Fe/Co-MOF Nanocatalysts: Greener Chemistry Approach for the Removal of Toxic Metals and Catalytic Applications

This study describes the preparation of new bimetallic (Fe/Co)–organic framework (Bi-MOF) nanocatalysts with different percentages of iron/cobalt for their use and reuse in adsorption, antibacterial, antioxidant, and catalytic applications following the principles of green chemistry. The prepared catalysts were characterized using several techniques, including X-ray powder diffraction, Fourier transform infrared spectroscopy, transmission electron microscopy, and scanning electron microscopy. These techniques proved the formation of MOFs, and the average crystallite sizes were 25.3–53.1, 27.6–67.2, 3.0–18.9, 3.0–12.9, and 3.0–23.6 nm for the Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 50%Fe:50%Co-MOF, and 90%Fe:10%Co-MOF samples, respectively. The nanoscale (Fe/Co) Bi-MOF catalysts as efficient heterogeneous solid catalysts showed high catalytic activity with excellent yields and short reaction times in the catalytic reactions of quinoxaline and dibenzoxanthene compounds, in addition to their antioxidant and antibacterial activities. Furthermore, the nanoscale (Fe/Co) Bi-MOF catalysts efficiently removed toxic metal pollutants (Pb2+, Hg2+, Cd2+, and Cu2+) from aqueous solutions with high adsorption capacity.

INTRODUCTION
Recently, the design, crystal engineering, and synthesis of metal−organic coordination frameworks have attracted great attention as promising functional materials. Metal−organic frameworks (MOFs) are a family of nonporous, crystalline solid-state materials with one-, two-, or three-dimensional structures. 1−3 MOFs are also known as coordination polymers, 4 and they are crystalline, microporous to mesoporous functional materials consisting of inorganic clusters interconnected by organic linkers. 5 Their inorganic components may be single or have more metal atoms coordinated by ligand atoms, and their organic components (ligand) are generally small organic molecules that can bind to metal ions or clusters via functional groups. Metal-linker clusters contain more than one metal atom or extended arrays of metal atoms linked by single ligand atoms. 6,7 Because of their unique characteristics (e.g., high specific surface area, high porosity, controllable pore size, low density, and great functionality), intriguing architectures, and intricate entangled motifs, MOFs have a wide range of applications in gas storage, gas separation, heterogeneous catalysis, microelectronics, nonlinear optics, sensing, drug delivery, membranes, and pollutant adsorption. 8−14 MOFs have many advantages over the more widely known zeolites. 15 Compared with traditional porous materials, MOFs have high specific surface areas, reaching about 10,000 m 2 /g. 13 Although MOFs are relatively new, their basic structures, known as the secondary building unit, 16 very closely resemble those of metal complexes as organic linkers can donate multiple lone pairs of electrons to metal cations, whereas metal cations are made up of vacant orbital shells that can accept these lone pairs of electrons to form MOFs. 17 A general strategy for the utilization of heterometallic− organic frameworks, especially bimetallic−organic frameworks (Bi-MOFs), is the use of two or more different metals. 18,19 On the subject of the catalytic activities of coordination polymers (MOFs), the types and properties of incorporated metal ions are important. 20−23 In principle, the active sites of MOF catalysts can be the metals at the nodes of a framework (metal centers with unsaturated coordination environments). 23, 24 In MOF compounds, the metal ions and organic linker (ligand) play important roles in catalytic reaction mechanisms. The inorganic metal moieties are Lewis acids; that is, they can act as electron pair acceptors, especially when uncoordinated metal sites are generated because of the removal of coordinated solvent molecules that remain after the preparation of MOFs. 10,25 The synthesis of multi-MOFs as nanomolds for hosting functional inorganic nanoparticles and their wide application have attracted attention in the last few years. 19,22,26 In this regard, heavy metals are of great concern because of the continuous industrial development. However, the intensive spread and accumulation of contaminants in water are considered serious environmental problems affecting human, animal, and plant life when released into the environment. 27−30 For instance, lead, mercury, copper, and cadmium ions are considered the worst heavy metals for humans and animals. 31−33 Quinoxaline and its derivatives are fragments of several biologically active and pharmacologically important compounds. 34 It exhibits a broad spectrum of biological activities (e.g., antibacterial, antiviral, anticancer, antifungal, antitubercular, and anti-inflammatory). 35 −38 In recent years, the use of recoverable and reusable catalysts has attracted special attention because of their nontoxicity and economic, environmental, and industrial advantages. 39 For instance, xanthene and its derivatives are important because of their wide range of biological and pharmaceutical characterizations and have agricultural bactericidal, antiviral, anti-inflammatory, and antiviral activities. 40 −44 They have been widely used in dyes, fluorescent materials for the visualization of biomolecules, and laser technologies. However, many of the synthesis methods of these catalysts suffer from one or more disadvantages, such as long reaction times (from 16 h to 5 days), harsh conditions, expensive reagents, use of toxic organic solvents, low product yields, and the need for excess reagents and catalysts.
This work aims to further understand new (Fe/Co) Bi-MOF nanocatalysts to determine their potential use in adsorption, antibacterial, antioxidant, and catalytic applications. Particular focus was devoted to the synthesis, structural characterization, and different applications of new inorganic−organic frameworks consisting of two different metals to investigate their diversity and potential as functional materials. We used the prepared nanocatalysts for the adsorption removal of four different heavy metals, catalytic activity of quinoxaline and dibenzoxanthene compounds, and antibacterial and antioxidant activities. Following the principles of green chemistry, the design of easily separable, reusable, nontoxic, low-cost, and insoluble acidic (Fe/Co) Bi-MOF nanocatalysts is the most important aim of this work. The pH values of the pollutant solutions were adjusted to 2−7 or 8 for heavy metals using 0.1 M HCl and 0.1 M NaOH solutions. The heavy metal equilibrium concentrations were determined using the spectrophotometric method at a certain wavelength for each pollutant. The adsorbed amount of heavy metals at equilibrium (q e in mg/g) was calculated using the following equation

EXPERIMENTAL
and 2.2.2. Catalytic Activity. Quinoxaline was synthesized using a mixture of 0.1 mol o-phenylenediamine and 0.1 mol diethyl oxalate in a round flask. Then, the flask was placed in an oil bath and refluxed for 2 h at 80°C in the presence of 0.03 g of activated (Fe/Co) Bi-MOF nanocatalysts. 37,39,54 Dibenzoxanthene was synthesized using a mixture of 2 mmol β-naphthol with 1 mmol benzaldehyde in a round flask. The flask was then placed in an oil bath and refluxed for 30 min at 70°C in the presence of 0.03 g of activated (Fe/Co) Bi-MOF nanocatalysts. 41,44,55,56 The formed solid was filtered, washed with water, and dried to give a pure compound. It could be also purified by crystallization from ethanol. The product was characterized by thin-layer chromatography (TLC) and Fourier transform infrared (FTIR) spectroscopy.
2.2.3. Antibacterial and Antioxidant Activities. We studied the antibacterial and antioxidant activities of all the nanocatalyst samples, where the antibacterial activities of the prepared nanocatalysts were individually tested against a panel of Gram-positive Bacillus subtilis, Gram-negative Escherichia coli, and Candida barbicans.

RESULTS AND DISCUSSION
3.1. X-ray Diffraction Analysis. Figure 1 shows the X-ray diffraction (XRD) patterns of the different percentages of (Fe/ Co) Bi-MOF catalysts.  47,49 The XRD patterns of Co-MOF showed two main peaks at 9.29 and 20.19°, which corresponds to a standard pattern. 49 The main peaks of the 10%Fe:90%Co-MOF, 50%Fe:50%Co-MOF, and 90%Fe:10% Co-MOFs appeared at 5. 35 22,[50][51][52][53]57 indicating that all (Fe/Co) Bi-MOF samples were crystallized. We observed that the peak at 9.63°o f the Fe-MOF catalyst was stronger than that of the Co-MOF catalyst. The XRD patterns show that the peak at 9.29°i ncreased with the increase in the iron percentage from 10 to 90% in the (Fe/Co) Bi-MOFs. The peak at 29.00°was clearer in the Co-MOF catalyst, which slightly decreased with the increased percentage of iron in the (Fe/Co) Bi-MOF catalysts. Generally, the peaks of the Fe-MOF increased with the increasing percentages of iron in the (Fe/Co) Bi-MOF catalysts, whereas the peak of the Co-MOF decreased with the decreasing percentages of cobalt in the (Fe/Co) Bi-MOF catalysts (i.e., the peaks of the Co-MOF catalysts decreased with the increasing percentages of iron). A sharp peak below 10°indicated that the MOF materials are crystalline. These results confirm the formation of (Fe/Co) Bi-MOFs and are in good agreement with the results of Wang et al. 50 The crystallite sizes were 25.3−53.1, 27.6−67.2, 3.0−18.9, 3.0−12.9, and 3.0−23.6 nm for Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 50%Fe:50%Co-MOF, and 90%Fe:10%Co-MOF, respectively.
3.2. FTIR Spectroscopy. The prepared catalysts were characterized using their FTIR spectra. Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 30%Fe:70%Co-MOF, 50%Fe:50%Co-MOF, 70%Fe:30%Co-MOF, and 90%Fe:10%Co-MOF samples were observed at room temperature to confirm the presence of BTC (to confirm the formation of MOFs), as shown in Figure 2. The peaks at 1663 and 1579 and 1478 and 1384 cm −1 indicate asymmetric and symmetric carboxylate groups, respectively. 2,48,58 Additionally, the absence of a strong absorption band near 1710 or 1714 cm −1 indicates that all carboxyl groups of the terephthalic ligand are deprotonated. 11,48,59 The broad and medium-intensity peak at 3420−3432 cm −1 was produced by the (OH) stretching vibrations of coordinated water molecules. 59,60 The Δν [ν as (COO) − ν s (COO)] values were 185 and 195 cm −1 , which are characteristic of the bridge coordination mode. 50,60 The relatively weak absorption band located at 3063 cm −1 was due to the stretching vibrations of the aromatic ring −(CH) ar , which confirms the presence of ligand molecules in the MOFs. 53,61 Additionally, the peak at 1105−1111 cm −1 was related to the C−O stretching vibration. 50,61 Generally, most of the bands in the region from about 1600−1605 to 700−800 cm −1 were due to aromatic ring stretching. The weak and  narrow bands at 1065−1083 and 750−782 cm −1 could be attributed to the γ(C−H) and δ(C−H) vibration of aromatic rings, respectively. 62 The band with medium strength at 538− 557 cm −1 was due to M−O vibrations, which proved that MOFs were indeed formed. 61 3.3. Transmission Electron Microscopy. Figure 3 shows the transmission electron microscopy (TEM) images of the morphology and microstructure of Fe-MOF, Co-MOF, and (Fe/Co)-MOF catalysts. The results showed irregular and similar hexagonal particles with nanosizes reaching nearly 28.8 and 27.6 nm for the Fe-MOF and Co-MOF, respectively. The TEM images of the 10%Fe:90%Co-MOF, 50%Fe:50%Co-MOF, and 90%Fe:10%Co-MOF catalysts showed the presence of two different shapes with defects in the (Fe/Co)-MOF samples, with nanosizes reaching nearly 23.0, 22.9, and 23.6 nm for 10%Fe:90%Co-MOFs, 50%Fe:50%Co-MOFs, and 90% %Fe:10%Co-MOFs, respectively. These results are in accordance with the XRD results. The 50%Fe:50%Co-MOFs showed an enhanced crystallinity with a crystallite size of 22.9 nm, making it the best nanosize sample with high catalytic and adsorption activities.
3.4. Scanning Electron Microscopy. Figure 4 shows the scanning electron microscopy (SEM) images of the MOF samples, which show irregular and similar hexagonal particles with smooth and large surface areas and nanosizes reaching nearly 28.8 and 27.6 nm for the Fe-MOF and Co-MOF, respectively. They also show large surface areas and the presence of two different shapes for the (Fe/Co) Bi-MOF samples with defects for 10%Fe:90%Co-MOF, 50%Fe:50%Co-MOF, and 90%%Fe:10%Co-MOF with nanosizes reaching nearly 23.0, 22.9, and 23.6 nm, respectively, which are in accordance with the XRD and TEM results.
3.5.1. Effect of pH on Adsorption. The pH of pollutant solutions is one of the major factors that influence the adsorption of heavy metals. 63 We studied the influence of pH on the removal of Pb(II), Hg(II), Cd(II), and Cu(II) using (Fe/Co) Bi-MOFs, presented in Figure 5. Generally, the uptake of Pb(II), Hg(II), Cu(II), and Cd(II) ions on absorbents by 0.03 g of Bi-MOF catalysts placed in bottles containing 100 mL of aqueous solutions of heavy metals as inorganic pollutants was significantly high in basic solutions compared with that in acidic solutions at an initial concentration (100 mg/L). The pH values selected in the experiments were prior to the precipitation limit of metals: pH = 5 for Pb 2+ , pH = 7 for Cd 2+ , and pH = 6 for Cu 2+ and Hg 2+ .  In Figure 5, the maximum adsorption capacities were 98.89, 95.98, 93.89, and 92.6% for lead, mercury, cadmium, and copper ions, respectively, on the Bi-MOF samples after 120 min. At low pH values, low removal efficiencies were observed because of hydrogen ions competing with metal ions on active sites (at ion-exchangeable sites). 64,65 The protonation of the surface of the catalysts at a lower pH led to the extensive repulsion of metal ions, which may be another reason for the decrease in the adsorption of metals at lower pH ranges.
3.5.2. Effect of Initial Metal Concentrations on Adsorption. We studied the effects of the initial concentrations of lead, mercury, cadmium, and copper ions on the removal rate at initial concentrations of 25, 50, 100, 150, 200, and 300 mg/ L after 120 min using 0.03 g of the 50%Fe:50%Co-MOF nanocatalyst. Figure 6 shows the adsorption capacities of heavy metals. The removal percentage decreased with the increase in initial concentrations of heavy metals from 25 to 300 mg/L. Evidently, at C o = 25, 50, 100, 150, 200, and 300 mg/L, the maximum adsorption capacities on the 50%Fe:50%Co-MOF nanocatalyst after 120 min were 100, 100, 98.37, 92. 16 70.50% for copper ions, respectively. At lower solution concentrations, the removal percentage was about 100% (i.e., approaching complete removal of heavy metals), which is highly appreciated in some cases such as the treatment of metal pollutants from wastewater. The decrease in the removal percentage of heavy metals at C 0 = 300 mg/L may be attributed to the saturation of available active sites on the  surface of the catalysts with the increase in the initial concentrations of heavy metals. 30,31,66,67 Additionally, the metal salt solutions contained nitrate or chloride anions, which increased in the metal solutions with increasing metal concentrations. In other words, the increase in the nitrate or chloride ions may block the active sites of the catalyst surfaces.
3.5.3. Effect of Adsorption Kinetics on the Adsorption of Heavy Metals. The adsorption efficiency as a function of contact time was monitored by varying the equilibrium time between the adsorbate (heavy metals) and the adsorbent (the prepared catalysts) in the range of 0−120 min for Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 30%Fe:70%Co-MOF, 50% Fe:50%Co-MOF, 70%Fe:30%Co-MOF, and 90%Fe:10%Co-MOF samples. The effect of the content time on the adsorption of lead, mercury, cadmium, and copper ions onto the prepared catalysts was determined at an initial concentration of 100 mg/L at 25°C using 0.03 g of the nanocatalysts. The time required to reach equilibrium is 6 h for the Fe-MOF and Co-MOF nanocatalysts and 120 min for the Bi-MOF nanocatalysts. Figures 7−12 show the effect of contact time on the adsorption of heavy metals. The removal of metal ions on the Fe-MOF and Co-MOF nanocatalysts observed a quick increase during the first 4 h and then a gradual increase within 6 h. Thereafter, no or a slight increase in metal adsorption was observed. For the removal of metal ions on the Bi-MOF nanocatalysts, a quick increase within the first 60 min, then a gradual increase within 120 min, and no able increase in the adsorption of heavy metals thereafter were observed. The experimental data confirmed that the 50%Fe:50%Co-MOF was the best sample for the adsorption of heavy metals. The uptake expressed in terms of the amount of the metal removed from the pollutant solutions was as follows: Pb (II) > Hg(II) > Cd(II) > Cu(II).
3.5.4. Effect of Catalyst Dosage. The adsorption of lead, mercury, cadmium, and copper cations by the (Fe/Co) Bi-MOF nanocatalysts was studied by changing the amount of the  catalyst (0.01, 0.03, 0.05, and 0.07 g) but keeping the initial metal concentration at 100 mg/L for 120 min on the 50% Fe:50%Co-MOF sample. The effect of the catalyst weight on the removal of Pb 2+ , Hg 2+ , Cd 2+ , and Cu 2+ ions is presented in Figure 13. The results reveal that the removal percentage of the toxic heavy metals from aqueous solutions increased with the increase in the catalyst weight from 0.01 to 0.07 g. The maximum adsorption efficiencies for 0.07 g of the 50%Fe:50% Co-MOF nanocatalyst after 120 min were 100, 99.37, 98.45, and 98.11% for lead, mercury, cadmium, and copper ions, respectively. This indicates that as the amount of the (Fe/Co) Bi-MOF nanocatalysts is increased, their ability to adsorb metals increases. This is because surfaces are finite at low catalyst weights, whereas they are infinite at high weights. 68 3.5.5. Effect of the Repeated Use of Catalysts. Figure 14 shows the effect of reusing catalysts on heavy metal adsorption using 0.03 g of the 50%Fe:50%Co Bi-MOF. The adsorption activity slightly decreased with catalyst reuse from 98.37, 95.98, 93.89, and 92.60% at the first use to about 60.11, 52.09, 46.89, and 50.31% at the fourth use for lead, mercury, cadmium, and copper ions, respectively. This clearly indicates the slight deactivation of the catalysts possibly because of the presence of nitrates or chlorides originating from the solutions of the metal salts or the gradual loss of weight of the catalysts during filtration and washing. 68 Our recyclability studies suggest that the prepared catalysts can be repeatedly used as efficient adsorbents in water treatment.  3.5.6. Adsorption Isotherms. The study of adsorption isotherms is very important in determining the correlations between adsorbates and adsorbents at equilibrium conditions. Thus, we used several adsorption isotherm models with distinguished parameters to understand the adsorption mechanism behavior of the adsorbent surface during adsorption and its adsorption capacity to adsorb adsorbates (pollutants). 66 We used in this study two widely used adsorption isotherm models: the Langmuir 69 and Freundlich 70 models.
The Langmuir and Freundlich isotherms equations can be written as follows Langmuir equations where q m (mg/g) is the maximum amount of pollutants (heavy metals) that can be adsorbed, K L is the Langmuir constant related to adsorption rates, K F is the Freundlich constant, and n indicates the favorability of the adsorption process. For the Langmuir isotherm, a straight line with slope =   can be expressed by a separation factor (R L ), which is defined as The R L value implies whether the adsorption is linear (R L = 1), favorable (0 < R L < 1), unfavorable (R L > 1), or irreversible (R L = 0). The obtained R L values were 0.0020, 0.00261, 0.00290, and 0.0035 for lead, mercury, cadmium, and copper cations, respectively, showing that the adsorption of these metal cations on the (Fe/Co) Bi-MOF nanocatalysts was favorable.
K F and n are Freundlich constants related to adsorption capacity and adsorption intensity, respectively. Generally, n > 1 illustrates that the adsorbate is favorably adsorbed on the adsorbent (catalysts), whereas n < 1 demonstrates that the adsorption process is chemical in nature. The plot of ln q e versus ln C e gave a straight line with a slope of 1/n and an intercept of ln K F . The values found for n were 3.127, 3.097, 1.851, and 2.51 for lead, mercury, cadmium, and copper cations, respectively, which proves that the adsorption of heavy metals on the (Fe/Co) Bi-MOF nanocatalysts is favorable and physical in nature.
The results show that the R 2 values obtained from the Langmuir isotherm for lead, mercury, cadmium, and copper ion adsorption on the (Fe/Co) Bi-MOF nanocatalysts were higher than those obtained from the Freundlich isotherm: the R 2 values from the Langmuir isotherm were 0.9998, 9997, Figure 13. Effect of the weight of the catalyst on the adsorption of Cd 2+ , Pb 2+ , Cu 2+ , and Hg 2+ after 120 min using 0.03 g of the 50%Fe:50%Co-MOF nanocatalyst.     Figure 15, whereas the Langmuir and Freundlich isotherms are described in Figure 16. 3.5.7. Kinetic Study. Adsorption kinetics is an important characteristic when determining the pollutant adsorption process efficiency. The adsorption of a solute by solid catalysts in aqueous pollutant solutions is a phenomenon that usually has complex kinetics. In this study, adsorption kinetics describes the rate of adsorbate (heavy metals) uptake onto the (Fe/Co) Bi-MOF catalyst samples. Several kinetic models are available to help us understand the behavior of adsorbents, examine the controlling mechanisms of the adsorption process, and test the experimental data. In the present investigation, the adsorption data were analyzed using two famous kinetic models: the pseudo-first-order model of heavy metal adsorption, which can be expressed by the equation 71 and the pseudo-second-order model of adsorption, which may be written as 72 The results confirmed that the pseudo-first-order kinetic model did not produce good linearity, and the predicted q e values obtained from this kinetic model did not display good agreement with the experimental ones. This suggested that heavy metal adsorption by (Fe/Co) Bi-MOF nanoparticles would not follow the pseudo-first-order kinetic model. Meanwhile, the pseudo-second-order kinetic model produced Figure 17. Pseudo-first-order and pseudo-second-order kinetic models for the adsorption of Pb 2+ by Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 30%Fe:70%Co-MOF, 50%Fe:50%Co-MOF, 70%Fe:30%Co-MOF, and 90%Fe:10%Co-MOF samples. good linearity, and the experimental and calculated q e values were obtained from this kinetic model as the value of R 2 of the pseudo-second-order kinetic model was higher than that of the pseudo-first-order kinetic model. In addition, the application of the pseudo-first-order model to the kinetic adsorption curves of lead, mercury, cadmium, and copper cations on the (Fe/Co) Bi-MOF nanocatalysts gave q e values that are significantly lower than the corresponding experimental ones. These suggested that the pseudo-second-order adsorption mechanism for the adsorption of lead, mercury, cadmium, and copper ions on the (Fe/Co) Bi-MOF nanocatalysts was dominant, showing that the adsorption process is controlled by chemisorption (the rate-limiting step in the adsorption of heavy metal ions is chemisorption), which involves valiancy forces through the sharing or exchange of electrons between the solvent and adsorbate. 73 The pseudo-first-order and pseudo-second-order models of heavy metal adsorption are described in Figures  17−20. 3.5.8. Mechanism of Adsorption of Heavy Metals. The infrared spectroscopy technique has proven to be a reliable technique for studying the interaction of a catalyst surface with pollutant species during the adsorption process of inorganic or organic pollutants from aqueous solutions. The change in the intensity of FTIR peaks after the adsorption process of heavy metals can be used to indicate interactions of the heavy metals with functional groups on the (Fe/Co) Bi-MOF nanocatalyst. Thus, FTIR spectra before and after metal ion adsorption were obtained and compared to identify the interaction of the 50% Fe:50%Co-MOF nanocatalyst surface with adsorbed heavy metals (Pb 2+ , Hg 2+ , Cd 2+ , and Cu 2+ ). Figure 21 shows that the intensities of characteristic FTIR peaks of the main M−O, M− OH, CO, and −C−O groups (as functional groups) were decreased after the adsorption of Pb 2+ , Hg 2+ , Cd 2+ , and Cu 2+ compared with those of the fresh 50%Fe:50%Co-MOF nanocatalyst, which indicates that these functional groups participated in the adsorption process of Pb 2+ , Hg 2+ , Cd 2+ , and Figure 18. Pseudo-first-order and pseudo-second-order kinetic models for the adsorption of Hg 2+ by Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 30%Fe:70%Co-MOF, 50%Fe:50%Co-MOF, 70%Fe:30%Co-MOF, and 90%Fe:10%Co-MOF samples. Cu 2+ . It was seen from the survey FTIR spectrum of the 50% Fe:50%Co-MOF nanocatalyst after the adsorption process that the FTIR peaks did not shift compared with those of the fresh (Fe/Co) Bi-MOF catalyst, indicating that the catalyst structure did not change after the adsorption of Pb 2+ , Hg 2+ , Cd 2+ , and Cu 2+ , which clearly indicates the presence of the adsorbed metals on the catalyst surface. This result might be related to the chemical interactions between (Fe/Co) Bi-MOFs and the heavy metal ions, as well as related to the formation of van der Waals forces, acid−base interactions, and diffusion on the catalyst surfaces. The magnetic properties of (Fe/Co) Bi-MOFs may lead to an improved adsorption rate. Based on the above discussion, it could be concluded that the main functional groups (M−O, M−OH, CO, and −C−O groups) were key active sites involved in the adsorption process of Pb 2+ , Hg 2+ , Cd 2+ , and Cu 2+ from aqueous solutions. The proposed adsorption mechanism on the (Fe/Co) Bi-MOF nanocatalyst has been illustrated in the Graphical Abstract.
3.6. Catalytic Activity. 3.6.1. Surface Acidity (Nonaqueous Titration). Nonaqueous titration was used to investigate the acidity of surfaces, including the total number of acid sites and acidic strength. According to the nonaqueous titration, the initial electrode potential (E i ) indicates the maximum acid strength of the surface sites, whereas the range where a plateau is reached (meq/g solid) indicates the total number of acid sites. The nonaqueous titration method was conducted using n-butyl amine in acetonitrile. The value for the electrode potential in this method is related to the surface acidity of the catalysts, as shown in Figure 22.
3.6.2. Catalysis of the Quinoxaline Synthesis by (Fe/Co) Bi-MOF Nanocatalysts. We studied the catalysis of quinoxaline synthesis at 80°C in the presence of catalytic amounts of (Fe/Co) Bi-MOF nanocatalysts (0.03 g). The progress of the reaction was monitored using TLC. After the completion of the reaction, the catalysts were filtered off, washed, and dried. The combined organic layers were then washed and dried. The Figure 19. Pseudo-first-order and pseudo-second-order kinetic models for the adsorption of Cd 2+ by Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 30%Fe:70%Co-MOF, 50%Fe:50%Co-MOF, 70%Fe:30%Co-MOF, and 90%Fe:10%Co-MOF samples. solvent was evaporated under reduced pressure, and pure products were obtained without any further purification. 37,74 Figure 23 shows the FTIR analysis of the quinoxaline compounds. A wide range of homogeneous and heterogamous catalysts such as Cu (BDC), Fe 3 O 4 , HCl, and iodine are wellsuited for such reactions. 54 Figure 24, the quinoxaline yield percentage increased gradually with the increasing percentages of iron up to the maximum at 50%Fe:50%Co-MOF and then decreased. This conforms with the XRD and TEM results, which showed that the nanosize of the (Fe/Co) Bi-MOF catalysts decreased with the increasing percentage of iron up to the maximum at 50%Fe:50%Co-MOF and then increased. When the nanosize decreased, the surface area increased, the number of active sites on the catalyst surface increased, and the catalytic activity was very high. Figure 25 and Table 1 show that the catalytic activity and number of acid sites increased with the increasing iron percentage up to 50%Fe:50%Co-MOF and then decreased. The increase in the surface acidity may have been due to the strong interaction between the greater amounts of iron in Bi-MOFs and cobalt.
3.6.2.2. Effect of the Molar Ratio of o-Phenylenediamine/ Diethyl Oxalate. The molar ratio of reactants in any catalytic reaction is one of the most important parameters that must be optimized. 78 The effect of the molar ratio of o-phenylenediamine to diethyl oxalate ranging from 1:1 to 1:5 on the quinoxaline synthesis at 80°C by the 0.03-g 50%Fe:50%Co-MOF nanocatalyst was investigated. The results are shown in Figure 26. It can be observed that the percentage yield of quinoxaline increased from 93.40 to 95.51 and 98.38% as the molar ratio of o-phenylenediamine to diethyl oxalate is increased from 1:1 to 1:2 and 1:3, respectively. Thereafter, the yield percentage decreased from 95.15 to 90.15% with the increase in the molar ratio of o-phenylenediamine to diethyl Figure 20. Pseudo-first-order and pseudo-second-order kinetic models for the adsorption of Cu 2+ by Fe-MOF, Co-MOF, 10%Fe:90%Co-MOF, 30%Fe:70%Co-MOF, 50%Fe:50%Co-MOF, 70%Fe:30%Co-MOF, and 90%Fe:10%Co-MOF samples. oxalate from 1:4 to 1:5. From these results, the 1:3 m ratio of o-phenylenediamine to diethyl oxalate was the optimum for the synthesis of quinoxaline. The decrease in the catalytic activity of the 50%Fe:50%Co-MOF nanocatalyst at molar ratios of 1:4 and 1:5 may be explained on the basis of the fact that the increase in the diethyl oxalate concentration hinders the reaction by blocking active sites on the catalyst surface as a result of the saturation of the active sites with diethyl oxalate molecules. 78,79 3.6.2.3. Reusability (Recyclability) Study of the Catalysts. The reusability of heterogeneous catalysts is an important aspect of any industrial process. Thus, we studied the reusability of the 50%Fe:50%Co-MOF catalyst by conducting repeated runs of the reaction at 80°C and keeping the reactant molar ratio of o-phenylenediamine to diethyl oxalate at 1:1.
The 50%Fe:50%Co-MOF catalyst was regenerated after every 2 h of reaction time, and it was separated through filtration, washed with ethanol and DMF, dried at 120°C for 2 h, and reused in another reaction under similar conditions. Figure 27 shows the results of reusing the catalyst four times. With the increase in the number of usages of the 50%Fe:50%Co-MOF nanocatalyst, its catalytic activity is slightly decreased from 93.40, 86.10, and 82.40, to 76.80 for the first, second, third, and fourth use, respectively, because of the gradual loss of the catalyst weight during filtration and washing. 80,81 3.6.2.4. Effect of the Weight of the Catalyst on the Synthesis of Quinoxaline. In this part, the amount of the 50% Fe:50%Co-MOF nanocatalyst was varied to 0.01, 0.03, 0.05, and 0.07 g while keeping the molar ratio of o-phenylenediamine to diethyl oxalate at 1:1 and the reaction temperature at  80°C for 2 h. Figure 28 shows that with the increase in the amount of the 50%Fe:50%Co-MOF nanocatalyst from 0.01, 0.03, and 0.05, to 0.07 g, the quinoxaline yield increased from 71.98, 93.40, and 96.00 to 97.23%, respectively. This was due to the availability of the surface area and acid sites (active sites), 79 where the catalyst weight increased (i.e., the number of nanoparticles increased), increasing the surface area and number of active sites.
3.6.2.5. Effect of the Reaction Time on the Synthesis of Quinoxaline. The influence of the reaction time on the quinoxaline synthesis by the (Fe/Co) Bi-MOF catalysts is given in Figure 29. Using 0.03 g of the 50%Fe:50%Co-MOF nanocomposite as the catalyst under other similar conditions, we studied the effect of various reaction time durations (0.5, 1.0, 1.5, 2.0, and 3.0 h). The results indicate that the quinoxaline yield increased with the increase in reaction time from 0.5 to 2 h. Thereafter, no notable increase or only a slight increase in the quinoxaline yield was observed, suggesting the influence of time on the yield. The percentage yield of quinoxaline increased from 18.00, 49.91, 74.44, and 94.30 to 94.81% after 0.5, 1.0, 1.5, 2.0, and 3.0 h, respectively. This means that the catalyst activity established equilibrium after 2 h.
3.6.2.6. Effect of the Used Solvent in the Reaction. Solvents are known to significantly affect the rates of catalytic activities. The effect of the used solvent is attributed to various factors, which include the solubility of the yield, thermodynamic interaction of the solvent with the reactants and products, and competitive adsorption of the solvent. Different reaction media were used to perform quinoxaline synthesis, such as those using H 2 O, C 2 H 5 OH, DMF, and dimethyl sulfoxide (DMSO) as solvents. Figure 30 clearly shows that under the same reaction conditions, when water and ethanol were used as solvents, the reaction led to relatively higher yields and shorter reaction times; the quinoxaline yields were 93.4 and 92.12% with ethanol and water, respectively. 34,82 In  comparison, under the same reaction conditions, when DMF and DMSO were used as solvents, the yield decreased to 71.71 and 45.03%, respectively. This decrease in the yield may have been caused by the solvent dissolving the formed yield. 34,83 3.6.3. Catalysis of the Dibenzoxanthene Synthesis via (Fe/ Co) Bi-MOF Catalysts. We studied the catalysis of dibenzoxanthene synthesis at 70°C in the presence of a catalytic amount of (Fe/Co) Bi-MOF nanocatalysts (0.03 g). The progress of the reaction was monitored using TLC. After the completion of the reaction, the catalyst was filtered off, washed, and dried. The combined organic layers were then washed and dried. The solvent was evaporated under reduced pressure, and a pure product was obtained without any further purification. 37,74 Figure 31 shows the FTIR analysis of the dibenzoxanthene compound.
3.6.3.1. Effect of the Percentages of Iron and Cobalt in (Fe/Co)-MOFs on the Catalytic Activity. The effect of the percentages of iron and cobalt in Bi-MOFs on catalytic activity using 0.03 g of different (Fe/Co) Bi-MOF catalysts for the catalysis of dibenzoxanthene synthesis processes was also studied. As shown in Figure 32, the percentage of   dibenzoxanthene yield increased gradually with the increase in the percentage of iron up to 50%Fe:50%Co-MOF and then decreased. This conforms with the XRD and TEM results, which showed that the nanosize of the (Fe/Co) Bi-MOF catalysts decreases with the increase in the percentage of iron up to 50%Fe:50%Co-MOF and then decreased. As the nanosize decreased, the surface area increased, the number of active sites on the catalyst surface increased, and thus, the catalytic activity is increased. 84 From Figure 33 and Table 1, the catalytic activity and the number of acid sites increased with the increase in iron percentage to 50%Fe:50%Co-MOF and then decreased. The increase in the surface acidity may have been due to the strong interaction between the greater amounts of iron in Bi-MOFs and cobalt.
3.6.3.2. Effect of the Molar Ratio of β-Naphthol/ Benzaldehdyde. We studied the effect of the molar ratio of β-naphthol to benzaldehyde (from 2:1 to 2:5) on the synthesis of dibenzoxanthene at 70°C using 0.03 g of the 50%Fe:50% Co-MOF catalyst. Figure 34 shows that the percentage yield of dibenzoxanthene increased from 73.11 and 85.20 to 87.04% with the increase in the molar ratio of β-naphthol to benzaldehyde from 2:1 and 2:2 to 2:3, respectively. Therefore, the yield percentage increased with the increase in the molar ratio of β-naphthol to benzaldehyde from 2:4 to 2:5. From these results, the molar ratio of 2:3 was found to be optimum for the synthesis of dibenzoxanthene. The absence of an increase in the catalytic activity of the 50%Fe:50%Co-MOF nanocatalyst at molar ratios of 1:3 to 1:4 may have been because β-naphthol was completely consumed in the reaction. Thus, both the benzaldehyde concentration and the amount of β-naphthol must be increased in the reaction.
3.6.3.3. Reusability (Recyclability) Study of the Catalysts. We also studied the reusability of the prepared nanocatalysts while keeping the reactant molar ratio of β-naphthol to benzaldehyde at 2:3. Figure 35 shows the result of reusing the catalysts four times. The catalytic activity of the 50%Fe:50% Co-MOF catalyst is slightly decreased from 87.45, 78.81, and 73.29 to 66.59 for the first, second, third, and fourth use, respectively, because of the gradual loss of the catalyst weight during filtration and washing. 17,81 The 50%Fe:50%Co-MOFs catalyst was regenerated after every 30 min of reaction time, and it was separated through filtration, washed with ethanol and DMF, dried at 70°C for 30 min, and reused in other reactions under similar conditions. 3.6.3.4. Effect of the Weight of the Catalyst on the Synthesis of Dibenzoxanthene. The amount of the 50% Fe:50%Co-MOF nanocatalyst was varied from 0.01, 0.03, and 0.05, to 0.8 g while keeping the molar ratio of β-naphthol to benzaldehyde at 2:3 and the reaction temperature at 70°C for 30 min. Figure 36 indicates that with the increase in the amount of the 50%Fe:50%Co-MOF nanocatalyst from 0.01, 0.03, and 0.05 to 0.08 g, the yield of dibenzoxanthene increased from 70.66, 87.45, and 89.98 to 92.33%, respectively. This is due to the availability of the surface area and acid sites (active sites). 67 As the catalyst weight increased (i.e., the number of nanoparticles increased), the surface area and number of active sites increased.
3.6.3.5. Effect of the Reaction Time on the Synthesis of Dibenzoxanthene. We studied the influence of the reaction time on the synthesis of dibenzoxanthene by (Fe/Co) Bi-MOF nanocatalysts. In this regard, 0.03 g of 50%Fe:50%Co was taken as the catalyst under other similar conditions. Figure 37 shows the effect of 10, 20, 30, and 40 min reaction times. The results indicate that the yield of dibenzoxanthene increased with the increase in the reaction time from 10 to 30 min. Thereafter, no notable increase or only a slight increase in the yield of dibenzoxanthene was observed, suggesting the influence of time on the yield. The yield percentage of dibenzoxanthene increased from 25  Fe:90%Co-MOF, 30%Fe:70%Co-MOF, 50%Fe:50%Co-MOF, 70%Fe:30%Co-MOF, and 90% Fe:10%Co-MOF) samples exhibited a remarkable antibacterial activity against the tested   bacterial strains, as shown in Table 2. The results show that the 50%Fe:50%Co-MOF nanocatalyst exhibited the highest antibacterial activity. The antibacterial activities of the     common standard antibiotic ampicillin and antifungal clotrimazole were also recorded using the same procedure at the same concentration and solvents to be considered as the reference drug.
3.8. Conclusions. (Fe/CO) Bi-MOF nanocatalysts have been successfully attained, obtaining a mixed-metal (Fe/Co) Bi-MOF material series with different weight ratios of Fe and Co. XRD, TEM, SEM and FTIR spectroscopy were used for the characterization of the prepared nanocatalysts. XRD, TEM, and SEM proved that (Fe/Co) Bi-MOFs were obtained with sizes in the nanoscale reaching 3 nm, and FTIR results confirmed that (Fe/Co) Bi-MOFs were formed. The (Fe/Co) Bi-MOFs as good adsorbents showed efficient removal of the toxic heavy metal cationic pollutants (Pb 2+ , Hg 2+ , Cd 2+ , and Cu 2+ ) from the aqueous solution with high adsorption capacity. The nanoscale (Fe/Co) Bi-MOF nanocatalysts as efficient heterogeneous solid acid catalysts showed high catalytic activity with excellent yields and short reaction time in the catalytic reactions of quinoxaline and dibenzoxanthene


Acceptorless dehydrogenative synthesis of primary amides from alcohols and ammonia

The highly desirable synthesis of the widely-used primary amides directly from alcohols and ammonia via acceptorless dehydrogenative coupling represents a clean, atom-economical, sustainable process. Nevertheless, such a reaction has not been previously reported, and the existing catalytic systems instead generate other N-containing products, e.g., amines, imines and nitriles. Herein, we demonstrate an efficient and selective ruthenium-catalyzed synthesis of primary amides from alcohols and ammonia gas, accompanied by H2 liberation. Various aliphatic and aromatic primary amides were synthesized in high yields, with no observable N-containing byproducts. The selectivity of this system toward primary amide formation is rationalized through density functional theory (DFT) calculations, which show that dehydrogenation of the hemiaminal intermediate into primary amide is energetically favored over its dehydration into imine.

Introduction
Ammonia is the simplest, most abundant precursor for the industrial preparation of nitrogen-containing compounds, but its application in selective homogeneously catalyzed organic synthesis by metal complexes has been challenging, partly due to the formation of stable Werner-type ammine complexes. 1 In 2008, our group reported an atom-and step-economical synthesis of primary amines from alcohols and ammonia, catalyzed homogeneously by an acridine-based PNP-type ruthenium pincer complex. 2a Since then, environmentally benign processes, which involve coupling of alcohols and ammonia, and generate no hazardous waste, have been developed to access amines, imines (including N-heterocycles) and nitriles. 2 The rst mechanistic steps common to these transformations are alcohol dehydrogenation into aldehyde or ketone, 3 followed by nucleophilic attack of ammonia on the newly-formed carbonyl, leading to a hemiaminal intermediate. The latter subsequently undergoes facile dehydration into imine, which typically reacts further to yield various Ncontaining products (Scheme 1a). In the case of primary alcohols, an alternative route can be envisioned, wherein the hemiaminal undergoes dehydrogenation rather than dehydration, 4 thereby affording a primary amide.
Primary amides feature extensively in organic synthesis, and their ubiquity in pharmaceuticals, natural products, agrochemicals and biologically-active molecules has made their preparation the focus of widespread attention. 5 Classical synthetic methodologies involve the amidation of carboxylic acid derivatives or hydration of nitrilesprocesses that either generate considerable waste or require the use of pre-prepared starting materials. 6 Consequently, there is high demand for more sustainable and efficient means of accessing primary amides from readily-available sources. 7 From this perspective, the synthesis of primary amides directly from alcohols and ammonia is highly desirable, since these precursors are abundant and inexpensive. Nevertheless, the existing methods for this transformation require stoichiometric amounts of additives, such as expensive or harmful oxidants or organic hydrogen acceptors, making them neither sustainable nor atom-economical. 8 For instance, Grützmacher and co-workers reported an efficient homogeneous rhodium-catalyzed coupling of alcohols and ammonia into primary amides, 8b but this required a 5-fold excess of methyl methacrylate as a hydrogen scavenger. Some catalytic systems employ O 2 as the hydrogen acceptor, with concomitant generation of water, thereby making them more environmentally friendly. 8j-l A notable example was reported by Mizuno and co-workers, who implemented a heterogeneous catalysis approach, using manganese-oxide-based molecular sieves under 3 bar of O 2 . 8j However, a fundamental drawback of all acceptor-based systems, apart from the inherent need for additives, is that hydrogen atoms from the substrate are transferred to the acceptor, rather than being released as hydrogen gasa valuable commodity chemical in its own right. An acceptorless process, whereby alcohols and ammonia are coupled into primary amides with liberation of H 2 as a byproduct, would therefore be highly advantageous, but this has never been reported. One of the challenges in developing such a process is that the hemiaminal intermediate can easily dehydrate into an imine upon heating, 8j thereby inducing side reactions. Furthermore, reactions involving H 2 evolution usually benet from an open system, whereas utilizing gaseous NH 3 as a reactant typically requires a closed system.
Herein, we report the unprecedented synthesis of primary amides directly from alcohols and gaseous ammonia, with concomitant evolution of H 2 , using a pyridine-based PNNruthenium pincer complex (Ru-1) as the catalyst, combined with catalytic amounts of base (Scheme 1b). This system exhibits excellent chemoselectivity toward the generation of primary amides, rather than other N-containing compounds, and enables the synthesis of various aliphatic and aromatic primary amides in generally high yields.

Establishment and optimization of the catalytic reaction conditions
To realize the synthesis of primary amides by dehydrogenative coupling of alcohols and ammonia, we reasoned that a pincer complex could constitute an appropriate catalyst, since such complexes have usually been stable in the presence of ammonia. 2 Several pyridine-based ruthenium-containing pincer complexes, especially those already known to generate secondary or tertiary amides from alcohols, 3b,4 were examined as catalysts in order to preferentially achieve the dehydrogenation of the ammonia-derived hemiaminal intermediate (Scheme 1a). 3-Phenyl-1-propanol (1a) was chosen as the model substrate for the preliminary experiments. In each of these trials, a toluene solution of 1a containing 1 mol% catalyst and 2 mol% t BuOK, was placed under 7 bar of ammonia inside a 90 mL Fischer-Porter tube at room temperature, and was then heated at 135 C by immersion in an oil heating bath (Table 1). Interestingly, using our PNN-Ru complex Ru-1 (ref. 3a) as the catalyst, 92% conversion was achieved aer heating for 18 h, with the target product 2a being observed in 33% yield, along with 30% of ester 3a (accounting for 60% of the consumed substrate). No other N-containing side products, such as amines, imines or nitriles, were observed by GC-MS analysis of the crude reaction mixture. This attests to the excellent activity of this catalytic system toward dehydrogenation rather than dehydration. By contrast, applying the PNP-Ru complex Ru-2 resulted in low conversion of 1a and a marginal yield of amide 2a. Complex Ru-3, 4e,f which is an analog of Ru-1 with an NHBn group (Bn ¼ benzyl), exhibited the same alcohol esterication reactivity as Ru-1, but was inferior with respect to primary amide synthesis. 9 Finally, using the new complex Ru-4, which is a PPh 2 -substituted variant of Ru-1, afforded no observable amide. It should be noted that a bipyridine-based PNN-Ru pincer complex, previously reported to catalyze the formation of tertiary amides from alcohols and amines, 4c was also employed as a catalyst for the coupling of alcohols and ammonia, but only secondary amines and imines were obtained. 2d In contrast to previously reported catalytic reactions wherein organoamines were used as coupling partners, the current system showed no improvement in ester-to-amide conversion upon prolonging the reaction time, possibly because ammonia is not nucleophilic enough to amidate esters under the applied conditions. 10 We therefore undertook systematic optimization of the catalytic conditions in order to improve the yields of primary amides, using 1a as substrate and Ru-1 as catalyst. Firstly, the effect of base was evaluated. No reaction was observed in the absence of base, but when KOH was used, alcohol conversion and product yield were slightly lower than with t BuOK ( Table 2, entries 1-3). These ndings clearly indicate that the base is necessary for catalysis to occur. Moreover, when the amount of t BuOK was reduced from 2 to 1 equiv relative to the catalyst, conversion of alcohol decreased from 92 Table 1 Examination of potential catalysts a a Conditions: 1a (0.5 mmol), cat. (1 mol%), t BuOK (2 mol%), toluene (2 mL), NH 3 (7 bar), 135 C, 18 h. Conversions and yields were determined by NMR spectroscopy using benzyl benzoate as internal standard.
to 80% (entry 4). We also examined alternative solvents, namely, THF and hexamethyldisiloxane (HMDSO), but no improvement was observed in amide yield (entries 5 and 6). The yield of 2a was found to be inversely related to temperature. Thus, increasing the temperature from 135 to 150 C slightly decreased the yield from 33 to 30% (entry 7), whereas lowering the temperature to 120 C raised the yield of 2a to 40%, while decreasing the yield of ester 3a to 20% (entry 8). The improvement in amide yield at 120 C might be due to enhanced ammonia solubility upon cooling. 11a We reason that the concentration of ammonia in solution directly affects the formation of hemiaminal from the in situ-generated aldehyde, which is further inuenced by the competition between ammonia and alcohol, the latter being responsible for the ester byproducts. Based on these assumptions, and the fact that ammonia solubility varies with solvent, 11b the effect of solvent composition was investigated at 120 C. Since ammonia miscibility is usually higher in ethers or alcohols than in benzene-derived solvents, we studied equivolume mixtures of toluene with THF, t AmOH (Am ¼ amyl) or t BuOH. As shown in Table 2, of the three toluene/cosolvent mixtures examined (entries 9-11), the one involving t AmOH provided the best results, with an enhanced amide yield of 56%, and only 11% of the ester (entry 10). It should be noted that using pure t AmOH resulted in low conversion (40%), and therefore incorporating toluene is crucial for ensuring signicant conversion. Interestingly, doubling the volume of the toluene/ t AmOH mixture, while halving the concentration of substrate 1a, improved the yield of primary amide to 71%, whereas the yield of ester dropped to only 3% (entry 13).
Considering that combining Ru-1 with a base also affords an efficient hydrogenation catalyst, 12 we reasoned that under the current conditions, which involve the use of a sealed reaction vessel, the generated hydrogen gas that accumulates inside it might counteract the forward dehydrogenation reaction. Thus, a technically-modied procedure was employed, whereby the gaseous contents of the Fischer-Porter tube were vented aer 24 h of heating at 120 C, followed by relling the ask with 7 bar of ammonia, and heating for another 12 h. As expected, this improved the yield of 2a to 83%, while the yield of ester remained at 3% (entry 14). Repeating the same procedure, but raising the temperature to 150 C during the nal 12 h, increased the conversion to 95%, with 2a isolated in 87% yield (entry 15).
With the optimal catalytic conditions in hand, we explored the substrate scope of the amidation reaction catalyzed by Ru-1 (Table 3). Linear primary alcohols of different chain lengh, namely, 1-heptanol, 1-butanol, and 1-propanol, gave the respective products, 2b-d, in 81-87% isolated yield. The compatibility of the catalytic system with various functional groups was also examined. Increased steric hindrance near the -CH 2 OH moiety, i.e., replacing H by CH 3 at the b-position, decreased the yield of the target product (2e, 51%). By contrast, tertiary amine and methoxy groups were well-tolerated (2f, 81%; 2g, 85%). Remarkably, C]C double bonds of olenic substrates were also preserved despite the presence of H 2 in the reaction vessel. Thus, using 4-hexen-1-ol as starting material, a mixture of primary amides was obtained due to alkene isomerization, with the internal alkenes being the major products (2h, 61% in total). In the case of citronellol, the trisubstituted double bond was unaffected, and the corresponding primary amide (2i) was isolated in 61% yield. Benzyl alcohol and its para-methyl, -methoxy, -chloro and -uoro derivatives were also studied as substrates under the same reaction conditions, and the corresponding benzamides were produced in high yields without a signicant substituent effect (2j-n, 82-91%). Finally, in addition to alcohols, 3-phenylpropionaldehyde was also examined as a starting material, but its self-condensation was the primary outcome.

Mechanistic investigation
In an attempt to elucidate the underlying mechanism of the amidation process, we studied the roles played by the various components of the catalytic system, employing both experimental and computational tools. As part of these efforts, the organic species present in the catalytic reaction mixture were monitored by a set of parallel catalytic experiments, carried out under identical reaction conditions (as per entry 13, Table 2), but sampled at different time intervals. The time-variation of the residual alcohol substrate 1a, as well as the yields of primary amide 2a and ester 3a, are presented in Fig. 1. These data clearly indicate that alcohol consumption is correlated with amide formation, and that the reaction rate gradually decreases with time.
The ester byproduct reached a yield of 3% within the rst 2 h, and this remained roughly constant throughout the catalytic run. This cannot simply be ascribed to the gradual decrease in alcohol concentration, since an appreciable amount of alcohol remained available in the mixture well aer the initial 2 h, thereby allowing further esterication to take place. Instead, considering the known chemistry of Ru-1 and related complexes, it is likely that the continuous generation of ester was offset by its hydrogenation back to the alcohol, 12 induced by the accumulation of H 2 gas in the closed reaction vessel, or by direct amidation under the pressurized NH 3 . The ability of our catalyst to promote facile ester hydrogenation was demonstrated by the fact that hexanol was obtained in 27% yield aer heating hexyl hexanoate at 120 C under only $0.3 bar of H 2 ($20 mL H 2 , representing $80% H 2 yield in the actual catalytic  a Conditions: alcohol 1 (0.5 mmol), Ru-1 (2 mol%), t BuOK (4 mol%), toluene (4 mL), t AmOH (2 mL), NH 3 (7 bar), 120 C. Pressure was released aer 24 h at 120 C, and then NH 3 was relled to 7 bar, and the reaction mixture was heated at 150 C for 12 h. Residual alcohol substrates and homocoupled ester byproducts were detected in all cases.
reaction; Scheme 2a). By contrast, when the ester was exposed to 7 bar of NH 3 under similar conditions, but in the absence of H 2 , only trace amounts of the corresponding amide were detected by GC-MS aer completion of the catalytic procedure (Scheme 2b). In a third experiment, the ester was treated simultaneously with $0.3 bar of H 2 and 7 bar of NH 3 under the optimal catalytic conditions (as per entry 15, Table 2), and this afforded the amide in 48% yield (Scheme 2c). Furthermore, we exposed primary amide 2a to $0.3 bar of H 2 and 7 bar of NH 3 under similar catalytic conditions, but observed only trace amounts of alcohol ($1%) aer 18 h of heating (Scheme 2d). 13 Taken together, these results imply that the ester byproduct formed during the catalytic process is funneled back into the catalytic cycle by its hydrogenation into alcohol, rather than direct amidation. Moreover, while ester formation is reversible under the applied catalytic conditions, amide formation is virtually irreversible, eventually leading the primary amide to become the dominant product. To supplement the above experiments, which addressed the species observed within the catalytic reaction mixture, we also explored the involvement of nitriles. These have not been observed in the current system, but they are nonetheless plausible reaction intermediates, because they can be generated by hemiaminal dehydration into imines, followed by dehydrogenation, and may subsequently undergo rehydration to produce the primary amides. Moreover, the current catalyst has recently been shown to promote nitrile hydration into amides. 6d To probe this alternative pathway, 3-phenylpropionitrile was subjected to the catalytic conditions (as per entry 13, Table 2) in the presence of 1 equiv. of added water (Scheme 2e). However, primary amide 2a was generated in only 37% yield, compared to 71% when alcohol 1a was used as substrate, indicating that the rate of nitrile hydration is signicantly lower than that of amide generation from alcohol and ammonia under the same conditions. This result, combined with the fact that no nitrile was observed in the current catalytic process, clearly demonstrate that nitrile hydration is not the dominant reaction pathway (see ESI, Section 3.2, † for further details).
As described above, utilizing t AmOH as a cosolvent to toluene was found to substantially improve the yield of amide. In order to probe the role of t AmOH, a catalytic experiment was carried out, employing the usual toluene/ t AmOH solvent mixture and 1a as the substrate, but in the absence of ammonia (Scheme 2f). Despite the high concentration of t AmOH, only a small portion of it coupled with 1a to afford the tertiary ester 3b in 9% yield, whereas the major product, obtained in 39% yield, was ester 3a. In fact, under the optimal catalytic conditions, in the presence of ammonia, and regardless of the identity of the primary alcohol substrate, esters of t AmOH were not observed, showing that any involvement of this cosolvent as a substrate is not competitive with the amidation reaction. Furthermore, when the amidation of 1a was carried out aer replacing t AmOH by 1,4-dioxane, which is both less polar and aprotic, only a small decrease in the yield of 2a was observed (80%, with 6% of 3a). This demonstrates that polarity and hydrogen-bonding are not decisive factors in the ability of t AmOH to promote alcohol amidation.
The above results led us to surmise that t AmOH enhances the solubility of ammonia in the reaction medium, thereby allowing amidation to outcompete the reversible esterication reaction. In the absence of sufficient literature data for ammonia solubility in the relevant solvents, we assessed it semi-quantitatively at room temperature by exposing a given solvent to 10 bar of NH 3 , and calculating the concentration of ammonia from the pressure drop. The obtained data, compiled in Fig. 2, indicate correlation between the ability of a given solvent to dissolve ammonia and the amide yield obtained in that solvent. Thus, NH 3 is less soluble in toluene than in the 4 : 2 toluene/ t AmOH mixture used in our catalytic experiments, and this is consistent with the observed improvement in amide yield upon going from toluene to toluene/ t AmOH. On the other hand, HMDSO, in which the amide yield was poor, exhibits the lowest ammonia solubility of all examined solvents. Furthermore, the solubility of NH 3 in dioxane is only a little lower than in t AmOH, and the same is true of their respective mixtures with toluene, and this is in line with the relatively small drop in amide yield when t AmOH was replaced by dioxane. Thus, the above ndings support our assertion that using t AmOH as a cosolvent improves the yield of amide by increasing ammonia solubility.
In order to clarify those mechanistic aspects that directly involve the coordination sphere of Ru-1, we also explored possible catalytically-related species (Scheme 3). As previously reported, the catalytic cycle commences when Ru-1 undergoes base-induced dearomatization to afford the well-documented catalytically-active complex Ru-5 (Scheme 3a). 3a When a C 6 D 6 solution of this complex was exposed to 3 bar of NH 3 , the ammonia-adduct Ru-6 immediately formed, as evidenced by a clear color change from purple to brown. Removal of ammonia under reduced pressure regenerated Ru-5, showing that formation of Ru-6 is reversible (see Fig. S1-S7 † for details). Interestingly, Ru-5 was found to be inert toward excess t AmOH (180 equiv.) at room temperature, indicating that this cosolvent does not directly interfere with the catalytic cycle. By contrast, Ru-5 is highly reactive toward primary alcohols, affording the ruthenium-alkoxide species Ru-7 (Scheme 3b) by adding their Fig. 2 Ammonia concentrations in the catalytically-relevant solvents at room temperature, after initial introduction of 10 bar of NH 3 . A, toluene/ t AmOH, 4 : 2; B, toluene/1,4-dioxane, 4 : 2 (volumetric ratios, mL/mL).
O-H bond across the metal-ligand framework, 12e as presently demonstrated by the reaction of Ru-5 with 3 equiv of 1-butanol. The resulting n-butoxide complex is unstable at room temperature in C 6 D 6 , gradually converting into dihydride Ru-8 (ref. 12c) (major species), as well as aldehyde adduct Ru-9 (ref. 14) (minor species), which can also be obtained directly from Ru-5 and the aldehyde, and is proposed to be an off-cycle species. The very generation of Ru-8 and Ru-9 strongly indicates that Ru-7 undergoes facile b-hydride elimination to produce the aldehyde, which under the catalytic conditions would subsequently react with ammonia. Based on the above results, it is proposed that conversion of alcohol into aldehyde is not the ratedetermining step in this system, since it readily occurs at room temperature.
A plausible catalytic cycle for the transformation mediated by Ru-5 is outlined in Scheme 3c. This reaction pathway was investigated by density functional theory (DFT) calculations, based on previously-reported works, 15 with ethanol as a minimal alcohol model, and toluene as an implicit solvent (Fig. 3). Given that the alkoxide complex Ru-7 was observed to form under the applied catalytic conditions (Fig. S17 †), 16 and based on a report involving a comparable osmium-alkoxide pincer complex, 15c we propose that Ru-7 promotes the coupling of ammonia and aldehyde through a concerted outersphere transition state (TS 7,10 , Fig. 3). This leads to a hydrogenbonded intermediate Ru-10, which releases one molecule of alcohol to generate the hemiaminalate complex Ru-11. This, in turn, undergoes b-hydride elimination to afford Ru-8 and liberate the primary amide product. According to the computed energy prole, the conversion of aldehyde to primary amide is thermodynamically favored by 6.4 kcal mol À1 , with an overall kinetic barrier of 26.8 kcal mol À1 , corresponding to the alcoholassisted extrusion of H 2 from Ru-8. 17 This indicates that the generation of primary amides from aldehydes and ammonia is energetically highly-feasible in the current system, in agreement with our experimental ndings (see Fig. S57-S63 † for other computed pathways, such as those involving pincer ligand armopening 18 and ammonia activation 19 ).
In an attempt to further rationalize the selectivity of the current catalytic system toward primary amide formation, we carried out a computational study of the alternative imineproducing pathway, involving aldehyde and ammonia, which could lead to other N-containing compounds. 20 The resulting energy proles, which are presented in Fig. S64 and S65, † show that this pathway is thermodynamically uphill by 3.9 kcal mol À1 , with an overall kinetic barrier of 37.0 kcal mol À1 , which is substantially higher than the direct amidation pathway shown in Fig. 3. 21 Therefore, primary amide generation by the present PNN-ruthenium system is both kinetically and thermodynamically more favorable than imine formation, and this explains the high selectivity of this system toward the synthesis of primary amides rather than other N-containing products.

Conclusion
We have presented an environmentally-benign and highly selective synthesis of primary amides through an unprecedented ruthenium-catalyzed acceptorless dehydrogenative coupling of alcohols and ammonia. The yield of primary amides was found to depend strongly on ammonia solubility in the reaction medium, and utilizing a toluene/ t AmOH mixture allowed high amide yields to be achieved, with only marginal yields of ester byproducts. The PNN-ruthenium catalyst used in our system efficiently promotes the dehydrogenation of the ammonia-derived hemiaminal intermediate, rather than its dehydration, thereby selectively forming primary amides instead of other N-containing compounds, such as amines, imines or nitriles, which are typically generated by other catalytic systems. Based on DFT calculations, we attribute this selectivity to both kinetic and thermodynamic preference for the generation of primary amides over that of imines, which are also the source of other N-containing products.

Data availability
Experimental details, GC trace, compound characterization data, NMR spectra, computational details.

Conflicts of interest
There are no conicts to declare.


Bis[(4-methylphenyl)diphenylphosphine-κP](nitrito-κ2 O,O′)silver(I)

The title silver(I) diphenyl-p-tolylphosphine complex crystallizes with one complete molecule in the asymmetric unit that features a bidentate nitrito, as well as two diphenyl-p-tolylphosphine ligands coordinated to a AgI center.

The title Ag I complex, [Ag(NO 2 )(C 19 H 17 P) 2 ], reveals a distorted pseudotrigonal-planar shape around the Ag I atom geometry resulting from the coordination of two phosphine ligands, as well as a nitrito-O,O 0 ligand coordinating to the silver(I) atom through the oxygen atoms; in this description, the two oxygen atoms are assumed to occupy one position, forming an acute O-Ag-O angle of 51.44 (9) . The plane resulting from the NO 2 coordination to Ag is nearly perpendicular to the plane from the coordination of the phosphine-P atoms to Ag [dihedral angle = 86.43 (9) ].

Synthesis and crystallization
Diphenyl-p-tolylphosphine (1 mmol) was dissolved in acetonitrile (10 ml). Silver nitrite (1 mmol) was dissolved in acetonitrile (5 ml). The diphenyl-p-tolylphosphine solution (10 ml) was added to the silver nitrite solution (5 ml), to give a 2:1 molar ratio reaction. The mixture was heated under reflux for 2 h after which the solution was left to crystallize.

Refinement
For full experimental details including crystal data, data collection and structure refinement details, refer to Table 1.

Figure 2
Packing diagrams as viewed along the (a) a and (b) c axes. Hydrogen atoms are omitted for clarity.

Figure 1
Perspective view of the molecular structure of the title compound showing displacement ellipsoids at the 50% probability level. Hydrogen atoms are omitted for clarity.
Special details Geometry. All esds (except the esd in the dihedral angle between two l.s. planes) are estimated using the full covariance matrix. The cell esds are taken into account individually in the estimation of esds in distances, angles and torsion angles; correlations between esds in cell parameters are only used when they are defined by crystal symmetry. An approximate (isotropic) treatment of cell esds is used for estimating esds involving l.s. planes. Refinement. All H atoms were placed in geometrically idealized positions and constrained to ride on their parent atoms.