4.09 ppm for {K⁺(4-mpipdtc)⁻}, 1, 3, 4, 5 and 6, respectively. For the equatorial protons, only the geminal coupling with the axial protons was large but in the case of the axial protons both the geminal coupling with the equatorial protons as well as the vicinal coupling with the adjacent axial protons are both large/significant. The deshielding of the equatorial and axial protons is due to the high electron density at the nitrogen atom,^28,29 however, the equatorial proton (H-1e) is more deshielded than the axial proton (H-1a).³⁰ The signals for the C₂H and C₃H protons appear as multiplets in the range of δ 1.20–1.75 and δ 1.68–1.79 ppm, respectively. A doublet for the methyl protons appeared in the range δ 0.93–0.99 ppm for all the compounds.³¹ The singlet for the SCH₂ protons is highly deshielded due to the presence of the sulfur atom around the protons and appeared at δ 5.37 ppm in compound 6.²⁴

The ¹³C NMR spectrum of the {K⁺(4-mpipdtc)⁻} ligand showed a signal at δ 208.76 ppm due to the CS₂ carbon. Complexes 1, 3 and 4 showed a signal for the NCS₂ carbon at δ 207.17, 203.55 and δ 200.81 ppm, respectively. This indicates the considerable double bond character in the nitrogen carbon bond of the dithiocarbamato group. A considerable δ⁺ surplus charge is localized on the nitrogen atom and a δ⁻ charge is delocalized via four-membered metallochelate ring –CS₂M.^30–32 The CS₂ carbon for compounds 5 and 6 appeared at δ 192.82 and 193.62 ppm, respectively, which are less downfield shifted than the free ligand as well as in complexes 1, 2, 3 and 4. The carbon atom (C1) adjacent to the nitrogen of dithiocarbamate was more deshielded than other carbon atoms in the piperidine ring (Scheme 4) and appeared at δ 51.29, 52.74, 45.48, 52.88, 47.43 and 52.04 ppm for {K⁺(4-mpipdtc)⁻}, 1, 2, 3, 4, 5 and 6, respectively. The C3 carbon is more deshielded than the C2 carbon and the signals for C3 and C2 were observed at δ 34.35, 30.89 {K⁺(4-mpipdtc)⁻}; 33.87, 30.39 (1); 33.13, 30.65 (3); 33.79, 30.25(4); 33.83, 30.85 (5) and 34.00, 30.83 (6) ppm.²⁹ The methyl carbon was observed in the range δ 21.19–21.56 ppm for all the compounds. In compound 6, a new signal observed at δ 54.52 ppm was due to the SCH₂S carbon.²⁴ The C=S signal was observed in the range of 200.81–207.17 ppm and were found to be shielded when compared to the free ligand (δ 208.76 ppm) due to the involvement of CS₂ in bonding.

Scheme 4 Numbering of the ligand moiety.

2.4 Photoluminescence studies

The emission spectra of compounds 1−6 were recorded in methanol and chloroform solutions at room temperature in the 10⁻⁵ M concentration range (Fig. 1). Compounds 3 and 5 were highly fluorescent in methanol. They exhibited photoluminescence properties and showed emission maxima at 373 and 413 nm (violet and violet-blue emission) upon excitation at 328 (Fig. 1(b) graph 3b) and 300 nm, respectively. Compounds 3 and 5 showed intense emission, which was slightly blue shifted when compared to the free ligand. The ligand exhibited emission maxima at 434 nm upon excitation at 300 nm. Compound 3 showed a blue shifted emission of 61 nm as it showed a maximum emission at 373 nm upon excitation at 300 nm and also showed an intense fluorescence emission at 367 nm upon excitation at 300 nm with a shift of 67 nm (Fig. 1(b)). Compounds 1, 2, 4 and 6 were relatively weak fluorescent materials and exhibited maxima at 350, 347, 347 and 348 nm respectively, upon excitation at the 300 nm. Compounds 1 and 4 showed strong blue shift emission when compared to the free ligand, probably due to the heavy metal effect or quenching behaviour or Ag⋯S/Hg⋯S secondary interactions. The violet and violet-blue emission of compounds 3 and 5 in the solution state implies that these compounds may be potentially applicable as materials for light emitting diode devices (LEDs).^33,34

2.5 Thermal studies

The thermal properties of complexes 1–4 were studied by TG-DTA under a nitrogen atmosphere in the temperature range 25–900 °C at a controlled heating rate of 10 °C min⁻¹ (Fig. 2). Thermogravimetric analysis of complex 1 showed it to be stable up to 210 °C and after that a single weight loss of 75.40% occurred in a single step between 210–320 °C to give silver sulfide (Ag₂S) as the final residue (obs.: 21.43% and calc.: 22.75%).³⁵ Complex 2 also showed almost the same pattern as that of complex 1. It was stable up to 213 °C and then lost its weight in a single step between 213–327 °C to give copper sulfide (CuS) as the end product (obs.: 21.85% and calc.: 23.19%).³⁶ Complex 3 was stable up to 150 °C. Its low thermal stability was due to the presence of a chloroform molecule in its crystal lattice. Thus, it loses chloroform before it melts (309 °C) and showed a first step weight loss at 150–310 °C (weight loss calc.: 17.02 and obs.: 16.40%).¹⁰ Thus, heating at ∼200 °C may be a good method for the conversion of [Co(4-mpipdtc)₃]·CHCl₃ to [Co(4-mpipdtc)₃]. Subsequently, in the second step between 310 and 355 °C it melts with decomposition to give cobalt sulfide (CoS₂) as the final residue (obs.: 17.70% and calc.: 17.55%).¹² Complex 4 completely decomposes in a three step process. It is stable up to 248 °C and the absence of any thermal change before this temperature indicated that sample restructuring does not take place before the degradation processes begins³⁷ and then, decomposes in the first step up to 290 °C (weight loss calcd: 48.53 and observed: 49.70%), probably forming mercuric sulfide (HgS), which is also not a stable product. The HgS also decomposes rapidly and up to 500 °C and undergoes near complete vaporization.³⁵ In the DTA curve, complex 4 showed two sharp endothermic curves around 250 °C near their melting temperatures. By comparing the TG and DTA curves, it was observed that the major weight loss process started after the melting temperatures of the compounds.³⁸ The broad exothermic hump observed in complex 4 at 500 °C implies slow decomposition leading to volatilization upon heating (ESI Fig. 9 and 10†).³⁷

Fig. 2 TGA curves showing the thermal degradation of complexes 1–4 at a heating rate of 10 °C min⁻¹ in a nitrogen atmosphere.

2.6 Electrochemical studies

The electrochemical behaviour of the ligand and its complexes were studied using cyclic voltammetry and are depicted as Fig. 3. The cyclic voltammograms were recorded with a 10⁻³ M concentration of the compounds at a platinum electrode using 0.05–0.1 M TBAP as supporting electrolyte in the potential range of −1.0 to +1.0 V with a scan rate of 20 mV s⁻¹. The ligand and complexes 1, 3 and 4 showed irreversible behaviour while complex 2 showed reversible redox behaviour. Complex 2 may be applicable in fields such as electrochemical materials and electrochemical solar cells. The ligand showed cathodic and anodic peaks at 0.140 and −0.775 V, respectively, characterizing the irreversible redox behaviour (Fig. 3(a)). Complexes 1, 3 and 4 exhibited cathodic peaks at 0.080, 0.990 and 0.120 V and the corresponding anodic peaks at −0.260, 0.285 and −0.230 V, respectively. The separation between the cathodic and anodic peak potentials (D_Ep = E_pa − E_pc) of 0.340, 0.705 and 0.350 V for complexes 1, 3 and 4, respectively, indicate that these complexes show irreversible redox behaviour.³⁹ Complex 2 exhibited the cathodic and anodic peaks at 0.545 and 0.485 V, respectively and the separation between the cathodic and anodic peak potentials (D_Ep = E_pa − E_pc) of 0.060 V indicates a reversible redox behaviour, which was assigned to the Cu(II)/Cu(I) couple with a formal redox potential E = (E_pa + E_pc)/2 = 0.515 V (Fig. 4).⁴⁰ There was no shift in the peak potentials on scanning the voltammogram at different scan rates, which indicated that the complex exhibited completely reversible behaviour. Complex 2 also exhibited another cathodic and anodic peak at −0.380 and −0.520 V, respectively, due to the ligand moiety. The free ligand showed cathodic and anodic peaks at 0.140 and −0.775 V, respectively, which are shifted to −0.380 and −0.520 V due to formation of the complex.

Fig. 3 The cyclic voltammograms obtained for the ligand and complexes: (a) the ligand in acetonitrile, (b) complex 1 in DCM-DMSO, (c) complex 2 in acetonitrile, (d) complex 3 in acetonitrile and (e) complex 4 in DMSO at Pt electrode using 0.1 M TBAP as the supporting electrolyte.

Fig. 4 The cyclic voltammograms obtained for [Cu(mpipdtc)₂] (2) at different scan rates.

2.7 Crystal structure description

The crystal structures of compounds 1–6 were obtained using single crystal X-ray diffraction data. The ORTEP diagrams of compounds 1–6 with atom numbering schemes are shown in Fig. 5–10. The details of the data collection, structure solution and refinement are listed in Table 2. Selected bond lengths and angles are included in Tables 3–8. Hydrogen bonding parameters for compounds 1–6 are given in Table 9.

Table 2 Crystallographic data, structure solution and refinement for compounds 1–6

a R1 = Σ||Fo| − |Fc||Σ|Fo|.b R2 = [Σw (|Fo^2| − |Fc^2|)^2/Σw|Fo^2|^2]^1/2.
Parameters	1	2	3	4	5	6
Empirical formula	C50H54Ag2N2P2S4	C14H24CuN2S4	C22H37Cl3CoN3S6	C13H17HgNS2	C14H24N2S4	C15H26N2S4
Formula weight	1088.87	412.13	701.18	451.99	348.59	362.62
Crystal system	Triclinic	Monoclinic	Triclinic	Monoclinic	Monoclinic	Triclinic
Space group	P1̄	P21/c	P1̄	P21/n	P21/c	P1̄
T (K)	293(2)	120(2)	120(2)	293(2)	120(2)	293(2)
λ, Mo Kα (Å)	0.71073	0.71073	0.71073	0.71073	0.71073	0.71073
a (Å)	10.538(5)	8.452(3)	10.266(2)	8.799(5)	12.715(8)	8.876(15)
b (Å)	10.829(6)	18.620(6)	12.091(4)	15.123(9)	6.421(4)	10.017(17)
c (Å)	11.215(6)	12.141(4)	14.720(5)	11.633(7)	22.315(14)	11.405(2)
α (°)	78.16(5)	90	113.48(3)	90.00	90	87.26(15)
β (°)	84.07(4)	106.57(4)	94.48(2)	108.23(9)	104.13(6)	74.87(16)
γ (°)	77.60(5)	90	105.77(2)	90.00	90	76.56(14)
V, (Å^3)	1221.2(11)	1831.4(11)	1576.3(9)	1470.2(15)	1767.1(2)	952.0(3)
Z	1	4	2	4	4	2
ρcalcd (g cm^−3)	1.481	1.495	1.477	2.042	1.310	1.265
μ (mm^−1)	1.074	5.890	10.462	10.732	4.867	0.495
F (000)	556	860	728	856	744	388
Crystal size (mm^3)	0.30 × 0.27 × 0.23	0.30 × 0.27 × 0.23	0.55 × 0.25 × 0.21	0.30 × 0.24 × 0.21	0.49 × 0.31 × 0.04	0.30 × 0.27 × 0.23
θ range for data collections (°)	3.38–28.95	3.79–76.34	3.39–76.66	2.55–28.32	3.58–76.95	3.41–28.84
Index ranges	−13 ≤ h ≤ 14	−10 ≤ h ≤ 9	−12 ≤ h ≤ 7	−11 ≤ h ≤ 11	−15 ≤ h ≤ 15	−15 ≤ h ≤ 15
	−14 ≤ k ≤ 14	−20 ≤ k ≤ 23	−14 ≤ k ≤ 15	−20 ≤ k ≤ 20	−4 ≤ k ≤ 7	−4 ≤ k ≤ 7
	−15 ≤ l ≤ 15	−15 ≤ l ≤ 15	−18 ≤ l ≤ 18	−15 ≤ l ≤ 15	−28 ≤ l ≤ 27	−28 ≤ l ≤ 27
No. of reflections collected	6511	3796	6532	3666	3616	5001
No. of independent reflections (Rint)	4196	3578	6186	2245	2888	2469
No. of data/restrains/parameters	6511/0/271	3796/0/193	6532/0/320	3666/0/154	3616/15/215	5001/0/190
Goodness-of-fit on F^2	0.968	1.076	1.027	0.962	1.011	1.128
R1^a, wR2^b [I > 2σ(I)]	0.0445, 0.1118	0.0563, 0.1541	0.0471, 0.1269	0.0408, 0.0987	0.0508, 0.1305	0.0824, 0.2245
R1^a, wR2^b (all data)	0.0646, 0.1299	0.0582, 0.1570	0.0490, 0.1288	0.0679, 0.1111	0.0647, 0.1448	0.1115, 0.1695
Largest difference in peak/hole (e Å^−3)	0.678, −0.696	1.458, −1.042	1.348, −0.861	1.480, −3.331	0.658, −0.558	0.475, −0.301

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Table 3 The interatomic distances (Å) and angles (°) obtained for [Ag₂(4-mpipdtc)₂(PPh₃)₂] (1)

Bond lengths (Å)			Bond angles (°)
	Exp.	Cal.		Exp.	Cal.
Ag1–P1	2.427	2.550	S1^i–Ag1–S2^i	67.143	65.1
Ag1–S1	2.582	2.707	S1–Ag1^i–S2	67.143	65.1
Ag1–S2^i	2.587	2.678	S1–Ag1–S2^i	110.744	112.1
Ag1–S1^i	2.787	2.678	S1–Ag1^i–S1^i	105.023	112.5
Ag1–Ag1^i	3.272	3.133	P1–Ag1–S1^i	117.073	113.5
S2–C1	1.706	1.732	P1–Ag1–S1	120.353	118.3
S1–C1	1.736	1.756	P1–Ag1–S2	123.334	123.3
C1–N1	1.327	1.353	Ag1–S1–Ag1^i	74.983	67.5

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Table 4 The interatomic distances (Å) and angles (°) obtained for [Cu(4-mpipdtc)₂] (2)

Bond lengths (Å)			Bond angles (°)
	Exp.	Cal.		Exp.	Cal.
Cu–S1A	2.295	2.398	S1B–Cu–S2B	77.302	75.3
Cu–S1B	2.295	2.398	S1A–Cu–S2A	77.592	75.3
Cu–S2A	2.305	2.398	S1B–Cu–S2A	101.192	104.6
Cu–S2B	2.307	2.398	S1A–Cu–S2B	103.502	104.6
S1A–C1A	1.729	1.738	C1A–S1A–Cu	84.598	84.8
S2A–C1A	1.724	1.738	S2A–C1A–S1A	113.201	179.9
N1A–C1A	1.325	1.339	S1A–Cu–S1B	177.253	179.9

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Table 5 The interatomic distances (Å) and angles (°) obtained for [Co(4-mpipdtc)₃]·CHCl₃ (3)

Bond lengths (Å)			Bond angles (°)
	Exp.	Cal.		Exp.	Cal.
Co–S1A	2.264	2.331	S1A–Co–S2A	76.562	75.9
Co–S2A	2.274	2.334	S2B–Co–S1B	76.532	76.0
Co–S1B	2.278	2.331	S1C–Co–S2C	76.182	75.9
Co–S2B	2.255	2.334	S1A–Co–S1B	94.562	94.9
Co–S1C	2.261	2.331	S1C–Co–S1A	94.873	94.8
Co–S2C	2.286	2.334	Cl2–C1S–Cl1	110.971	—
S1A–C1A	1.719	1.728	S2B–Co–S1A	167.413	167.3
N1A–C1A	1.321	1.344	S2A–C1A–S1A	109.901	112.3

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Table 6 The interatomic distances (Å) and angles (°) obtained for [PhHg(4-mpipdtc)] (4)

Bond lengths (Å)			Bond angles (°)
	Exp.	Cal.		Exp.	Cal.
Hg1–C4	2.059	2.237	C4–Hg1–S1	172.1	154.4
Hg1–S1	2.393	2.692	C4–Hg1–S2	115.9	133.7
Hg1–S2	2.854	2.846	S1–Hg1–S2	68.6	68.2
S1–C7	1.751	1.747	C7–S1–Hg1	93.4	89.6
S2–C7	1.720	1.753	C7–S2–Hg1	79.3	84.6
N1–C7	1.318	1.341	S2–C7–S1	118.5	119.6

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Table 7 The interatomic distances (Å) and angles (°) obtained for {4–mpipdtc₂} (5)

Bond lengths (Å)			Bond angles (°)
	Exp.	Cal.		Exp.	Cal.
S1A–C1A	1.827	1.860	C1A–S1A–S1B	102.09	104.5
S2A–C1A	1.657	1.659	N1A–C1A–S2A	126.02	126.0
C1A–N1A	1.341	1.351	N1A–C1A–S1A	112.22	111.5
S1A–S1B	2.008	2.037	S2A–C1A–S1A	121.70	122.3

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Table 8 The interatomic distances (Å) and angles (°) obtained for {CH₂4-mpipdtc₂} (6)

Bond lengths (Å)			Bond angles (°)
	Exp.	Cal.		Exp.	Cal.
S1–C7	1.7907	1.812	C7–S1–C15	102.34	102.8
S3–C14	1.783	1.812	C14–S3–C15	102.14	102.8
S2–C7	1.656	1.678	S3–C15–S1	116.84	116.5
S4–C14	1.659	1.678	S2–C7–S1	121.54	122.5
S1–C15	1.796	1.820	S4–C14–S3	122.04	122.5
S3–C15	1.791	1.819	N1–C7–S2	125.15	124.3
N1–C7	1.333	1.354	C4–N1–C6	110.56	111.9
N2–C14	1.342	1.354	C7–N1–C6	122.26	122.1

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Table 9 The hydrogen bond parameters [(Å) and (°)], C–H⋯π and S⋯S interactions obtained for compounds 1–6

Compounds	D–H⋯A	dD–H	dH⋯A	dD⋯A	<DHA	Symmetry equivalent operators
1	C11–H11⋯S11	0.930	2.786	3.696	166.17	2 – x, 1 − y, –z
	C6–H6B⋯Cg	0.970	2.775	3.646	149.80	−x, −y, −z
2	C6B–H6BA⋯S2B	0.991	2.911	3.653	132.43	x, 0.5 − y, 0.5 + z
	S1A⋯S1A			3.473
3	C4A–H4AA⋯S2C	1.000	2.994	3.895	150.52	2 − x, 1 − y, 1 − z
	C6A–H6AA⋯S2C	0.990	2.920	3.827	151.84	2 − x, 1 − y, 1 − z
4	C8–H3B⋯Cg	0.970	2.920	3.692	137.32	0.5 − x, 0.5 + y, 0.5 − z
5	C6A–H6AB···S2A	0.990	2.806	3.719	153.68	1 − x, 0.5 + y, 1.5 − z
	S1A⋯S2A			3.547
6	C4–H4B⋯S4	0.970	2.791	3.674	151.69	2 − x, 2 − y, 1 − z

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2.7.1 Crystal structure description of [Ag₂(4-mpipdtc)₂(PPh₃)₂] (1). Fig. 5 shows ORTEP diagram of complex 1 along with the atom numbering scheme. Complex 1 forms a dinuclear molecule in which each silver ion is bonded through two bridging uninegative bidentate dithiocarbamato ligands via sulfur and one triphenylphosphine co-ligand. The ligand dithiocarbamato shows a mixed structural function that includes formation of chelate rings and the sulfur bridges in the dimeric Ag(I) complex 1. The Ag–S bond lengths at the sulfur bridge (Ag(1)–S(1) = 2.582(11) Å) are shorter than the non-bridging Ag–S bond distances (Ag(1)–S(1) = 2.787(12) and Ag(1)–S(2) = 2.587(12) Å) indicating that the bridging Ag–S bonds are stronger than the non-bridging Ag–S bonds. The ligand having mixed structural function coordinates to Ag(I) ion and forms tricyclic fragments [Ag₂S₄C₂]. The geometry of this [Ag₂S₄C₂] ring can be approximated to a chair conformation.⁴¹ The bridging bond angles with Ag and S atoms of different ligand units (S(1)–Ag(1)–S(2) = 110.74(4)°) are slightly greater than a normal tetrahedral bond angle of 109.47°. The bond length P(1)–Ag = 2.4275(9) Å is quite similar to the reported Ag(I) complexes of triphenylphosphine.⁴² The carbon–nitrogen bond length of 1.327(6) Å in the carbodithioate moiety indicates a substantial delocalization of the π-electron density attributing some double bond character.³⁰ The crystal structure of the complex is stabilized by weak intermolecular C–H⋯S interactions occurring between the dithio sulfur and C–H hydrogen of the phenyl ring (ESI Fig. 11†) (Table 9). In addition, the structure is stabilized by intramolecular C–H⋯π interactions (2.775 Å) between the C–H of the piperidine ring and π electrons of a phenyl ring of PPh₃, which is well within the reported range (2.75–2.89 Å) (ESI Fig. 12†).⁴³ The intermolecular interactions are clearly hydrogen bond interactions while the intramolecular interactions deal with the C–H⋯π interactions only.

Fig. 5 The ORTEP diagram of [Ag₂(4-mpipdtc)2(PPh₃)₂] (1).

2.7.2 Crystal structure description of [Cu(4-mpipdtc)₂] (2). Fig. 6 shows the ORTEP diagram of complex 2 with the atom numbering scheme. In complex 2 the copper(II) ion is bonded through four sulfur atoms of two dithiocarbamato ligands. The CuS₄ core exhibits a distorted square planar geometry. The four Cu–S bond lengths are unequal and lie in the range 2.295(6)–2.307(6) Å. The bite angles of two chelate rings are slightly different {S(1A)–Cu–S(2A) = 77.59(2) and S(1B)–Cu–S(2B) = 77.30(2)°} and less than normal bond angle (90°), which indicate that the complex has a distorted square planar geometry. The other S(1A)–Cu–S(2B) and S(1B)–Cu–S(2A) bond angles are 103.50(2) and 101.19(2)°, respectively. The N–C bond lengths of the two NCS₂ moieties {N(1A)–C(1A) = 1.325(3) and N(1B)–C(1B) = 1.321(3) Å} are shorter than a single N–C bond length attributed to partial double bond character, which suggests the delocalization of the π-electrons in the NCS₂ moiety.³⁰ The crystal structure of complex 2 is stabilized by weak intermolecular C–H⋯S interactions occurring between the dithiosulfur atom and C–H hydrogen in the piperidine ring (ESI Fig. 13†) (Table 9). The crystal structure is further stabilized by S⋯S interactions, which lead to a wave-like architecture (ESI Fig. 14†).

Fig. 6 The ORTEP diagram of [Cu(4-mpipdtc)₂] (2).

2.7.3 Crystal structure description of [Co(4-mpipdtc)₃]·CHCl₃ (3). Fig. 7 shows the ORTEP diagram of complex 3 along with the atom numbering scheme. The coordination sphere of complex 3 is fulfilled by six dithio-sulfur atoms from three bidentate dithiocarbamato moieties of the ligand. The formation of three four membered CS₂Co chelate rings, with a bite angle in the range of 76.18(2)–76.56(2)° represent a minor deviation from an ideal octahedral geometry. The Co–S bond distances are unequal being 2.264(6), 2.274(7), 2.278(7), 2.255(6), 2.261(7) and 2.286(7) Å and are comparable to the bond distances reported earlier for other similar cobalt dithio complexes.^44–46 The carbon–sulfur distances within the chelate rings are intermediate between single and double bond lengths. These bond lengths, C(1A)–S(1A) = 1.719(3), C(1A)–S(2A) = 1.716(2), C(1B)–S(1B) = 1.714(2), C(1B)–S(2B) = 1.714(3), C(1C)–S(1C) = 1.713(3) and C(1C)–S(2C) = 1.715(2) Å (Table 5) suggest a considerable delocalization of charge in the chelate ring.⁴⁶ In the solid state, complex 3 is stabilized via two types of intermolecular C–H⋯S interactions that occur between the dithio sulfur atoms and the hydrogen atoms in the piperidine ring/hydrogen atoms of chloroform molecules (Table 9) (ESI Fig. 15†).

Fig. 7 The ORTEP diagram of [Co(4-mpipdtc)₃]·CHCl₃ (3).

2.7.4 Crystal structure description of [PhHg(4-mpipdtc)] (4). Fig. 8 shows the ORTEP diagram of complex 4 together with the atom numbering scheme. Hg(II) is bonded with phenyl carbon and the dithiosulfur atoms of the ligand, and the complex has an almost linear geometry. The C–S bond distances of 1.751(6) and 1.720(6) Å in complex 4 agree well with those reported for other similar complexes,^47,48 being an intermediate between the C–S single (1.82 Å) and C=S double (1.56 Å) bond distances.⁴⁹ The Hg–S1 bond length is 2.393(2) Å, whilst that of Hg(1)–S(2) is 2.854(2) Å, which suggests that the bond between Hg and S(2) is longer than Hg(1)–S(1), thereby reflecting the propensity of a linear geometry around Hg(II) ion. The bond angle of 172.18(16)° indicates a deviation from the ideal linear geometry. The intermolecular distance between Hg(II) ion and S(2) is 3.166 Å, which suggests a strong interaction between metal ion and sulfur atom. In the solid state, the crystal structure is stabilized by C–H⋯π interactions (2.920 Å) occurring between C–H of piperidine and π electrons of the phenyl rings (ESI Fig. 16†).

Fig. 8 The ORTEP diagram of [PhHg(4-mpipdtc)] (4).

2.7.5 Crystal structure description of (4-mpipdtc)₂ (5). Fig. 9 shows the ORTEP diagram of compound 5 along with the atom numbering scheme. The bond lengths and angles are slightly different in the two halves of (4-mpipdtc)₂ (5) with an S(2A)–S(1A)–S(1B)–S(2B) torsion angle of 90.20°, which indicates that they are nearly perpendicular to each other. In addition, the S(1)–S(2) bond length of 2.0081(10) Å is found in the normal range for a disulfide bond. The C(1A)–S(1A), C(1B)–S(1B) and S(1A)–S1(B) bond lengths of 1.827(3), 1.832(3) and 2.0081(10) Å, respectively show C–S and S–S single bond character and indicate that compound 5 is present in the disulfide form.²⁰ This is further supported by the adjacent C(1A)–S(2A) and C(1B)–S(2)B bond lengths of 1.657(3) and 1.646(3) Å, which are shorter than those in complexes 1–4 (Tables 3–7), thus suggesting double bond character. The crystal structure is stabilized by intermolecular C–H⋯S hydrogen bonding and S⋯S interactions, which leads to a linear chain architecture (ESI Fig. 17†) (Table 9).

Fig. 9 The ORTEP diagram of (4-mpipdtc)₂ (5).

2.7.6 Crystal structure description of {CH₂(4-mpipdtc)₂} (6). Fig. 10 shows the ORTEP diagram of {CH₂(4-mpipdtc)₂} (6) with the atom numbering scheme. Crystal structure of compound 6 indicates that the methylene group is bonded through the two sulfur atoms of two units of dithiocarbamate ligand at distances of 1.796(7) Å {C(15)–S(1)} and 1.791(7) Å {C(15)–S(3)}. The bond angle of 116.8(4)° {S(1)–C(15)–S(3)} is larger than the ideal tetrahedral value of 109.47°, which may be due to repulsion between the two C=S bonding electron pairs. The bonds C(7)–N(1) {1.333(8) Å} and C(14)–N(2) {1.342(8) Å} are shorter than the reported C–N average bond lengths (1.462 Å).²² The averages of the C–S and C=S bond distances are 1.790 and 1.657 Å, respectively, which are well within the range of the reported single and double bond lengths.^17,22 In the solid state, the crystal structure is stabilised by weak intermolecular C–H⋯S interactions occurring between hydrogens in the piperidine ring and sulfur atom in the dithiocarbamate ligand (ESI Fig. 18†) (Table 9).

Fig. 10 The ORTEP diagram of CH₂(4-mpipdtc)₂ (6).

2.8 Frontier molecular orbital analysis

The HOMO and LUMO are very important parameters for quantum chemistry. The most important frontier molecular orbitals (FMOs) such as the highest occupied molecular orbital (HOMO) and lowest unoccupied molecular orbital (LUMO) play a crucial part in the chemical stability of a molecule.⁵⁰ The HOMO represents the ability to donate an electron and LUMO as an electron acceptor represents the ability to accept an electron. The energy gap between the HOMO and LUMO also determines the chemical reactivity, optical polarizability and chemical hardness-softness of a molecule.⁵¹ HOMO and LUMO analysis is also used to determine the charge transfer within a molecule. Considering the chemical hardness, if a molecule has a large HOMO–LUMO gap, it is a hard molecule or small HOMO–LUMO gap it is a soft molecule. One can also relate the stability of a molecule to hardness, which means that the molecule with the least HOMO–LUMO gap is more reactive. Soft systems are large and highly polarizable, while hard systems are relatively small and much less polarisable.

The analysis of the wavefunction indicates that the electron absorption corresponds to the transition from the ground state to the first excited state and is mainly described by one-electron excitation from the HOMO to LUMO.⁵² Fig. 11 and 12 show the shapes and energy levels of the HOMO and LUMO for complexes 1–6. The HOMO of complexes 1, 2 and 3 show that the electron density is mainly localized on the chelate rings while the LUMO of complex 1 shows that the electron density would be localized over the piperidine ring. The LUMO of complexes 2 and 3 show that the electron density would be localized over piperidine ring and chelate ring. The HOMO of complex 4 shows that electron density is localized on the dithio sulfur atoms and bonded mercury(II) ion whereas the LUMO of complex 4 shows that the electron density would be localized over the piperidine ring as well as the mercury(II) ion and bonded dithio sulfur atoms. The HOMO of compound 5 shows that the electron density is mainly localized on the disulfide sulfur atoms while the LUMO of compound 5 shows that the electron density is going to be localized over piperidine ring as well as the disulfide sulfur atoms. The HOMO of compound 6 shows that electron density is localized on the both dithio sulfur, methylene sulfur atoms and piperidine nitrogen atom, whereas LUMO of compound 6 shows that the electron density would be localized over both dithio sulfur, methylene sulfur atoms and piperidine ring. The HOMO–LUMO energy gaps for complexes 1–4 are 3.94946, 4.20471, 3.96279 and 4.55900 eV, respectively. It is pertinent to mention here that the energy gap is smaller in complexes 1 and 3 than in complexes 2 and 4. Therefore, complexes 1 and 3 can serve as better targets for photophysical studies and electronics applications when compared to complexes 2 and 4.^53,54 In addition, complexes 1 and 3 are softer and hence more reactive when compared to complexes 2 and 4. The HOMO–LUMO energy gaps for compounds 5 and 6 are 4.71982 and 4.79764 eV, respectively (Table 10). The electronic transition from the ground state to the first excited state for these molecules is due to the transfer of electrons from the HOMO to LUMO level being mainly the π⋯π* type.

Fig. 11 The frontier molecular orbitals of complexes 1, 2, 3 and 4.

Fig. 12 The frontier molecular orbitals of molecules 5 and 6.

Table 10 The calculated frontier molecular orbital energies (eV) obtained for compounds 1–6

Molecules	Energies (eV)
	HOMO	LUMO	Energy gap
1	−4.67573	−0.72627	3.94946
2	−5.09425	−0.88954	4.20471
3	−5.48364	−1.52085	3.96279
4	−5.66868	−1.10968	4.55900
5	−5.49752	−0.77770	4.71982
6	−5.48990	−0.69226	4.79764

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3. Conclusions

In conclusion, we report the syntheses, structural investigation and physicochemical properties (photoluminescence, thermal analyses and cyclic voltammetry) of Ag(I), Cu(II), Co(III) and Hg(II) complexes bearing the 4-methyl-piperidine-carbodithioate ligand. Additionally, Mn(II) assisted formation of a disulfide compound and the Ag(I) mediated synthesis of a carbothioylsulfanyl derivative have also been discovered. The solid state structures of the Ag, Cu, Co complexes and the disulfide and carbothioylsulfanyl derivatives revealed intermolecular C–H⋯S interactions including S⋯S interactions in the Cu complex and the disulfide derivative. Further, intramolecular C–H⋯π and the Hg⋯S secondary interactions were found in the Hg(II) complex. The photoluminescence properties of the Co(III) complex and the disulfide derivative make them potential materials to be useful in light emitting diodes (LED). When compared to the free ligand, a large blue shift emission in its Ag(I) and Hg(II) complexes was a consequence of the quenching behaviour and/or an effect of the Hg⋯S/Ag⋯S secondary interactions. The reversible redox behaviour of the Cu(II) complex makes it a potential probe in electrochemical materials and electrochemical solar cells. The thermal degradation of the metal complexes, as analysed by TG-DTA, showed strong evidence for the formation of their respective metal sulfides as the final thermally stable chemical entities.

4. Experimental

4.1 Materials and methods

4-Methyl-piperidine was purchased from Sigma Aldrich. Other chemicals were of reagent grade and used as purchased without further purification. All synthetic manipulations were carried out open to the ambient atmosphere and at room temperature. The solvents were distilled before use following standard procedures. The carbon, hydrogen, nitrogen and sulfur contents were determined on a CHN Model CE-440 Analyzer and on an Elementar Vario EL III Carlo Erba 1108. Infrared spectra were recorded in the 4000–400 cm⁻¹ region as KBr pellets on a Varian Excalibur 3100 FT-IR spectrophotometer. ¹H and ¹³C NMR spectra were recorded in CDCl₃ on a JEOL AL 300 FT NMR spectrometer using TMS as an internal reference. The fluorescent data were collected at room temperature on a Varian CARYECLIPSE spectrophotometer in CHCl₃ (1 and 4) and MeOH (2, 3, 5 and 6). Thermogravimetric analysis (TG-DTA) of complexes 1–2 was performed using a Perkin Elmer-STA 6000 thermal analyzer, TA Instrument and complexes 3–4 were performed using a thermal analyzer Model No-TGA/DSC1 STAR System Germany, in a nitrogen atmosphere at a heating rate of 10 °C min⁻¹.

4.2 Cyclic voltammetry

Electrochemical studies were performed on a Metrohm (Netherlands) Instrument (Autolab PGSTAT204) potentiostat/galvanostat attached with NOVA 1.11 software. A three electrode, one compartment cell with solid platinum electrode as the working electrode, platinum wire as the counter electrode and Ag/AgCl as the reference electrode was used. All electrochemical studies were carried out at room temperature under a nitrogen atmosphere.

4.3 X-ray crystallography

Structural measurements of compounds 1–6 were performed on a computer-controlled Oxford Gemini diffractometer equipped with a CrysAlis CCD software using a graphite mono-chromated Mo Kα (λ = 0.71073 Å) radiation source at 293 K. Multi-scan absorption correction was applied to the X-ray data collection for all the compounds. The structures were solved using direct methods (SHELXL-2013) and refined against all data by full matrix least-square on F² using anisotropic displacement parameters for all non-hydrogen atoms. All hydrogen atoms were included in the refinement at geometrically ideal positions and refined with the Riding model.⁵⁵ The MERCURY package and ORTEP-3 for Windows program were used for generating the molecular graphics.^56,57

4.4 Quantum chemical calculations

To understand the binding sites and to verify the composition of the complexes, density functional theory (DFT) calculations were performed. In these calculations, the B3LYP density functional theory method was used, which is a hybrid version of the DFT and Hartree–Fock (HF) methods.⁵⁸ In this method, the exchange energy from Becke's exchange functional was combined with the exact energy obtained from the Hartree–Fock theory.⁵⁹ The three parameters define the hybrid functional, specifying how much of the exact exchange is mixed in along with the component exchange and correlation functionals.⁶⁰ The 6-31G** basis set for all the atoms except metal ions was used. The LanL2DZ basis set with an effective core pseudo potential was used for the metal atoms.⁶¹ All the geometry optimizations and frequency calculations (to verify a genuine minimum energy structure) were performed using the Gaussian 09 program package.⁶² It worth pointing out here that the starting geometries were chosen based on the X-ray crystallographic structures with required modification for the computations if needed.

4.5 Synthesis

4.5.1 Synthesis of potassium 4-methyl-piperidine-carbodithioate {K⁺(4-mpipdtc)⁻}. Potassium 4-methyl-piperidine-carbodithioate was prepared by the dropwise addition of carbon disulfide (1.5 mL, 20 mmol) to an ethanolic solution of 4-methyl-piperidine (2.4 mL, 20 mmol) and potassium hydroxide (1.2 g, 20 mmol). The reaction mixture was stirred continuously for 2 h. The solid potassium 4-methyl-piperidine-carbodithioate {K⁺(4-mpipdtc)⁻} obtained was filtered, washed with ethanol and dried under reduced pressure (Scheme 2). Yield: 75%; mp 185 °C. Anal. calc. for C₇H₁₂NS₂K (213.40): C, 35.78; H, 5.66; N, 6.56; S, 30.05. Found: C, 35.95; H, 5.90; N, 6.30; S, 29.70%. IR (KBr, cm⁻¹): 2951 ν(C–H), 1585 ν(C–N), 1015 ν(C=S). ¹H NMR (δ, ppm): 0.95 (d, 3H, CH₃), 1.20–1.59 (m, 4H, H2), 1.70–1.75 (m, 1H, H3), 3.76 (t, 2H, H1-axial), 5.67 (d, 2H, H1-equatorial). ¹³C NMR (δ, ppm): 21.56 (CH₃), 30.89 (C2), 34.35 (C3), 51.29 (C1), 208.76 (C=S) (Scheme 4). UV-Vis. (CHCl₃, λ_max/nm, ε/M⁻¹ cm⁻¹): 265 (1.2 × 10⁴), 297 (1.2 × 10⁴).

4.5.2 Synthesis of [Ag₂(4-mpipdtc)₂(PPh3)₂] (1). AgNO₃ (0.169 g, 1 mmol) was added to a 10 mL methanol solution of {K⁺(4-mpipdtc)⁻} (0.213 g, 1 mmol) and stirred for 1 h at room temperature. A white precipitate was obtained. Triphenylphosphine (0.262 g, 1 mmol) solution in chloroform was added to the above suspension. A colourless clear solution was obtained, which was filtered and kept for crystallization. Colourless crystals suitable for X-ray analysis were obtained after the slow evaporation of the solvent over a period of 15 days (Scheme 2). Yield: 58%; mp 196 °C. Anal. calc. for C₅₀H₅₄Ag₂N₂P₂S₄ (1088.87): C, 55.14; H, 4.96; N, 2.57; S, 11.76. Found: C, 55.00; H, 5.20; N, 2.40; S, 11.52%. IR (KBr, cm⁻¹): 3058 ν(C–H, phenyl) (calc. 3177), 2989 ν(C–H) (calc. 2948), 1469 ν(C–N) (calc. 1462), 962 ν(C=S) (calc. 960), ν(Ag–S) (calc. 320). ¹H NMR (δ, ppm): 0.93 (d, 3H, CH₃), 1.32–1.57 (m, 4H, H2), 1.68 (m, 1H, H3), 3.07 (t, 2H, H1-axial), 5.30 (d, 2H, H1-equatorial), 7.34–7.49 (m, 15H, phenyl H). ¹³C NMR (δ, ppm): 21.27 (CH₃), 30.39 (C2), 33.87 (C3), 52.74 (C1), 126.21–134.09, 207.17 (C=S) (Scheme 4). UV-Vis. (CHCl₃, λ_max/nm, ε/M⁻¹ cm⁻¹): 263 (9 × 10³).

4.5.3 Synthesis of [Cu(4-mpipdtc)₂] (2). Cu(OAc)₂·H₂O (0.200 g, 1 mmol) was added to a solution of {K⁺(4-mpipdtc)⁻} (0.426 g, 1 mmol) in 20 mL of a methanol-chloroform mixture and stirred for 3 h at room temperature. A clear green solution was obtained, which was filtered and kept for crystallization. Prism shaped dark blue crystals suitable for X-ray analysis were obtained over a period of 10 days. Yield: 78%; mp 213 °C. Anal. calc. for C₁₄H₂₄CuN₂S₄ (412.13): C, 40.79; H, 5.86; N, 6.80; S, 31.04. Found: C, 40.25; H, 5.60; N, 6.75; S, 31.33%. IR (KBr, cm⁻¹): 2948 ν(C–H) (calc. 2995), 1504 ν(C–N) (calc. 1492), 965 ν(C=S) (calc. 966), 429 ν(Cu–S) (calc. 405). UV-Vis. (MeOH, λ_max/nm, ε/M⁻¹ cm⁻¹): 272 (8 × 10³), 439 (3 × 10²).

4.5.4 Synthesis of [Co(4-mpipdtc)₃]·CHCl₃ (3). CoCl₂ (0.130 g, 1 mmol) was added to a solution of {K⁺(4-mpipdtc)⁻} (0.639 g, 3 mmol) in 20 mL of a methanol-chloroform mixture and stirred for 3 h at room temperature. A clear green solution was obtained, which was filtered and kept for crystallization. Dark green crystals suitable for X-ray analysis were obtained over a period of 10 days (Scheme 2). Yield: 80%; mp 309 °C. Anal. calc. for C₂₂H₃₇Cl₃CoN₃S₆ (701.18): C, 37.68; H, 5.31; N, 5.60; S, 18.28. Found: C, 37.88; H, 5.15; N, 5.40; S, 18.53%. IR (KBr, cm⁻¹): 2949 ν(C–H) (calc. 2993), 1493 ν(C–N) (calc. 1493), 966 ν(C=S) (calc. 875), 428 ν(Co–S) (calc. 358). ¹H NMR (δ, ppm): 0.96 (d, 3H, CH₃), 1.23 (m, 4H, H2), 1.71 (m, 1H, H3), 2.89 (t, 2H, H1-axial), 4.66 (d, 2H, H1-equatorial). ¹³C NMR (δ, ppm): 21.32 (CH₃), 30.65 (C2), 33.13 (C3), 45.48 (C1), 203.55 (C=S) (Scheme 4). UV-Vis. (MeOH, λ_max/nm, ε/M⁻¹ cm⁻¹): 271 (3 × 10⁴), 328 (2 × 10⁴).

4.5.5 Synthesis of [PhHg(4-mpipdtc)] (4). A methanol-chloroform solution of PhHg(OAc) (0.337 g, 1 mmol) was added to a 20 mL of a methanol-chloroform mixture (1 : 1) containing {K⁺(4-mpipdtc)⁻} (0.213 g, 1 mmol) and stirred for 3 h at room temperature. A colourless clear solution was obtained, which was filtered and kept for crystallization. Rod shaped crystals suitable for X-ray analyses were obtained over a period of 7 days (Scheme 2). Yield: 75%; mp 248 °C. Anal. calc. for C₁₃H₁₇HgNS₂ (451.99): C, 34.54; H, 3.78; N, 3.10; S, 14.18. Found: C, 34.35; H, 3.60; N, 3.01; S, 14.33%. IR (KBr, cm⁻¹): 3062 ν(C–H phenyl) (calc. 3120), 2947 ν(C–H) (calc. 2999), 1484 ν(C–N) (calc. 1484), 968 ν(C=S) (calc. 960), 430 ν(Hg–S) (calc. 428). ¹H NMR (δ, ppm): 0.93 (d, 3H, CH₃), 1.32–1.57 (m, 4H, H2), 1.68 (m, 1H, H3), 3.10 (t, 2H, H1-axial), 4.95 (d, 2H, H1-equatorial), 7.34–7.49 (m, 5H, phenyl H). ¹³C NMR (δ, ppm): 21.19 (CH₃), 30.25 (C2), 33.79 (C3), 52.88 (C1), 128.23–154.46, 200.81 (C=S) (Scheme 4). UV-Vis. (CHCl₃, λ_max/nm, ε/M⁻¹ cm⁻¹): 295 (5 × 10³).

4.5.6 Synthesis of (4-mpipdtc)₂ (5). Mn(OAc)₂·4H₂O (0.245 g, 1 mmol) was added to a solution of {K⁺(4-mpipdtc)⁻} (0.426 g, 2 mmol) in 20 mL of methanol and stirred for 3 h at room temperature. A clear yellow solution was obtained, which was filtered and kept for crystallization. Rod-shaped crystals suitable for X-ray analysis were obtained over a period of 15 days (Scheme 3). Yield: 82%; mp 102 °C. Anal. calc. for C₁₄H₂₄N₂S₄ (348.59): C, 48.23; H, 5.31; N, 6.93; S, 36.78. Found: C, 48.43; H, 5.23; N, 6.76; S, 36.58%. IR (KBr, cm⁻¹): 2950 ν(C–H) (calc. 2966), 1481 ν(C–N) (calc. 1479), 959 ν(C=S) (calc. 902), 816 ν(C–S) (calc. 821), 556 ν(S–S) (calc. 564). ¹H NMR (δ, ppm): 0.99 (d, 3H, CH₃), 1.39–1.57 (m, 4H, H2), 1.79 (m, 1H, H3), 3.32 (t, 2H, H1-axial), 5.00 (d, 2H, H1-equatorial). ¹³C NMR (δ, ppm): 21.20 (CH₃), 30.85 (C2), 33.83 (C3), 47.43 (C1), 192.82 (CS₂) (Scheme 4). UV-Vis. (MeOH, λ_max/nm, ε/M⁻¹ cm⁻¹): 283 (4 × 10⁴).

4.5.7 Synthesis of CH₂(4-mpipdtc)₂ (6). AgNO₃ (0.169 g, 1 mmol) was added to a solution of {K⁺(4-mpipdtc)⁻} (0.426 g, 2 mmol) in 20 mL of methanol and stirred for 1 h at room temperature. A white precipitates was obtained, to which dichloromethane (10 mL) was added dropwise and stirred for 2 h, followed by filtration and kept for crystallization. Rod-shaped colourless crystals of 6 suitable for X-ray analysis were obtained over a period of 15 days (Scheme 3). Yield: 80%; mp 140 °C. Anal. calc. for C₁₅H₂₆N₂S₄ (362.62): C, 49.68; H, 7.22; N, 7.72; S, 35.36. Found: C, 49.42; H, 7.23; N, 7.81; S, 35.53%. IR (KBr, cm⁻¹): 2924 ν(C–H), 1434 ν(C–N), 998 ν(C=S), 714 ν(C–S). ¹H NMR (δ, ppm): 0.95 (d, 3H, CH₃), 1.24–1.43 (m, 4H, H2), 1.65 (m, 1H, H3), 2.95 (t, 2H, H1-axial), 4.09 (d, 2H, H1-equatorial), 5.37 (s, 2H, SCH₂), ¹³C NMR (δ, ppm): 21.19 (CH₃), 30.83 (C2), 34.00 (C3), 52.04 (C1), 54.52 (SCH₂S) 193.62 (CS₂), UV-Vis. (MeOH, λ_max/nm, ε/M⁻¹ cm⁻¹): 260 (1 × 10⁴).

Acknowledgements

Paras Nath thanks the UGC, India for the award of JRF and Dr M. K. Bharty is thankful to the Science and Engineering Research Board, India for the award of a Project (No. SB/EMEQ-150/2014; Diary No. SERB/F/372/2015-16).

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Footnote

† Electronic supplementary information (ESI) available: Crystallographic data and additional figures. CCDC 1052191, 1412092, 1412089, 1407803, 1412091 and 1402772. For ESI and crystallographic data in CIF or other electronic format see DOI: 10.1039/c6ra15186h

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