Thiocarbohydrazone-Derived Metal-Free Electrocatalyst for Hydrogen Evolution In An Aqueous Medium

Summary

This protocol presents a reliable and reproducible protocol for designing, preparing, and evaluating an electrocatalyst for the hydrogen evolution reaction (HER). It enables undergraduate, postgraduate, and Ph.D. students to connect synthetic chemistry, electrode fabrication, and the application of electrode fabrication in catalytic electrochemical hydrogen production in a neutral aqueous medium.

Abstract

Most electrocatalysts reported for the hydrogen evolution reaction (HER) are based on transition-metal composites, nanoparticles, metal–organic frameworks, alloys, sulfides, and metal–nitrogen–carbon frameworks. These materials often require harsh synthetic conditions and generally perform optimally in acidic or alkaline media. In contrast, organic electrocatalysts for HER are relatively rare and are predominantly reported to function under acidic conditions. Thiocarbohydrazones (TCHs) constitute an important class of compounds with applications in coordination and bioinorganic chemistry. The presence of proton-donating and proton-accepting sites, such as NH, SH, and C=N groups, enables them to act as proton relays. Consequently, TCH derivatives can be considered promising candidates for electrocatalytic hydrogen evolution. This work reports the three-step synthesis and characterization of a thiocarbohydrazone derivative, P₃. The synthetic procedure is straightforward, affords high yields, and does not require expensive reagents, making it well-suited for educational demonstrations. All compounds were characterized by ¹H-NMR, HRMS, UV-Visible, fluorescence, and FTIR spectroscopy. The HER activity of P₃ was evaluated by immobilizing it on a 3 mm glassy carbon electrode (GCE) at a loading of 0.4 mgcm⁻² using perfluorosulfonic acid ionomer binder. Electrochemical measurements were carried out in 100 mM phosphate buffer (pH 7.4) using an Ag/AgCl reference electrode. Linear sweep voltammetry of the P₃-modified GCE displayed a pronounced cathodic response at −1.5 V, corresponding to hydrogen evolution, which was confirmed by gas chromatographic analysis. Chronoamperometric measurements performed for more than 2 h showed a stable current response, confirming the stability of P₃ on the GCE surface. These results demonstrate the potential of organic thiocarbohydrazone derivatives as promising electrocatalysts for HER in a neutral aqueous medium.

Introduction

The continuous growth in global energy demand has intensified the need for sustainable, renewable energy technologies to reduce dependence on fossil fuels and mitigate their environmental impact¹. Consequently, significant research efforts are being directed toward the development of environmentally benign energy sources and efficient energy storage technologies^2,3. Among the various alternatives, hydrogen (H₂) has emerged as one of the most promising sustainable energy sources due to its high energy density and pollution-free combustion, which produces only water as a by-product^4,5. Hydrogen can be generated through electrocatalytic or photocatalytic processes, both of which require efficient catalysts to lower activation barriers and enhance the overall rate of hydrogen production⁶.

Nature provides an excellent blueprint for hydrogen evolution catalysis via hydrogenase enzymes, which contain either [NiFe] or [FeFe] active sites; [FeFe]-hydrogenases exhibit exceptional catalytic activity for the hydrogen evolution reaction (HER) at remarkably low overpotentials^7,8,9. Structural studies based on X-ray crystallography have provided valuable insights into the origin of this high catalytic efficiency under ambient conditions^10,11. It has been established that the diiron active site is connected to a 4Fe-4S cluster through a cysteine thiolate bridge. During H₂ generation, the 4Fe-4S cluster serves both as an electron reservoir and as a source of electrons required for catalysis. Furthermore, the dithiolate bridge linking the two iron centers contains an amine functionality that promotes rapid proton-coupled electron transfer (PCET)^7,8,9, resulting in high turnover frequencies ranging from 6000–21,000 s⁻¹. Inspired by these structural and functional features, numerous synthetic mimics of the hydrogenase active site have been developed following the elucidation of the H-cluster structure^12,13,14. The importance of proton relays in facilitating efficient HER has also been demonstrated in several synthetic molecular catalysts. DuBois and co-workers¹⁵ established the critical role of pendant amine groups in achieving high catalytic turnover frequencies in a family of Ni(II)-based HER catalysts. Similarly, Sakai and co-workers¹⁶ reported that the [Ni(dcpdt)₂]²⁻ catalyst contains four nitrogen donor atoms that function as proton relay sites, lowering the HER overpotential to 330–400 mV at pH 4–6. These proton-relay functionalities complement electron-storage features in a manner analogous to the proton-transfer and electron-reservoir components found in hydrogenase enzymes^10,11. These studies have demonstrated that HER overpotentials can be systematically lowered by modulating PCET-driven ligand reduction potentials in a series of Ni(II) complexes¹⁷.

Most HER electrocatalysts reported to date are based on noble metals such as platinum, palladium, ruthenium, rhodium, and iridium, as well as non-noble transition metals possessing favorable hydrogen adsorption energies. These catalysts have been developed in various forms, including nanoparticles, alloys, metal–organic frameworks, composites, selenides, nitrides, carbides, sulfides, phosphides, and metal–nitrogen–carbon materials^18,19,20. Although many of these systems exhibit excellent electrocatalytic activity, their widespread application is often limited by high cost, complex synthetic procedures, and restricted operational stability. Moreover, their performance is generally optimized in strongly acidic or alkaline media, whereas HER under neutral conditions remains comparatively less explored. These limitations have stimulated interest in the development of metal-free organic electrocatalysts, which offer advantages such as structural tunability, mild synthetic procedures, lower cost, and environmental compatibility.

Among metal-free systems, organic hydrides have attracted considerable attention as potential electrocatalysts for hydrogen generation. It has been proposed that hydrogen atoms capable of exhibiting hydride-donor character under suitable conditions may facilitate HER through heterogeneous pathways. Natural redox processes frequently employ organic hydride donors such as flavin adenine dinucleotide (FADH₂) and nicotinamide adenine dinucleotide phosphate (NADPH)²¹, providing inspiration for the development of metal-free HER catalysts. More recently, frustrated Lewis pairs comprising sterically hindered proton and hydride acceptors have emerged as a new class of organic systems capable of activating hydrogen²².

Several classes of metal-free organic materials have been investigated for HER, including an organic hydride of 4,4′-dihydropyridine-1,1′-(2,6-dimethyl-1,4-dihydropyridine-3,5-diyl)bis(ethan-1-one)²³, hydrogen-bonded organic frameworks (HOFs)^24,25 and covalent organic frameworks (COFs)²⁶. In addition, a variety of small organic molecules, such as benzothiadiazole²⁷, benzodithiazole-caffeine²⁸, benzoselenonates²⁹, 9-phenyl-10-methylacridinium iodide³⁰, 2,4,6-triphenylpyridine derivatives³¹, phosphorus-based corrole complexes³², 2,2′-bipyridyl³³, 4,4′-bipyridyl³⁴, phenanthroline³⁵, and bis-imidazolium-embedded heterohelicene³⁶ have been explored as homogeneous HER catalysts. While these systems demonstrate catalytic activity under acidic or basic conditions, reports of immobilized metal-free organic electrocatalysts operating efficiently in neutral aqueous media without the addition of external acids or bases remain scarce.

The design of new organic electrocatalysts can benefit from structure-activity relationships established for existing HER systems, particularly regarding the presence of protonatable and deprotonatable functional groups that facilitate proton transfer and PCET processes. Thiocarbohydrazide (TCH), thiosemicarbazide (TSC), and their derivatives represent a promising class of compounds in this context (Figure 1).

Figure 1: Basic structures of thiosemicarbazone (TSC) and thiocarbohydrazone (TCH).​ Representative structures of TSC and TCH frameworks showing schematic sites for protonation and deprotonation. Please click here to view a larger version of this figure.

These molecules contain, C=N, S-H and multiple N-H functionalities that can participate in proton-transfer reactions and proton-relay mechanisms^17,37. Furthermore, the free amino groups present in TCH and TSC enable facile structural modification, allowing the introduction of additional proton-donating and proton-accepting sites. Such modifications may influence ligand-centered reduction processes and provide a means to tune HER overpotentials.

To date, TSC-based complexes of Co(II), Ni(II), Cu(II), Zn(II), and Pd(II) have been investigated as HER catalysts primarily in non-aqueous media and generally require the addition of external acids or bases^{38,39,40,41,42}. In contrast, studies exploring metal-free thiocarbohydrazone derivatives as electrocatalysts for HER under neutral aqueous conditions, particularly in immobilized form, remain largely unexplored.

In the present work, the synthesis and characterization of the thiocarbohydrazone derivative P₃ are reported. The intrinsic electrochemical behavior of P₃ was investigated using cyclic and linear sweep voltammetry, revealing distinct redox processes. To evaluate its catalytic potential toward HER, P₃ was immobilized in a perfluorosulfonic acid ionomer binder and deposited onto a glassy carbon electrode surface to enable heterogeneous electrocatalysis under neutral aqueous conditions. The observed catalytic activity and stability of P₃ in the absence of an external proton source demonstrate the potential of thiocarbohydrazone derivatives as a promising class of metal-free organic electrocatalysts for hydrogen evolution.

Protocol

1. Synthesis of thiocarbohydrazone derivative P₃

1. Perform the synthesis of thiocarbohydrazone derivative P₃ in a laboratory safety hood. Arrange the experimental setup as shown in Figure 2.
2. Synthesis of (E)-N′-(pyridine-2-ylmethylene) hydrazinecarbothiohydrazide (P₁)
   1. Weigh about 100.00 mg thiocarbohydrazide in a clean 50 mL beaker and dissolve it in 10 mL of distilled ethanol.
   2. Weigh 100.90 mg 2-pyridyl carboxaldehyde, transfer it to a 100 mL round-bottom flask, and dissolve it completely in 5 mL of distilled ethanol.
   3. Add one drop of concentrated hydrochloric acid to the above ethanolic solution and stir it for one min using a magnetic stir bar.
   4. Carefully add the previously prepared ethanolic solution of thiocarbohydrazide to the solution prepared in step 1.2.3 using a glass dropper and reflux the reaction mixture for 3 h in an oil bath. Cool it to room temperature.
   5. Filter the pale-yellow solid obtained through qualitative filter paper, Grade 1, wash twice with 5–7 mL distilled ethanol and once with diethyl ether. Recrystallize the crude product from methanol and dry under vacuum. Yield: 90%
3. Synthesis of (E)-N′-((E)-2-hydroxybenzylidene)-2-(pyridine-2-ylmethylene) hydrazine-1-carbothiohydrazide (P₂)
   1. Weigh about 100.00 mg of the recrystallized dry P₁ in a clean 50 mL beaker and dissolve it in 10 mL of distilled ethanol.
   2. Take 62.50 mg salicylaldehyde in a 100 mL round-bottom flask equipped with a stir bar and dissolve it in 5 mL of distilled ethanol with continuous stirring on a magnetic stirrer.
   3. Add one drop of concentrated hydrochloric acid to the solution in a round-bottom flask and stir it for 1 min.
   4. Carefully add the solution of P₁ to the acidified solution of salicylaldehyde in the round-bottom flask using a glass dropper and reflux the reaction mixture for 3 h.
   5. Separate the lemon-yellow precipitate from the reaction mixture by filtration through qualitative filter paper, Grade 1. Wash the precipitate twice with 5–7 mL of distilled ethanol and once with diethyl ether. Recrystallize the crude product from methanol and dry under vacuum. Yield: 80%
4. Synthesis of (E)-N′-((E)-2-((4-nitrobenzyl)oxy)benzylidene)-2-(pyridine-2-ylmethylene) hydrazine-1-carbothiohydrazide (P₃)
   1. Weigh 100.00 mg of the recrystallized, dry P₂ and dissolve it in 50 mL dry acetone. Add 46.15 mg potassium carbonate to it and stir the solution for 10 min at room temperature.
   2. Subsequently, add 25.00 mg sodium iodide to the above stirring solution, followed by 72.15 mg p-nitrobenzyl bromide (PNB).
   3. Reflux the above reaction mixture for 4 h. Bring it to room temperature, then filter the reaction mixture to remove the potassium carbonate precipitate. Collect the filtrate in a round-bottom flask and concentrate under reduced pressure to obtain crude P₃.
   4. Purify the crude product by column chromatography using silica gel as a stationary phase and n-hexane:ethyl acetate (1:1; v/v) as a mobile phase. Collect 10 mL fractions eluted from the column into 15 mL test tubes.
   5. Concentrate the solution from each test tube to approximately 2 mL and place a small drop from each test tube on a precoated silica gel plate with a capillary, along with the spots of P₂ and PNB (Figure 3). After spotting, carefully place the TLC plate in a TLC chamber containing n-hexane/ethyl acetate (1:1, v/v) as the mobile phase, holding it with tweezers. Cover the chamber. Allow the solvent to rise to three-quarters of the height of the TLC plate. Mark the solvent front with a pencil, dry the plate with a hot-air gun, and visualize under the UV visualization chamber.
   6. Combine all the fractions containing only the product and concentrate under vacuum to obtain a pale-yellow pure product. Yield: 70%.

Figure 2: Reaction setup in a safety hood. Schematic representation of the reaction arrangement inside a safety hood used for synthetic reactions. Please click here to view a larger version of this figure.

Figure 3: Illustration of a TLC plate. Schematic representation of a TLC plate showing the separation of P₃ from impurities formed during the reaction. Please click here to view a larger version of this figure.

2. Investigating solution-state electrochemical behavior of P₃ in dry DMSO solvent

1. For electrochemical measurements, keep the following items ready as suggested.
   1. A clean glassy carbon electrode as a working electrode, Ag/AgCl as the reference electrode, platinum wire as the counter electrode, and O-rings to hold electrodes.
   2. An electrochemical cell with a fitting lid having holes for holding electrodes and purging with N₂ gas. Fix the cell to a stand with a clamp.
   3. A nitrogen cylinder with a regulator, glass dropper, and silicone tubing.
   4. Prepare a potentiostat and connecting wires with crocodile clips.
2. A supporting electrolyte solution
   1. Weigh 3.87 g of tetrabutylammonium hexafluoro phosphate (TBAPF₆) and transfer it to a clean 100 mL volumetric flask with a stopper.
   2. Carefully add about 50 mL dry DMSO to this volumetric flask. Mix well to dissolve all particles of TBAPF₆ and make the volume up to the mark to obtain a 0.1 M solution. Label this volumetric flask as 0.1 M TBAPF₆ in dry DMSO and keep it aside.
3. Solution of P₃
   1. Weigh 4.34 mg P₃ and transfer it to a 10 mL clean volumetric flask with a stopper.
   2. Add 0.1 M TBAPF₆ in dry DMSO and make the volume up to the mark. This produces a 1 mM (i.e., 0.001 M) solution of P₃.
4. 100 mM (i.e., 0.1 M) phosphate buffer, pH 7.4⁴³
   1. Weigh 1.74 g K₂HPO₄ in a clean 50 mL beaker and add 10 mL deionized water (DIW) or 18.2 MΩ·cm water to it. Label this as ‘Solution A’.
   2. Weigh 1.36 g KH₂PO₄ in another clean 50 mL beaker and add 10 mL deionized water. Label this as ‘Solution B’.
   3. Using a 10 mL serological pipette, transfer 8 mL of ‘Solution A’ and 2 mL of ‘Solution B’ into a 100 mL clean volumetric flask and make the volume up to the mark with DI water. Use separate pipettes for each solution.
   4. Check the pH of this solution on a pH meter and adjust the pH to 7.4 with careful addition of either 0.1 M HCl or 0.1 M NaOH solution as required. Once the pH is adjusted to 7.4, the phosphate buffer (0.1 M) is ready for use.
5. Cyclic voltammetric measurements
   1. Switch on the potentiostat and wait for 10 min to stabilize the instrument. Open the cyclic voltammetry (CV) program.
   2. Set the potential limits from +1.0 V to −1.5 V; start and end potentials to 0 V; scan rate to 50 mV/s, and number of cycles to four.
   3. Organize the electrochemical assembly as shown in Figure 4.
   4. Add 10 mL 0.1 M TBAPF₆ in dry DMSO solution (electrolyte) to an electrochemical cell and purge with N₂ for 15 min. Record the voltammogram of the electrolyte (blank) in the potential range of +1.0 to −1.5 V at a scan rate of 50 mV/s. Clean the electrochemical cell and dry it after use.
   5. Similarly, in the cleaned, dry electrochemical cell, take 1 mM P₃ solution (10 mL), purge N₂ for 15 min, and record the voltammogram using the same parameters as those used above.
   6. Record the voltammogram of P₃ at each scan rate separately, i.e., at 20, 50, 100, 150, 200, and 250 mV/s.
   7. Calculate the E_pa, i_pa and i_pc for NO₂•–, as shown in equation (1), for each scan rate and prepare a table as shown in Table 1.

Figure 4: Electrochemical setup. Schematic illustration of a three-electrode system showing the potentiostat connections, nitrogen purging, and electrode arrangement. Please click here to view a larger version of this figure.

Scan Rate (mV/s)	(scan rate)1/2 (υ1/2)	Epa (V)	Epc (V)	ipc (μA)	ipa (μA)
20	4.47	−1.13	−1.23	−0.4	0.45
50	7.07	−1.13	−1.23	−0.59	0.72
100	10	−1.13	−1.23	−0.81	1
150	12.24	−1.13	−1.23	−0.99	1.24
200	14.14	−1.12	−1.24	−1.13	1.38
250	15.81	−1.12	−1.24	−1.28	1.62

Table 1: Peak potential and current values of P₃ at different scan rates in dry DMSO solution. E_pa, E_pc, i_pa and i_pc values obtained from cyclic voltammogram of P₃ in 0.1 M TBAPF₆ in dry DMSO recorded at different scan rates.

3. Preparation of P₃ modified glassy carbon electrode

1. Weigh 1.00 mg P₃ in a glass vial. Add 500 µL of DMSO/acetonitrile (1:1, v/v) to it. Gently heat the vial to obtain a clear solution. Only 10 µL of this solution is required for loading on the glassy carbon electrode (GCE).
2. Polish a 3 mm glassy carbon electrode⁴⁴. Place two drops of alumina paste on a polishing pad and trace a figure-of-eight motion on the pad for 1 min while firmly holding the electrode. Wipe the paste after polishing. Rinse it thoroughly with distilled water and dry it with a clean, lint-free tissue.
3. First, carefully place 3 µL of P₃ on the 3 mm area of GCE. Dry it with a hot-air gun. After drying the first layer, place 3 µL on top and dry with a hot-air gun. Repeat this one more time by adding 4 µL of P₃ solution and drying with a hot-air gun.
4. Then place 4 µL perfluorosulfonic acid ionomer binder (1%) solution on it. Allow the ionomer layer to dry at room temperature; the electrode is then ready for use.

4. Electrochemical hydrogen evolution studies using linear sweep voltammetry

1. Open linear sweep voltammetry (LSV) program. In the LSV parameters, enter the start and end potentials as 0 and −1.5 V, respectively, and set the scan rate to 100 mV/s.
2. Fill the electrolytic cell with 10 mL 0.1 M phosphate buffer (pH 7.4) and connect all three electrodes to the potentiostat using crocodile clips.
3. Use P₃-modified GCE, i.e., P₃-GCE, as a working electrode; platinum wire as a counter electrode, and Ag/AgCl as a reference electrode.
4. After connecting the electrodes, purge N₂ for 15 min. Then seal the electrolytic cell using PTFE tape. Record a linear sweep voltammogram.

5. Stability and durability test of P₃-GCE using chronoamperometry

1. Open chronoamperometry program. Set the applied potential to −1.22 V (vs Ag/AgCl) and the time to 2 h.
2. Use the same electrochemical setup for LSV measurements. Click the Start Measurement, and the measurement will stop automatically after 2 h.

Results

P₃ was synthesized using a three-step procedure. In the first step, the reaction of TCH with 2-pyridyl carboxaldehyde in a 1:1 molar ratio gave compound P₁ (Supplementary Figure 1). In the second step, the reaction between P₁ and salicylaldehyde gave P₂. Finally, the reaction of P₂ with p-nitrobenzyl bromide yielded compound P₃. Both P₁ and P₂ were purified by recrystallization from methanol and dried before proceeding to the next steps in synthesis. The final product, P₃, was purified by column chromatography as described in the protocol section. After purification, P_1, P_2, and P₃ were characterized by ¹H-NMR (Supplementary Figure 2, Supplementary Figure 3, Supplementary Figure 4), FTIR (Supplementary Figure 5, Supplementary Figure 6, Supplementary Figure 7), HRMS (Supplementary Figure 8, Supplementary Figure 9, Supplementary Figure 10), UV-visible spectroscopy (Supplementary Figure 11), powder X-ray diffraction (Supplementary Figure 12), and field-emission scanning electron microscopy (FE-SEM, Supplementary Figure 13). ¹H-NMR, FTIR, and HRMS are useful for identifying the compounds; details are provided in the Supplementary File 1. The PXRD of P₃ showed sharp peaks between 2θ = 5° and 30°, indicating a crystalline nature, which was further confirmed by the FE-SEM image showing long crystalline needles. The crystallite size calculated using the Debye–Scherrer equation is 12.88 nm. The relatively small crystallite size of P₃ is expected to expose more catalytically active sites, thereby facilitating charge transfer and enhancing interactions with the electrolyte. Furthermore, the porous morphology of P₃ immobilized in a perfluorosulfonic acid ionomer binder provides a high surface area, promoting efficient proton adsorption, increasing the accessibility of active sites, and facilitating the release of generated H₂ gas from the electrode surface.

The inherent redox properties of P₃  in solution were evaluated by cyclic voltammetry in dry DMSO containing 0.1 M TBAPF₆ as the supporting electrolyte (Figure 5, Table 2) at a scan rate of 50 mV/s. The voltammogram of the electrolyte blank solution alone did not show any redox peaks (black line), whereas P₃ (blue line) showed a redox couple with an anodic peak, A, appearing at E_pa1 = −1.13 V and the corresponding cathodic peak, A’, at E_pc = −1.23 V (Figure 5A). This redox couple is assigned to the reduction of the NO₂ group in P₃ to NO₂•⁻, as shown in equation (1)^45,46,47,48 below. The peak-to-peak separation for this couple was calculated as ΔE_p = E_pa1 – E_pc = 0.1 V, which is larger than 0.059 V, suggesting a quasi-reversible process⁴⁹.

R - NO₂ + e^- → R - NO₂^-  -------- (1)

	Epa (V)	Epc (V)	ΔE (V)	E1/2 (V)
NO2/NO2.–​	−1.13	−1.23	0.1	−1.18
N-H	0.8	-	-	-

Table 2: Peak potential values obtained from the cyclic voltammogram of P₃ in dry DMSO solution. Peak potential values of P₃ obtained from the cyclic voltammogram of the solution of P₃ prepared in 0.1 M TBAPF₆ as a supporting electrolyte in dry DMSO at a scan rate of 50 mV/s.

In addition an anodic peak at E_pa2 = +0.80 V (denoted as B in Figure 5A) is attributed to the irreversible oxidation corresponding to proton-coupled electron transfer PCET, from the N–H in the thiocarbohydrazone moiety in P₃^50,51,52. Further, the scan rate dependence of the NO₂ redox couple (A/A’) and oxidation of TCH (B) was studied by scanning the CV at different scan rates between 50 mV/s and 250 mV/s (Figure 5B). From the plot of peak current (i_pa and i_pc) versus the square root of scan rate (ν^1/2), the peak current for both the peaks varied linearly with the square root of the scan rate (Figure 5C,D), indicating the diffusion-controlled nature of both processes in P₃.

Figure 5: Electrochemical behavior of P₃ in dry DMSO. (A) Cyclic voltammogram of 1 mM P₃ recorded in 0.1 M TBAPF₆ in dry DMSO at 50 mV/s scan rate; (B) Cyclic voltammogram of 1 mM P₃ in 0.1 M TBAPF₆ in dry DMSO at different scan rates; Plots of current vs. (ν^1/2) for redox peaks (C) A–A’ and (D) B. Please click here to view a larger version of this figure.

P₃ possesses several notable features that make it a candidate electrocatalyst for HER. It contains heteroatoms (N, O, and S), protonatable and deprotonatable basic pyridine nitrogen and -NH groups, a redox-active NO₂ group, and several hydrogen-bond donor and acceptor functionalities. Since P₃ is water-insoluble, the electrocatalytic behavior of P₃  toward HER was evaluated in 0.1 M phosphate buffer (pH 7.4) and 0.1 M KNO₃ (pH 7.2) in aqueous medium by immobilizing it on GCE using a perfluorosulfonic acid ionomer binder.

The quasi-reversible NO₂/NO₂^•– redox couple, with ΔEp = +0.1 V vs Ag/AgCl electrode, is observed only in the solution-state cyclic voltammogram of P₃ in dry DMSO. This couple is not observed in the solid-state voltammogram of P₃ immobilized in a perfluorosulfonic acid ionomer binder (P₃-GCE) in 0.1 M phosphate buffer at pH 7.4. In phosphate buffer, P3-GCE gives peak-to-peak separation of Epa = +0.277 V and E_pc = −0.007 V, with a peak-to-peak separation of 0.284 V vs Ag/AgCl electrode. This is attributed to slow, heterogeneous electron-transfer kinetics, which require a larger overpotential to drive the redox process, indicating that the electron transfer is followed by a slow proton transfer that requires structural rearrangement. This is also seen in the Tafel slope (−370 mV/dec).

To gain further insight into the role of the NO₂ group in the proton-coupled electron transfer (PCET) process, the pH-dependent redox behavior of P₃-GCE was investigated in phosphate buffer solutions over the pH range 5.0–9.0 (pH 5.0, 5.5, 6.0, 6.5, 7.0, 7.4, 8.0, 8.5, and 9.0). The corresponding Pourbaix diagram is presented in Figure 6. In the anodic region, a pH-independent one-electron oxidation process is observed at nearly constant potential, indicating that the oxidation is centered on the sulfur donor atoms of the thiocarbohydrazone framework⁵³. In contrast, the reduction process exhibits a clear pH dependence. The slope of the first reduction peak potential versus pH plot is approximately +67 mV pH⁻¹, which is close to the theoretical value expected for a one-electron/one-proton coupled process. This observation confirms that the reduction of P₃ proceeds through a PCET mechanism involving the transfer of one electron and one proton. These results provide additional evidence for the participation of the NO₂ functionality in the electron-transfer process and further support the quasi-reversible redox behavior observed in the cyclic voltammetry studies.

Figure 6: Pourbaix diagram of P₃-GCE. Plots of the first oxidation and reduction potentials (vs RHE) of P₃-GCE as a function of pH (Pourbaix diagram), where potentials were determined by observing cyclic voltammograms of P₃-GCE in aqueous phosphate buffer solutions under various pH conditions (pH 5–9). Anodic potentials are presented as red circles with solid black lines, whereas cathodic potentials are presented as green circles with solid black lines. Please click here to view a larger version of this figure.

The linear sweep voltammograms (LSVs) of bare GCE, ionomer-coated GCE, and P₃-GCE were recorded in 0.1 M phosphate buffer (pH 7.4) at 100 mV/s (Supplementary Figure 14). The LSVs of the bare GCE and perfluorosulfonic acid ionomer binder-coated GCE showed a linear response, whereas the P₃-GCE showed a prominent cathodic current characteristic of HER. The onset potential (Figure 7A) for this cathodic current was −0.370 V (vs RHE), and an overpotential of −0.690 V (vs RHE) was required to reach a current density of 0.1 mA/cm². The Tafel plot (Figure 7B) analysis yielded a slope of −370 mV/dec, indicating slow HER kinetics by P₃-GCE at neutral pH 7.4. The band gap of 2.24 eV was calculated from the Tauc plot (Figure 7C).

Figure 7: Electrocatalytic HER studies using P₃-GCE in 0.1 M phosphate buffer of pH 7.4: (A) Linear sweep voltammogram of P₃-GCE in 0.1 M phosphate buffer (pH 7.4) at a scan rate of 100 mV/s; (B) Tafel plot; (C) Tauc plot for estimation of optical band gap of P₃; (D) Cyclic voltammogram of P₃-GCE in non-Faradaic region (+0.4 to −0.4 V vs Ag/AgCl) at different scan rates; (E) Plot of capacitive current at 0 V (vs Ag/AgCl) versus scan rate, and (F) EIS Nyquist plot in 0.1 M phosphate buffer. Please click here to view a larger version of this figure.

The current is normalized to the geometric area of the working electrode (0.07 cm²) to compare the HER performance of P₃-GCE in two different electrolytes. The double-layer capacitance (C_dl) was calculated from cyclic voltammograms recorded in the non-Faradaic region. The double-layer capacitance is directly proportional to the electrochemically active surface area for catalysis, and this is reflected in the C_dl values obtained in 0.1 M phosphate buffer (14.59 µF) and 0.1 M KNO₃ in DIW (5.73 µF) (Figure 7D,E). The C_dl value in phosphate buffer is approximately 2.5 times that of 0.1 M KNO₃ in DIW, as reflected in the ECSA. In addition, the electric double-layer enhances electrode–electrolyte interactions and may increase the availability of protons near the electrode surface.

The electrochemically active surface area of 0.029 cm² was calculated assuming a specific capacitance of 35 µF/cm², corresponding to six times the roughness of the bare electrode. The solution resistance and charge transfer resistance (R_ct) of the catalyst film were calculated using electrochemical impedance spectroscopy (EIS). The EIS was fitted to a Nyquist plot (Figure 7F) using various equivalent circuit models, yielding a solution resistance of 113 Ω and R_ct = 2492 Ω. These high values indicate, resistive nature of P₃-GCE. The turnover frequency was 265.6 x 10^-6 s^-1, and an exchange current density of 986.5 x 10^-6 A/cm² was achieved at zero overpotential.

Additionally, the electrocatalytic behavior of P₃-GCE was studied under identical conditions in an aqueous solution of 0.1 M KNO₃ (pH 7.2) as an electrolyte (Figure 8A–F). Interestingly, in 0.1 M KNO₃, P₃-GCE produced a higher cathodic current for HER with an onset potential of −0.329 V (vs RHE; Figure 8A). An overpotential of −0.519 V (vs RHE) was required to reach a current density of 0.1 mA/cm². The Tafel plot (Figure 8B) yielded a slope of −348 mV/dec. The charge transfer resistance, R_ct, for this system was 5522.8 Ω, the turnover frequency was 409.6 (× 10⁻⁶) s^-1, and the exchange current density was the same as that of the phosphate buffer.

Figure 8: Electrocatalytic HER studies using P₃-GCE in 0.1 M KNO₃ in DIW: (A) Linear sweep voltammogram of P₃-GCE in 0.1 M KNO₃ in DIW at a scan rate of 100 mV/s; (B) Tafel plot; (C) Tauc plot for estimation of optical band gap of P₃; (D) Cyclic voltammogram of P₃-GCE in non-faradaic region (+0.2 to −0.2 V vs Ag/AgCl) at different scan rates; (E) Plot of capacitive current at 0 V (vs Ag/AgCl) versus scan rate, and (F) EIS Nyquist plot in 0.1 M KNO₃. Please click here to view a larger version of this figure.

To study the stability of P₃ film on the GCE surface, chronoamperometric measurements (Figure 9A) were carried out at −1.22 V (vs Ag/AgCl) for 2 h, and a stable current response was observed. In contrast, in the 0.1 M KNO₃ solution, the cathodic current increased continuously up to 45 min and then began to decrease (Figure 9B), thereby limiting the measurement time to 1 h. The chronoamperometric measurements for the phosphate buffer system were also studied in short on-off catalytic cycles at 10 min intervals (Figure 9C), indicating a stable response of the catalyst.

Figure 9: Stability and durability tests of P₃-GCE. Chronoamperometry response of P₃-GCE for 2 h at −1.22 V (vs Ag/AgCl) in (A) 0.1 M phosphate buffer (pH 7.4) and (B) 0.1 M KNO₃. (C) Chronoamperometric recyclability response of P₃-GCE at on- and off-potentials in 0.1 M phosphate buffer (pH 7.4) (on-potential = −1.22 V (vs Ag/AgCl), off-potential = 0 V (vs Ag/AgCl)). Please click here to view a larger version of this figure.

Gas chromatography (Figure 10A) analysis of the headspace after 2 h of chronoamperometry confirmed the presence of hydrogen, with a retention time matching that of the standard H₂ gas sample analyzed under identical conditions. Bubble formation (Figure 10B) was noticed at the electrode-electrolyte interface during prolonged electrolysis. All electrochemical parameters for the performance of P₃-GCE in both electrolyte systems are summarized in Table 3.

Figure 10: Detection of H₂ formed by gas chromatography: (A) Gas chromatographic (GC) analysis of hydrogen generated during controlled potential electrolysis using P₃-GCE in 0.1 M phosphate buffer and 0.1 M KNO₃; (B) Image showing bubble formation at the P₃-GCE electrode surface during chronoamperometric measurements. Please click here to view a larger version of this figure.

Supplementary Figure 1: Synthesis of P₃. A scheme showing the steps involved in the synthesis of P₃. Please click here to download this file.

Supplementary Figure 2: ¹H-NMR spectrum of P₁. ¹H-NMR spectrum of P₁ recorded in DMSO-d₆ solvent. Please click here to download this file.

Supplementary Figure 3: ¹H-NMR spectrum of P₂. ¹H-NMR spectrum of P₂ recorded in DMSO-d₆ solvent. Please click here to download this file.

Supplementary Figure 4: ¹H-NMR spectrum of P₃. ¹H-NMR spectrum of P₃ recorded in DMSO-d₆ solvent. Please click here to download this file.

Supplementary Figure 5: Infrared spectrum of P₁. FTIR spectrum of a finely powdered sample of P₁. Please click here to download this file.

Supplementary Figure 6: Infrared spectrum of P₂. FTIR spectrum of a finely powdered sample of P₂. Please click here to download this file.

Supplementary Figure 7: Infrared spectrum of P₃. FTIR spectrum of a finely powdered sample of P₃. Please click here to download this file.

Supplementary Figure 8: HRMS spectrum of P₁. HRMS spectrum of P₁ recorded in positive ion mode showing m/z = 196.05 corresponding to [M+H]⁺ ion. Please click here to download this file.

Supplementary Figure 9: HRMS spectrum of P₂. HRMS spectrum of P₂ recorded in positive ion mode with m/z = 299.09 corresponding to [M+H]⁺ ion. Please click here to download this file.

Supplementary Figure 10: HRMS spectrum of P₃. HRMS spectrum of P₃ recorded in positive ion mode with m/z = 435.12 corresponding to [M+H]⁺ ion. Please click here to download this file.

Supplementary Figure 11: UV-visible spectra of P₁, P_2, and P₃. UV-Visible spectra of P₁ (black solid line), P₂ (red line), and P₃ (solid blue line) recorded at 10 µM concentration in neat DMSO. Please click here to download this file.

Supplementary Figure 12: Powder XRD pattern of P₃. X-ray powder diffraction pattern of P₃. Please click here to download this file.

Supplementary Figure 13: FE-SEM image of P₃ and perfluorosulfonic acid ionomer binder-coated P₃. FE-SEM images showing (A) Needle-like crystals of P₃ (scale: 5 µm) and (B) perfluorosulfonic acid ionomer binder-coated P₃ showing porous surface (scale: 1 µm). Please click here to download this file.

Supplementary Figure 14: LSV of P₂ and P₃ modified GCE. LSV of bare GCE (brown line with solid stars), GCE coated with perfluorosulfonic acid ionomer binder (green solid line), P₂-GCE (blue line with crosses), and P₃-GCE (red line with hollow squares) recorded in 0.1 M phosphate buffer (pH 7.4) at a scan rate of 100 mV/s. Please click here to download this file.

Supplementary Figure 15: Cyclic voltammogram of P₃-GCE. Cyclic voltammogram of P₃-GCE in (A) 0.1 M KNO₃ in DIW and (B) 0.1 M phosphate buffer (pH 7.4) at a scan rate of 100 mV/s. Please click here to download this file.

Supplementary Figure 16: LSV profile of P₃-GCE. LSV profile of P₃-GCE in 0.1 M phosphate buffer (red line with hollow circles) and 0.1 M KNO₃ (black solid line) in DIW at a scan rate of 100 mV/s. Please click here to download this file.

Supplementary Figure 17: Stability of P₃-GCE. Performance of P₃-GCE in 0.1 M phosphate buffer (pH 7.4) for 5 days. Please click here to download this file.

Supplementary File 1: Supplementary Experimental Details and Calculations. The details of the synthesis and instrumentation used for characterization of P₁-P₃; interpretation of HRMS, ¹H-NMR, FTIR, UV-Visible spectra and electrochemical measurements along with the equations used for conversion of redox potential value from Ag/AgCl to RHE (eq. 1); overpotential (η, eq. 2); electrochemically active surface area (ECSA, eq. 3); turnover frequency (TOF, eq. 4 to 6); calculation of conductivity and film thickness (eq 7 to 10), Faradic efficiency (eq. 11-14) can be found in the supplementary file. Please click here to download this file.

Discussion

This protocol describes the three-step synthesis of a thiocarbohydrazone derivative, P₃, and the preparation of a P₃-modified glassy carbon electrode, P₃-GCE, for investigating the hydrogen evolution reaction (HER) in a neutral aqueous medium. The straightforward synthesis and purification procedures, together with the use of readily available and inexpensive reagents, make P₃ a suitable candidate for exploring HER in a typical laboratory setting and for educational demonstrations.

The electrochemical behavior of P₃ was initially investigated under homogeneous conditions in DMSO to evaluate its intrinsic redox properties. The presence of a quasi-reversible redox couple associated with the nitro group and an irreversible oxidation process arising from the thiocarbohydrazone backbone indicates the existence of electroactive sites within P₃. Catalytic studies were subsequently performed under heterogeneous conditions using P₃-GCE. Electrochemical investigations were carried out in two electrolyte systems, namely 0.1 M phosphate buffer (pH 7.4) and 0.1 M KNO₃ (pH 7.2) (Supplementary Figure 15 and Supplementary Figure 16 ).

S.No.	Parameters	Phosphate Buffer	KNO3
1	Onset potential (vs RHE) (mV)	−376.3	−329
2	Overpotential (0.1 mA/cm2)(mV)	−690	−519
3	Tafel slope (mV/dec)	−370	−348
4	Exchange current density (mA/cm2) (X10-3)	0.9865	0.9866
5	Rct (Charge transfer resistance) (ohm)	2492	5522.8
6	Conductivity (x X10-4)	9.85	4.43
7	Band gap (eV)	2.24
8	TOF  (s-1) (X10-6)	265.6	409.6
9	ECSA (cm2)	0.029	0.011
10	Double layer capacitance (mF) (X10-3)	14.59	5.73
11	Faradaic efficiency (%)	0.27	0.43

Table 3: Summary of electrochemical parameters of P₃-GCE. Summary of all electrochemical parameters for the performance of P₃-GCE in 0.1 M phosphate buffer (pH 7.4) and 0.1 M KNO₃ electrolyte systems.

The electrochemical parameters presented in Table 3 reveal that the performance of P₃-GCE is strongly influenced by the electrolyte medium. In 0.1 M KNO₃, P₃-GCE exhibits a lower onset potential (−329 mV vs. RHE) and a smaller overpotential (−519 mV at 0.1 mA/cm²) compared to those observed in 0.1 M phosphate buffer (−376.3 mV and -690 mV, respectively), indicating that hydrogen evolution is thermodynamically more favorable in the KNO₃ electrolyte. Furthermore, the slightly lower Tafel slope in KNO₃ (−348 mV/dec) relative to phosphate buffer (−370 mV/dec) suggests somewhat improved HER kinetics. The exchange current density values are nearly identical in both electrolytes (~0.99 × 10^-3 mA/cm²), indicating similar intrinsic catalytic activity at the active sites. However, phosphate buffer exhibits significantly lower charge-transfer resistance (2492 Ω) and higher conductivity (9.85 × 10^-4) than KNO₃ (5522.8 Ω and 4.43 × 10⁻⁴, respectively), demonstrating more efficient interfacial electron transport. This improved charge-transfer behavior is further supported by the larger electrochemically active surface area (0.029 cm²) and higher double-layer capacitance (14.59 µF) observed in phosphate buffer, suggesting greater accessibility of catalytic sites and enhanced electrode–electrolyte interaction. Despite the higher charge-transfer resistance in KNO₃, P₃-GCE displays a higher turnover frequency (409.6 × 10^-6 s^-1) and Faradaic efficiency (0.43%) than in phosphate buffer (265.6 × 10^-6 s^-1 and 0.27%, respectively), indicating more efficient utilization of active sites for hydrogen generation once the reaction is initiated. These observations suggest that while phosphate buffer promotes favorable electrode kinetics and charge transfer, KNO₃ is more effective at facilitating hydrogen evolution on the P₃-GCE under the conditions investigated.

The electrocatalytic behavior of P₃ was first examined in phosphate buffer. Analysis of the linear sweep voltammetry (LSV) profile revealed lower hydrogen generation than that of many organic electrocatalysts reported in the literature (Table 4). However, most reported organic electrocatalysts have been evaluated in homogeneous, non-aqueous media, whereas P₃-GCE operates as a heterogeneous catalyst in a neutral aqueous environment. Furthermore, many reported organic electrocatalysts are sacrificial in nature^{23,24,25,26,27,28,29,30,31,32,33,34,35,36,37,38,39,40,41,42}, which is a significant limitation. In contrast, P₃-GCE exhibited excellent stability and reusability. Once prepared, the electrode delivered a consistent response in 0.1 M phosphate buffer over five days when used for at least 5 h each day (Supplementary Figure 17). Although the response of P₃-GCE in 0.1 M KNO₃ was less stable, the current density was higher than that observed in phosphate buffer, reflecting the influence of electrolyte composition on catalytic performance.

Sr. No.	Material		Type	Solvent	Acid/ Base used	Current range	Catalysis Mode	Reference
1	Organic hydride (dihydropyridine)		Organic	Acetonitrile	CH3COOH, CF3COOH	μA	Homogeneous	23
2	Benzothiadiazole (BTDN)		Organic	Acetonitrile	salicylic acid	μA	Homogeneous	27
3	Benzothiadiazole–caffeine (BT-DCAF)		Organic	DMSO	CF3COOH	μA	Homogeneous	28
4	9-phenyl-10-methylacridinium iodide		Organic	Acetonitrile	HClO4	μA	Homogeneous	30
5	2,4,6-Triphenylpyridine		Organic	Acetonitrile	CH3COOH	μA	Homogeneous	31
6	Phosphorus corrole		Organic (non-metallic complex)	DMF	TFA	μA	Homogeneous	32
7	2,2`-Bypyridyine		Organic	Acetonitrile	CH3COOH	μA	Homogeneous	33
8	4,4`-Bypyridyine		Organic	Acetonitrile	HBF4, HClO4, TsOH, CF3COOH	μA	Homogenous	34
9	1,10-Phenanthroline		Organic	Acetonitrile	HClO4, CF3COOH	μA	Homogeneous	35
10	Bis-imidazolium-embedded heterohelicene		Organic	Acetonitrile	CH3COOH	μA	Homogeneous	36
11	P3-GCE		Organic	i) 0.1 M Phosphate buffer ii) 0.1 M KNO3 in DIW	pH 7.4	μA	Heterogenous (immobilized on electrode)	This work

Table 4: Comparison of reported organic HER electrocatalysts with P₃-GCE. The table provides a comparison of reported organic electrocatalysts for HER with P₃-GCE.

Most previously reported organic HER systems, including acridinium salts, benzothiadiazole derivatives, triphenylpyridine, phosphorus corrole, bipyridine- and phenanthroline-based catalysts, NADPH-like heterohelicenes, organic hydrides, and benzothiadiazole-caffeine conjugates, operate as homogeneous molecular catalysts in organic solvents such as acetonitrile, DMF, or DMSO in the presence of strong proton sources including HClO₄, trifluoroacetic acid (TFA), acetic acid, and HBF₄. These reaction conditions provide a high concentration of readily available protons, thereby facilitating efficient hydrogen evolution. In contrast, P₃-GCE functions as a heterogeneous electrocatalyst immobilized on a glassy carbon electrode and operates in environmentally benign aqueous electrolytes, namely 0.1 M phosphate buffer (pH 7.4) and 0.1 M KNO₃. The significantly reduced proton availability under these near-neutral conditions is expected to limit the HER rate, thereby reducing hydrogen generation. Nevertheless, the ability of P₃-GCE to catalyze hydrogen evolution without corrosive acids, organic solvents, or dissolved molecular catalysts highlights its practical advantages in operational simplicity, catalyst recovery, and environmental sustainability. Therefore, although the hydrogen evolution performance is lower than that of many reported homogeneous organic electrocatalysts, P₃-GCE demonstrates the feasibility of employing a metal-free organic molecule for HER under mild aqueous conditions.

The observed behavior can be rationalized using the HER mechanism illustrated in Figure 11. Hydrogen evolution generally involves two key processes: adsorption of water molecules or protons onto the cathode surface (Volmer step), followed by hydrogen desorption through either the Heyrovsky or Tafel step. Efficient HER requires an appropriate balance between these processes. As HER is a two-electron-transfer process, its kinetics are strongly influenced by the electrolyte pH. In acidic media, the high proton concentration facilitates rapid reaction kinetics. In contrast, under neutral or alkaline conditions, proton generation depends on the slower dissociation of water molecules, resulting in a higher kinetic barrier^54,55. “Despite the measurable Tafel response of −370 mV/dec, which suggests relatively fast reaction kinetics, the HER performance of P₃-GCE appears to be limited by charge-transfer resistance. Electrochemical impedance spectroscopy (EIS) revealed a charge-transfer resistance of 2492 Ω, indicating restricted electron transport to the catalyst surface. Additionally, P₃ is likely stabilized within the ionomer matrix through hydrogen-bonding interactions. Upon exposure to water, this hydrogen-bonding network may become more extensive. Consequently, the strong interaction between P₃ and its surroundings may hinder the desorption of hydrogen from the electrode surface after its formation. This effect contributes to the high overpotential observed for HER and partially explains the lower catalytic activity of P₃-GCE compared with both organic and inorganic electrocatalysts reported in the literature.

Figure 11: Hydrogen-bonding network at the P3-GCE through a hydrogen-bonding network involving water and the ionomer matrix. Diagram showing concerted H2 evolution by stabilization of P3 in perfluorosulfonic acid ionomer binder via a hydrogen-bonding network with water and the matrix. Please click here to view a larger version of this figure.

Tafel slopes of approximately 120, 40, and 30 mV/dec are generally associated with the Volmer, Heyrovsky, and Tafel rate-determining steps, respectively⁵⁵. For P₃-GCE, the rate-determining step is likely the Volmer process, in which water molecules first dissociate, and the resulting hydrogen species adsorb onto the catalyst surface. The presence of N–H, sulfur, and imine functionalities in P₃ facilitates proton adsorption and subsequent hydrogen evolution. The HER mechanism in neutral media consists of three elementary reaction steps:

P₃ + H₂O + e^- → P₃H_ads + OH^- (Volmer Step) -------- (1)

P₃H_ads + e^- + H₂O ⇆ P₃ + H₂ + OH^- (Heyrovsky step) -------- (2)

P₃H_adsH_ads ⇄ P₃ + H₂ (Tafel step) -------- (3)

2H₂O + 2e^- → H₂ + 2OH^- (Overall Reaction)

Mechanistic Discussion and Electrocatalytic Performance

The Volmer step (1) involves the dissociation of a water molecule and the adsorption of a hydrogen atom on an active site on the electrode or catalyst surface. Hydrogen evolution subsequently proceeds through either an electrochemical pathway, the Heyrovsky step (2), or a chemical pathway, the Tafel step (3). Several studies have reported that at low overpotentials, HER proceeds through the Volmer step, followed by the parallel Heyrovsky and Tafel pathways, whereas at high overpotentials the Tafel step becomes negligible and the reaction predominantly follows the Volmer–Heyrovsky mechanism^56,57,58. Given the relatively high overpotential observed for P₃, the HER is proposed to proceed mainly via the Volmer–Heyrovsky pathway. The proposed multiple pathways for HER on P₃-GCE are illustrated in Figure 12.

Figure 12: Plausible HER mechanism pathways. Schematic illustration of pathways showing a possible mechanism for HER on P3-GCE. Please click here to view a larger version of this figure.

Another factor influencing the HER performance is the buffering capacity of the 0.1 M phosphate buffer at pH 7.4, which is approximately 0.057 M. In this system, competition exists between P₃, embedded in a perfluorosulfonic acid ionomer binder, and the buffer base for the protons generated during the initial HER step. This competition reduces the availability of H⁺ ions at the catalyst surface, thereby lowering hydrogen production. Gas chromatography analysis of the headspace in the electrocatalytic reaction assembly obtained from P₃-GCE showed a distinct peak corresponding to H₂ gas, with a retention time matching that of a standard hydrogen sample. In addition, visible bubble formation at the electrode-electrolyte interface during prolonged electrolysis confirmed gas evolution. These observations collectively demonstrate that the cathodic current at P₃-GCE is attributable to electrocatalytic hydrogen evolution. Only a small difference was observed in the amount of hydrogen generated by P₃-GCE in 0.1 M phosphate buffer (25.28 µM) and in 0.1 M KNO₃ solution (21.31 µM). After the HER process in 0.1 M KNO₃, the solution's pH increased to 7.95. The initial increase in current in chronoamperometric measurements can be attributed to a higher local deprotonation rate near the electrode surface, driven by OH⁻ ions produced during the first catalytic step. However, as electrolysis continues, the depletion of available water molecules and protons reduces the efficiency of the P₃-GCE in KNO₃ solution after 1h.

Advantages and Comparison with Other Catalysts

P₃-GCE offers several advantages: (i) it functions as a heterogeneous catalyst immobilized as a film on the electrode surface, (ii) it operates effectively at neutral pH (7.4) without the need for an external proton source, (iii) the catalyst remains stable for several days, and (iv) it exhibits durability over multiple catalysis on/off cycles. Table 4 compares the performance of P₃-GCE with various reported organic electrocatalysts. Inorganic metal complexes derived from thiocarbohydrazone ligands generally require an external proton source for HER and typically operate in non-aqueous media. For example, Co(II)⁵⁹ and Cu(II)⁶⁰ complexes of salen-type ligands reported by Dutta et al. function efficiently at neutral pH; however, the catalysis remains homogeneous and is dependent on the metal center.

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Sr. No.	Material	Electrode substrate	Electrolyte	Over potential at 10 mA/cm2 (mV)	Tafel Slope (mV)	Reference
1	C60-SWCNT15 (single walled carbon nano tube)	SiO2/Si film	Phosphate buffer (pH = 7)	330	166.1	61
2	N-MPG (Nitrogen doped mesoporous graphene)	Glassy carbon electrode	0.5 M H2SO4	239	109	62
3	B-SuG (Boron substituted graphene)	Glassy carbon electrode	0.5 M H2SO4	201a	99	63
4	NS 500 (N,S co-doped graphene)	Nano porous nickel substrate	0.5 M H2SO4	−280	80.5	64
5	N-CQDs-17h (Nitrogen-doped carbon quantum dots)	Glassy carbon electrode	0.5 M H2SO4	341	126	65
6	PA-900 (P- doped carbon)	Glassy carbon electrode	0.5 M H2SO4	290	−112.7	66
7	PDHCI (carbon-based pentalene-rich nanoribbons)	Au(111)	0.1 M HClO4	At 0.4 mA/cm2 −258	82	67
8	CD2 (Sulfanilamide/Citric Acid Derived N,S‑ Doped Carbon Dot)	Nickel foam	1 M KOH	216	124.1	68
9	PG (Poly guanine polymer DNA inspired)	Carbon felt	0.5 M H2SO4	411	243	69
10	BSA-FCNT (Bovine serum albumin functionalized multi walled carbon nano tube)	Glassy carbon electrode	0.5 M H2SO4	-0.44	161	70
11	P3-GCE	Glassy carbon electrode coated with P3 in perfluorosulfonic acid ionomer binder	0.1 M phosphate buffer (pH = 7.4)	At 0.1 mA/cm2 −690	−370	This work
			0.1 M KNO3 in DIW	At 0.1 mA/cm2 −519	−348	This work
*a = Current density not available

Table 5: Comparison of reported carbon and protein-based electrocatalysts. The table provides a comparison of reported carbon- and protein-based electrocatalysts for HER with P₃-GCE.

To provide a broader perspective on the catalytic performance of P₃-GCE, its HER activity was compared with representative carbon-, protein-, and DNA-based electrocatalysts reported in the literature (Table 5). Most carbon-based electrocatalysts, including C₆₀-SWCNTs^15,61, N-MPG⁶², B-SuG^63,64, N-CQDs-17h⁶⁵, PA-900⁶⁶, and CD2^67,68, require relatively high overpotentials in the range of 201–341 mV to achieve current densities close to 10 mA/cm² with Tafel slopes between 80 and 166 mV/dec. DNA-inspired polyguanine (PG)⁶⁹ exhibits an even higher overpotential of 411 mV and a Tafel slope of 243 mV/dec, indicating comparatively sluggish HER kinetics. The protein-based catalyst BSA-FCNT⁷⁰ also exhibits a higher Tafel slope (161 mV/dec), indicating slower charge transfer and proton reduction processes. In contrast, P₃-GCE exhibits catalytic activity in neutral media, delivering hydrogen evolution at an overpotential of -690 mV in phosphate buffer (pH 7.4) and -519 mV in 0.1 M KNO₃, both measured at a current density of 0.1 mA/cm². Although direct comparison of overpotential values is complicated by differences in electrolyte composition, pH, electrode substrates, and operating current densities, the measured Tafel slopes obtained for P₃-GCE (–370 and –348 mV/dec in magnitude) demonstrate its ability to facilitate HER under environmentally benign, near-neutral conditions without the need for highly acidic or alkaline electrolytes. These results highlight the potential of the synthesized organic molecule P₃ as a metal-free electrocatalyst and justify its inclusion alongside established carbon-, protein-, and DNA-based HER systems.

Methodological Considerations and Reproducibility

A critical step in the protocol is thoroughly polishing and cleaning the glassy carbon electrode prior to catalyst immobilization. Inadequate polishing can result in uneven catalyst distribution and an irregular electroactive surface area, which in turn affects current density measurements obtained from linear sweep voltammetry (LSV). Uniform catalyst immobilization is therefore essential for obtaining consistent electrochemical responses. Residual contaminants on the electrode surface may increase the background current, shift the onset potential, and obscure the catalyst's intrinsic catalytic activity.

Additional factors that strongly influence the reliability of the protocol include electrolyte purity and the effective removal of oxygen from the solution. The presence of precipitates in the buffer or dissolved oxygen can significantly affect electrochemical measurements.

Reliable electrical contact between the electrodes and the potentiostat is equally important. Loose or unstable connections, as well as accidental physical contact between electrodes during measurements, can increase resistance and distort voltammograms at cathodic potentials. Firm mechanical connections and stable electrical contacts help minimize noise and improve the accuracy of overpotential measurements.

The reproducibility of the protocol is influenced by operator-dependent factors such as electrode polishing and drop-casting of P₃ onto the glassy carbon surface. Variations in these steps can alter the electrochemically active surface area and affect the observed cathodic response.

Further studies are required to gain a deeper understanding of the catalytic HER behavior of P₃-GCE.

Materials

List of materials used in this article

Name	Company	Catalog Number	Comments
Acetone	Qualigens	Q33516	Flammable liquid, severe eye irritation
Concentrated Hydrochloric Acid	Qualigens	Q29145	Corrosive; handle with gloves and eye protection and use in fume hood
Diethyl Ether	Avra	ASD1520	Highly flammable; avoid ignition sources and inhalation
Dipotasssium hydrogen phosphate	Sisco Research Lboratories Pvt. Lts.	90654	Avoid contact with skin and eyes
Ethanol	Qualligens	Q21535	Flammable liquid, keep away from heat and open flame
Ethyl acetate	Qualigens	Q2345C	Highly flammable, Keep away from heat, sparks and open flames
Hexane	Qualigens	Q13177	Highly flammable liquid and vapor. Causes skin and eye irritation.
p-Nitrobenzyl bromide	Sigma Aldrich	N13054	Skin corrosion, avoid direct exposure
Potassium Dihydrogen Orthophosphate	Sisco Research Lboratories Pvt. Ltd.	52403	Avoid contact with skin, eyes and clothing
Potassium nitrate	Thermo Fisher Scientific	A14527	Keep away from heat, hot surfaces, sparks, open flames and other ignition sources
2-Pyridyl carboxaldehyde	Avra	ASP2301	Skin irritation, use only outdoors or in a well-ventilated area
Salicylaldehyde	Sisco Research Lboratories Pvt. Lts.	24922	Harmful if swallowed avoid inhalation and skin contact
Silica Gel (230-400 mesh) for column chromatography	Sisco Research Lboratories Pvt. Lts.	96671	Avoid direct contact with skin, eyes and inhalation
Thiocarbohydrazide	Alfa Aesar	207530050	Toxic when in contact with skin, acute inhalation toxicity
Potentiostat	Metrohm, Ivium	PGSTAT204, IVIUM VERTEX.1A	Handle electrical equipment with dry hands