s2fieldsofstudy:
Chemistry
Environmental Science

text:
Determination of Triiodide Ion Concentration Using UV-Visible Spectrophotometry†

The concentration of iodine species can be measured using several different analysis methods. One of the most wellknown methods involves iodine titration with a thiosulfate solution. Total iodine can be analyzed using spectrometry, such as inductively coupled plasma-atomic emission spectrometry (ICP-AES), atomic absorption spectrometry (AAS) and inductively coupled plasma-mass spectrometry (ICP-MS). As elemental iodine has various valance states, various stable iodine species are present in aqueous solutions. Stable iodine species can include iodide ions, triiodide ions, iodine and iodate ions. For the quantitative analysis of these species, ion chromatography (IC) has been used in analytical laboratories. In addition, ultraviolet-visible spectrophotometry and Raman spectrometry have been used for the speciation and quantitative analyses of iodine species. Recently, after the severe accident at the Fukushima nuclear power plant in Japan, the chemical behaviour of radioactive iodine received increased attention as a result of interest in evaluating the volatility of iodine species, which depends on the chemical conditions of the exposed environments. The use of γ-ray spectroscopy to measure radioactivity is the best method for determining radioactive iodine concentrations. However, to use radioactive iodine, a licensed Determination of Triiodide Ion Concentration Using UV-Visible Spectrophotometry†

INTRODUCTION
The concentration of iodine species can be measured using several different analysis methods [1][2][3] . One of the most wellknown methods involves iodine titration with a thiosulfate solution 4,5 . Total iodine can be analyzed using spectrometry, such as inductively coupled plasma-atomic emission spectrometry (ICP-AES), atomic absorption spectrometry (AAS) and inductively coupled plasma-mass spectrometry (ICP-MS) 3,6,7 . As elemental iodine has various valance states, various stable iodine species are present in aqueous solutions. Stable iodine species can include iodide ions, triiodide ions, iodine and iodate ions. For the quantitative analysis of these species, ion chromatography (IC) has been used in analytical laboratories 3 . In addition, ultraviolet-visible spectrophotometry and Raman spectrometry have been used for the speciation and quantitative analyses of iodine species 8,9 . Recently, after the severe accident at the Fukushima nuclear power plant in Japan, the chemical behaviour of radioactive iodine received increased attention as a result of interest in evaluating the volatility of iodine species, which depends on the chemical conditions of the exposed environments. The use of γ-ray spectroscopy to measure radioactivity is the best method for determining radioactive iodine concentrations. However, to use radioactive iodine, a licensed The concentration of triiodide ions was determined using UV-visible spectrophotometry. The maximum absorption λmax for the analysis of triiodide ions were observed at 288 and 351 nm, with molar absorptivity values of 2.65 × 10 4 and 1.72 × 10 4 L cm -1 mol -1 , respectively. The absorbance of triiodide ions had little effects on the solution pH in the range of 2.0 to 6.5. However, the absorbance strongly depended on the iodide ion concentration. Consequently, we confirmed that triiodide ion concentrations in the range of 0.005-0.1 mM could be determined using UV-visible spectrophotometry.
Keywords: Triiodide ion, UV-visible spectrophotometry, Iodide ion, Solution pH. †Presented at 5 th International Symposium on Application of Chemical and Analytical Technologies in Nuclear Industries (Nu-ACT 2013), Daejeon, Korea laboratory equipped with numerous radioactivity safety devices is generally required. Therefore, many researchers have performed their experiments using non-radioactive iodine under γ-ray radiation instead of radioactive iodine 10 . In this case, the analysis of iodine species should be performed by chemical methods. Similar to radioactive iodide ions, which are stably soluble in aqueous solution, non-radioactive iodide ions can also be oxidized into volatile iodine under γ-ray radiation.Under the selected experimental conditions, the iodide ion concentration is relatively high, allowing the oxidized species, molecular iodine,to react with iodide ions and form triiodide ions, which are the more stable soluble species. Therefore, it is important to measure the concentration of triiodide ions to evaluate the volatility of iodine species.
In the present work, we established a simple method to analyze the triiodide ion concentration using UV-visible spectrophotometry. First, we identified distinct absorption wavelengths and determined the maximum absorbance for the analysis of triiodide ions. Second, the absorbance of triiodide ions was confirmed to be independent of the solution pH. Third, we evaluated the effect of iodide ions on the formation of triiodide ions by analyzing the UV-visible absorption spectra. Lastly, the detection range of triiodide ions was determined for the UV-visible spectrophotometric method. An iodide solution (100 mM NaI) was obtained by the dissolution of sodium iodide and an iodine (I 2 ) solution was obtained by the dissolution of solid iodine in distilled water. A triiodide solution was prepared by adding the iodine solution to the sodium iodide solution and mixing. Thus, the triiodide ion concentration was mainly determined by the concentration of iodine (I 2 ) in the solutions. Iodide ions were added in excess to stable form the triiodide ion in the solutions. A perchloric acid solution (0.1 M HClO 4 ) was used to control the solution pH. All chemicals used were of analytical grade and all experiments were carried out at ambient temperature. A WPA Lightwave II UV-visible spectrophotometer was used (Biochrom Ltd).

RESULTS AND DISCUSSION
UV-visible absorption spectrum of triiodide ion: Fig. 1a shows the UV-visible absorption spectra of 0-0.15 mM triiodide ions. The solution pH was maintained at 6.5. The large absorption partially observed between 200 and 250 nm is attributed to the addition of iodide ions to the solutions in excess. The triiodide ions exhibited strong light absorption (ε >10 4 L cm -1 mol -1 ) at 288 and 351 nm. The maximum absorbance values at these two wavelengths increased with the triiodide ion concentration. The molar absorptivity values at 288 and 351 nm were 2.65 × 10 4 and 1.72 × 10 4 L cm -1 mol -1 , respectively.  Using the absorbance values presented in Fig. 1a, we obtained the relationship between the maximum absorbance value at 288 nm and triiodide ion concentration, as shown in Fig. 1b. A linear relationship was observed in the concentration range of 0.005-0.10 mM I3 -, corresponding to the range of triiodide ion concentrations that could be determined using UV-visible spectrophotometry.
Effect of solution pH on the UV-visible absorption spectrum of triiodide ions: Fig. 2 shows the UV-visible absorption spectra of 0.1 mM triiodide ion solutions at three different pH values: 2.0, 4.0 and 6.5. Above a pH of 6.5, it is well-known that the triiodide ion isdisproportionated into iodide and iodate ions according to reactions (1,2). Therefore, in our experiments, we limited the solution pH below 6.5. Interference of iodide ions in the UV-visible absorption spectrum of triiodide ions: Molecular iodine is stably soluble in its triiodide ion form in iodide solution,as shown in reaction (3). Although the equilibrium constant of the reaction is high, the concentration of triiodide ions dissolved in the iodide solution may depend on the concentration of iodide ions.
We obtained triiodide solutions by adding iodine (I 2 ) to the sodium iodide solutions in four different concentrations and mixing. As the substantial concentration of iodine is 0.01 mM, the concentration value of triiodide ion is approaching to 0.01 mM with an increase in the iodide ion concentration.  In Fig. 3b, the maximum absorbance λ max of triiodide ion at 288 nm was plotted against the concentration of iodide ion. As the concentration of iodide ion increased, the absorbance value increased with approaching a maximum value. Thus, the absorbance value increased rapidly when the initial concentration of iodide ion was below 10 mM. However, above the concentration 10 mM, the absorbance value increased slowly and approached a certain value. The result indicates that the concentration of iodide ion in triiodide samples should be constant or relatively high, when the UV-visible spectroscopic analysis is used for a triiodide concentration determination method.
Triiodide ion determination range using UV-visible spectrophotometry: Considering the equilibrium between triiodide and iodide ions, we recorded UV-visible absorption spectra of various concentrations of triiodide ions with a constant iodide concentration of 5 mM. Fig 4 shows the relationship between the maximum absorbance λ max at 288 nm and the initial concentration of molecular iodine added to the iodide solution. It was observed that the absorbance of triiodide ion increased linearly with the initial concentration of iodine without any interference of solution pH. The result indicates that iodine concentrations of up to 0.1 mM are completely converted into triiodide ions in a 5 mM iodide solution. Based  Fig. 4, we also found that the detection limit of triiodide ions was 0.005 mM. Therefore, it was concluded that triiodide ion concentrations of 0.005 to 0.1 mM could be determined using UV-visible spectrophotometry as long as the solution pH was 2.0-6.5.

Conclusion
The triiodide ion was observed to have two distinct absorptions at 288 and 351 nm with molar absorptivity values of 2.65 × 10 4 and 1.72 × 10 4 L cm -1 mol -1 , respectively. The absorbance of triiodide ions was independent of the solution pH from 2.0-6.5. However, the absorbance was strongly affected by the iodide ion up to 5.0 mM. Consequently, it was confirmed that triiodide ion concentrationsof 0.005-0.1 mM could be determined using UV-visible spectrophotometry.


extfieldsofstudy: Chemistry Medicine

s2fieldsofstudy:
Chemistry
Agricultural And Food Sciences

text:
Flavonoids from Machilus japonica Stems and Their Inhibitory Effects on LDL Oxidation

Stems of Machilus japonica were extracted with 80% aqueous methanol (MeOH) and the concentrated extract was successively extracted with ethyl acetate (EtOAc), normal butanol (n-BuOH), and water. Six flavonoids were isolated from the EtOAc fraction: (+)-taxifolin, afzelin, (−)-epicatechin, 5,3'-di-O-methyl-(−)-epicatechin, 5,7,3'-tri-O-methyl-(−)-epicatechin, and 5,7-di-O-methyl-3',4'-methylenedioxyflavan-3-ol. The chemical structures were identified using spectroscopic data including NMR, mass spectrometry and infrared spectroscopy. This is the first report of isolation of these six compounds from M. japonica. The compounds were evaluated for their diphenyl picryl hydrazinyl scavenging activity and inhibitory effects on low-density lipoprotein oxidation. Compounds 1 and 3–6 exhibited DPPH antioxidant activity equivalent with that of ascorbic acid, with half maximal inhibitory concentration (IC50) values of 0.16, 0.21, 0.17, 0.15 and 0.07 mM, respectively. The activity of compound 1 was similar to the positive control butylated hydroxytoluene, which had an IC50 value of 1.9 µM, while compounds 3 and 5 showed little activity. Compounds 1, 3, and 5 exhibited LDL antioxidant activity with IC50 values of 2.8, 7.1, and 4.6 µM, respectively.

Introduction
The family Lauraceae includes about 32 genera and 2500 species that are distributed in tropical and subtropical regions, especially in Southeast Asia [1]. In Korea, members of the Lauraceae family are found in the southern regions and in Jeju Island. In this study, we describe the chemical constituents and activity of Machilus japonica extracts. M. japonica is an evergreen tree with yellow-green flowers and oval leaves. A few studies on the chemical composition of this plant have resulted in the identification of lignans [2] and terpene compounds [3]. In addition, the biological activities of M. japonica extracts such as insecticidal activity [4], elastase inhibition activity [5], and inhibition of matrix metalloproteinase-9 activity [6] have been investigated; however, all of these activities have been identified from extracts of M. japonica leaves. Taking a different approach, the aim of this study was to explore the active components of the stem of M. japonica and to screen their biological activities.
A total of six flavonoids were isolated and identified from the stem of M. japonica. Because many flavonoids have been reported to exhibit antioxidant activities [7], the isolated flavonoids were also evaluated for DPPH radical scavenging activity. low-density lipoprotein (LDL) is susceptible to oxidative damage, and oxidized LDL (oxLDL) plays a key role in the development of atherosclerotic lesions [8]. oxLDL in vessel walls is subjected to rapid uptake by scavenger receptors on monocyte derived macrophages, leading to the formation of foam cells that accumulate cholesterol [9]. Thus, we also evaluated the isolated compounds for their ability to inhibit LDL oxidation. Our results suggest that extracts of M. japonica stems and the specific flavonoids found in these extracts may prove useful for preventing or treating hypercholesterolemia and atherosclerosis.
Compound 5 was obtained as colorless needles and a molecular ion peak [M] + was observed at m/z 332 in the EI/MS spectrum. The IR spectrum showed absorbance bands for hydroxyl (3412 cm −1 ) and aromatic (1616 cm −1 ) groups. The 1 H-NMR and 13 C-NMR spectra of compounds 4 and 5 were analogous, showing (−)-epicatechin skeleton signals. Due to the m/z of 332 and observation of three methoxy groups, compound 5 was presumed to be (−)-epicatechin-tri-O-methylate. To confirm the position of the methoxy groups, a gHMBC experiment was conducted. Judging by the correlation between the three methoxy proton signals at δH 3.69 (6H, s) and 3.66 (3H, s) with three oxygenated olefin quaternary carbon signals at δC 160.0 (C-7), δC 159.6 (C-5), and δC 148.3 (C-3') in the gHMBC spectrum, compound 5 was identified as 5,7,3'-tri-O-methyl-(−)-epicatechin, which was confirmed by comparison with spectroscopic data in the literature [15].

Evaluation of Radical Scavenging Activity and Inhibitory Effect on LDL Oxidation
Compounds 1 and 3-6 exhibited DPPH radical scavenging with IC50 values of 0.16, 0.21, 0.17, 0.15, and 0.07 mM, respectively, which were equivalent with that of ascorbic acid (0.18 mM) ( Table 1). The lack of scavenging activity by compound 2 was speculated to be due to the glucosyl moiety on the C ring [17]. Compound 6 exhibited the highest antioxidant activity, which we attributed to the dioxymethylene moiety on the B ring. Highly reactive molecules called free radicals can cause tissue damage by reacting with polyunsaturated fatty acids in cellular membranes, nucleotides in DNA, and critical sulfhydryl bonds in proteins. In addition to cellular damage, cataract formation, photodermatoses, aging, and inflammatory diseases such as arthritis are associated with free radicals [18]. Therefore, the search for compounds from natural sources that can protect against free radicals by endogenous and exogenous antioxidants is of special significance for human health.
Antioxidants act by donating hydrogen atoms to lipid radicals. Radicals obtained from antioxidants with molecular structures such as phenols are stable species that can halt the oxidation chain reaction [19]. To determine whether these compounds might be effective in the development of hypercholesterolemic or antiatherogenic agents, their potential for inhibiting LDL oxidation was evaluated.
Compounds 1, 3, and 5 demonstrated LDL antioxidant activity with IC50 values of 2.8, 7.1, and 4.6 µM, respectively (Table 2). Compound 1 was similar to the positive control, BHT, which had an IC50 value of 1.9 µM, while compounds 3 and 5 also showed significant activity. Compound 4 exhibited a low level of LDL antioxidant activity with an IC50 value of 79.1 µM. Compounds 2 and 6 were not effective as LDL antioxidants, which was attributed to the glucosyl moiety at C-3 for compound 2 and steric hindrance of the hydroxyl groups at C-3' and C-4' by dioxymethylene moieties in compound 6. Table 2. Inhibitory effects of flavonoids isolated from the stems of Machilus japonica on low-density lipoprotein (LDL) oxidation. a The IC50 value of each compound was defined as the concentration (μM) that resulted in 50% inhibition of LDL oxidation. The results are averages of three independent experiments, and the data are expressed as the mean ± SD. LDL oxidation is regarded as a key step in the formation of atherosclerotic lesions [20,21]. Experimental evidence demonstrating an association between oxLDL cholesterol and both the presence of atherosclerotic lesions and progression of carotid artery atherosclerosis support this hypothesis [22,23]. Vitamin E, one of the most popular natural antioxidants, inhibits atherogenesis by inhibiting LDL oxidation (IC50: 2.4 μM) [24]. Several other effective natural dietary antioxidants comprising phenolic compounds and carotenoids have been identified as well, in addition to vitamins and enzymes [25]. The inhibitory activity of compounds 1, 3, and 5 was very similar to that of the well-known antioxidant BHT (IC50: 1.9 μM), and showed significant antioxidant capacity relative to several other naturally-occurring antioxidants such as nectandrin B (IC50: 4.1 μM) from tabu (Machilus thunbergii) [26] and (+)-lariciresinol (IC50: 11.9 μM) from Rousa dogwood (Cornus kousa) [27]. Therefore, the flavonoids isolated from Machilus japonica may be a good natural source of antiatherogenic agents.

Plant Materials
Dried stems of Machilus japonica were supplied by GFC Co., Ltd., Suwon, Korea in January 2010, and were identified by Dae-Keun Kim, College of Pharmacy, Woosuk University, Jeonju, Korea. A voucher specimen (KHU2010-0103) has been reserved at the Laboratory of Natural Product Chemistry, Kyung Hee University, Yongin, Korea.

General Experimental Procedures
Melting points were determined using a Fisher-John's Melting Point apparatus (Fisher Scientific, Miami, FL, USA) with a microscope and the values obtained were uncorrected. Optical rotations were measured using a JASCO P-1010 digital polarimeter (Tokyo, Japan). NMR spectra were recorded on a 400 MHz FT-NMR spectrometer (Varian Inova AS 400, Palo Alto, CA, USA). IR spectra were obtained from a Perkin Elmer Spectrum One FT-IR spectrometer (Buckinghamshire, England). FAB-MS data were recorded on a JEOL JMS-700 (Tokyo, Japan), EI-MS on a JEOL JMSAX 505-WA (Tokyo, Japan), and ESI-MS on a Finnigan LCQ Advantage spectrometer (Thermo Scientific, Waltham, MD, USA). The UV lamp used was a Spectroline Model ENF-240 C/F (Spectronics Corporation, Westbury, NY, USA). Kiesel gel 60 silica gel resin was used for column chromatography (c.c.) (Merck, Darmstadt, Germany) and the ODS was a LiChroprep RP-18 (Merck). TLC analysis was carried out using Kiesel gel 60 F254 and RP-18 F254S (Merck). Deuterated solvents were purchased from Merck Co. Ltd. and Sigma Aldrich Co. Ltd. (St. Louis, MO, USA).

DPPH Radical Scavenging Activity
The DPPH radical scavenging activity assay was based on the capacity of a substance to scavenge stable DPPH radicals. Briefly, reaction mixtures containing test samples (100 µL) and an ethanolic DPPH solution (100 µL of 0.06 mM) were placed in 96-well microplates and incubated at 37 °C for 30 min. Absorbance values were measured at 517 nm.

LDL Isolation and Oxidation Assay
Plasma was obtained from fasted healthy normalipidemic volunteers. LDL isolation and TBARS assays were performed as previously described with slight modification [28]. Briefly, an LDL solution (250 μL, 50-100 μg protein) in 10 mM PBS (pH 7.4) was supplemented with 10 μM CuSO4. Oxidation was performed in screw-capped 5-mL glass vials at 37 °C in the presence of either compound 1, 3, or 5. After incubation for 4 h the reaction was terminated by the addition of 1 mL 20% TCA. Following precipitation, 1 mL 0.67% TBA in 0.05 N NaOH was added and the mixture was vortexed, after which the final mixture was heated for 5 min at 95 °C, cooled on ice, and centrifuged for 2 min at 1000× g. The optical density of MDA generated in the assay was measured at 532 nm. Calibration was performed using an MDA standard prepared from tetramethoxypropane.


s2fieldsofstudy: Chemistry Biology

text:
Modular o-Quinone Catalyst System for Dehydrogenation of Tetrahydroquinolines under Ambient Conditions

Quinolines are common pharmacophores present in numerous FDA-approved pharmaceuticals and other bioactive compounds. Here, we report the design and development of new o-quinone-based catalysts for the oxidative dehydrogenation of tetrahydroquinolines to afford quinolines. Use of a Co(salophen) cocatalyst allows the reaction to proceed efficiently with ambient air at room temperature. The utility of the catalytic method is demonstrated in the preparation of a number of medicinally relevant quinolines.

C opper amine oxidases contain a tyrosine-derived oquinone in their active site that mediates aerobic oxidation of primary amines to aldehydes (e.g., topaquinone, Scheme 1A). 1 Biomimetic o-quinones such as Q1 and Q2 (Scheme 1A; Q2 red is proposed to form an o-quinone in situ) have been shown to be effective synthetic catalysts for aerobic dehydrogenation of primary amines, typically affording homocoupled imines. 2 Both topaquinone and the biomimetic quinone catalysts mediate amine oxidation via a "transamination" pathway, initiated by formation of an imine adduct of the substrate with the quinone. This mechanism accounts for the highly selective oxidation of primary over secondary and tertiary amines. We recently reported that 1,10-phenanthroline-5,6-dione (phd, Scheme 1A) promotes amine oxidation by a non-biomimetic "addition−elimination" pathway involving a hemiaminal intermediate (Scheme 1B). 3 This novel mechanism enabled the substrate scope to be expanded to include secondary amines. Aerobic dehydrogenation of a number of different nitrogen heterocycles was achieved by using phd in combination with ZnI 2 and pyridinium p-toluenesulfonate (PPTS) as a cocatalyst (Scheme 1C).
This phd/ZnI 2 catalyst system demonstrated the feasibility of aerobic secondary amine dehydrogenation, but reactions often required up to 48 h to reach completion and certain product classes were not accessible. For example, quinolines are an important class of heterocycles, but even the parent tetrahydroquinoline underwent dehydrogenation to quinoline in only 18% yield (Scheme 1C). Here, we describe an octahedral [Ru(phd) 3 ] 2+ catalyst that shows considerably higher activity for amine oxidation, including successful aerobic dehydrogenation of diverse tetrahydroquinolines at room temperature with ambient air as the source of O 2 . 4 This work highlights the modular nature of the phd o-quinone catalyst that makes it readily amenable to optimization and adaptation to different applications. Replacement of iodide with Co-(salophen) (salophen = N,N′-bis(salicylidene)-1,2-phenylenediamine) as a redox cocatalyst contributes significantly to the efficiency of the reactions.
In our initial studies, we compared the previously optimized phd/ZnI 2 catalyst system with simple octahedral [Fe(phd) 3 ] 2+ and [Ru(phd) 3 ] 2+ complexes in the oxidation of tetrahydroquinoline to quinoline ( Figure 1). The time course traces ( Figure 1) show the low activity and conversion of the previously reported phd/ZnI 2 catalyst (red trace); the catalyst loses activity ∼6−7 h into the reaction after reaching ≤20% conversion to the quinoline product. The Fe and Ru complexes (2.5 mol %) were also tested (green and blue traces, respectively). The use of Bu 4 NI (1 mol %) as a cocatalyst reflected previous observations showing that the I − /I 3 − redox couple promotes aerobic oxidation of the reduced, hydroquinone form of the phd catalyst. 5 [Fe(phd) 3 ] 2+ showed a similar initial rate to the ZnI 2 catalyst, but it exhibited somewhat improved stability. In contrast, [Ru(phd) 3 ] 2+ exhibited a significant increase in activity and a 93% yield of quinoline was obtained after 24 h. On the basis of this result, we characterized [Ru(phd) 3 ](ClO 4 ) 2 via X-ray crystallography ( Figure 2). 6,7 This [Ru(phd) 3 ] 2+ /Bu 4 NI catalyst was tested with a series of challenging N-heterocyclic substrates that had required 48 h to reach completion with the phd/ZnI 2 catalyst (Scheme 2). Improved yields and significantly decreased reaction times were observed in each case, with the most dramatic improvement observed in the dehydrogenation of tetrahydroquinoline.
Iodide was previously shown to mediate aerobic oxidation of the reduced hydroquinone form of the catalyst, and a catalytic sequence for the present dehydrogenation reactions is depicted in Scheme 3, where Co-Cat red/ox = 3I − /I 3 − . (It is not known whether dehydrogenation of the intermediate dihydroquinoline involves the catalyst.) We speculated that alternative cocatalysts could lead to even better catalytic reactivity. Backvall and others have highlighted the role of cocatalysts for aerobic oxidation of benzoquinone in multicomponent catalytic reactions, 8,9 and molecular catecholase mimics have been identified for aerobic oxidation of hydroquinones. 10 Drawing on these precedents, we tested a number of possible cocatalysts as replacements for Bu 4 NI, including Cu(pc), Fe(pc), Co(salophen), and Co(salpr) (pc = phthalocyanine; salpr = bis(salicylideneiminato-3propyl)methylamine). 11 Co(salophen) proved to be particularly effective, enabling full conversion within 3 h (Figure 3). 12 Subsequent studies showed that Co(salophen) enabled the reactions to proceed efficiently under ambient conditions (at room temperature with ambient air as the oxidant). The [Ru(phd) 3 ] 2+ catalyst structurally resembles Ru-polypyridyl complexes commonly used as photoactive catalysts, but control experiments show that the reactions exhibit identical behavior in the presence and absence of light. 13 In addition, no reaction was observed in the absence of [Ru(phd) 3 ] 2+ , suggesting that Co(salophen) is not a competent dehydrogenation catalyst under these conditions.
When probing the reactivity of polycyclic substrate 20 (Scheme 4), both dehydrogenation and benzylic oxygenation occurred to afford product 21 in 68% isolated yield. This reaction provides concise access to the indeno[2, 1-c]quinoline substructure present in numerous biologically active compounds, 17 including antiprotozoal agent 22 18 and phase II topoisomerase inhibitor  In conclusion, these results demonstrate the utility of [Ru(phd) 3 ] 2+ as a novel o-quinone catalyst for dehydrogenation of N-heterocycles. The results show that the substitutionally inert Ru 2+ ion is more effective than Zn 2+ in activating phd toward secondary amine dehydrogenation. Replacement of iodide with Co(salophen) as a redox cocatalyst to promote aerobic oxidation of the hydroquinone catalyst leads to substantial improvement in catalyst activity and enables the reactions to proceed under ambient conditions. The modular nature of the catalyst system described here has important implications for future studies targeting other aerobic quinonemediated oxidation reactions.

* S Supporting Information
Full experimental procedures and characterization data for all products, and additional screening data. This material is available free of charge via the Internet at http://pubs.acs.org.

Notes
The authors declare no competing financial interest.


s2fieldsofstudy:
Chemistry
Materials Science

text:
Theoretical study of the reaction mechanism Ba, Ti, O in the early growth of BTO thin films

s: We use the DMol program based on density functional theory(DFT) to clarify the processes of microcosmic reactions and the particle states in the early growth of BaTiO3 (BTO) thin films. This research is important for optimal preparation and structure control of BTO thin films. After designing and optimizing some possible intermediate states, we find that the molecules may serve as nucleation sites for BTO growth. This occurs by combining TiO2 with BaO molecules to form BaTiO3 with the G conformation, which is similar to a quarter of a BTO unit cell. By virtue of understanding these mechanisms, perovskite-structured BTO thin films are formed.

Introduction
A revolution is occurring in science and technology based on the recently developed ability to measure, manipulate, and organize matter on ferroelectric thin films. These thin films are widely exploited for microwave phase shifters and dynamic random access memory due to their advantageous physical properties (piezoelectricity, ferroelectricity, pyroelectricity, and photoelectricity-the latter is used in nonlinear optics). As a typical representative of the perovskite structure thin films, BaTiO 3 thin films are used not only to make microelectronic devices, but also to develop ferroelectric/superconductor devices by combining them with superconducting thin films [1][2][3]. As a result, progress in understanding BTO thin films will have a very far-reaching impact.
Currently, BaTiO 3 (BTO) and SrTiO 3 (STO) thin films are prepared primarily by Laser Molecular Beam Epitaxy (LMBE), not only because it allows in situ monitoring, but also because the chemical composition of thin films can be controlled. Using LMBE to prepare BTO and STO thin films, Li, et al [4][5][6] demonstrated the feasibility of unit cell migration with surface self-diffusion by in situ reflective high energy electron diffraction and atomic force microscopy. The experiments revealed "unit cell migration", which formed the basis for a growth model for films. In addition, they measured the activation energy of unit cell migration in BTO and STO thin films to be 0.29 ± 0.01 eV [4,5].
However, these experimental investigations are inadequate to describe physical chemistry data such as interface structure, adsorption, diffusivity, and particle trajectories in the growth of thin films. Additionally, experimental information is lacking regarding the formation of the BTO perovskite structure unit cell and the physicochemical processes involving Ba, Ti, and O atoms. Thus, developing a theoretical picture is very important for advancing research in this area. We had investigated possible reaction mechanisms in initial growth process of SrTiO 3 (STO) by a first-principles calculation in previous research [7]. To this end, we present a first-principles calculation using density functional theory (DFT) [8,9] to investigate both the possible combination reactions and the initial process of migration growth of BTO to the cell stage.
Regarding unit cell formation with Ba, Ti, and O microcosmic particles formed by flash evaporation in the unit cell migration growth of BTO thin films, we hypothesize two cases. The first is that Ba, Ti, and O atoms form intermediates (O 2 , BaO, BaO 2 , TiO, TiO 2 ) by colliding, after which the unit cell is formed. The second is the formation of a perovskite structure unit cell by direct deposition of Ba, Ti, and O atoms on the surface of the substrate. This paper is organized as follows, a brief introduction about the calculation method shows in section 2, the main results and discussion of the paper are described in section 3 and the end of the paper is conclusion.

Methodology
We use the local orbital density functional method DMol 3 program [10,11], which evolved from a gas phase molecular method to a method that allows for state-of-the-art calculations on non-conducting and metallic solids [12,13]. We used the PW91 function of the Generalized Gradient Approximation [14,15] and convergence conditions with a large basis set and high accuracy to optimize each structure of this system. During the calculations, the wave function of the valence electrons is expanded using the Double Numeric Polarization basis set in DFT, which is similar to 6-311++G** of Gaussian theory, but more precise. Its selfconsistent convergence precision is set above 10 -6 a.u. The energy convergence accuracy reaches 10 -5 a.u., and the convergence accuracy of grads and displacements surpassed 10 -4 a.u./nm and 5 × 10 -4 nm, respectively. In the DMol 3 program, all the electrons of atoms with atomic number less than 21 are involved in the calculation. Thus, all the electrons of O atoms and the valence electrons of Ba and Ti atoms are treated explicitly, but the core electrons of Ba and Ti atoms are replaced with effective core potential approximations [15]. Besides, we chose the most accurate transient state (TS) search methods to obtain the TS structure, namely the complete linear synchronous transit and the complete quadratic synchronous transit [17,18]. The activation energy of reactions can then be obtained after analyzing the energy difference between the intermediate state and the TS [19,20]. Incidentally, the TS can be confirmed easily if it has a higher energy than that of the reactants or intermediates and if it exists only at an imaginary frequency [21][22][23].

Results and Discussion
Generally, in an atmosphere containing Ba, Ti, and O atoms, Ba and Ti are easily oxidized by O atoms of higher activity, so that O 2 , BaO, BaO 2 , TiO, TiO 2 , and BaTiO 3 molecules are formed preferentially prior to formation of the unit cell. Based on our analysis, we present some reaction models and their initial conditions, assuming a vacuum and non-equilibrium states with high-density reactants.  Figure 1.

The Formation of Molecular BaO
The same analysis may be applied to Ba atoms, which are oxidized and form BaO molecules due to interactions with O atoms or O 2 molecules. Here, we likewise use two models, model IIA (IIB) where the Ba and O atoms (O 2 molecule) interact to form BaO. The reaction process is quite simple within model IIA-our calculation indicates that there is no energy barrier in this process. Using the energies given in Table 1, we find that the stabilization energy of the BaO molecule is 572.03 kJ/mol, which is consistent with values reported in the literature; i.e. E s ≥ 562.8 ± 14.6 kJ/mol [24]. In addition, the bond length of BaO is 0.2034 ± 0.0053 nm [24].
Within model IIB, two Ba atoms form coordination bonds with the lone-pair electrons of an O 2 molecule, instantly (on the time scale of interest) transforming into an M2 compound. The formation of the Ba-O bond weakens the O-O bond, and it is gradually broken resulting in the decomposition of M2 toward the BaO molecule via the transient TS2 state. By analyzed the transient TS2 state, we find that the stabilization energy of M2 is 424.74 kJ/mol and that the activation energy of the reaction is 55.01 kJ/mol. These results imply that the reaction always favors the formation of BaO molecules.

The Formation of Molecular BTO
After BaO and TiO 2 molecules are formed, they may interact to form BaTiO 3 . BaO, TiO 2 , and BaTiO 3 molecules, providing precursors for forming the BTO unit cell.
To be sure of the geometric conformations of BaTiO 3 , we built seven different possible structures according to the principle of molecular symmetry. As shown in Fig 5, after optimization of the seven structures, we find that the A, B, and C conformations are similar to the F conformation and that the D and E conformations are similar to the G conformation. Furthermore, the system energies of the F and G conformations are the lowest, at -813585.88 kJ/mol and -813587.18 kJ/mol, respectively, indicating that these conformations are more stable.
To study the components and structures of BTO thin film crystals, we create a 2×2 BTO super-cell with four BTO molecules, corresponding to BaTiO 3 molecules. The total energy of the super-cell is -3257201.00 kJ/mol, and the average energy of the BTO molecule is -814300.25 kJ/mol, making it 714.37 and 713.07 kJ/mol less than the F and G conformations, respectively. In addition, we note that the most stable conformation, G, is similar to a quarter-corner of the BTO unit cell, so that a BTO molecule with the G conformation, formed by BaO and TiO 2 molecules, will be shaped into the BTO unit cell with a perovskite structure. Put another way, the G conformation, which is a BTO molecule, may be regarded as one component of a BTO unit cell.
To form a BTO molecule (see Figs. 6 and 7), the F conformation is first formed by collision between BaO and TiO 2 molecules, which is a barrierless reaction. Immediately thereafter, the Ba atoms moves toward the center of the three O atoms until the G conformation is formed because of the attractive forces of the lone-pair electrons of the three O atoms. This process involves a barrier. Upon optimization and frequency analysis, we find a new transient state TS3. The characteristic parameters of TS3 are given in Table 2, and indicate that the reaction favors the G conformation and that it occurs easily due to the low activation energy (10.55 kJ/mol).
Our calculation and optimization indicate that there is little probability that Ba, Ti, and O atoms form a BTO unit cell directly but that some intermediates (such as BaO, TiO 2 , and BTO molecules) need to be formed beforehand. The activation energies for the formation processes of BaO, TiO 2 , and BTO are 55.01 kJ/mol, 11.81 kJ/mol, and 10.55 kJ/mol, respectively, with the lower value indicating that the reaction occurs more easily.

Analysis Using the Frontier Orbital Theory
The Frontier Orbital Theory considers that molecular orbitals (the highest occupied molecular orbital, or HOMO, and the lowest unoccupied molecular orbital, or LUMO) play a pivotal role during the chemical reaction. For lower E HOMO , the electronic capacity to offer molecules is lower because the electrons in the orbital are steadier. We use E LUMO to describe the ability of a molecule to accept electrons. The lower it is, the better electron-acceptor it is.
Our calculations indicate that the energy of the BaO HOMO is -288.44 kJ/mol and its LUMO energy is -154.02 kJ/mol. For TiO 2 the HOMO and LUMO energies are -558.31 kJ/mol and -352.97 kJ/mol, respectively, so that E HOMO (BaO) > E HOMO (TiO 2 ), and the E LUMO (BaO) > E LUMO (TiO 2 ). As a result, the LUMO of TiO 2 can react with the HOMO of BaO because BaO can offer electrons and TiO 2 can accept electrons. As shown in Figure 8, the main HOMO of BaO comes from the lone-pair electrons of the O 2p orbital, and the main part of the TiO 2 LUMO comes from the 3d orbital of the Ti atom, indicating that the O atoms of BaO molecules will combine with the Ti atoms of TiO 2 molecules. This conclusion is the same as the conclusion reached above. In addition, the calculated BTO HOMO and LUMO orbits are shown in Figure 8. The BTO HOMO distributed around the O atoms mainly form its lone-pair electrons in the 2p orbital, and the LUMO distributed around the Ba and Ti atoms mainly form their 5s and 3d orbital.

Conclusion
In our calculation, we find that BaO, TiO 2 , and BaTiO 3 molecules are the primary particle states in the early growth of BTO thin films by LMBE, with activation energies of 55.01 kJ/mol, 11.81 kJ/mol, and 10.55 kJ/mol, respectively. We find that TiO 2 molecules are formed first in vacuum where Ba, Ti, and O atoms coexist and become nucleation sites for BTO unit cell growth. Furthermore, the BaTiO 3 molecule, a part of the BTO unit cell, is formed by TiO 2 and BaO molecules. The BTO primary unit cell with a perovskite structure then evolves and grows with BTO unit cell migration growth on the substrate surface.


Molecular Interaction study of Multi-Functional Group of Alcohols through Dielectrics Using TDR

The complex permittivity spectra for tert-butyl alcohol(TB) with n-propanol(nP) were determined over the frequency range of 10 MHz to 20 GHz using the time domain reflectometry (TDR) in the temperature range of 25°C to 55°C for 11 different concentrations of the system. The static dielectric constant (ε0) and relaxation time (τ) have been determined these spectra using the Debye model. Excess properties and Kirkwood correlation factor of the mixtures have been determined. The excess permittivity is found to be positive in the tart-butyl alcohol rich region and negative in the n-Propanol rich region. However, the excess inverse relaxation time has different trend. The static dielectric constants for the mixtures have also been fitted with the modified Bruggeman model by assuming an additional parameter in the model.

Introduction
The measurement and characterization of the response of a material to an applied electric field is called dielectric analysis. Dielectric analysis measures changes in the properties of a material as it is subjected to a periodic electric field. The dielectric relaxation of binary polar liquids provides information regarding their interactions. In this work, we report the dielectric study of tert-butyl alcohol (TB) with n-propanol (nP) mixture; both are hydroxyl group molecules.

Experimental
TB and nP (AR grade, Merck Pvt. Ltd., West Mumbai, India) were used without further purification. The solutions were prepared at 11 different volume percentage of nP in TB from 0 % to 100 % in steps of 10 %. Using these volume percents the mole fraction is calculated as where m i , v i , and ρ i represent the molecular weight, volume percent, and density of the i th (i=1, 2) liquids, respectively. The density and molecular weight of the liquids are as follows: TB: Density=0.8093 gmcm -3 , Mol. Wt. = 74.12 nP : Density= 0.7997 gmcm -3, Mol. Wt. = 60. 11 The complex permittivity spectra were studied using the time domain reflectometry [1]- [6] method. The Hewlett Packard HP 54750 sampling oscilloscope with HP 54754A TDR plug in module has been used. The temperature controller system with water bath and a thermostat has been used to maintain the constant temperature within the accuracy limit of ±1 ο C.

Data Analysis
The time dependent data were processed to obtain complex reflection coefficient spectra ρ * (ω) over the frequency range from 10 MHz to 20 GHz using Fourier transformation [7], [8] as where p (ω) and q (ω) are Fourier transforms of [R 1 (t) -R x (t)] and [R 1 (t) + R x (t)] respectively, c is the velocity of light, ω is angular frequency, d is the effective pin length and j= − 1 .
The experimental values of ε * (ω) are fitted with the Debye equation [9] ε with ε 0, ε ∞ and τ as fitting parameters. A nonlinear least-squares fit method [10] was used to determine the values of dielectric parameters.

Results and Discussion
The obtained values of ε 0, ε ∞ and τ are listed in Table 1. The excess permittivity ε E is defined as where x-mole fraction and suffices m, 1, 2 represents mixture, liquid 1(TB) and liquid 2(nP) respectively. The excess permittivity may provide qualitative information about multimers formation in the mixture as follows: ε E = 0 , the liquid 1 and 2 do not interact ; ε E < 0 , the liquid 1 and 2 interaction in such a way that the total effective dipoles get reduced and ε E > 0, indicates the liquid 1 and 2 interaction such that the total effective dipole moment increases.
Similarly, the excess inverse relaxation time is defined as where (1/τ) E is excess inverse relaxation time which represents the average broadening of dielectric spectra. The inverse relaxation time analogy is taken from spectral line broadening in the resonant spectroscopy [12]. The information regarding the dynamic of liquid 1 and 2 interactions from this excess property is as: (1/τ) E = 0; there is no change in the dynamics of liquid 1 and 2 interaction, (1/τ) E <0 ; the liquid 1 and 2 interaction produces a field such that the effective dipoles rotate slowly and (1/τ) E >0 ; the liquid 1 and 2 interaction produces a field such that the effective dipoles rotate quickly .
The experimental values of both the excess parameters were fitted to the Redlich-Kister equation [13], [14] as where A is either ε E or (1/τ) E . By using these B n values, A E values were calculated. The 'g' values for pure liquid may be obtained by the expression where µ is dipole moment in gas phase, ρ is density, M is molecular weight, k is Boltzmann's constant and N is Avogadro's number. The dipole moments for TB and nP in gas phase are taken as 1.69 D and 1.68 D respectively [15]. Temperature dependent g eff and g f for the system is also shown in Fig. 1, (a) and (b).

Conclusion
The study has provided the permittivity parameters for the mixture of binary polar liquids. The molecules of these compounds have several hydrogen bonding sites and enter into intra or inter molecular hydrogen bonding giving rise to different conformations. This provides a unified approach to understand the interaction between two constituents in the liquid mixtures, using different models.


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ω-Heteroarylalkylcarbamates as inhibitors of fatty acid amide hydrolase (FAAH)

Please note that technical editing may introduce minor changes to the text and/or graphics, which may alter content. The journal’s standard Terms & Conditions and the Ethical guidelines still apply. In no event shall the Royal Society of Chemistry be held responsible for any errors or omissions in this Accepted Manuscript or any consequences arising from the use of any information it contains. Accepted Manuscript MedChemComm

ω ω ω ω-Heteroarylalkylcarbamates as inhibitors of fatty acid amide hydrolase (FAAH) Tobias Terwege, Helmut Dahlhaus, Walburga Hanekamp and Matthias Lehr* Fatty acid amide hydrolase (FAAH) is a serine hydrolase that terminates the analgetic and antiinflammatory effects of endocannabinoids such as anandamide. Herein we describe structure-activity relationship studies on a new series of ω-heteroarylalkylcarbamate inhibitors of FAAH. The most active compounds exhibit IC 50 values in the low nanomolar range. Investigations on selectivity and metabolic stability of these inhibitors are also presented.
Anandamide is one principal endocannabinoid in the mammalian organism. 1,2 It is formed "on demand" during several pathological disorders and mediates analgetic and anti-inflammatory effects by activation of the cannabinoid receptors CB 1 and CB 2 . These two receptors are differentially distributed within the various cells. [3][4][5] While the CB 1 receptor is preferentially expressed in neuronal tissues, the CB 2 receptor occurs mainly in cells of the immune system such as macrophages.
An important enzyme in the endocannabinoid metabolism is fatty acid amide hydrolase (FAAH), which rapidly cleaves the lipid mediator anandamide into arachidonic acid and ethanolamine. [6][7][8][9] Inhibition of FAAH is supposed to potentiate the action of anandamide. Therefore, inhibitors of FAAH may represent new agents against pain and inflammation. In the past years many potent inhibitors of FAAH have been found. [10][11][12][13][14] The majority of these compounds are substances, which form covalent binding interactions with the catalytic serine of the active site of FAAH such as carbamates like 1 (URB 597) 15 and 2, 16 activated ketone derivatives like the α-ketoheterocycle 3 (PHOP) 17 and the propan-2-one 4 18 (Figure 1), and urea derivatives like 5 (PF-04457845). 19 The latter compound is being examined in clinical studies for fear response and cannabis withdrawl. 20 Several years ago we have described dual inhibitors of cyclooxygenase and 5-lipoxygenase, which contain a pyrrole scaffold with vicinal phenyl substituents (e.g. 6) 21 ( Figure 2) showing structural simi-larities with the heterocyclic part of the FAAH inhibitor 2. 16 This initiated our interest in the question, in which way inhibitory potency of 2 is influenced when the 2-methyl-4,5-diphenylimidazole system of 2 is replaced by a 2,3-diphenyl-5-methylpyrrole substituent. Moreover, this pyrrole ring system offered the opportunity to introduce further substituents into the ring at the free carbon site, which could lead to additional binding interactions with FAAH. Therefore, we have synthesized the pyrrole analogue of 2 and used it as starting point for structure-activity relationship studies.  Scheme 1 Reagents and conditions: (a) tert-Butyl N-(6-bromohexyl)carbamate/tert-butyl N- (6-bromoethyl) The pyrrolyl-substituted N-alkylphenylcarbamates investigated were synthesized using one of the three routes outlined in Scheme 1. In the first instance, the appropriate nitrogen heterocycle was alkylated with a BOC-protected ω-bromoalkylamine followed by cleaving of the BOC group with trifluoroacetic acid and reaction of obtained amine with phenyl chloroformate. Alternatively, the heterocycle was alkylated with a dibromoalkane. The obtained bromoalkyl compound was treated with potassium phthalimide yielding a phthalimideprotected alkylamine. Removal of the protecting group was accomplished using hydrazine hydrate and the formed amine was directly reacted with phenyl chloroformate to yield the target compound. Instead, the phthalimide intermediate was directly prepared by reaction of the nitrogen heterocycle with ω-bromoalkyphthalimide as shown for the synthesis of 27 and 30.
Inhibition of FAAH hydrolase by the test compounds was determined with an assay using microsomes from rat brain as enzyme source and N-(2-hydroxyethyl)-4-pyren-1-ylbutanamide as fluorogenic substrate. 23 Inhibitory potencies of the tested compounds were assessed by comparing the amount of 4-pyren-1-ylbutanoic acid released from the substrate in their absence and presence after an incubation time of 60 min by reversed-phase HPLC with fluorescence detection (see Supplementary Information). With this assay for the diphenylimidazole 2 an IC 50 of 0.35 µM was measured (Table 1). Replacement of the imidazole scaffold of 2 by a pyrrole (9) led to an about 7-fold loss of activity. Shortening of the alkyl chain connecting the carbamate group and the pyrrole heterocycle from six to four and two carbon atoms further decreased inhibitory potency. With an IC 50 of 6.2 µM the butyl derivative 13 was only about half as active as the hexyl compound 9. The derivative with an ethyl linker (15) even was inactive at the highest test concentration of 10 µM. Because removing of the methyl group in position 2 of the imidazole ring of 2 did not change the activity of the compound, 24 the impact of the corresponding methyl group of the pyrrole 9 on inhibitory potency was not determined. Introduction of polar ethoxycarbonyl, acetyl and cyano substituents at the free position of the pyrrole ring of 9 was also detrimental. The inhibition values of obtained compounds 18, 21 and 24 were in the magnitude of 10 µM (Table 2). Interestingly, moving the phenyl group from position 3 to position 4 of the pyrrole ring, to give compound 27, as well as completely removing the phenyl ring, to give compound 30, significantly increased FAAH inhibition. The activity of these two inhibitors was comparable to that of the diphenylimidazole 2. Similar results were obtained with analogous indole derivatives (Table 3). With IC 50 values of about 5 µM the 2,3-diphenylindole 32 and the 3-phenylindole 34 possessed the lowest potency of the series. The 2-phenylindole 36 was about tenfold more active than the 3-phenylderivative 34. Thus, a phenyl substituent is much better tolerated in position 2 than in position 3 of the indole scaffold. The indole derivative 38, which has no phenyl-substituents in 2-and 3position, possessed the highest activity (IC 50 0.25 µM) showing that phenyl substituents are even unfavourable.
Next the effect of an introduction of nitrogen atoms in position 2 and 3 of the indole ring of 38 was studied. The indazole and benzotriazole derivatives 40 and 42 bearing a nitrogen in position 2 of the heteocycle were about three-to fourfold less active than the indole 38 (Table 4). In contrast, the imidazole derivative 44 possessing an additional nitrogen only in position 3 was equipotent to the indole 38.   ., 2012, 00, 1-3 This journal is © The Royal Society of Chemistry 2012 Taken together, replacement of the diphenylpyrrole heterocycle from the starting compound 9 by an unsubstituted indole (38) led to an about tenfold increase of FAAH inhibitory potency. At this point we wanted to evaluate the role of the alkyl linker of 38 on activity in more detail. Both shortening and elongation of the hexyl chain by one carbon increased activity about twofold: The IC 50 values of the pentyl and heptyl derivatives 47 and 49 were about 0.09 µM (Table  5). A more pronounced variation of the chain length to four (45) or eight (51) carbons resulted in a reduction of FAAH inhibition.  In literature a series of highly potent 3-pyridyl carbamate inhibitors of FAAH is described. 25 Therefore, we finally wanted to investigate the effect of a replacement of the phenyl group of the carbamates 47 and 38 by a 3-pyridyl residue. This structural variation led to a drastic increase of activity (Table 5). With IC 50 values of 0.0036 µM and 0.0053 µM obtained compounds 52 and 54 were about as active as the α-ketoheterocycle 3 (PHOP), which counts to the most active inhibitors of FAAH known today.
Besides FAAH, monoacylglycerol lipase (MAGL) is another important enzyme during the process of endocannabinoid inactivation. Just like FAAH, MAGL contains a catalytic serine in the active site. The main substrate of MAGL is 2-arachidonoyl glycerol. To get some information about the specificity of the developed substance class, some of the new FAAH inhibitors were also tested for MAGL inhibition. 26 As shown in Table 5, the evaluated phenyl ω-(indol-1-yl)alkylcarbamates did not inhibit MAGL at the highest test concentration of 10 µM. Contrary, with IC 50 values in the submicromolar range, the corresponding pyridyl carbamates 52 and 54 inhibited MAGL significantly. However, the inhibitory potency of these substances against FAAH is still 60-to 90-fold higher than against MAGL.
Another enzyme of the serine hydrolase family is cytosolic phospholipase A 2 α (cPLA 2 α, which catalyzes the first step of the so called arachidonic acid cascade by cleaving membrane phospho-lipids in arachidonic acid and lysophospholipids. At the highest test concentration of 10 µM cPLA 2 α was not inhibited 27 by the indole compounds listed in Table 5. These results indicate that the inhibitors investigated display a certain specificity with regard to their ability to inhibit serine hydrolase enzymes. Since carbamates are known to be susceptible to hydrolysis in biological environments [28][29][30] leading to the inactivation of the compounds, we also measured the stability of some of the new substances in rat liver S9 fractions in presence of the co-factor NADPH 31 and in porcine blood plasma (Table 5). After incubation in the liver S9 fractions for 30 min, still about 50% of the selected indole compounds were present in the incubation mixture. For comparison, the metabolism rate of the known FAAH inhibitor URB597, which also exhibits a carbamate structure, was significantly higher. Here, only about 13% of the parent compound could be detected after the metabolic reactions. Incubation in blood plasma for 30 min revealed striking different results. The pyridin-3-yloxy carbamates 52 and 54 had totally disappeared at the end of the trial most likely due to an extensive hydrolysis or transcarbamoylation. The stability of the phenoxycarbamates depended on the chain length of the alkyl spacer. The pentyl-substituted compound 47 exhibited high stability, with 91% parent remaining after incubation. In contrast, only 39% of the heptyl derivative 49 could be detected under the same conditions.


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Transition Metal Complexes and Radical Anion Salts of 1,10-Phenanthroline Derivatives Annulated with a 1,2,5-Tiadiazole and 1,2,5-Tiadiazole 1,1-Dioxide Moiety: Multidimensional Crystal Structures and Various Magnetic Properties

Advances in the molecular variety and the elucidation of the physical properties of 1,10-phenanthroline annulated with 1,2,5-thiadiazole and 1,2,5-thiadiazole 1,1-dioxide moieties have been achieved, and are described herein. A 1,2,5-thiadiazole compound, [1,2,5]thiadiazolo[3,4-f][1,10]phenanthroline (tdap), was used as a ligand to create multidimensional network structures based on S•••S and S•••N intermolecular interactions. A 1,2,5-thiadiazole 1,1-dioxide compound, [1,2,5] thiadiazolo[3,4-f][1,10]phenanthroline, 1,1-dioxide (tdapO2), was designed to create a stable radical anion, as well as good network structures. Single crystal X-ray structure analyses revealed that transition metal complexes of tdap, and radical anion salts of tdapO2 formed multidimensional network structures, as expected. Two kinds of tdap iron complexes, namely [Fe(tdap)2(NCS)2] and [Fe(tdap)2(NCS)2]•MeCN exhibited spin crossover transitions, and their transition temperatures showed a difference of 150 K, despite their similar molecular structures. Magnetic measurements for the tdapO2 radical anion salts revealed that the magnetic coupling constants between neighboring radical species vary from strongly antiferromagnetic (J = −320 K) to ferromagnetic (J = 24 K), reflecting the differences in their π overlap motifs.

Introduction
The electrical and magnetic properties of molecular crystals have been studied extensively, and various molecule-based conductors [1][2][3][4][5][6][7], superconductors [6,[8][9][10][11][12][13][14], and magnetic materials [15][16][17][18][19][20][21][22][23][24][25][26][27] have been synthesized to date. The research has been characterized by the enhancement of dimensionality in intermolecular interactions. As a result of the major efforts to enhance the dimensionality of intermolecular interactions, heterocyclic sulfur nitrogen (SN) compounds, which consist of nitrogen atoms and sulfur atoms, have arisen, and their solid-state properties have been investigated in relation to their crystal structures. The heterocyclic chalcogen-nitrogen compounds have attracted much attention due to their potential to form a multidimensional network structure in their crystals, which involve X•••N and X•••X (X = S, Se, Te) contacts and π-overlaps. Furthermore, some of these compounds are known to be radical species without any protecting groups. These features are believed to bring about the unique chemical and physical properties of these materials, such as the high-T c magnetic orderings, high conductivities and drastic structural phase transitions (see [28] and references therein). Although the intermolecular interactions are weaker than those of other heterocyclic chalcogen-nitrogen compounds, such as SeN and TeN compounds, various structures of heterocyclic SN compounds have been reported due to the convenience of their synthesis [28][29][30][31][32][33].
The synthesis of a closed-shell SN compound, [1,2,5]thiadiazolo [3,4-f] [1,10]phenanthroline (tdap, Scheme 1a), was reported in 2006 [34]. Using this compound as a ligand, the intermolecular interactions and electron spins which are the remarkable features of SN radicals can be separated into the 1,2,5-thiadiazole moieties and the transition metal ions, respectively. By this method, the stability can be greatly improved, whilst simultaneously maintaining the open shell structure. In this study, a 1,2,5-thiadiazole moiety is used to introduce the intermolecular interaction to transition metal ions which have a large quantum spin number.
In the present paper, we describe the crystal structures and spin-crossover transitions of tdap transition metal complexes, and the crystal structures and magnetic properties of radical anion salts of tdapO 2 .

Transition Metal Complexes of [1,2,5]thiadiazolo[3,4-f][1,10]phenanthroline
The synthesis of [1,2,5]thiadiazolo [3,4-f] [1,10]phenanthroline (tdap, Scheme 1a) was reported in 2006 by Conte et al. [34]. They reported the luminescent properties of the Eu III complex with the ligand in 2008 [49], and the photoluminescence and electroluminescence of the Tb III complex in 2011 [50]. In addition, they reported the photonuclease activity of the Fe II complex in 2010 [51]. In these papers, the crystal structures of these complexes were also reported. Thus there was no intermolecular interaction found, and the properties were discussed as single molecule features. Despite this report, we expected that the thiadiazole moiety of this ligand would have potential to form intermolecular interactions and would be a good building block for molecule-based magnets. In this chapter, we describe the crystal structures and intermolecular interactions of first-row transition metal complexes of tdap. Among them, the crystal structures and magnetic properties of two spin crossover complexes are discussed with emphasis on the intermolecular interactions and molecular structures. In this chapter, 10 [1,2,5]thiadiazolo [3,4-f] [1,10]

phenanthroline (tdap)
The electrochemical measurement of tdap in dimethyl sulfoxide (DMSO) has been reported, and revealed a redox peak in the reduction range (E 1/2 = −1.46 V vs. a normal hydrogen electrode (NHE)) [51]. To compare the acceptor ability to the related compound tdapO 2 described in a later section, the electrochemical measurement in acetonitrile (MeCN) was carried out (E 1/2 = −1.93 V vs. Fc/Fc + Figure 1, where Fc is ferrocene). These results suggest that tdap has poor acceptor ability.  [48]. Copyright 2013 American Chemical Society.

Synthesis of Transition Metal Complexes of tdap
Single crystals of transition metal complexes of tdap were obtained by the slow diffusion method using an H-shaped glass cell purged with N 2 gas [52]. Crystal solvent-free transition metal complexes [Fe(tdap) 2 (NCS) 2 ] (black blocks) [53] and [Co(tdap) 2

Spin-Crossover Complexes of tdap
The behavior of SCO transitions is known to be greatly affected by intermolecular interactions. The magnetic properties and crystal structures of two SCO complexes of tdap are discussed in this study. The non-solvated [Fe(tdap) 2 (NCS) 2 ] is reported to exhibit a SCO transition at 320 K [53].  Table 1. The transition was accompanied by an expansion of the Fe-N distances in the same manner as for previously reported SCO complexes [55].  Both in the HS and LS states, the space groups were P2 1 /n, and there was a 5.5% difference in unit cell volume between the two states. This change results not only from the thermal expansion of the crystal lattice, but also from the expansion of the Fe-N distances upon the SCO transition. The crystal structure of [Fe(tdap) 2 (NCS) 2 ]•MeCN, measured at 120 K, is shown in Figure 2b. In the LS state, the molecular packing exhibits S(1)•••N(11) contacts (3.017(4) Å: blue broken lines in Figure 2b), which are shorter than the sum of van der Waals (VDW) radii (3.35 Å), and weak S(2)•••S(4) (3.5911(7) Å: red broken lines in Figure 2b) contacts, which are almost the same as the sum of VDW radii (3.60 Å). Hereafter, the VDW radii given by Bondi were used (C: 1.75 Å; Cl: 1.75 Å; N: 1.55 Å; O: 1.50 Å; S: 1.80 Å) [56]. This weak S(2)•••S(4) contact forms a one-dimensional (1D) structure for the [Fe(tdap) 2 (NCS) 2 ] complexes, while the S•••N short contacts quench the intermolecular interactions between complexes. After the transition to the HS state, the S(2)•••S(4) contacts show a 2.5% expansion, while the S(1)•••N(11) distance shows a 2.4% reduction. As a result, the S(2)•••S(4) distances elongated to 3.6805(8) Å, which is longer than the sum of the VDW radii, and the complexes showed less interaction in the HS state.
Magnetic susceptibility measurements of [Fe(tdap) 2 (NCS) 2 ]•MeCN were performed under 0.5 T, in the temperature range of 2-400 K. Both the heating and cooling processes were performed to confirm the non-existence of structural change due to the evaporation of crystal solvent. The fact that the same profiles were observed in both processes suggests that no evaporation occurred even at 400 K. The paramagnetic susceptibility (χ p ) of the complex was calculated with a diamagnetic susceptibility (χ dia ) that was obtained by assuming that the temperature (T) dependence of χ p followed the Curie law at low temperature. The temperature dependence of χ p of [Fe(tdap) 2 (NCS) 2 ]•MeCN is shown in Figure 3. Since there is no observable thermal hysteresis in the SCO transition, only the heating process is shown. At low temperature, below 120 K, the value of χ p T is nearly constant (0.13 emu mol −1 K), though the spin state is expected to be S = 0. This is probably partially due to the presence of the HS state and/or contamination of the paramagnetic complex. Above 150 K, χ p T shows a gradual increase, and reaches 3.57 emu mol −1 K at 400 K. This value is larger than the theoretical spin-only value for a 5 T 2g ground state (3.0 emu mol −1 K with g = 2, where g is a g-factor), indicating spin-orbital coupling.  [53]. Adapted by permission of The Royal Society of Chemistry).

[Fe(tdap) 2 (NCS) 2 ]
The solvent-free [Fe(tdap) 2 (NCS) 2 ] was crystallized in the Pna2 1 space group. The molecular structure and selected bond lengths of solvent-free [Fe(tdap) 2 (NCS) 2 ] are shown in Figure 4a and Table 1   The comparison of the crystal structures in the HS state and LS state reveals several characteristic differences. It is notable that the N(10)-C(26) distance (for the NCS ligand) in the HS state is extremely short (see Figure 4a), presumably due to chemical pressure caused by the intermolecular contacts. The N-C(S) bond lengths at 400 K (1.117(9) Å for N(9)-C(25) and 1.055(11) Å for N(10)-C(26)) were 3.6% and 9.0% shorter than those at 173 K (1.159(5) Å for N(9)-C (25) and N(10)-C(26)), while [Fe(phen) 2 (NCS) 2 ] [54] and [Fe(tdap) 2 (NCS) 2 ]•MeCN exhibit a slight increase of 1.6% and a decrease of 0.6% in the N-C(S) distances, respectively. A comparison between the molecular structures of the solvated and non-solvated forms indicates that the large shrinkage in the latter would be caused by its characteristically short S•••S contacts. It is hard to clearly rationalize this difference, but the expansion of the FeN6 core could be balanced by such shrinkages in the non-solvated [Fe(tdap) 2 (NCS) 2 ].
Magnetic susceptibility measurements of [Fe(tdap) 2 (NCS) 2 ] were performed under 0.5 T, in the temperature range of 2-400 K. The χ p of the complex was calculated with a χ dia that was obtained by assuming that the temperature dependence ofχ p followed the Curie law. The temperature dependence of χ p for [Fe(tdap) 2 (NCS) 2 ] is shown in Figure 3. Below 150 K, the value of χ p T is nearly constant (0.21 emu mol −1 K), which is probably extrinsic. Above 180 K, χ p T shows a gradual increase, and reaches 2.56 emu mol −1 K at 400 K. This value is smaller than the theoretical spin-only value for a 5 T 2g ground state (3.0 emu mol −1 K with g = 2), indicating an incomplete transition. It is notable that there is no thermal hysteresis in the SCO transition in [Fe(tdap) 2 (NCS) 2 ], despite the intermolecular contacts in this compound. The unsubstituted 1,10-phenanthroline complex [Fe(phen) 2 (NCS) 2 ] is known to exhibit a steep SCO transition at 177 K without hysteresis [54]. Since the FeN6 geometry in [Fe(tdap) 2 (NCS) 2 ] is very similar to that in [Fe(phen) 2 (NCS) 2 ], the intermolecular contacts in [Fe(tdap) 2 (NCS) 2 ] are considered to bring about a high temperature shift of more than 100 K in the SCO transition. As described above, the intermolecular interactions of [Fe(tdap) 2 (NCS) 2 ] disturb the molecular expansion of the SCO transition. As a result, the expanded structure in the HS state is destabilized, and the LS state is maintained to a high temperature. Light-induced excited spin-state trapping (LIESST) experiments were performed under 0.5 T, in the temperature range of 5-120 K. An Hg-Xe lamp (550 nm, 1.5 mW cm −2 ) was used as a light source.
The light, passing through infrared (IR) and green filters, was guided via an optical fiber into the superconducting quantum interference device (SQUID) magnetometer. The sample was placed on the edge of the optical fiber. The thermal relaxation of the paramagnetism induced by photo-irradiation at 5K (Δχ p T) is shown in Figure 5. Light irradiation brings about a magnetization increase, but is saturated to a small value that is only 1.3% of the theoretical value for a complete LS to HS transition. Since [Fe(tdap) 2 (NCS) 2 ] appears as black crystals, only the surface molecules exhibit the spin transition. Upon heating, the photo-induced magnetization disappears at 80 K; this temperature is higher than that of the unsubstituted phenanthroline complex (55 K) [54].

Crystal Structures of tdap Transition Metal Complexes
The crystal structure of the crystal solvent-free cobalt complex [Co(tdap) 2 (NCS) 2 ] is shown in Figure 6. The [Co(tdap) 2 (NCS) 2 ] complexes are located on a 2-fold rotation axis. The unit cell contains two right-handed and two left-handed enantiomers of this chiral complex. The packing motif of [Co(tdap) 2 (NCS) 2 ] is much different from that of the crystal solvent-free iron complex [Fe(tdap) 2 (NCS) 2 ] ( Figure 4). There are no short intermolecular contacts, although this crystal does not contain crystal solvents which could quench the intermolecular interaction between neighboring complexes as is observed in [Fe(tdap) 2 (NCS) 2 ]•MeCN. The molecular arrangement of this crystal seems to be governed only by VDW forces and a small π-interactions (interplanar distances of ~3.47 Å) between tdap moieties.   Figure 8. In the crystal of the Ni complex, no short intermolecular interaction is found, either between complexes or between a complex and an MeCN molecule.
Magnetic susceptibility measurements of [M(tdap) 2 (NCS) 2 ]•MeCN (M = Mn, Co, Ni, Zn) were performed under 0.5 T, in the temperature range of 2-300 K. Curie paramagnetic behavior was found in the Mn, Co and Ni complexes, and the Zn complex exhibited diamagnetism. The measured Curie constants of the Mn, Co, Ni and Zn complexes at 300 K were 4.37, 2.95, 1.22 and 0.0057 emu mol −1 K, respectively, which are comparable to (Mn) or higher than (Co, Ni) the theoretical spin-only value. The higher value in the Co and Ni complexes was due to the spin-orbit coupling. Furthermore, the stronger magnetic anisotropy of the Co and Ni ions brought about a gradual decrease of the χ p T value of the Co complex, and a small increase of the χ p T value of the Ni complex.  There are two crystallographically nonequivalent thiocyanate ions. While one of the thiocyanate ions has a positional disorder, the other ion is fixed on a certain position, maybe due to the intermolecular interaction described above. It appears that the disordered thiocyanate ion simply fills a vacant space.  [53]. Adapted by permission of The Royal Society of Chemistry.

Dinuclear Copper Complexes
The crystal structure of the monovalent copper complex [Cu 2 (tdap) 2 (NCS) 2 ] is shown in Figure 11. Two monovalent Cu ions are bridged by two thiocyanate ions. The coordination geometry around Cu ions is distorted tetrahedral, reflecting the d 10 electron configuration of Cu I . The tdap moieties form a π stacked column, with an interplanar distance of 3.50 Å, in the direction of the a axis. These columns are bridged by the coordination bonds of copper-thiocyanate-copper, and the [Cu 2 (tdap) 2 (NCS) 2 ] complexes form a zig-zag sheet structure.
The crystal structure of the divalent copper complex [Cu 2 (tdap) 2 (NCS) 4 ] is shown in Figure 12. Two Cu ions are bridged by two thiocyanate ions. In contrast to [Cu 2 (tdap) 2 (NCS) 2 ], an additional thiocyanate ion is coordinated to a Cu ion. The geometry around a Cu II ion is square pyramidal, which is formed by four squarely coordinated nitrogen atoms from a tdap and two thiocyanate ions, and a sulfur atom from a thiocyanate ion on a perpendicular axis to the square. There are short contacts between the sulfur atoms of thiocyanate ions and tdap (3.3632(7) and 3.4073 Å) which form a ladder structure. Two Cu ions face each other across the ladders. The distance between intermolecular Cu ions, which is shown as broken lines in Figure 12, is 3.5812(14) Å.

Radical Anion Salts of [1,2,5]thiadiazolo[3,4-f][1,10]phenanthroline 1,1-dioxide
The molecule tdapO 2 is designed to possess a stable radical anion species, due to the strong electron-withdrawing property of the thiadiazole dioxide moiety [36,[57][58][59], which enhances the acceptor ability of tdapO 2 . In addition, the phenanthroline and thiadiazole dioxide moieties at the molecular ends are expected to operate as a ligand and/or a proton acceptor, which can realize multi-dimensional interactions (Scheme 1b). In this chapter, we describe the electrochemical properties of tdapO 2 first. Then we describe the multi-dimensional network structures and magnetic properties of the tdapO 2 radical anion salts, which were reported previously [47,48], and a novel radical anion salt [CoCp 2 ]•tdapO 2 .
The crystal structure of tdapO 2 consists of a 2D layer formed by C-H•••N hydrogen bonds, as shown by the broken lines in Figure 13c. Figure 13d shows a side view of the layered structure, where the interlayer arrangement includes a S-O•••H hydrogen bond, and very partial π-overlap. Figure 13b suggests the bending of the tdapO 2 molecule (ca. 15°), which may have been due to the repulsive force between neighboring molecules caused by the out-of-plane S=O bonds. Figure 13. Molecular structure (a,b) and crystal structure (c,d) of tdapO 2 . Adapted with permission from [47]. Copyright 2011 American Chemical Society. Figure 14 shows the CV of tdapO 2 , which indicates two reversible reduction peaks at -0.520 and -1.32 V vs. Fc + /Fc; these processes suggest that the mono and dianionic species of tdapO 2 are stable and the acceptor ability of tdapO 2 are comparable to those of the halogen-substituted p-benzoquinones [61][62][63][64]. It is notable that the related compound tdap exhibits very weak acceptor ability (see Figure 1). This means that the oxidation of the sulfur in the thiadiazole ring notably reduces the reduction potential. Solution EPR spectrum measurement of the reduced tdapO 2 was carried out on a CH 2 Cl 2 solution of mono-anionic tdapO 2 , which was prepared by the following method. A purple solution of reduced tdapO 2 in MeCN was prepared by potentiostatic reduction (−0.6 V vs. Ag/Ag + ) of 1 mM tdapO 2 in an MeCN solution of 100 mM TBAClO 4 . The resulting MeCN solution was then evaporated to dryness and dissolved in CH 2 Cl 2 . This CH 2 Cl 2 solution was degassed by a frozen degassing technique, and vacuum-sealed in a quartz cell. The measured and simulated EPR spectra are shown in Figure 15. The complex hyperfine structure suggests the existence of an unpaired electron that is well delocalized over the tdapO 2 skeleton. Figure 15. Experimental (red) and simulated (gray) solution EPR spectra of mono-anionic tdapO 2 . The inset shows the hyperfine coupling constants in units of mT used for the simulation, and in parentheses, the estimated spin density on the nitrogen and carbon atoms. Adapted with permission from [48]. Copyright 2011 American Chemical Society.
Theoretical fitting was carried out, and a best fit was obtained with the following parameters: g = 2.0032, two N (I = 1) with a N = 0.349 mT, two N (I = 1) with a N = 0.025 mT, two H (I = 1/2) with a H = 0.152 mT, two H (I = 1/2) with a H = 0.148 mT, and two H (I = 1/2) with a H = 0.05 mT, where I is the nuclear spin quantum number, and a N and a H are the hyperfine coupling constants for nitrogen and hydrogen, respectively. The spin distribution on the tdapO 2 radical anion was estimated from the McConell relation, a H = Q C  C and a N = Q N  N . Q C and Q N are the proportionality constants for the hydrogen-neighboring carbon, and for nitrogen, respectively [65,66].  C and  N are the spin distribution on the carbon and nitrogen, respectively. Q constants of Q C = −2.45 mT [67] for carbon of azaaromatic rings and Q N = +2.5 mT [68] for nitrogen were applied. The calculated spin densities are 0.14 and 0.01 for each of the two distinct nitrogen atoms, and 0.06 and 0.02 for the carbon atoms neighboring a hydrogen, at position 2 and 4, and at position 3 (Figure 15 inset), respectively. The positions of each spin density are proposed by comparing the obtained value with the calculated spin density of mono-anionic tdapO 2 [48].

Nonsolvated and Solvated Potassium Salts: K•tdapO 2 and KtdapO 2 •0.5MeCN
The crystal structure of nonsolvated K•tdapO 2 is shown in Figure 17a,b. Because the tdapO 2 molecules are located on a mirror plane, the molecular structure of tdapO 2 is ideally planar. The crystal structure of K•tdapO 2 consists of a 2D layer, formed by a coordination polymer chain along the a + c direction (Figure 17a), and regular π-stacking along the b axis with interplanar distances of 3.30 Å (Figure 17b). It thus seems that 2D layers are connected only by VDW force. Figure 17c,d show the crystal structure of K•tdapO 2 •0.5MeCN, which consists of a 2D layer formed by π-stacking along the a axis and a coordinate bonding chain along the c axis (Figure 17c). The 2D layers are connected by coordination bonds of MeCN molecules which are located between two alkali metal ions. The MeCN molecules of K•tdapO 2 •0.5MeCN are disordered between two K ions as shown in Figure 17c. The π-stacking chain includes a very weak alternation with interplanar distances of 3.36 and 3.42 Å. The π-overlaps between tdapO 2 molecules in these salts (Figure 18b,d) are much larger than that in the neutral tdapO 2 (Figure 13c), probably due to the exchange interactions between the radical anions.   5MeCN (c,d). Temperature dependence of χ p for K•tdapO 2 and K•tdapO 2 •0.5MeCN (e). Adapted with permission from [47] and [48]. Copyright 2011 and 2013 American Chemical Society.
The coordination geometry around the potassium ion and π-overlaps in K•tdapO 2 and K•tdapO 2 •0.5MeCN are shown in Figure 18. The π-overlaps between neighboring tdapO 2 molecules of these two salts are much larger than that of neutral tdapO 2 , probably due to the exchange interaction between radical anions. The π-overlap motifs of the nonsolvated form (Figure 18b) and the solvated form (Figure 18d) are significantly different from each other, presumably due to the geometrical difference around the potassium ions induced by MeCN coordination, as shown in Figure 18a,c.   Figure 18. Coordination geometry around the potassium ion (a,c) and π-overlaps (b,d) in K•tdapO 2 (top) and K•tdapO 2 •0.5MeCN (bottom). Adapted with permission from [48]. Copyright 2011 American Chemical Society.
The measurements of the temperature dependences of the magnetic susceptibility for K•tdapO 2 and K•tdapO 2 •0.5MeCN were performed under 3 T, in the temperature range of 2-300 K (Figure 17e). The temperature dependences of χ p of these radical anion salts were calculated with diamagnetic susceptibilities that were obtained as fitting parameters of the Bonner-Fisher model and Bleaney-Boweres model, respectively.
The χ p of K•tdapO 2 (green circles in Figure 17e) shows a gradual increase with a decrease in temperature from 300 K. After reaching a maximum at 13.5 K, χ p shows a decrease, followed by another quick increase below 6.5 K. The latter behavior is probably caused by the Curie spins on the lattice defects. The χ p values in the range of 16-300 K can be well fitted to the antiferromagnetic Bonner-Fisher model (Equation (1)), assuming the spin Hamiltonian in the zero field to be H = -2J S i •S i+1 , where J is the antiferromagnetic coupling constant, and S i and S i+1 are the monoelectronic spin operators [69][70][71]. The numerical expression is 0.25 + 0.14995x + 0.30094x 2 1 + 1.9862x + 0.68854x 2 + 6.0626x 3 (1) where x = |J|/k B T. The solid curve on the open circles in Figure 17e is the theoretical best fit of this model with g = 2.00 ± 0.05 and J/k B = −13 ± 1 K, where N A is the Avogadro's number,  B is the Bohr magneton and k B is the Boltzmann constant. The χ p of K•tdapO 2 •0.5MeCN (blue circles in Figure 17e) was smaller than that of K•tdapO 2 , reflecting the alternating nature of the π-stack column. The temperature dependence of χ p for K•tdapO 2 •0.5MeCN decreases in temperature from 300 K and forms a broad maximum around 150 K. Further cooling induces a quick increase below 10 K due to Curie defects. The behavior at the high temperature can be fitted to the Bleaney-Bowers model (Equation (2)) with the spin Hamiltonian in zero field: H = −2J S 1 •S 2 [69,72]. The numerical expression is The χ p values in the range of 200-300 K can be well fitted with following fitting parameters: g = 2.0 ± 0.1 and J/k B = −110 ± 10 K.

Mixed Valent Salts: K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2
The crystal structures of K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 were isostructural. Figure 19a,b shows the crystal structure of K•(tdapO 2 ) 2 , which consists of two crystallographically nonequivalent tdapO 2 molecules 1 (blue) and 2 (red). The bond lengths of these two tdapO 2 molecules (1 and 2) are sufficiently distinguishable (green circles in Figure 16), and clearly indicate that one molecule is neutral (1) and the other molecule is anionic (2). They form an alternating π-stacking arrangement along the a + c axis, with very similar interplanar distances of 3.29 Å but different intermolecular arrangements (i) and (ii). In a similar way, the tdapO 2 molecules in Rb•(tdapO 2 ) 2 form an alternating π-stacking arrangement with interplanar distances of 3.29 (i) and 3.31 Å (ii). As shown in Figure 19a, a potassium ion is tweezered by the phenanthroline moieties of the coplanar molecules 1 and 2 in different π-stacking columns, and is capped by the two thiadiazole dioxide moieties from both sides of the plane. These salts are in a class I mixed valence state [73], or, in other words, in a charge-ordered state-a state that has attracted much attention in recent years. Figure 19b indicates that the molecules 1 and 2 exhibit a checker-board-type charge ordering [74]. Figure 19. Crystal structure (a,b) of K•(tdapO 2 ) 2 and temperature dependence of χ p (c) for K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 . Adapted with permission from [47] and [48]. Copyright 2011 and 2013 American Chemical Society. Figure 19c depicts the temperature dependence of χ p for K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 , where the molar unit contains two tdapO 2 molecules. The behavior of K•(tdapO 2 ) 2 (green circles in Figure 19c) in the temperature range from 33 to 350 K can be fitted to the Bonner-Fisher model (Equation (1)) with the following fitting parameters: g = 1.95 ± 0.05 and J/k B = −20 ± 1 K for K•(tdapO 2 ) 2 . The behavior of Rb•(tdapO 2 ) 2 (blue circles in Figure 19c) in the temperature range from 54 to 300 K can be fitted to the Bonner-Fisher model (Equation (1)) with the following fitting parameters: g = 1.94 ± 0.06 and J/k B = −29 ± 1 K.
The |J| values of K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 are larger than that of K•tdapO 2 , though the π-stacking column in K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 includes the non-magnetic molecule 1. In general, the intermolecular magnetic exchange interaction results from cancellation between the ferromagnetic and antiferromagnetic contributions. Therefore, it is not entirely surprising that the antiferromagnetic interactions in K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 are stronger than that in K•tdapO 2 , if the latter involves a stronger ferromagnetic contribution. Below ca. 15K, the χ p values of K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 exhibit a sudden jump (Figure 20a), followed by a Curie-like behavior. The field dependence of the magnetization was measured at various temperatures. The magnetization curves (Figure 20b,c) indicate the presence of a small spontaneous magnetization below these temperatures, which is probably caused by spin canting. These are some of the highest transition temperatures among molecule-based magnetic materials that do not include transition metals [18].  (Figure 16) suggest that the valence of tdapO 2 is −1 and the composition formula is Cs 7 •(tdapO 2 ) 6 •ClO 4 . A Cs ion is surrounded by five oxygen atoms belonging to five tdapO 2 and four nitrogen atoms belonging to three tdapO 2 . The other Cs ion is surrounded by six oxygen atoms belonging to six tdapO 2 and eight oxygen atoms (whose occupancy is 0.5) belonging to two ClO 4 ions. A notable point is that six tdapO 2 radical anions are located on a so-called "Kagomé lattice," in which antiferromagnetically interacted spins form a frustrated system [75]. They also form a strongly alternating π-stacking arrangement along the c axis, with interplanar distances of 3.13 (i) and 3.29 Å (ii). Figure 21. Crystal structure (a,b) and temperature dependence of χ p (c) for Cs 7 •(tdapO 2 ) 6 •ClO 4 . Adapted with permission from [48]. Copyright 2013 American Chemical Society. Figure 21c depicts the temperature dependence of χ p for Cs 7 •(tdapO 2 ) 6 •ClO 4 , where the molar unit contains one tdapO 2 molecule. With the decrease in temperature from 300 K, χ p shows a gradual decrease. After reaching a minimum around 100 K, the value of χ p is maintained down to 30 K, and then shows another quick increase. The latter behavior is probably caused by the Curie spins on the lattice defects, whose Curie constant is calculated as 0.0037 emu mol −1 K. The behavior of Cs 7 •(tdapO 2 ) 6 •ClO 4 in the temperature range from 30 to 300 K can be fitted to the Bleaney-Bowers model. The best fit was obtained with the following fitting parameters: g = 2.0013 (fixed) and J/k B = −310 ± 10 K for Cs 7 •(tdapO 2 ) 6 •ClO 4 . The g value was fixed at 2.0013, which was obtained from solid EPR measurement for a microcrystalline sample of Cs 7 •(tdapO 2 ) 6 •ClO 4 . Figure 22a,b shows the crystal structure of (NH 4 ) 2 •tdapO 2 •I, which consists of an alternating π stacking of tdapO 2 with interplanar distances of (i) 3.35 and (ii) 3.44 Å (see Figure 22b). These are longer than those observed in the alkali metal salts. The intermolecular π-overlaps, (i) and (ii), in the stacking chain are shown in Figure 23b. Both the phenanthroline and thiadiazole dioxide moieties form hydrogen bonds to different ammonium ions, indicating the proton-accepting ability of the tdapO 2 radical anion (Figure 23a). Between the stacking columns, a 2D hydrogen-bonding network is formed by NH 4 + , I − and the proton-accepting moieties of tdapO 2 . Adapted with permission from [48]. Copyright 2013 American Chemical Society.
The crystal structure of Hppda•tdapO 2 •MeCN is shown in Figure 22c,d. Because it is difficult to determine the existence of the protons by XRD analysis, due to their small electron density, the valence of tdapO 2 was confirmed to be monovalent from the bond lengths of the crystal structure ( Figure 16). This result implies that a ppda molecule is protonated and becomes a 4-aminoanilinium ion (Hppda). The geometry of the hydrogen bonds around N(5), shown in Figure 23c, also implies a protonation of N(5). The crystal structure of Hppda•tdapO 2 •MeCN consists of a regular π-stacking of tdapO 2 with interplanar distances of 3.42 Å (Figure 22d). This distance is comparable to those of (NH 4 ) 2 •tdapO 2 •I. The intermolecular π-overlap in the stacking chain is shown in Figure 23d. The π-stack columns are connected by hydrogen bondings between tdap radical anions and Hppda cations, and form a 3D hydrogen-bonding network structure.
The blue circles in Figure 22e show the temperature dependence of χ p T for (NH 4 ) 2 •tdapO 2 •I. When the temperature decreases from 300 K, χ p T exhibits a gradual increase down to 30 K. After passing a broad maximum, χ p T decreases quickly. This behavior indicates the coexistence of ferromagnetic and antiferromagnetic interactions, with the former being stronger than the latter. Since the crystal structure consists of alternated π-stacking, this temperature dependence can be interpreted in terms of a ferromagnetic dimer, with weak antiferromagnetic interdimer coupling ( [76] using where J 1 and J 2 express the intradimer ferromagnetic coupling constant, and the interdimer antiferromagnetic coupling constant, respectively, and z is the number of the nearest neighbors connected by J 2 . The solid curve on blue circles in Figure 22e is the theoretical best fit with g = 1.94 ± 0.10, J 1 /k B = 24 ± 1 K and zJ 2 /k B = −7.9 ± 0.1 K. It is clear that the tdapO 2 radical anion has the potential to yield both ferromagnetic and antiferromagnetic intermolecular interactions. The green circles in Figure 22e show the temperature dependence of χ p T for Hppda•tdapO 2 •MeCN. When the temperature decreases from 300 K, χ p T exhibits a gradual increase down to 10 K. After passing a broad maximum, χ p T decreases quickly. This behavior is similar to that of (NH 4 )•tdapO 2 •I. Since the crystal structure consists of regular π-stacking, this temperature dependence can be interpreted in terms of a ferromagnetic 1D chain with weak antiferromagnetic interchain coupling (H = −2J 1 ΣS i •S i+1 and a molecular mean-field correction) [69,71,77] using where χ p = (N A g 2 μ B 2 )/(4k B T)[N/D] 2⁄3 , N = 1.0 + 5.7979916x + 16.902653x 2 + 29.376885x 3 + 29.832959x 4 + 14.036918x 5 , D = 1.0 + 2.7979916x + 7.0086780x 2 + 8.6538644x 3 + 4.5743114x 4 , and x = J 1 ⁄(k B T). The χ' p and χ p are the magnetic susceptibility with and without mean-field correction, respectively. The parameters J 1 and J 2 express the intrachain ferromagnetic coupling constant and the interchain antiferromagnetic coupling constant, respectively, and z is the number of the nearest neighbors connected by J 2 . The solid curve on green circles in Figure 22e is the theoretical best fit with g = 1.84 ± 0.10, J 1 /k B = 5.6 ± 0.  Table S1 (Supplementary Materials). The valence of tdapO 2 could be determined as −1 from its bond length ( Figure 16). The π-overlap arrangements of [CoCp 2 ]•tdapO 2 are shown in Figure 24. The intermolecular arrangement of the tdapO 2 radical anions in this system is much different from those in the other radical anion salts. The molecular planes of tdapO 2 are not parallel. The overlapped part shows a closer interplanar distance, maybe due to the exchange interaction through π orbitals.   The temperature dependence of χ p for [CoCp 2 ]•tdapO 2 is shown in Figure 25c. With a decrease in temperature from 300 K, χ p shows a gradual increase. After reaching a maximum around 50 K, χ p shows a slight decrease, followed by another quick increase below 20 K. The latter behavior is probably caused by the Curie spins on the lattice defects. The structural study revealed that the tdapO 2 π-stack column is well separated by the diamagnetic cation [CoCp 2 ] + . The χ p values in the range of 60-300 K can be well fitted to the antiferromagnetic Bonner-Fisher model (Equation (1)). The solid curve on the circles in Figure 25c is the theoretical best fit of this model with g = 1.92 ± 0.10 and J/k B = −49 ± 5 K.

Theoretical Calculations
It is quite remarkable that the tdapO 2 radical anion salts exhibit exchange-coupling constants that vary significantly from antiferromagnetic J/k B = −310 K to ferromagnetic J/k B = 24 K values (see Table 2). Wave function-based ab initio difference-dedicated configuration interaction (DDCI) calculations for the exchange interactions of 6 different radical dimers that were found in four kinds of tdapO 2 salts were carried out [48]. The exchange coupling constants J, effective hopping integral t, on-site repulsion U, and direct exchange K are summarized in Table 2. The ferro-and antiferromagnetic contributions can be estimated by following the 2J = 2K − 4t 2 /U partitioning, with the 2K term representing the ferromagnetic contribution and the −4t 2 /U term the antiferromagnetic contribution. The U values for tdapO 2 salts are almost constant, in agreement with the finding of Rota et al. that the U values are mostly governed by building units [78]. The magnetic orbitals are the in-phase and out-of-phase linear combinations of the singly occupied MOs (SOMOs) localized on each radical ( Figure 26). The relative position of the radicals (i.e., the three-dimensional arrangement) becomes a determining factor governing the magnetic interactions reflecting the localized character of the spin densities. Table 2 shows the evolution of the J value with respect to the interplanar distance between the radical dimers. The significant dependence of the exchange coupling constant on this structural parameter is very instructive and indicates that the overlap between the monomers plays a major role in defining the magnetic character of these materials. The calculated exchange coupling constant values are in good agreement with the experimental estimates (Table 2), despite some qualitative deviations observed in weakly interacting systems.
The comparison between Hppda•tdapO 2 •MeCN and the hypothetical Hppda•tdap•MeCN (lacking oxygen atoms on the thiadiazole ring) is rather instructive. Removal of the oxygen atoms shifts the magnetic behavior from ferro-(13 K) to antiferromagnetic (−4 K). It is remarkable that removal of the electron-withdrawing oxygen atoms leads to a ca. 3800 K increase of U in the hypothetical Hppda•tdap•MeCN radical as compared to Hppda•tdapO 2 •MeCN. Such variation of U is overtaken by a significant increase of the hopping integral t (from 85 to 791 K), which drives the hypothetical Hppda•tdap•MeCN into the antiferromagnetic regime. The presence of oxygen atoms reduces the superexchange contribution (absolute value of the antiferromagnetic contribution) with competitive effects on the parameters t and U. Table 2. Interplanar distances (d int ) and magnetic exchange coupling constants (J 1 and J 2 ) from the theoretical fits of experimental data, and J calc and corresponding t, K, and U parameters from DDCI calculations performed on radical anion pairs. Adapted with permission from [48]. Copyright 2013 American Chemical Society.  [48]. Copyright 2013 American Chemical Society.

Summary
The 1,2,5-thiadiazole compound tdap was induced to transition metal complexes in an effort to enhance intermolecular interactions and form highly networked systems. As a result, some of these complexes exhibited S•••S, S•••N, or S•••X intermolecular interactions, and formed a dimer, 1D ladder, 2D layer, and three-dimensional (3D) network structure. These interactions could be classified into four types: S(tdap)•••S(tdap), S(tdap)•••X (X = S, Cl belonging to coordinated anion), S(tdap)•••N(tdap), and S(tdap)•••X (X = N, Cl, S belonging to packing molecule or anion). The packing molecules and anions tended to quench the network structure in transition metal complexes of tdap. The effect of the intermolecular interaction on the physical properties was shown in the spin-crossover complex by the shift in their transition temperatures. [Fe(tdap) 2 (NCS) 2 ]•MeCN and [Fe(tdap) 2 (NCS) 2 ], underwent an SCO transition whose critical temperatures (T c ) were 210 and 315 K, respectively. A significant shift of transition temperature was found in solvent-free [Fe(tdap) 2 (NCS) 2 ], and may have been due to the chemical pressure caused by strong intermolecular interactions.
The 1,2,5-thiadiazole dioxide compound tdapO 2 was designed as a stable electron acceptor molecule with coordination sites. The electrochemistry of this compound indicated its good acceptor ability. The chemical and electrochemical reduction successfully gave 8 kinds of radical anion salts. These salts formed highly networked structures governed by effective coordination or hydrogen bonds. Most of the radical anion salts exhibited 1D magnetic behavior with various magnetic coupling constants that ranged from strong antiferromagnetic (−320 K) to ferromagnetic (24 K). Mixed-valent salts K•(tdapO 2 ) 2 and Rb•(tdapO 2 ) 2 showed magnetic ordering at low temperature, indicating the existence of considerable intercolumnar magnetic interactions in these salts.
Whereas the intermolecular interaction and physical properties of 1,2,5-thiadiazole compounds have been extensively reported on, for many systems, the solid state physical properties of 1,2,5-thiadiazolo 1,1-dioxide compounds have just started to be investigated. As shown in this review, 1,2,5-thiadiazole 1,1-dioxide compounds exhibit unique properties, such as coordination ability, proton acceptor ability, electron acceptor ability, and stability of radical anion species, with properties very different from those of 1,2,5-thiadiazole compounds. We expect the 1,2,5-thiadiazole 1,1-dioxide unit to form a new class of compounds which show exotic properties similar to those achieved for other heterocyclic chalcogen-nitrogen compound, such as thiadiazole, dithiazole, dithiadiazole etc.