Review Reports - Nickel Complexes as Artificial Urea Lyases with 1012-Fold Rate Enhancement

Round 1

Reviewer 1 Report

Comments and Suggestions for Authors

The paper is devoted to the investigation of new polydentate ligands and nickel complexes with these ligands as the catalysts of urea decomposition. The catalytical cleavage of urea N-C bonds in relatively mild conditions was demonstrated that could be of future interest. Still the synthetic part is extremely speculative and become the main objection to the publication.

Despite the difference of side radicals all the ligands demonstrate the same coordination pattern binding to nickel. And the difference between complexes is hardly explained with the nature of side radicals. Especially in pairs Ni₂L3 and Ni₂L3``. Was both complexes prepared simultaneously? If so – how was the separation? If the complexes were prepared separately – no description of separate synthesis was presented. In pairs Ni₂L1 and Ni₂L2 the alkyl chain is the same except of the end and no electronic effects could be expected. Still the final products again differ in nickel coordination sphere. Why? Why complex Ni₂L1 with two chlorides and Ni₂L3 with one chloride have similar conductivity? Why complex 3 is non-electrolyte completely? If authors could not prepare crystals suitable for XRD the NMR of complexes could gave some answers about the structure. And simple Ag titration could provide the direct result on the chlorine content.

Before answering the questions above and clear experimental description the paper could not be published.

Author Response

We would like to thank the reviewer for the comments, as we believe they improved the manuscript. The complexes Ni₂L3 and Ni₂L3``were prepared with different protocols. We have added the following description to the manuscript “During the synthesis of the complexes containing the Phenyl-radical containing ligand, two distinct products were obtained in independent experiments. The complex isolated in the first synthesis was green, whereas the product obtained in the second synthesis was orange. Upon repeating the syntheses, it was observed that, despite using the same drying time, the color of the hydrated nickel chloride salt varied after vacuum drying. In some experiments, the salt became pale yellow, yielding a green complex. In others, the salt developed a yellow color with a slight orange tint, and the resulting complex was orange. These observations suggest that the extent of dehydration of the nickel salt may influence the color of the resulting complexes and, potentially, their coordination geometries, as water molecules can also coordinate to the metal center.

The hydrated nickel chloride salt was subjected to high vacuum and gentle heating using a heat gun to remove the water of hydration. Because the heating was not temperature-controlled, the actual temperature reached during this process was unknown. The dehydration of NiCl₂.6H₂O may occur in multiple stages. The first mass loss takes place between 24 and 120 °C and corresponds to approximately 30% of the initial mass, equivalent to the loss of four water molecules. Complete dehydration requires temperatures above 250 °C. Therefore, the nickel salt was likely not fully dehydrated under the conditions employed. Nevertheless, the differences in salt coloration observed during the syntheses of Ni₂L3 and Ni₂L3``suggest that the last may have involved a greater degree of dehydration. This difference in hydration state may have affected the coordination environment of the nickel centers, potentially accounting for the distinct colors observed in the resulting complexes.”

In addition to this discussion, we have added to the protocol a phrase indicating the color observed for the nickel chloride salt during the dehydration process.

In pairs Ni₂L1 and Ni₂L2 the alkyl chain is the same except of the end and no electronic effects could be expected. Still the final products again differ in nickel coordination sphere. Why? Why complex Ni₂L1 with two chlorides and Ni₂L3 with one chloride have similar conductivity? Why complex 3 is non-electrolyte completely? If authors could not prepare crystals suitable for XRD the NMR of complexes could gave some answers about the structure. And simple Ag titration could provide the direct result on the chlorine content.

We thank the reviewer for these important comments. The different numbers of chloride ligands proposed for these complexes are based primarily on the elemental analysis. The CHN data are consistent with the respective molecular formulations containing the different numbers of chloride ligands, despite the structural similarity of the corresponding ligands. We therefore do not attribute this difference to a specific electronic effect of the terminal substituent, but rather report the formulations that are consistent with the experimentally determined compositions.

The conductivity measurements provide additional information regarding the behavior of these complexes in solution. In dichloromethane, a non-coordinating solvent, the complexes behave as non-electrolytes, consistent with the absence of dissociated counterions. In contrast, increased conductivity is observed in methanol, a coordinating solvent, consistent with solvent-induced displacement of labile coordinated ligand(s) and formation of ionic species in solution. Consequently, similar conductivity values for complexes containing different numbers of chloride ligands are not necessarily contradictory, since conductivity reflects the solution species generated upon dissolution rather than directly reporting the total number of chloride ligands present in the isolated compound.

We also attempted NMR characterization of the complexes. However, consistent with the magnetic susceptibility measurements (Figure 2), the complexes do not behave as straightforward diamagnetic Ni species. Their NMR spectra displayed substantial broadening and chemical-shift changes that prevented reliable assignment, and therefore these data could not provide meaningful additional structural information. We have added the 1H NMR spectra of all four complexes to the SI (S52). Since complex (Ni,Ni)L3” is significantly less paramagnetic, its spectrum is better resolved, and we could attempt a signal assignment (S53).

Finally, addition of AgNO₃ to the complexes results in formation of a white precipitate, consistent with the presence of chloride that can become labile in solution. Taken together, the elemental analysis, mass spectrometry, magnetic susceptibility, solvent-dependent conductivity, and qualitative AgNO₃ test provide complementary evidence supporting the proposed compositions and the presence of labile chloride ligands.

Before answering the questions above and clear experimental description the paper could not be published.

Author Response File: Author Response.docx

Reviewer 2 Report

Comments and Suggestions for Authors

The peer‑reviewed manuscript “Nickel Complexes as Artificial Urea Lyases with 10¹²‑Fold Rate Enhancement” by Letícia K. Sebastiany et al. is devoted to the development of a series of binuclear nickel complexes based on bis‑triazole ligands, which act as artificial urea lyases — catalysts for the decomposition of urea via the β‑elimination pathway, leading to the formation of isocyanate and ammonia. The authors report achieving a reaction acceleration of up to 10¹² times compared to the uncatalyzed process, which, according to them, is a record value for abiotic ureases and approaches the efficiency of natural enzymes. Remarks

The article is devoted to complexes, but the characterization is incomplete. The composition of the complexes has been confirmed using indirect methods (HRMS (MALDI-TOF), UV-Vis, conductometry, magnetic measurements, thermogravimetry, FTIR), but not their structure, which can be unambiguously confirmed only by X‑ray crystallography. If the authors were unable to grow single crystals for X‑ray crystallography, it is necessary to add a discussion of thermal analysis, correlating mass loss with the organic component.

The abstract is too general. It needs to be shortened and more data on nickel complexes and ligands should be added.
In the experimental part, the authors write about the dehydration of the initial nickel chloride and conducting the reaction under dry conditions, yet the diagram shows an aqueous chloride, and the complex also contains coordinated water and hydroxyl groups. This discrepancy needs to be explained.

In the description of the structure of the complexes, it is also necessary to provide data on the structurally confirmed coordination of similar ligands in nickel complexes from the literature.
The limited evidence of inner‑sphere coordination.

FTIR data on urea coordination under anhydrous conditions (band at 1770 cm⁻¹) and in experiments with photolysis of “caged” urea (band at 1640 cm⁻¹) are interpreted as evidence of urea binding via nitrogen and via carbonyl oxygen, respectively. These two coordination modes differ, and the authors do not discuss which one is realized under catalytic conditions. Furthermore, the bands in the region of 1500–1650 cm^-1may overlap with ligand vibrations, which makes unambiguous assignment difficult.

The question of catalyst stability. The paper does not discuss whether the complex retains its structure during catalysis, whether nickel precipitates, or whether nanoparticle formation occurs. The control with NiCl₂ did not result in ammonia formation, which is important, but it does not rule out the formation of other active forms. Heterogeneity tests (for example, a filtration test) would be appropriate.

Comparison with the literature. The statement about “the highest rate enhancement reported for an abiotic urea lyase” requires a more detailed comparison with previous works. It would be convenient to present a table comparing the conditions, substrates, k₂, and accelerations for known systems. Formatting and technical details.

In the text there are typos and inconsistencies: "respectivelly", "comparisson", "stoped"

The numbering of the compounds (L1, L2, L3, L3") is not always consistent; in the experimental part for the (Ni,Ni)L1 given molecular formula C56H75Cl2N8Ni2O7, whereas in HRMS indicated C75H75N10Ni2O7 is a different formula that requires explanation.

The caption for Figure 3 states “7.5 % of water” in one place and “15.8 % water” in another — this needs to be reconciled..

The manuscript requires significant revision.

Author Response

The article is devoted to complexes, but the characterization is incomplete. The composition of the complexes has been confirmed using indirect methods (HRMS (MALDI-TOF), UV-Vis, conductometry, magnetic measurements, thermogravimetry, FTIR), but not their structure, which can be unambiguously confirmed only by X‑ray crystallography. If the authors were unable to grow single crystals for X‑ray crystallography, it is necessary to add a discussion of thermal analysis, correlating mass loss with the organic component.

First of all we would like to thank you for the comment, as we believe it helped us improve the manuscript. We have added a discussion of the thermal analysis to the Results section: “The thermogravimetric profiles and residual masses are consistent with the formulations derived from elemental analysis, supporting the different ligand and chloride contents assigned to the complexes (Table 1). The mass loss between 35 and 120 °C is attributed to loosely bound solvent: two ethanol molecules for (Ni,Ni)L1 (7.1%), six water molecules for (Ni,Ni)L2 (9%), and three water molecules for (Ni,Ni)L3 (5%) and (Ni,Ni)L3” (4%). Between 120 and 280 °C, more tightly bound water and chloride are lost. (Ni,Ni)L1 loses three water molecules (4.3%, 120–200 °C), followed by 2 Cl⁻, one OH⁻ and one H₂O (8%, 200–260 °C). (Ni,Ni)L2 loses three H₂O and three Cl⁻ (12%), (Ni,Ni)L3 loses two OH⁻ and one Cl⁻ (6%), and (Ni,Ni)L3” loses three Cl⁻ (9%). At higher temperatures, the organic ligand decomposes (55–75% of the mass), leaving residues of about 21% for (Ni,Ni)L1, 24% for (Ni,Ni)L2, 14% for (Ni,Ni)L3, and 15% for (Ni,Ni)L3”. Because thermal analysis cannot establish connectivity, these data support the composition of the complexes but not their structure.

Table 1. Tentative assignment of the TGA/DTG experiment for each complex.

Complex

Weight loss

Temperature range

Assignment

(Ni,Ni)L1

»7.1%

»4.3%

»8%

»30%

35-120°C

120-200°C

200-260°C

260-380°C

380-800°C

1 EtOAc

3 H₂O

2Cl^-+ OH^- + H₂O

Ligand

(Ni,Ni)L2

»9%

»12%

»30%

»25%

35-120°C

120-280°C

280-380°C

380-800°C

6 H₂O

3H₂O + 3Cl^-

Ligand

(Ni,Ni)L3

»5%

»6%

»44%

»31%

35-120°C

120-220°C

220-380°C

380-800°C

3 H₂O

2OH^- + Cl^-

Ligand

(Ni,Ni)L3”

»4%

»9%

»42%

»30%

35-120°C

120-280°C

280-360°C

360-800°C

3 H₂O

3 Cl^-

Ligand

The abstract is too general. It needs to be shortened and more data on nickel complexes and ligands should be added.

Thank you for your comment. We have modified the abstract and believe that it is now less general.

In the experimental part, the authors write about the dehydration of the initial nickel chloride and conducting the reaction under dry conditions, yet the diagram shows an aqueous chloride, and the complex also contains coordinated water and hydroxyl groups. This discrepancy needs to be explained.

Thank you for your comment. We have added this sentence to the manuscript: Even though the reactions were carried out under anhydrous conditions, all isolated complexes contain water according to TGA, indicating that water is retained or taken up during isolation and purification, and that residual hydration of the nickel salt may contribute to the variability between batches.

In the description of the structure of the complexes, it is also necessary to provide data on the structurally confirmed coordination of similar ligands in nickel complexes from the literature.

Thank you for your comment. We have added this paragraph in the discussion of the complexes structure. “Dinuclear nickel(II) complexes supported by phenolate-bridged compartmental ligands have been structurally characterized extensively as urease models, typically showing hexacoordinate high-spin Ni(II) centers with Ni···Ni distances of 2.9-3.5 Å, a μ-phenolate bridge, and additional exogenous bridges or terminal solvent and chloride ligands [37] [38] [39] [40]. Likewise, 1,2,3-triazoles coordinate Ni(II) through N[3] with Ni-N distances of 2.0-2.2 Å [41] [42]. However, we have evidenced the N[2] coordination in similar mononuclear triazole complexes. [43] Our spectroscopic, magnetic, and conductometric data are consistent with this coordination mode: the Ni(II) centers are high-spin and octahedral, the triazole bands in the FTIR shift upon coordination (ex: 1675cm^-1 to 1653 cm^-1), and the chloride content determined by elemental analysis and conductivity agrees with the proposed chlorides in the inner sphere. Without a crystal structure, we propose the structures in Scheme 1 on this basis.”

The limited evidence of inner‑sphere coordination.

Thank you for your comment. We have added a solution-based experiment under anhydrous conditions using caged urea and the control of caged-urea only assay in 100% acetonitrile. Analysis of the result with these new data reveal that urea coordinates to the nickel complex in anhydrous solutions. However, the assay in ACN/H2O mixture only evidences urea consumption. We thus changed the manuscript accordingly:

To determine whether urea coordination to the complex could also be detected in solution, we investigated the reaction between the complex and a ruthenium-caged urea[52], which upon photolysis generates free urea. As acetonitrile is a coordinating solvent, it can easily displace the coordinated urea, which results in not optimal dark subtractions. However, it is evident that the control reaction of the caged urea in acetonitrile reveals the emergence of bands at 1728, 1602, 1517 and 1497 cm^-1, which can be attributed to δ_s NH₂, ν(C=O), [53] bpy (ν(C=N) and ν(C=C) stretching modes) [54] and ν(C-N) [53]. The blueshift of the carbonyl bands from the solid state is mostly due to the disruption of standard hydrogen-bonding aggregation. [55] In the presence of the complex ((Ni,Ni)L3) we can analyze the bands that emerge relative to the control experiment performed in the absence of caged urea. In an anhydrous condition, the 1728cm^-1 and the 1497 cm^-1 bands are easily visible, revealing that free urea is in solution. Interestingly, bands at 1702, 1636 and 1594 cm^-1 are distinct from the complex-only control. These bands could be related to urea coordination to the complex in anhydrous conditions. For instance, O-coordination normally weakens the C=O bond and shifts ν(C=O) to lower wavenumber(1728"1702 cm^-1), while the C-N/NH₂ modes shift higher (1602"1636 cm^-1). [37] An unambiguous assignment is not possible, however, because the 1636 cm^-1 band could also arise from release of water from the complex.

In the presence of water, all urea-related bands except 1728 cm^-1 appear as negative features, indicating that urea is consumed. The absence of a negative band at 1728 cm^-1 may reflect overlap with a new band at a similar wavenumber. The most prominent new band is at 1640 cm^-1. Although this is consistent with O-bound urea at nickel[56, 57], it more likely reflects water release (δ(H2O) falls in this region), given that the complex is highly active in water and no intermediate could be isolated at room temperature. We were unable to detect isocyanate or other potential reaction intermediate because the strong absorption of acetonitrile obscures the nitrile/cyanate region of the spectrum. These solution data therefore support that in anhydrous conditions, urea coordinates to the nickel centers, corroborating to the solid-state assay.

Figure 5. Photolysis experiments of Ru-caged urea and (Ni,Ni)L3 using a diode laser (l- 532nm). The Ru-caged urea[52] (A) and the complex (B) were dissolved in acetonitrile. The anhydrous reaction between photocleaved urea and the complex is shown in (C), whereas the reaction in the presence of water is shown in (D). Conditions: Caged urea (10mM), complex (5mM). The dark spectrum is the subtraction of a spectrum in the dark versus the initial spectrum before photolysis. The time os irradiation of each experiment is shown in the insets.

The question of catalyst stability. The paper does not discuss whether the complex retains its structure during catalysis, whether nickel precipitates, or whether nanoparticle formation occurs. The control with NiCl₂ did not result in ammonia formation, which is important, but it does not rule out the formation of other active forms. Heterogeneity tests (for example, a filtration test) would be appropriate.

We thank the reviewer for raising the question of catalyst stability and agree that the NiCl₂ control excludes free Ni(II) but not other nickel-containing species. We would like to point out some important information that we have added to the manuscript as the following “Before examining the identity of the catalytically competent species, we considered whether the observed activity could instead arise from heterogeneous nickel species. The strong ligand dependence of the catalytic activity argues against a common heterogeneous nickel species, such as Ni(0) nanoparticles or Ni(OH)₂ precipitates, as the origin of the observed reactivity. First, the pH changes by only ~0.5 unit during the reaction, owing to the equilibrium between isocyanate (pKa = 3.5) and ammonium (pKa = 9.35). Second, the reaction proceeds at room temperature, well below the hydrothermal conditions typically associated with Ni(OH)₂formation. Additionally, no reducing agent is present, making the formation of Ni(0) particles unlikely. Lastly, catalytic activity can be restored by addition of further urea solution, indicating that the active species persists in solution.”

Comparison with the literature. The statement about “the highest rate enhancement reported for an abiotic urea lyase” requires a more detailed comparison with previous works. It would be convenient to present a table comparing the conditions, substrates, k₂, and accelerations for known systems.

Thank you for the suggestion. We have softened our conclusion on rate enhancement and added a table with literature data.

Formatting and technical details.

In the text there are typos and inconsistencies: "respectivelly", "comparisson", "stoped"

Fixed.

Added “The higher mass observed for the complexes is due to the coordination of the matrix to the complex.”

The caption for Figure 3 states “7.5 % of water” in one place and “15.8 % water” in another — this needs to be reconciled..

Fixed.

The manuscript requires significant revision.

Author Response File: Author Response.docx

Reviewer 3 Report

Comments and Suggestions for Authors

This is a very interesting paper on the synthesis of three triazole ligands and their complexation to Ni(II) to form four novel dimeric complexes. The ligands are well characterized using standard techniques, while the paramagnetic nickel complexes are characterized using magnetic susceptibility measurements, high resolution mass spectrometry, elemental analysis, conductivity and UV-vis spectroscopy. The complexes acted as highly efficient artificial urea lyases catalyzing the decomposition of urea to ammonia and isocyanate. Complexes (Ni,NI)L1 and (Ni,Ni)L3 exhibited very high efficiencies for urea-decomposition, amongst the highest values reported for synthetic catalysts. Interestingly (Ni,Ni)L2 and (Ni,NI)L3” showed lower catalytic activity, indicating that moderate changes around the coordination sphere of the Ni(II) centres can have significant effect on how the complexes behave as catalysts. The authors showed that water in the solution was important in the formation of the active catalyst as no ammonia was formed when the reaction was carried out under anhydrous conditions. Studies were carried out to understand the speciation of the complexes upon addition of water using conductivity measurements and HRMS. The work is novel, well studied and the findings will be of interest to researchers working in this area.

My main query would be concerning the FTIR data as evidence for the coordination of urea to (Ni,Ni)L3. This reaction was carried out in acetonitrile. No information as far as I could see is given on the reaction conditions e.g. mol ratios of the reactants, or workup to remove unreacted urea. No spectrum is given for urea itself in KBr. Is all unreacted urea removed from the sample and how was that confirmed? KBr studies can only give qualitative information, as the amount of compound and the pathlength of the pellet will vary each time. Why did the authors not do a solution-based study like the one given in Figure 5 recording spectra in absorbance mode and creating difference spectra?

For the experiment given in Figure 5 it would be helpful to provide concentrations, and mol ratios for the reactants. Also, I think the authors need to show a few more control experiments. What did the spectra look like for the caged urea irradiated without (Ni,Ni)L3? I think it would be useful to give ordinary FTIR spectra in absorbance mode for the region of interest for each of (Ni,Ni)L3, the (RuBpy-urea) and Urea. This would allow the reader to more confidently assign the negative features observed in the difference spectra, and determine which species are decreasing in concentration. I assume that the spectrum designated as "dark" (t = 0) in Figure 5A, and the corresponding t = 0 spectrum (not currently shown) in Figure 5B, should be essentially featureless, since they represent the subtraction of a spectrum from itself. If this is not the case, additional explanation would be helpful.

The authors did an interesting study using HRMS to look for different species when (Ni,Ni)L1 was in the presence of water. Was there reason that they could not do a similar study on the complexes in acetonitrile upon the addition of urea to directly look for urea coordination? Such data would provide more direct evidence for the coordination of urea than the FTIR data alone.

Author Response

We thank the reviewer for this comment. We have added the synthetic conditions (urea:complex ratio, solvent, reaction time) and the workup procedure to the SI, together with reference spectra of free urea and of (Ni,Ni)L3 recorded as KBr pellets. The absence of a band at 1770 cm-1 in the free urea sample and in free (Ni,Ni)L3 confirms that this band is due to urea-complex interaction. We agree that KBr measurements are not suited to quantitative comparison, since intensities depend on pellet thickness and loading. Our conclusion, however, rests on the position of the bands, which does not depend on these factors.

As suggested by the reviewer we have added the control experiment of caged urea and the anhydrous reaction.

“To determine whether urea coordination to the complex could also be detected in solution, we investigated the reaction between the complex and a ruthenium-caged urea[52], which upon photolysis generates free urea. As acetonitrile is a coordinating solvent, it can easily displace the coordinated urea, which results in not optimal dark subtractions. However, it is evident that the control reaction of the caged urea in acetonitrile reveals the emergence of bands at 1728, 1602, 1517 and 1497 cm^-1, which can be attributed to δ_s NH₂, ν(C=O), [53] bpy (ν(C=N) and ν(C=C) stretching modes) [54] and ν(C-N) [53]. The blueshift of the carbonyl bands from the solid state is mostly due to the disruption of standard hydrogen-bonding aggregation. [55] In the presence of the complex ((Ni,Ni)L3) we can analyze the bands that emerge relative to the control experiment performed in the absence of caged urea. In an anhydrous condition, the 1728cm^-1 and the 1497 cm^-1 bands are easily visible, revealing that free urea is in solution. Interestingly, bands at 1702, 1636 and 1594 cm^-1 are distinct from the complex-only control. These bands could be related to urea coordination to the complex in anhydrous conditions. For instance, O-coordination normally weakens the C=O bond and shifts ν(C=O) to lower wavenumber(1728"1702 cm^-1), while the C-N/NH₂ modes shift higher (1602"1636 cm^-1). [37] An unambiguous assignment is not possible, however, because the 1636 cm^-1 band could also arise from release of water from the complex.

We have added these information to the caption of the figure.

We thank the reviewer for this suggestion. We attempted ESI-HRMS of (Ni,Ni)L3 in acetonitrile after addition of urea. Unfortunately, we do not see any changes in mass spectrum that could be associated with urea coordination. We attribute this to labile binding and possibly in-source dissociation due to the competition by acetonitrile. For example, we have shown that the caged-urea in acetonitrile suffers ligand substitution quite easily (Figure 5). We have added this data to the SI.

Author Response File: Author Response.docx

Round 2

Reviewer 1 Report

Comments and Suggestions for Authors

All my comments were taken into discussion. Thank you.

Reviewer 2 Report

Comments and Suggestions for Authors

The authors responded to the comments and made the necessary revisions. The article may be accepted.