Spontaneous chlorine production from chloride-containing brines

Domain: physics
Published: 2026-01-07
Editorial outcome: major_revision

Abstract
Chlorine, a crucial basic chemical, is primarily produced by the electrolysis of chloride-containing brines, a highly energy-intensive process with a substantial carbon footprint. Notably, concentrated chloride-containing brines, e.g., acidic wastewater, desalination wastewater, seawater, possess significant osmotic energy, which can be harnessed using membrane-based diffusion cells. Considering this, we here present a spontaneous chlorine production method by using the inherent energy and chloride ions present in these brines. The method is first demonstrated with simulated acidic wastewater because in industry, diffusion cells are already widely used to recycle waste acid. Sulfonated covalent-organic framework membranes are employed to facilitate the diffusion of protons and reject multi-valent cations, purifying acid and avoiding side reactions on the anodes. Consequently, our method simultaneously recovers acid, produces hydrogen and chlorine without consuming external energy. We also validate the general applicability of the method with simulated desalination wastewater. Since our method is compatible with the diffusion-based industrial processes, it holds significant promise for facile, scalable implementation. We also expect the method to be extended for the spontaneous production of other crucial chemicals such as ammonia from nitrate-containing brines.

Introduction
Chlorine (Cl2) is a base chemical for our modern industry. Each year, tens of millions of tons of chlorine are consumed1,2. Over 95% of chlorine is produced by the electrolysis of chloride-containing brines, consuming around 100–200 TWh of electricity each year, which accounts for a few percent of the world’s electricity generation3. During the last decades, great efforts have been made to increase the energy efficiency of the chlorine production process, e.g., by developing anode materials to improve the charge yield4. Such efforts experienced limited success, since the charge yield of the current chlorine production process is already >95%.

We note that the two crucial elements of the chlorine production process are the energy and chloride ions, and both elements naturally exist in many brines, e.g., seawater, desalination wastewater, industrial waste acid, and pharmaceutical wastewater. The concentrated Cl− ions in these brines create high osmotic potential when discharged into lower salinity water (river water, lake water, etc.), which can be harvested to generate electricity5–7. For example, over 50 billion tons of desalination wastewater are discharged annually, typically with a Cl− concentration exceeding 1 M8. Each ton releases osmotic energy of over 0.2 kWh when mixed with seawater, meaning that the desalination wastewater could generate ~10 TWh of electricity per year9. This is sufficient to produce up to 4.5 million tons of chlorine using the Cl− ions in the brine. The mixing of seawater and river water releases even more energy, although the harvesting is also more challenging5,6. In addition, millions of tons of industrial waste acid containing concentrated Cl− ions are generated each year10. Although these brines contain much less energy in total, they present a high possibility for scalable chlorine production, since in our current industry, waste acid is often recycled using the diffusion dialysis process11, whose set-up is identical to that of the osmotic power generation that also utilizes the diffusion of protons. In fact, previous studies in our group and other groups have demonstrated the integration of osmotic energy harvesting and water treatment or resource (e.g., acid) recovery processes12–17. All these facts suggest that chlorine could be massively produced from brines without consuming external energy.

Considering this, we here demonstrate how chlorine can be spontaneously produced from chloride-containing brines, including simulated industrial waste acid and desalination wastewater (schematically illustrated in Fig. 1). We noticed that when industrial waste acid is used, the chlorine production may suffer from the reduction of heavy metal ions on the anode. This problem is addressed by using a covalent-organic framework (COF) membrane, which allows fast proton transport and rejects heavy metal ions. To sustain the chlorine production and couple the transport of the protons, we leveraged the reversible electrochemical reaction of Ag/AgCl electrodes to selectively extract Cl− ions from the waste acid12,13,18. The system not only produces chlorine but also recycles acid as a valuable chemical product, as well as generating hydrogen during the electrolysis of HCl. The hydrogen and chlorine production rates reached approximately 150 L m−2 h−1 and could be continuously sustained for at least 7 days. Furthermore, we demonstrated the general applicability of our method by showing the chlorine production from simulated desalination waste brines.

Results
Design of the system
As shown in Fig. 1a, the conventional chlor-alkali industry uses a cation ion exchange membrane to electrolyze concentrated brine for chlorine production. This process relies on a continuous supply of electrical energy, leading to significant power consumption. In contrast to the conventional chlor-alkali industry, which requires external electrical power, our system operates spontaneously and does not rely on energy input (Fig. 1b).

The system consists of serially connected units for power generation. Each unit is a two-compartment electrochemical cell whose feed reservoir was filled with waste acid, and the permeate reservoir was filled with fresh acid recovered from the waste acid. For practical relevance, the waste acid contained various heavy metal ions (e.g., Fe3+, Co2+, Ni2+) and pollutants (tetracycline hydrochloride, levofloxacin hydrochloride). The reservoirs were connected to Ag/AgCl electrodes and separated by proton-selective COF membranes. As protons diffuse from the waste acid reservoir to the fresh acid reservoir, the same number of electrons would flow through the electrodes to maintain charge balance. This induces the reversible electrochemical reaction, AgCl + e− ⇌ Ag + Cl−. The net outcome is that both protons and Cl− are ions that are selectively transported from the waste acid reservoir to the fresh acid reservoir, generating electricity and recovering acid. By periodically switching the electrodes and refreshing the reservoirs, the system could be continuously operated12,13,18.

To produce chlorine, the left-most electrode illustrated in Fig. 1 was switched to a graphite electrode, and the right-most electrode was a Pt electrode. Note that when the waste acid contains contaminants that can be easily oxidized, the waste acid in the left-most reservoir should be exchanged with the recycled pure acid. For many industrial waste acids, there is no such problem since the heavy metal ions are typically already oxidized. During the chloride production, the recycled protons are reduced on the Pt electrodes, generating hydrogen, which is also a high-value product. In this system, the chlorine evolution reaction (CER) and the hydrogen evolution reaction (HER) are spatially isolated, avoiding the safety problem caused by the mixing of gases.

High proton selectivity of the membrane
To ensure the purity of the recycled acid and thus suppression reduction of the metal ions on the HER electrode, the membrane should have high proton selectivity. Commercial ion-exchange membranes used to recycle acid typically only have a selectivity of up to a few tens. Our previous work demonstrated that a device made of TpPa–SO3H, a covalent-organic framework material with dense sulfonic groups, exhibits high proton selectivity against all other chemical species19. The sulfonic groups form hydrogen bonding networks with the help of water molecules and thus enable fast proton hopping. Here, we fabricated the material as a membrane (Fig. 2a) by casting its monomer solutions (1,3,5-triformylphloroglucinol in n-methyl-2-pyrrolidone and 2,5-diaminobenzenesulfonic acid in dimethyl sulfoxide) on a glass substrate and allowed them to evaporate (Supplementary Fig. 1).

Scanning electron microscopy (SEM) image showed the crack-free surface (Supplementary Fig. 2). The thickness of the membrane was typically 10 μm (Fig. 2a, inset). The homogeneity of the surface was also suggested by the uniform distribution of C, N, O, and S elements (Supplementary Fig. 3). Fourier transform infrared spectrometer (FTIR), Raman, and solid-state carbon-13 nuclear magnetic resonance (13C NMR) confirmed the chemical structure of the TpPa–SO3H (Supplementary Fig. 4). Transmission electron microscope (TEM) revealed the 3.49 Å interlayer spacing (Fig. 2b, Supplementary Fig. 5). The X-ray diffraction (XRD) pattern of the membrane exhibited a strong peak at 4.9° (Fig. 2c). Thanks to the abundant sulfonic groups on the COF pores, the membrane carries negative surface charge and exhibited hydrophilicity (Supplementary Fig. 6), which is favored for fast proton transport. Additionally, it displayed high thermal stability and high resistance to both acidic and alkaline conditions (Supplementary Fig. 7), some other properties important for practical use.

The selectivity of the membrane was assessed with diffusion dialysis experiments using a two-compartment electrochemical cell, since the osmotic power was generated by the diffusion of protons in our work. As shown in Fig. 2d, the feed reservoir was filled with a complex solution containing protons (0.1 M) and various cations, including Al3+, Mn2+, Fe3+, Co2+, Ni2+, Cu2+, and Zn2+, each cation having a concentration equal to that of the protons. The permeate reservoir was initially filled with deionized (DI) water. The permeability of protons was measured by continuously recording the pH value of the permeate reservoir, and the those of metal cations were measured by recording their concentrations with inductively coupled plasma mass spectrometry (ICP-MS). During 150 min of diffusion, the proton concentration in the permeate reservoir increased almost linearly with time, allowing us to easily calculate the permeability by fitting the data points with a line (Fig. 2e). The permeabilities of metal cations were dramatically lower than that of the protons. Magnified view in Fig. 2f suggests that the diffusion of metal cations also followed the linear behavior. The H+/Al3+, H+/Mn2+, H+/Fe3+, H+/Co2+, H+/Ni2+, H+/Cu2+, and H+/Zn2+ selectivities were calculated to be 673, 248, 3767, 224, 228, 201, and 225, respectively (Fig. 2g), far superior to commercial ion-exchange membranes.

To understand the proton/heavy metal ion selectivity performance, we further conduct characterization of the TpPa–SO3H membrane’s pore size or its high heavy metal ion adsorption capacity (Supplementary Figs. 8, 9). The low N2 adsorption volume and limited heavy metal ion adsorption capacity suggest the presence of small pores. In real industrial waste acid, the concentration of metal cations is typically much lower than that of protons. This means that in the permeate reservoir, the metal cation concentration would be negligible, avoiding side reactions on the HER electrode, such as metal ion deposition (Supplementary Fig. 10).

In practice, the wastewater may also contain organic pollutants. Our membrane can almost completely block them since they have a much larger size than the metal cations. Using tetracycline hydrochloride and levofloxacin hydrochloride as model pollutants (100 ppm for each in the feed reservoir), we conducted diffusion experiments and detected none of them in the permeate reservoir after 12 h using a ultraviolet-visible spectrophotometer (Fig. 2h, i).

Co-transport of Cl− ions and protons
To sustain the chlorine production, the Cl− ions and protons must transport at the same rate from the waste acid reservoir to the recycled acid reservoir, since otherwise the charge balance would break down. However, they should not transport together through the membranes so as not to cancel the net ionic current. The effective transport of Cl− ions was realized by the reversible electrochemical reaction of Cl− ions with the electrodes (Fig. 3a). The membrane, on the other hand, only allows the transport of protons while rejecting Cl− ions. This is crucial to increasing the performance of the system. The Cl− reaction creates a redox potential while the selective diffusion of protons creates an additional diffusion potential, and these potentials can add up, enhancing the power generation performance and increasing the transport rate of protons and Cl− ions (Supplementary Fig. 11).

As a verification of the effective Cl− transport through the electrodes, we inserted a pair of Ag/AgCl electrodes in a two-compartment cell. The feed reservoir contained 5 M HCl, and the permeate reservoir contained DI water. The concentration of protons and Cl− in the permeate reservoir was then measured with time. Fig. 3b shows that Cl− concentration increased linearly with time and maintained the same with that of the protons throughout the experiment to satisfy the charge balance requirement.

As the redox reactions proceed, the Cl on the cathode would be gradually depleted, and the anode would be covered by excess AgCl, hindering further Cl− transport. To sustain the reaction, the anode and cathode Ag/AgCl can be switched. As a verification, we used a two-compartment cell whose feed reservoir was filled with 5 M HCl and the permeate reservoir with 0.5 M HCl. Electrode B (Fig. 3c, now inserted in the permeate reservoir) turned white after 30 min, due to the consumption of AgCl on the electrode (Supplementary Fig. 12). However, after we switched the electrodes for the two reservoirs, Electrode B turned black again in 30 minutes, while the other electrode (Electrode A) turned white, suggesting the reversibility of the reaction. The XRD patterns confirmed that electrode A was entirely converted to Ag, while Electrode B displayed strong AgCl peaks (Fig. 3d).

Indeed, after each electrode switching, the ion transport performance can be recovered. As a verification, we tested two-compartment cells containing 5 M/0.5 M NaCl or 5 M/0.5 M HCl. For both cases, we switched the electrodes every 600 s and measured the transferred charges by recording the ionic current with time (Supplementary Fig. 13). Fig. 3e shows that the transferred charge in each cycle kept almost the same, confirming the stability in repeated tests. For HCl, the amount of the transferred charge is larger than that for NaCl due to the higher permeability of protons. Furthermore, in a 100-cycle test, the performance can still be robustly maintained with periodic electrode regeneration (Supplementary Fig. 14).

The diffusion of protons generates a diffusion potential, Ediff, and the electrochemical reaction of Cl− generates a redox potential Eredox. These potentials can add up, which in turn also facilitates the transport of protons and Cl−. To confirm so, we used the 5 M HCl and 0.5 M NaCl solutions to fill the feed and permeate reservoirs, respectively. Fig. 3f shows that without Ag/AgCl electrodes inserted, the proton permeability was 29.6 mol m−2 h−1. However, with electrodes inserted into the reservoirs, it increased significantly to 90.4 mol m−2 h−1. Such permeability is at least 1 to 2 orders of magnitude higher than common commercial ion-exchange membranes20, and is enabled by the fast proton hopping on hydrogen-bonded networks of the membrane, as we demonstrated in our previous work19. The Cl− would have the same permeability as required by the charge balance.

Based on the Nernst equation, we can calculate the total electrical potential as,
E = Eredox + Ediff = RT/F ln(afeed,Cl−/aperm,Cl−) + 2t+ RT/F ln(afeed,H+/aperm,H+) (1)
where R, T, F, t+ and a represent the gas constant, temperature, Faraday constant, cation transference number and activity of ions, respectively. The subscript feed and perm denote the feed reservoir and permeate reservoir, respectively. E can be measured as the open-circuit voltage (VOC) of the power generator, and the maximum E is reached when the proton selectivity of the membrane against anions is perfect, i.e., t+ = 1. We then varied the afeed/aperm by varying the concentration ratio of the feed reservoir and permeate reservoir from 5 to 100 and measured the VOC by recording sweeping current-voltage curves. Fig. 3g, h show that the measured values agreed well with the theoretical maximum values, confirming the synergistic role of Ediff and Eredox, as well as the almost perfect proton selectivity (Supplementary Fig. 15).

In many previous osmotic power generation works, the selective transport of Cl− ions was carried out by anion-exchange membranes (set-up illustrated in Supplementary Fig. 16)21. However, these membranes have low Cl− selectivity against protons, significantly limiting the osmotic power generation performance. Supplementary Fig. 17 shows that when a commercial anion-exchange membrane was used to transport Cl−, the generated potential was much lower than that when Ag/AgCl electrodes were used.

Power generation performance
Obviously, the osmotic power generation performance determines the chlorine generation performance. A high osmotic power density requires high proton permeability and high proton/anion selectivity, which our membrane already has. To assess the power density, we used a two-compartment cell and established a variety of concentration ratio, ranging from 5 to 500, by filling feed reservoir with 5 M HCl and the permeate reservoir with different concentrations of HCl (Supplementary Fig. 18). The test area (S) of the membrane was set at 7 mm2, larger than most of the previous osmotic power generation membranes. Then we measured current (I) with an external load of varying resistance (R) to calculate the output power density (P)22–24, P = I2R/S (Fig. 4a, Supplementary Fig. 19). The maximum P (Pmax) is typically reached at a moderate R. Fig. 4b shows that Pmax increased with the decrease of the concentration ratio when the concentration ratio is larger than 10, due to the increasing current (Supplementary Fig. 20) induced by the decreasing reservoir resistance of the permeate reservoir. For a lower concentration ratio, Pmax decreased due to the lower VOC (Supplementary Fig. 21). The maximum P reached 55 W m−2 at the moderate concentration ratio of 10. Meanwhile, the high output power density could be maintained for over 15 days without obviously decreasing (Supplementary Fig. 22).

We note that previous studies often tested a microscopic area23, which obviously has limited practical relevance. It is known that an extremely small area usually results in artificially increased Pmax23,25. For a fairer comparison with previous studies, we tested a range of membrane areas (Supplementary Fig. 23). In all cases, the Pmax was significantly higher than those of the previous osmotic generators obtained at the same test area22,23,26–50 (Fig. 4c, Supplementary Data 1). This highlights the advantage of the high proton permeability and selectivity of the membrane, as well as the advantage of generating osmotic power with acidic water.

To put the results into the content of practical application, we added heavy metal ions, including Al3+, Mn2+, Fe3+, Co2+, Ni2+, Cu2+, and Zn2+, 100 mM for each cation, into the feed reservoir to simulate real waste acid. We found that Pmax did not decrease. Instead, it increased to 67 W m−2 possibly due to the elevated conductivity of the feed reservoir (Fig. 4d, Supplementary Fig. 24). Importantly, the concentration of the heavy metal ions in the permeate reservoir was at least 4 orders of magnitude lower than that of the protons, confirming the negligible metal ion leakage (Fig. 4e). We also added tetracycline hydrochloride and levofloxacin hydrochloride (100 ppm for each) into the feed reservoir. Still, Pmax did not decrease (Fig. 4f, Supplementary Fig. 25). Two positively charged organic pollutants, Rhodamine B and Crystal Violet, were used to further evaluate the anti-fouling capability of the membrane (Supplementary Fig. 26). The transport of pollutants was also almost completely blocked, making the membrane anti-fouling (Supplementary Fig. 27).

Spontaneous chlorine production
In this work, we used a Pt electrode for the HER and a graphite electrode for the CER due to its high acid resistance. Before the chlorine production experiments, we assessed the performance of the electrodes using a conventional three-electrode system. As shown in Fig. 5a, the Pt electrode required -0.15 V vs. the reversible hydrogen electrode (RHE) to reach a current density of 100 mA cm−2 for the HER in 0.5 M HCl, whereas the graphite electrode required 1.54 V vs. RHE to drive the CER at the same current density in 5 M HCl. We further performed a two-electrode electrolysis test, and the I-V characteristic curve indicates that a voltage of 1.68 V is required to attain a current density of 100 mA cm−2, which aligns with the polarization curve obtained from the three-electrode setup (Fig. 5b). The slight deviation maybe attributed to the osmotic potential and membrane resistance contributed by the TpPa–SO3H membrane.

Chlorine production often suffers from the competition between oxygen evolution reaction (OER) and CER, which affects the faradaic efficiency (FE). However, we note that OER is highly unfavorable in the extremely low-pH environments (Supplementary Fig. 28). Indeed, we quantified the produced Cl2 using the iodometric titration method at various current densities to calculate the FE51,52. The results, presented in Fig. 5c, showed a high FE of 98.4% for the CER at 10 mA cm−2. Even at a very high current density of 100 mA cm−2, the FE reached 90.7%.

Equation 1 tells us that each osmotic power generation unit can produce a maximum of 0.184 V when 5 M/0.5 M HCl was used for chlorine production. This mean that 8 units are required to initiate the spontaneous production of chlorine. Considering the electrochemical overpotential, internal resistance voltage drops, we tested serially connected units from 8 to 18 (Fig. 5d, Supplementary Fig. 29). Then, the terminal electrodes of the serial osmotic generators were replaced with a graphite electrode and a Pt electrode. The Ag/AgCl electrodes of the feed and permeate reservoirs were switched every 600 s for continuous operation. The hydrogen generated on the Pt electrode was collected using the water drainage method, while the chlorine production rate was calculated from the ionic current data. Obviously, the VOC of the system increased linearly with the number of units in series, reaching 2.95 V when the unit number was 18 (Fig. 5e), resulting in an enhanced chlorine production rate. When the number of units were lower, the reaction was either slower or could not be initiated. As shown in Supplementary Fig. 30, the average current density was 36.8 mA cm−2, meaning that the FE would be larger than 97% (the FE at 50 mA cm−2). Figure 5f, g show that the hydrogen and chlorine production rate reached ~154 L m−2 h−1 along with nearly 100% HER Faraday efficiency (Supplementary Fig. 31). Such a production rate can be maintained for at least 7 days (Fig. 5h).

Extending to other chloride-containing brines
Obviously, all chloride-containing brines can be used to harvest the osmotic energy and produce chlorine if a brine with lower salinity is readily available. For the waste acid, the lower salinity brine is the recycled acid. For other brines, natural water could be used as the low salinity brine. For example, the waste brines of the desalination plant have higher salinity (typically >1 M Cl−) than seawater, and they are often discharged into seawater, providing the foundation for osmotic power generation (Supplementary Fig. 36).

We found that a high power density of >6 W m−2 can be achieved using simulated waste desalination brine generated by the membrane distillation method (5 M NaCl)9 and simulated seawater (0.5 M NaCl) using a TpPa–SO3H membrane (Supplementary Fig. 37). When combined with industrial waste heat, the power density could be further improved to 8.4 W m−2 (Supplementary Fig. 38). The power density is lower than that in Fig. 4c, due to the fact that our membrane was optimized for proton transport. For chlorine production, we used the dimensionally stable anode (DSA) electrode as the anode, which was widely used in the chlor-alkali industry because of its good stability and chlorine selectivity. The high selectivity of the electrode was further confirmed in Supplementary Fig. 39. The chlorine and hydrogen production rates of the integrated system reached approximately 12 L m−2 h−1, which is lower than that produced using waste acid, but confirms the versatility of our spontaneous chlorine production method (Supplementary Fig. 40).

Methods
Preparation of the TpPa–SO3H membrane
0.064 g (0.304 mmol) 1,3,5-triformylphloroglucinol (Tp) and 0.086 g (0.456 mmol) 2,5-diaminobenzenesulfonic acid (Pa-SO3H) were dissolved in 3 mL N-methylpyrrolidone (NMP) and dimethyl sulfoxide (DMSO), respectively. After sonication at room temperature for 20 min, the two solutions were mixed and stirred thoroughly. Then the mixture was sonicated for 20 min to remove the solution bubble. The mixed solution was cast onto a clean glass slide (10 cm × 10 cm) and kept on the heating plate at 60 °C for 3 days to form the TpPa–SO3H. The self-standing COF membranes could peel off from the glass slides by put into water. The membrane was washed thoroughly with NMP, DMSO and deionized water to remove residual monomer fragments before it dried in the oven at 60 °C. Then the membrane was cut into various sizes for further characterization and tests.

Preparation of the Ag/AgCl electrode
The Ag/AgCl electrode was fabricated via a chemical oxidation method. Silver plates (8 mm × 30 mm) were immersed in sodium hypochlorite solution for 30 min, during which hypochlorite reacted with the silver surface to form a layer of silver chloride. The initially bright silver plates turned glossy gray upon conversion. Following the reaction, the electrodes were thoroughly rinsed with deionized water to remove residual hypochlorite. The resulting Ag/AgCl electrodes were then used for subsequent tests.

Material characterization
Scanning electron microscopic images were taken with a Hitachi S-4800 field emission scanning electron microscope (SEM) with an accelerating voltage of 10 kV. The nanosheets were characterized by transmission electron microscopy (TEM, JEM F200). Powder X-ray diffraction (XRD) patterns were carried out in the 2θ range of 3° to 40° at room temperature using a Bruker D8 Advance X-ray diffraction diffractometer equipped with Cu Kα radiation (λ = 1.5147 Å). The surface chemical properties were characterized by Fourier transform infrared spectrometry (FTIR, Bruker Vertex 70v), Raman spectrometry (Nicolet 5700) and X-ray photoelectron spectrum (XPS, Thermo Scientific K-Alpha). The solid-state 13C NMR spectra were recorded on an AVANCE Ⅲ HD 500 MHz NMR spectrometer with cross-polarization magic angle-spinning (CP-MAS). The absorption spectra were recorded on Ultraviolet-visible (UV-Vis) spectrophotometer (Hitachi U-3900, Japan). The contact angles of the membranes were measured by a Contact Angle Surface Analyzer (LSA100) at room temperature. The Anton Paar SurPASS is used to determine the Zeta potential of the membrane surface. Thermogravimetric analysis (TGA) was performed on a STA 8000 thermogravimetric analyzer under a N2 atmosphere at a heating rate of 10 °C min−1.

Diffusion experiment
To investigate the ion transport behavior, the Ion selectivity tests were performed by concentration-driven diffusion using a self-made H-type device. The COF membrane was tightly mounted between two reservoirs, with one reservoir filled with 10 mL electrolyte feed solution and the other reservoir filled with 10 mL deionized water as the starting permeate solution. The concentration of cations was measured by Inductively Coupled Plasma mass spectrometry (ICP-MS, Agilent 7850), the proton concentration was measured with a pH meter (LAQUA F-71) and the other anion concentrations were determined by Ion Chromatography (IC, ECO).

Adsorption of heavy metal ions and positively charged pollutants
The adsorption of heavy metal ions by TpPa–SO3H membranes was carried out using a 50 mL mixed solution containing 0.001 M of each metal ion: Al³⁺, Mn²⁺, Fe³⁺, Co²⁺, Ni²⁺, Cu²⁺, and Zn²⁺. A known mass of the TpPa–SO3H membrane was immersed in the initial metal ion solution. The system was maintained under magnetic stirring at 600 rpm for 12 h to facilitate the adsorption process. Following this, the concentration of metal ions in the solution was measured to determine the extent of metal ion adsorption onto the membrane. The amount of metal ions adsorbed was calculated using the following formula:
Q = (c0 − ct) × V / m
where Q represents the adsorption capacity (mg/g), V is the volume of the solution (mL), c0 is the initial concentration of metal ions (mg/L), ct is the concentration after adsorption (mg/L), and m is the mass of the TpPa–SO3H membrane (g).

The adsorption experiments for positively charged organic pollutants were performed similarly, using 50 ppm solutions of Rhodamine B and Crystal Violet, respectively, as model pollutants. The adsorption process was conducted under the same experimental conditions as described for the heavy metals, with a stirring speed of 600 rpm for 12 h. After adsorption, the concentrations of Rhodamine B and Crystal Violet in the solution were quantified by a UV-Vis spectrophotometer, allowing for the determination of the adsorption capacity of the TpPa–SO3H membrane.

Osmotic energy conversion setup
The COF membrane was placed between two reservoirs. A pair of Ag/AgCl electrodes (8 mm × 40 mm) were employed as the work electrode. The effective test area for energy harvesting was 7.06 mm2 if without further explanation. The VOC and ISC were measured from the I-V curve tests, and the sweep voltage was from −1 V to 1 V. The equivalent circuit of the experimental setup can be represented by Supplementary Fig. 11, VOC, Eredox, Ediff, and Rmem represent the measured potential, the redox potential generated by the unequal potential drop at the electrode-solution interface and the diffusion potential contributed by the ion-selective membrane, and the internal resistance of the membrane, respectively. The measured VOC actually consists of two parts (i.e., Eredox and Ediff). In our work, the salt bridges were used to eliminate the redox potential of the system. Then, the corresponding output power density can be calculated as P = I2R/S, where I, R, and S represent ion current, load resistance and effective testing membrane area, respectively. All the osmotic power generation tests were carried out by Source Meter (Keithley 2450).

For a given concentration ratio, the cation transference number (t+) can be calculated as
t+ = 1 / (2 * (R * T * ln(γi * cL / γf * cH) / (z * F * E) + 1))
where R, T, z, F, γ, cH and cL represent the gas constant, temperature, charge valent, Faraday constant, activity coefficient of ions, high and low ion concentrations, respectively. The activity coefficients used in the calculation are provided in Supplementary Table 4.

Measurement of redox potential and diffusion potential
The experimental setup is shown in Supplementary Fig. 16. The measurements were performed using Ag/AgCl reference electrodes under a controlled environment. The system comprised two compartments separated by either a solution interface (for redox potential measurement) or a commercial anion exchange membrane (AEM) (for diffusion potential measurement). The compartments were filled with electrolyte solutions of varying concentrations to establish concentration gradients.

The redox potential of the Ag/AgCl electrode and the diffusion potential of the AEM were determined by recording the I–V curve. The potential was scanned over a range of −1 V to +1 V using a Source Meter (Keithley 2450), with the Ag/AgCl reference electrodes serving as both the working and reference electrodes in high and low-concentration electrolyte solutions. The intercept of the I–V curve on the voltage axis was recorded as the redox potential of the electrode or the osmotic potential of the AEM.

Electrochemical hydrogen evolution and chlorine evolution evaluation
The performance of electrodes was tested on an electrochemical workstation instrument (Biologic SP-300). The experimental setup is based on a conventional three-electrode system, with the working electrode (WE) (Pt, graphite or DSA), while the counter electrode (CE) was a graphite rod. The reference electrode (RE) used in this study was an Ag/AgCl electrode with a salt bridge contain the statured KCl solution. The effective area of the working electrode is 7.06 mm2. The polarization curves were obtained by employing a scanning rate of 10 mV s−1. All the recorded potentials of the polarization curves in this study were without conducting iR compensation and converted to reverse hydrogen electrode (RHE) values based on equation:
ERHE = E + 0.197 V + 0.059 pH

Two-electrode electrolysis and faradaic efficiency evaluation
The two-electrode electrolysis and faradaic efficiency evaluation were performed on the Source Meter (Keithley 2450). Within the H-type electrolytic cell, the working electrode (Graphite or DSA) was connected with the anode, while the counter electrode (Pt) was connected with the cathode. The “Voltage sweep” method was conducted to obtain the I–V characteristic curve with the scanning rate of 10 mV s−1. The faradaic efficiency of chlorine was determined through iodometric titration. First, 50 mL of electrolyte was added to the H-type electrolytic cell. Then the electrolyte was purged with N2 for 10 min. Following this, the “Current bias” method was conducted to maintain the current density from 10 to 100 mA cm-2. After 100 s of tests, the obtained Cl2 was trapped by the immediate addition of an excess amount of potassium iodide (Cl2 + 3I− → I3− + 2Cl−), and the solution was kept to attain equilibrium. The obtained amount was determined by performing Iodometric titration versus 0.1 mM Na2S2O3 (I3− + 2S2O32− → 3I− + S4O62−) to calculate the chlorine production in the experiment. The calculation method for chlorine faradaic efficiency was:
FE (%) = Experimental yield / Theoretical yield = (0.1 mM * volume of Na2S2O3 / 2) / (i / 2F) * dt

Integrated device performance evaluation
In the spontaneous chlorine and hydrogen production system with simultaneous acidic wastewater recovery, a varying number of pairs of units, each equipped with TpPa–SO3H membranes, were connected in series with Ag/AgCl electrodes. Graphite and Pt electrodes were used at both terminals, forming a fully integrated osmotic energy-driven device for chlorine and hydrogen production. 5 M HCl aqueous solution and 0.5 M HCl aqueous solution were injected into each pair of cells to simulate an acidic wastewater recovery system.

For the desalination brine-seawater system, 30 pairs of units were connected similarly, with DSA and Pt electrodes at the respective terminals to generate chlorine and hydrogen. The electrolytes at both ends consisted of 5 M NaCl (pH=2) and 1 M NaOH, respectively. Additionally, each pair of cells was injected with a 5 M NaCl aqueous solution and a 0.5 M NaCl aqueous solution to simulate a membrane distillation brine-seawater system.

Peer review rounds
[{"round": 1, "reviewer_comments": {"reviewer_1": "The authors present an innovative method for spontaneous chlorine production from chloride-containing brines using sulfonated COF membranes. This approach, which requires no external energy input, offers a significant advantage over conventional electrolysis methods. The manuscript is well-written, logically structured, and presents promising results with potential for substantial impact in low-carbon chemical synthesis and wastewater treatment. However, several critical points require clarification prior to publication.\n1. The authors demonstrate that the membrane maintains structural stability in acidic environments over a 15-day period. However, the long-term performance of osmotic energy harvesting under acidic conditions, which is essential for sustained chlorine production, was not addressed. The authors are recommended to conduct additional experiments to evaluate this performance.\n2. While platinum is known for its high catalytic activity in the hydrogen evolution reaction (HER), the Faradaic efficiency of hydrogen evolution in hydrochloric acid should be quantified or calibrated, as side reactions may impact overall efficiency.\n3. The porous structure and metal ion adsorption capacity are critical determinants of heavy metal ion retention performance. The authors should provide comprehensive characterization of these properties to strengthen their findings.\n4. Building on the above, the authors should discuss whether proton/heavy metal ion selectivity arises from the membrane’s small pore size or its high heavy metal ion adsorption capacity. If adsorption drives selectivity, the potential for adsorption saturation and its impact on selectivity should be clarified, given its importance to chlorine production performance.\n5. As industrial chlor-alkali processes often operate at elevated temperatures, the authors are encouraged to evaluate the influence of temperature on ion transport behavior and overall energy conversion efficiency.", "reviewer_2": "The authors presented a spontaneous chlorine production method by using the inherent 18 energy and chloride ions present in Chloride-Containing brines. The idea is sort of interest. My concern are shown as following:\n1. Preparation of the TpPa-SO3H membrane and Preparation of the Ag/AgCl electrode were checked by the reviewer and not too much novelty of preparation of materials was founded. So what is the main novelty of this paper? Just the processes?\n2. Author Contributions, L.J. supervised the project, Does L.J. really part part in the project? maybe he is really busy.\n3. How about the cost of the spontaneous chlorine production method? maybe a rough cost should be given.\n4. How long does the electrode last with the consideration of futher application?", "reviewer_3": "This manuscript presents a novel and potentially impactful electrochemical system that integrates osmotic energy harvesting with acid recovery to drive chlorine production from chloride-rich brines. The process design is innovative and technically robust, bridging electrochemical engineering with advanced membrane materials. The use of proton-selective COF membranes is a key enabler, allowing selective proton transport while excluding other ions. Overall, the study is thorough, and the results are well-analysed. The manuscript is suitable for publication in Nature Communications, pending minor revisions. The following points should be addressed:\n1. As a central schematic, Figure 1 should be improved for clarity. Electrochemical reactions and ion transport directions across compartments should be clearly labelled. The concept of continuous operation, via electrode switching and reservoir refreshing, should be provided and briefly explained. Including a comparative schematic contrasting the proposed process with conventional chlor-alkali systems would help highlight advantages and limitations.\n2. A diagram or discussion summarizing redox potentials of the involved reactions would be helpful. Specifically, the authors should clarify how other metal ions (e.g., heavy metals) in the brine are prevented from undergoing undesirable side reactions such as deposition.\n3. The long-term stability of the Ag/AgCl electrodes is not sufficiently addressed. Prior studies have noted issues with AgCl layer buildup increasing resistance. While images and SEM/EDX analysis are provided (Fig. 3c, Fig. S9), additional discussion or experimental data on electrode durability over extended operation would strengthen the manuscript.\n4. The authors predominantly use 5 M HCl in their experiments. However, acid concentration in real feedstocks may vary. A sensitivity analysis exploring system performance across a range of acid concentrations would enhance the relevance of the study to practical applications.\n5. While the TpPa-SO₃H COF membrane exhibits high proton selectivity and stability, testing membranes with lower selectivity or across a range would help define the operational limits and robustness of the system.\n6. Some data (e.g., Fig. 2g) lack error bars, and the selectivity values should be carefully reviewed for accuracy. The authors are encouraged to ensure all reported values are statistically supported.\n7. The study examines two model organic molecules (TCH and LFH), but real waste acid streams may contain a broader range of contaminants, including positively charged organics. Potential membrane fouling through adsorption (especially positively charged molecules) should be discussed, and membrane integrity after long-term exposure should be evaluated.\n8. The system design confines H₂ and Cl₂ generation to the terminal electrodes, which may constrain scalability. The authors should comment on potential design adaptations to improve scalability for industrial applications.\n9. The rationale for using 18 cells (achieving ~2.95 V) to drive spontaneous Cl₂ evolution should be more clearly linked to the electrochemical requirements and electrode polarization behaviour.\n10. Generating Cl₂ in an open system raises safety concerns. While HER and CER are spatially separated, the authors should discuss gas handling and containment strategies to support practical implementation."}, "author_response": "Response to Reviewer 1\n(Q=question/comment; A=answer; changes to the manuscript or supplementary information\nhighlighted in red)\nComments: The authors present an innovative method for spontaneous chlorine production from\nchloride-containing brines using sulfonated COF membranes. This approach, which requires no\nexternal energy input, offers a significant advantage over conventional electrolysis methods. The\nmanuscript is well-written, logically structured, and presents promising results with potential for\nsubstantial impact in low-carbon chemical synthesis and wastewater treatment. However, several\ncritical points require clarification prior to publication.\nResponse: Thank you for the careful review and constructive comments concerning our manuscript.\nWe have revised our manuscript carefully and thoroughly according to your suggestions.\nQ1: The authors demonstrate that the membrane maintains structural stability in acidic environments\nover a 15-day period. However, the long-term performance of osmotic energy harvesting under acidic\nconditions, which is essential for sustained chlorine production, was not addressed. The authors are\nrecommended to conduct additional experiments to evaluate this performance.\nA1: Thank you for the important suggestion! Following the suggestion, we performed a long-term\nstability test of power generation performance over 15 daysunder 5 M/0.5 M HCl condition. As shown\nin Supplementary Fig. 22, the maximum power density remains essentially unchanged from 1 to 15\nday with no obvious decrease. These results indicate that the TpPa-SO H membrane could maintain\n3\nhigh proton selectivity and stable power output, thereby supporting sustained chlorine production in\nour system.\nSupplementary Fig. 22. Long-term stability of power generation performance over 15 days.Error\nbars represent the standard deviation from at least three independent tests.\nWe revised the first paragraph of Power generation performance section,\n“...The maximum P reached 55 W m-2 at the moderate concentration ratio of 10. Meanwhile, the\nmax\nhigh output power density could be maintained for over 15 days without significant decrease\n(Supplementary Fig. 22).”\nQ2: While platinum is known for its high catalytic activity in the hydrogen evolution reaction (HER),\nthe Faradaic efficiency of hydrogen evolution in hydrochloric acid should be quantified or calibrated,\nas side reactions may impact overall efficiency.\nA2: Thank you for your helpful suggestion! We quantified the HER Faradaic efficiency under the\nhydrochloric-acid conditions used in this work by integrating the current–time traces (Supplementary\nFig. 30) to obtain the theoretical H yield. We then compared this calculated volume with the\n2\nexperimentally collected H . As shown in Supplementary Fig. 31, the measured points closely overlap\n2\nthe calculated trace over the entire time, indicating ~100% Faradaic efficiency within the test.\nSupplementary Fig. 31. Comparison between measured and Calculated hydrogen volumes.\nWe revised the third paragraph of Spontaneous chlorine production section,\n“...Fig. 5f and g show that the hydrogen and chlorine production rate reached ~154 L m-2 h-1 along\nwithnearly 100% HER Faraday efficiency (Supplementary Fig. 31).”\nQ3: The porous structure and metal ion adsorption capacity are critical determinants of heavy metal\nion retention performance. The authors should provide comprehensive characterization of these\nproperties to strengthen their findings.\nQ4: Building on the above, the authors should discuss whether proton/heavy metal ion selectivity arises\nfrom the membrane’s small pore size or its high heavy metal ion adsorption capacity. If adsorption\ndrives selectivity, the potential for adsorption saturation and its impact on selectivity should be clarified,\ngiven its importance to chlorine production performance.\nA3&A4: Thank you for your raising concerns regarding the porous structure and metal ion adsorption\ncapacity of the TpPa-SO H membrane, which are critical for its heavy metal ion retention performance.\nthe results of which are presented in Supplementary Figs. 8 and 9, respectively.\nThe nitrogen adsorption-desorption isotherm (Supplementary Fig. 8) reveals an extremely low\nadsorption quantity, suggesting a small surface area for the TpPa-SO H membrane. The adsorption and\n3\ndesorption curves exhibit typical behavior with minimal hysteresis, indicative of a compact structure\nwith limited pore volume. We hypothesize that the disordered surface structure of the membrane\nhinders nitrogen molecule penetration, contributing to this low adsorption.\nFurther evidence comes from the heavy metal adsorption experiments (Supplementary Fig. 9).\nOptical images show negligible changes in the membrane’s color and structure before and after\nexposure to a 0.001 M feed solution of various heavy metals (Al, Mn, Fe, Co, Ni, Cu, Zn).\nQuantification of metal ion adsorption (Q, mg/g) confirms very low uptake, with values consistently\nbelow 25 mg/g across all tested metals. This low adsorption capacity, corroborated by additional\nexperiments with positively charged pollutants showing similarly minimal uptake (Supplementary Fig.\n26), suggests that the membrane’s selectivity for protons over heavy metal ions is not driven by\nadsorption capacity but rather by its small pore size. The compact structure and restricted pore\naccessibility likely act as a physical barrier, preferentially allowing proton migration while rejecting\nlarger or multi-valent cations and pollutants.\nRegarding the potential for adsorption saturation, the low adsorption capacity indicates that\nsaturation is unlikely to significantly impact selectivity under the conditions tested, as the membrane\ndoes not rely on adsorption-driven retention. This is advantageous for chlorine production, as it\nminimizes the risk of performance degradation due to pollutant buildup.\nAccording to these findings, we revised the second paragraph of High proton selectivity of the\nmembrane section,\n“To understand the proton/heavy metal ion selectivity performance, we further conduct\ncharacterization of TpPa-SO H membrane’s pore size or its high heavy metal ion adsorption capacity\n3\n(Supplementary Fig.8-9). The low N adsorption volume and limited heavy metal ion adsorption\n2\ncapacity suggest the presence of small pores.”\nAdsorption of heavy metal ions and positively charged pollutants\nThe adsorption of heavy metal ions by TpPa-SO H membranes was carried out using a 50 mL mixed\n3\nsolution containing 0.001 M of each metal ion: Al³⁺, Mn²⁺, Fe³⁺, Co²⁺, Ni²⁺, Cu²⁺, and Zn²⁺. A known\nmass of the TpPa-SO H membrane was immersed into the initial metal ion solution. The system was\n3\nmaintained under magnetic stirring at 600 rpm for 12 hours to facilitate the adsorption process.\nFollowing this, the concentration of metal ions in the solution was measured to determine the extent\nof metal ion adsorption onto the membrane. The amount of metal ions adsorbed was calculated using\nthe following formula:\n𝑐0 −𝑐𝑡\nwhere Q represents the adsorption capacity (mg/g), V is the volume of the solution (mL), c is the\n𝑄 = ×𝑉 0\n𝑚\ninitial concentration of metal ions (mg/L), c is the concentration after adsorption (mg/L), and m is the\nt\nmass of the TpPa-SO H membrane (g).\nusing 50 ppm solutions of Rhodamine B and Crystal Violet, respectively, as model pollutants. The\nadsorption process was conducted under the same experimental conditions as described for the heavy\nmetals, with a stirring speed of 600 rpm for 12 hours. After adsorption, the concentrations of\nRhodamine B and Crystal Violet in the solution were quantified by UV-Vis spectrophotometer,\nallowing for the determination of adsorption capacity of the TpPa-SO H membrane.\nand desorption curves reveal typical behavior, with minimal hysteresis, suggesting a compact structure\nwith limited pore volume.\nSupplementary Fig. 9. Adsorption of heavy metals onto the TpPa-SO H membrane.\n3\na, Optical images of the membrane before and after exposure to the metal ions, showing noticeable\nchanges in the appearance. b, Quantification of metal ion adsorption (Q, mg/g) for various heavy\nmetals (Al, Mn, Fe, Co, Ni, Cu, Zn) from a 0.001 M feed solution. Error bars represent the standard\ndeviation from at least three independent tests.\nQ5: As industrial chlor-alkali processes often operate at elevated temperatures, the authors are\nencouraged to evaluate the influence of temperature on ion transport behavior and overall energy\nconversion efficiency.\nA5: Thank you for this constructive suggestion, which highlights a key factor in bridging our\nspontaneous chlorine production system with industrial chlor-alkali processes, where elevated\ntemperatures (typically > 60 °C) are common to enhance kinetics and efficiency. To address this, we\nconducted additional experiments evaluating the temperature dependence of ion transport and energy\nharvesting performance using simulated desalination wastewater. Specifically, we measured current-\nvoltage (I-V) curves and corresponding areal power densities at temperatures ranging from 298 K to\nrealistic industrial conditions.\nAs shown in the newly added Supplementary Fig. 39 (panels a, c, e), the current density increases\nmarkedly with rising temperature across all gradients, with the I-V slopes steepening at higher\ntemperatures. This enhancement is attributed to accelerated ion diffusion rates, as governed by the\nArrhenius relationship, where thermal energy reduces the activation barrier for ion mobility through\nthe sulfonated COF membrane channels. Consequently, the peak power densities (panels b, d, f) also\nrise significantly: relative gains from 298 K to 333 K are approximately 138% for 1 M/0.5 M, 45% for\nthat operating at elevated temperatures could substantially boost the overall energy conversion\nefficiency and chlorine production rates in our system, particularly for higher-salinity brines where\nmass transport limitations are more pronounced at ambient conditions.\nImportantly, the TpPa-SO H COF membrane exhibited thermal stability up to 333 K (as confirmed\n3\nby prior TGA data in Supplementary Fig. 7a, showing minimal weight loss below 300 °C), with no\ndegradation in selectivity or structural integrity observed during testing. This temperature resilience\nsupports the system's compatibility with industrial heat sources, such as waste heat from desalination\nplants, potentially further reducing the carbon footprint. For instance, desalination processes like multi-\neffect distillation (MED) or membrane distillation (MD) often discharge wastewater at temperatures\nof 40–70 °C or higher, which can be directly integrated to preheat the feed brines in our system,\nenhancing ion transport without additional energy input. This waste heat utilization not only aligns\nwith sustainable practices in desalination facilities—where up to 20–50% of input energy is lost as\nlow-grade heat—but also enables co-location of our chlorine production modules with existing plants\nfor efficient resource recovery.\nWe revised the second paragraph of Extending to other chloride-containing brines section,\n“...When combined with industrial waste heat, the power density could be further improved to 8.4 W\nm-2(Supplementary Fig. 39).”\nSupplementary Fig. 38 Effect of temperature on osmotic-energy harvesting performance.\na, c, e, I–V curves of the device measured at 298, 303, 313, 323 and 333 K under NaCl salinity gradients\nof 1 M/0.5 M (a), 3 M /0.5 M (c) and 5 M /0.5 M (e). b, d, f, Corresponding areal power density as a\nfunction of external load resistance. The peak power density increases with temperature; relative gains\nfrom 298 K to 333 K are ~138% (1 M/0.5 M), 45% (3 M/0.5 M) and 39% (5 M/0.5 M), respectively.\nResponse to Reviewer 2\n(Q=question/comment; A=answer; changes to the manuscript or supplementary information\nhighlighted in red)\nComments: The authors presented a spontaneous chlorine production method by using the inherent 18\nenergy and chloride ions present in Chloride-Containing brines. The idea is sort of interest. My concern\nare shown as following:\nResponse: Thank you for the important comments on our paper. We have revised our manuscript\naccording to your suggestions. Point-by point responses to the comments are enclosed.\nQ1: Preparation of the TpPa-SO3H membrane and Preparation of the Ag/AgCl electrode were checked\nby the reviewer and not too much novelty of preparation of materials was founded. So what is the main\nnovelty of this paper? Just the processes?\nA1: Thank you very much for the comment! We acknowledge that the novelty of this paper does not\nlie in the synthesis of these materials. TpPa-SO H membrane was reported, although previous studies\n3\nfocused on the synthesis of TpPa-SO H with high porosity and high crystallinity. Our membrane was\n3\nintentionally made amorphous to reduce the pore size and increase the proton/cation selectivity.\nThe novelty of our work lies in the concept and method of using wastewater or natural water\nfor spontaneous generation of chlorine.\nChlorine, as one of the most critical basic chemicals and a cornerstone of the chemical industry,\nis traditionally produced via the energy-intensive chlor-alkali process, which consumes substantial\nelectricity and carries a significant carbon footprint. Our method harnesses the inherent osmotic energy\nand chloride ions present in chloride-containing brines—such as acidic wastewater or desalination\nwastewater—to achieve spontaneous chlorine production without external power supply.\nIn this regard, using established material systems could make our work more impactful since it\ndemonstrates the general applicability of the concept better. We revised our Discussion section to\nhighlight this, “This work used a reported membrane for the demonstration of general applicability,\nbut novel membranes could enhance the performance.”\nQ2: Author Contributions, L.J. supervised the project, Does L.J. really take part in the project? maybe\nhe is really busy.\nA2: Thank you very much for asking! L.J. has rich experience in osmotic energy harvesting and\ncontributed to supervising the design of the energy harvesting experiments as well as the corresponding\ndiscussions.\nHowever, L.J. was not the primary supervisor in our work. This work was mainly supervised by\nJ. G.. Consequently, L.J. was not listed as a corresponding author. He was listed as the 6th author.\nTherefore, we believe the author list is proper. To avoid confusion, we further clarified the author\ncontribution, “The project was mainly supervised by J.G., with help from L.J.”\nQ3: How about the cost of the spontaneous chlorine production method? maybe a rough cost should\nbe given.\nA3: Thank you very much for the constructive question, which is indeed crucial for assessing its\neconomic viability. To address this, we have added a detailed cost analysis in the revised\nSupplementary Information as Supplementary Note 1: \"Chlorine production cost analysis,\" along with\nsupporting data in Supplementary Tables 3 and 4. This section provides a rough cost comparison\nbetween our approach and the conventional chlor-alkali process, focusing on key materials\n(membranes and electrodes), overall system expenditure, and the elimination of electricity costs.\nIn our analysis, for laboratory demonstrations, silver foil electrodes were used, but for large-scale\nproduction, silver can be loaded onto cost-effective carbon cloth substrates at a loading of 0.01 g/cm2,\nmaintaining stability over cycles while reducing electrode costs to approximately $500/m2\n(Supplementary Table 3) (Nature Water 2024, 1091-1101; Science 2024, 385, 1438-1444). The TpPa-\nSO H COF membrane, derived from material synthesis and fabrication expenses, is priced at around\n3\n$74/m2 due to the small thickness. For a system with 18 serially connected units (as optimized in our\nexperiments), the total upfront cost for membranes and Ag/AgCl electrodes is approximately\n$18300/m2, which is about 4-fold higher than the $4400/m2 for conventional chlor-alkali using Nafion\nmembranes.\nHowever, the economic advantage of our method results from eliminating electricity consumption\nduring electrolysis, the major cost driver in traditional processes that accounts for over 50% of total\nexpenses (typically ~$180 per ton of Cl at an electricity price of 0.075 $/kWh, assuming ~2400\n2\nkWh/ton based on industrial benchmarks; Supplementary Table 4) (Communications, E. C. The\nElectrolysis process and the real costs of production, Euro Chlor, 2018). In our method, osmotic energy\nfrom brines drives the process spontaneously, reducing the projected total production cost to ~$244\nper ton of Cl (assuming comparable brine and chemical treatment costs, and similar\n2\nmembrane/electrode stability to commercial systems). This represents an approximate one-third\nreduction compared to the conventional chlor-alkali cost of ~$360 per ton. The analysis assumes\nother operational costs (e.g., brine processing) remain similar, but our system's modular design could\nfurther lower them through waste integration.\nImportantly, our method is insulated from electricity price fluctuations, which are increasingly\nvolatile due to global energy market dynamics. As electricity prices escalate, our approach becomes\neven more economically attractive, amplifying the cost savings. For instance, if electricity costs double\nto $0.15/kWh, conventional processes could exceed $500/ton, widening the gap to over 50% savings\nin our system. This long-term economic benefit, combined with the potential for one-third cost\nreduction per ton of chlorine, underscores the scalability and sustainability of our innovation.\nWe have incorporated this analysis into the revised Supplementary Information and Discussion\nsection to highlight the method's practical advantages. Thank you for this feedback, which strengthens\nthe manuscript's emphasis on economic feasibility.\nWe revised our Discuss section,\n“...Compared to traditional chlorine production, our method requires no electricity input and thus has\nalmost no carbon emission considering that the carbon emission of chlorine production primarily\ncomes from the energy consumption (Fig. 5i, Supplementary Table 2),while also significantly\nreducing production cost (see Supplementary Note 1).”\nWe revised our Supplementary Information,\nSupplementary Note 1. Chlorine production cost analysis.\nTo assess the feasibility and economic advantage of our approach, we present a cost comparison\nbetween our system and the conventional chlor-alkali process. The analysis focuses on the cost of key\nmaterials, including membranes and electrodes, as well as the overall system expenditure.\nWhile laboratory experiments utilized silver foil electrodes, However, for large-scale production,\nsilver can be loaded onto carbon cloth or similar materials, significantly reducing the electrode cost\nwhile maintaining stability during cycles. We estimate the silver loading to be 0.01 g/cm², and\nconsidering the carbon cloth and processing costs, the price of electrode is approximately 500 $/m2.\nThe TpPa-SO H COF membrane, estimated from material synthesis and fabrication costs, is priced at\n3\naround $74/m². As shown in Supplementary Table 3, for a system comprising 18 units in series, the\ntotal upfront cost for the membrane and Ag/AgCl electrodes is approximately 18300 $/m², which is\nnotably higher than the 4,400 $/m² required by conventional chlor-alkali systems.\nHowever, a significant portion of the cost of chlorine production in conventional systems stems\nfrom the electricity consumed during electrolysis process (~2400 kWh/ton based on industrial\nbenchmarks; Supplementary Table 4). While the initial investment of our system exceeds that of\nconventional systems, it eliminates the need for electricity during the electrolysis process. If we assume\nthat other costs (such as brine and treatment chemicals) remain comparable to the traditional chlor-\nalkali industry, and that the stability of our membranes and electrodes is similar to commercial cation\nexchange membranes, we estimate that the production cost of chlorine could be reduced by\napproximately one-third. This cost reduction is a direct result of eliminating the electricity\nconsumption in the electrolysis stage, which contributes to over 50% of the costs in traditional chlorine\nproduction.\nSupplementary Table 3. Comparison of key materials used in chlorine production systems.\nMaterials Price ($/m2)\nCation exchange membrane (Nafion N424) * 4400\nTpPa-SO H membrane 74\n*https://ion-power.com/membranes/\n**https://www.fuelcellearth.com/fuel-cell-products/p50/\nhttps://clayrevolution.com/products/silver-powder-999?variant=40523555569746\nSupplementary Table 4. Cost comparison of producing 1 ton of chlorine using the conventional\nchlor-alkali system and the method presented in this work.\nUnit and price Conventional chlor-alkali system ($) This work ($)\nElectricity 0.075 $/kwh 180 ~ 0\nCation exchange membrane 4400 $/m2 20\nTpPa-SO H membrane and\n3\n1074 $/m2 84\nAg/AgCl electrode\nOthers ~ 160 ~ 160\nTotal 360 244\nQ4: How long does the electrode last with the consideration of further application?\nA4: Thank you very much for the question! In our original manuscript, Fig. 3e and Supplementary Fig.\n13 reported stable Ag/AgCl performance over 15 consecutive charge–discharge cycles. We have now\nperformed an extended 100-cycle test using 5 M/0.5 M HCl (Supplementary Fig. 14). The transferred\ncharge in each cycle remained largely stable during the early stages, but after ~25 cycles we observed\na ~20% decreasing, attributed to surface passivation/roughness changes of the Ag/AgCl layer\n(Corrosion Science 2021, 184; Electrochim Acta 2011, 56, 2781-2791). However, this decline can be\nreversed through periodic regeneration, i.e., e.g., immersing in sodium hypochlorite every 25 cycles\n(indicated by arrows in Supplementary Fig. 14b).\nWe revised the fourth paragraph of Co-transport of Cl- and proton, and its importance for\ncontinuous operation and high transport rate section,\n“...Furthermore, in a 100-cycle test, the performance can still be robustly maintained with periodic\nelectrode regeneration (Supplementary Fig. 14).”\nSupplementary Fig. 14. Performance of Ag/AgCl electrodes over 100 cycles in a 5 M/0.5 M HCl\nsolution.\na, I-t curve during the 100 cycles. b, Comparison of transferred charge (C) for each cycle, illustrating\nthe consistent charge transfer throughout the experiment. Arrows indicate the sodium hypochlorite\ntreatment every 25 cycles to maintain electrode performance.\nResponse to Reviewer 3\n(Q=question/comment; A=answer; changes to the manuscript or supplementary information\nhighlighted in red)\nComments: This manuscript presents a novel and potentially impactful electrochemical system that\nintegrates osmotic energy harvesting with acid recovery to drive chlorine production from chloride-\nrich brines. The process design is innovative and technically robust, bridging electrochemical\nengineering with advanced membrane materials. The use of proton-selective COF membranes is a key\nenabler, allowing selective proton transport while excluding other ions. Overall, the study is thorough,\nand the results are well-analysed. The manuscript is suitable for publication in Nature Communications,\npending minor revisions. The following points should be addressed:\nResponse: We appreciate your careful evaluation of our work. According to your thoughtful and\nvaluable suggestions, we have made the relevant modifications as shown below:\nQ1: As a central schematic, Figure 1 should be improved for clarity. Electrochemical reactions and ion\ntransport directions across compartments should be clearly labelled. The concept of continuous\noperation, via electrode switching and reservoir refreshing, should be provided and briefly explained.\nIncluding a comparative schematic contrasting the proposed process with conventional chlor-alkali\nsystems would help highlight advantages and limitations.\nA1: Thank you very much for the valuable suggestions. We have revised the figure to enhance clarity.\nIn the revised Fig. 1, we replaced the single-panel illustration with a two-panel schematic.\nPanel\na now depicts the conventional chlor-alkali process, explicitly labeling CER (2Cl⁻ → Cl + 2e⁻), HER\n2\n(2H O + 2e⁻ → H + 2OH⁻), Na+ transport across a cation-exchange membrane, and the product\n2 2\nstreams (Cl , H , NaOH). Panel b shows our spontaneous chlorine production system with an 18-unit\n2 2\nmodular stack, adding clear electron-flow arrows along the series connection and ion-transport arrows\nacross the TpPa-SO H proton-selective membrane to highlight H+ migration while pollutants/heavy\n3\nmetals remain on the feed side.\nThe concept of continuous operation is now incorporated through the annotation \"Switching\nperiodically\" in panel b, accompanied by bidirectional arrows. This indicates the periodic reversal of\nAg/AgCl electrodes to sustain the AgCl + e⁻ Ag + Cl⁻ reaction, ensuring continuous Cl⁻ transport.\nAdditionally, the flow lines from \"Concentrated acid\" and \"Acidic wastewater\" reservoirs, along with\n⇌\n\"Diluted acid\" output, suggest a cycle of reservoir refreshing to maintain the concentration gradient, a\ncritical aspect for sustained operation.\nWe revised the Design of the system section inour manuscript,\n“As shown in Figure 1a, the conventional chlor-alkali industry uses a cation ion exchange membrane\nto electrolyze concentrated brine for chlorine production. This process relies on a continuous supply\nof electrical energy, leading to significant power consumption. In contrast to the conventional chlor-\nalkali industry, which requires external electrical power, our system operates spontaneously and does\nnot rely on energy input (Fig. 1b).\nFig. 1 Schematic illustration of conventional chlor-alkali system and spontaneous chlorine production system.\na, In the conventional chlor-alkali system, the electrolysis of concentrated brine requires a continuous\nsupply of external electrical power to produce chlorine. b, The spontaneous chlorine production system,\nin contrast, operates without external electrical power. The integrated system consists of multiple units\nconnected in series. Acidic wastewater containing concentrated protons and Cl- ions was used. The\nprotons and Cl- ions selectively migrate from the feed reservoirs containing waste acid to the permeate\nreservoirs containing lower-concentration recovered acid through the proton-selective membranes and\nAg/AgCl electrodes, respectively. On the Ag/AgCl electrodes, electrochemical reactions occur,\nconsuming Cl- in the feed reservoirs and releasing Cl- in the permeate reservoirs, effectively\ntransporting Cl-. The electrochemical reactions and the proton diffusion generate electrical potential.\nAs a result, on the leftmost anode, chlorine is spontaneously produced and on the rightmost cathode,\nspontaneous hydrogen evolution occurs.\nQ2: A diagram or discussion summarizing redox potentials of the involved reactions would be helpful.\nSpecifically, the authors should clarify how other metal ions (e.g., heavy metals) in the brine are\nprevented from undergoing undesirable side reactions such as deposition.\nA2: Thank you very much for the helpful suggestion! We added Supplementary Fig. 10, which maps\nthe standard redox potentials of the reactions relevant to our system, including HER (H+/H ), CER\n2\n(Cl⁻/Cl ), OER (H O/O ), and representative metal-ion couples (e.g., Cu2+/Cu, Ni2+/Ni, Zn2+/Zn,\n2 2 2\nFe3+/Fe2+). This redox “landscape” shows that, if multivalent metal ions reached the cathode, their\nreduction would compete with HER (e.g., Cu2+/Cu at +0.34 V vs SHE).\nTo address this problem, our system employs a proton-selective TpPa-SO H COF membrane that\n3\nefficiently blocks the transport of multivalent heavy-metal ions (Fig. 2). As a result, the metal\ndeposition is prevented.\nWe revised the second paragraph of High proton selectivity of the membrane section in our\nmanuscript,\n“...In real industrial waste acid, the concentration of metal cations is typically much lower than that\nof protons. This means that in the permeate reservoir, the metal cation concentration would be\nnegligible, avoiding side reactions on the HER electrode, such as metal ion deposition (Supplementary\nFig. 10).”\nSupplementary Fig. 10. Standard electrochemical reduction potentials (E) of various metal ion\nredox reactions compared with the hydrogen evolution reaction (HER) and chlorine evolution\nreaction (CER).\nQ3: The long-term stability of the Ag/AgCl electrodes is not sufficiently addressed. Prior studies have\nnoted issues with AgCl layer buildup increasing resistance. While images and SEM/EDX analysis are\nprovided (Fig. 3c, Fig. S9), additional discussion or experimental data on electrode durability over\nextended operation would strengthen the manuscript.\nA3: Thank you very much for the constructive suggestion! We conducted extended tests, which indeed\nshowed AgCl layer build-up and increasing interfacial resistance, which in turn led to a decline in the\nper-cycle transferred charge. But the electrodes can be easily regenerated by immersing them in sodium\nhypochlorite.\nWe performed a 100-cycle durability study in 5 M/0.5 M HCl (Supplementary Fig. 14). After 25\ncycles, the transferred charge per cycle decreased by ~20%, suggesting the progressive AgCl\naccumulation that elevates the charge-transfer/transport resistance and reduces the electrochemically\nactive area(Corrosion Science 2021, 184; Electrochim Acta 2011, 56, 2781-2791). However, this\ndecline can be reversed through periodic maintenance—treating with sodium hypochlorite (indicated\nby arrows in Supplementary Fig. 14b).\nTo sum up, we revised the fourth paragraph of Co-transport of Cl- and proton, and its\nimportance for continuous operation and high transport rate section\n“...Furthermore, in a 100-cycle test, the performance can still be robustly maintained with periodic\nelectrode regeneration (Supplementary Fig. 14).”\nSupplementary Fig. 14. Performance of Ag/AgCl electrodes over 100 cycles in a 5 M/0.5 M HCl\nsolution.\na, I-t curve during the 100 cycles. b, Comparison of transferred charge (C) for each cycle, illustrating\nthe consistent charge transfer throughout the experiment. Arrows indicate the sodium hypochlorite\ntreatment every 25 cycles to maintain electrode performance.\nQ4: The authors predominantly use 5 M HCl in their experiments. However, acid concentration in real\nfeedstocks may vary. A sensitivity analysis exploring system performance across a range of acid\nconcentrations would enhance the relevance of the study to practical applications.\nA4: Thank you very much for this valuable suggestion. To address this, we have conducted a sensitivity\nanalysis by exploring the osmotic energy harvesting performance across different HCl concentration\nratios.\nWe set the low concentrations to be 0.5 M and 0.1 M HCl, and varied the high concentration from\nFor instance, with a fixed low concentration of 0.5 M HCl, the peak power density rises from ~2 W m-\n2 at 1 M/0.5 M to over 55 W m-2 at 5 M/0.5 M. Similarly, with a fixed low concentration of 0.1 M HCl,\nthe power density increases from ~6 W m-2 at 1 M/0.1 M to ~35 W m-2 at 5 M/0.1 M (Supplementary\nFig. 32d–f). This trend is attributed to the increased driving force for proton diffusion and reduced\nreservoir resistance.\nFor a fixed concentration ratio, the power density increases with the increase of concentration due\nto the reduced reservoir resistance. For example, the power density at 5 M/0.5 M was 55 W m-2 while\nthat at 1 M/0.1 M was merely 6 W m-2.\nThese findings demonstrate the system’s adaptability to varying acid concentrations and highlight\nthe importance of optimizing feed ratios for maximum efficiency. We have integrated this analysis into\nthe revised manuscript in Spontaneous chlorine production section,\n“To evaluate the system’s sensitivity, we tested various feed solution concentrations. As the\nconcentration ratio increases, the power density increases gradually. For fixed concentration ratio, the\npower density increases with the increase of concentration due to the reduced reservoir resistance\n(Supplementary Fig. 32-33).”\nSupplementary Fig. 32. Energy conversion performance of TpPa-SO H membrane with varying\na, d, I–V curves for 1 M/0.5 M, 3 M/0.5 M, and 5 M/0.5 M (a), and I–V curves for 1 M/0.1 M, 3 M/0.1\nM, and 5 M/0.1 M (d); b, e, power density versus external resistance; c, f, current density versus\nexternal resistance.\nSupplementary Fig. 33. Energy conversion performance of serially connected devices with\nvarying HCl concentration ratios.\na, Open-circuit voltage for different numbers of serially connected units at different concentration\nratios (5 M/0.5 M, 3 M/0.5 M, and 1 M/0.1 M).b, Open-circuit voltage and current density over time\nforsystems: 12 units for 5 M/0.5 M, 17 units for 3 M/0.5 M, and 17 units for 1 M/0.1 M.c, Average\ncurrent density (left axis) and Cl production rate (right axis) for different feed concentration ratios.\nQ5: While the TpPa-SO3H COF membrane exhibits high proton selectivity and stability, testing\nmembranes with lower selectivity or across a range would help define the operational limits and\nrobustness of the system.\nA5: Thank you very much for the suggestion! We have expanded our study to include three additional\nmembranes—Nafion, Graphene Oxide (GO), and Polyvinylidene difluoride (PVDF)—with lower\nselectivity. The proton/Cl- and proton/metal ion selectivities are both important for the performance of\nour system.\nWe first assessed the role of proton/Cl- selectivity. As shown in Supplementary Fig. 34a, the I-V\ncurves reveal distinct performance characteristics. Nafion (N117, ~183 μm) exhibits proton/Cl-\nselectivity comparable to TpPa-SO H, but its thicker structure results in higher internal resistance,\n3\nreducing ion transport efficiency compared to the thinner TpPa-SO H membrane. GO displays\n3\nmoderate proton selectivity, while PVDF, with the lowest selectivity, shows the weakest performance.\nThe power density (Supplementary Fig. 34b) and current density (Supplementary Fig. 34c) tested with\nexternal resistance further confirm that membranes with higher proton selectivity and lower internal\nresistance—such as TpPa-SO H—achieve higher osmotic energy harvesting performance.\nnumber of units (12 for TpPa-SO H and Nafion, 13 for GO, and 15 for PVDF) to achieve a consistent\n3\nopen-circuit voltage of ~2 V, as depicted in Supplementary Fig. 35a. Despite the similar open-circuit\nvoltages, the current density and Cl production rates vary significantly (Supplementary Fig. 35b, c).\n3 2\nfollowed by Nafion (~46 L m-2 h-1), GO (~33 L m-2 h-1), and PVDF (~27 L m-2 h-1). This trend\nunderscores the critical role of high proton selectivity in maximizing ion transport and chlorine\nproduction efficiency.\nFinally, to investigate the role of proto/heavy metal ion selectivity, we compared TpPa-SO H and\n3\n(Supplementary Fig. 36a), with no Cu deposition on the Pt HER electrode (inset photos) and high\nFaraday efficiency for H production (measured volumes closely matching theoretical values;\n3\npermeation over 1 h, unlike PVDF, which allows rapid Cu2+ crossover (Supplementary Fig. 36c). In\ncontrast, the 15-unit PVDF system exhibited rapid current decay due to cathodic Cu deposition\n(Supplementary Fig. 36d, inset photos), rendering it unstable.\nThese findings suggest the system's operational limits: membranes with low selectivity\ncompromise energy harvesting efficiency, Cl production rates, and long-term stability by permitting\n2\nside reactions (e.g., metal deposition). We updated the manuscript in Spontaneous chlorine\nproduction sectionwith these data,\n“...We also investigated the relationship between the membrane proton selectivity and system\nperformance. The results show that membranes with high proton/Cl- selectivity exhibit superior energy\nconversion performance and increased chlorine production rates (Supplementary Fig. 34–35).\nFurthermore, the high proton/heavy metal ion is critical for sustained operation, as inadequate rejection\nled to rapid side reactions at the cathode (Supplementary Fig. 36).”\nSupplementary Fig. 34. Energy conversion performance of membranes with different proton/Cl-\nselectivities.\na, I–V curves for TpPa-SO H, Nafion, Graphene oxide (GO) and Polyvinylidene difluoride (PVDF).\n3\nb, c,Power density (b) and current density (c) versus external resistance.\nSupplementary Fig. 35. Performance of systems with different proton selectivity membranes.\na, Open-circuit voltage for different numbers of serially connected units with different proton/Cl-\nselectivity membranes (TpPa-SO H, Nafion, GO and PVDF). b, Open-circuit voltage and current\n3\ndensity over time for systems. c, Average current density (left axis) and Cl production rate (right axis)\n2\nfor each membrane system. Error bars represent the standard deviation from at least three independent\ntests.\nSupplementary Fig. 36. Role of proton/heavy metal ion selectivity on sustained system operation.\na, I-t curves for a 12-unit system employing TpPa-SO H membranes with 5 M HCl (blue) or 5 M HCl\n3\ncontaining 1 M CuCl (orange) as feed solutions. Inset: photographs of the Pt electrode for hydrogen\n2\nevolution reaction before and after operation, showing no Cu deposition. b, Measured hydrogen\nproduction volume versus time for the system with Cu2+-containing feed solution, compared to the\ntheoretical volume (dashed line) calculated from the current curve, indicating high HER Faraday\nefficiency and minimal side reactions. c, Cu2+ rejection experiments using TpPa-SO H and PVDF\n3\nmembranes. The feed side (left, deep green) contained 5 M HCl and 1 M CuCl , while the permeate\n2\nside (right) contained 5 M HCl. Photographs at 0, 30 and 60 min show effective Cu2+ rejection by\nTpPa-SO H, in contrast to rapid permeation through PVDF. d, I-t curves for a 15-unit system\n3\nemploying PVDF membranes with 5 M HCl (pink) or 5 M HCl + 1 M CuCl (gray) as feed solutions,\n2\nexhibiting rapid current density decay in the presence of Cu2+ due to cathodic deposition. Inset:\nphotographs of the Pt electrode before and after operation, showing evident Cu deposition.\nQ6: Some data (e.g., Fig. 2g) lack error bars, and the selectivity values should be carefully reviewed\nfor accuracy. The authors are encouraged to ensure all reported values are statistically supported.\nA6:Thank you for pointing this out! We have re-processed the raw datasets, repeated the analyses,\nand updated the figures to include statistical descriptors. This update does not affect the conclusions.\nWe also re-checked all selectivity values against and verified their accuracy. The selectivities are now\nclearly reported in High proton selectivity of the membrane section,\n“...The H+/Al3+, H+/ Mn2+, H+/Fe3+, H+/Co2+, H+/Ni2+, H+/Cu2+, and H+/Zn2+ selectivities were\ncalculated to be 673, 248, 3767, 224, 228, 201, and 225, respectively (Fig. 2g), far superior to\ncommercial ion-exchange membranes.”\nFig. 2 High selectivity of the COF membrane. a, Optical image of the synthesized TpPa-SO H COF\n3\nmembrane. Inset: cross-sectional scanning electron microscopy image, showing a membrane thickness\nof approximately 10 µm. b, Transmission electron microscopy image of the membrane, showing the\ninterlayer spacing. c, Measured and simulated X-ray diffraction pattern of the membrane, confirming\nthe COF structure. Inset: molecular structure of the TpPa-SO H COF.\n3\nd, Schematic illustration of the\nset-up to measure the selectivity of the membrane using the diffusion dialysis method. e, Concentration\nof permeated protons and various metal ions (Al3+, Mn2+, Fe3+, Co2+, Ni2+, Cu2+, Zn2+) in response to\ndiffusion time. f, Magnified view of time-dependent concentration of metal ions. g, Calculated proton\nselectivity against various metal ions. h, i, UV-Vis absorbance spectra of levofloxacin hydrochloride\n(LFH) (h) and tetracycline hydrochloride (TCH) (i) in the feed and permeate reservoirs, suggesting\nthe almost perfect rejection of organic pollutants by the membrane. Error bars represent the standard\ndeviation from at least three independent tests.\nQ7: The study examines two model organic molecules (TCH and LFH), but real waste acid streams\nmay contain a broader range of contaminants, including positively charged organics. Potential\nmembrane fouling through adsorption (especially positively charged molecules) should be discussed,\nand membrane integrity after long-term exposure should be evaluated.\nA7: Thank you very much for the suggestion!\nTo evaluate potential fouling by positively charged organics, we selected two representative\ncationic dyes—Rhodamine B (RB) and Crystal Violet (CV)—as model pollutants. As shown in\nSupplementary Fig. 26, we conducted adsorption tests by exposing the TpPa-SO H membrane to 50\n3\nppm solutions of these dyes for 12 hours. The results indicate negligible changes in the membrane's\nmorphology and color, as evidenced by optical images and photographs of the dye solutions before\nand after exposure (Supplementary Fig. 26a-b). The concentration of dyes in the solution remained\nlargely unchanged, with an adsorption capacity (Q) of less than 10 mg/g for both RB and CV\n(Supplementary Fig. 26c-d). This low adsorption is likely due to the membrane's small pore size, which\nrestricts the diffusion of these larger organic molecules into the membrane.\nTo assess the impact on long-term performance, we evaluated the osmotic energy harvesting\nperformance of the TpPa-SO H membrane before and after 12 hours of exposure to these contaminants,\n3\nas presented in Supplementary Fig. 27. The open-circuit voltage (OCV) of the pristine membrane and\nthose soaked in CV and RB showed negligible differences. Similarly, the maximum power density\nremained consistent and the power density versus external resistance curves overlapped closely. These\nresults demonstrate that the membrane retains its proton selectivity and energy conversion efficiency\ndespite exposure to positively charged pollutants.\nIn fact, in separation applications, membrane fouling is typically caused by microorganisms. They\nstrongly adhere to grow on the membrane. In industrial wastewater, the corrosive environment makes\nmicroorganisms hard to live, mitigating the membrane fouling problem.\nWe incorporated these data into the Power generation performance section of the revised\nmanuscript,\n“Two positively charged organic pollutants, Rhodamine B and Crystal Violet, were used to further\nevaluate the anti-fouling capability of the membrane (Supplementary Fig. 26). The transport of\npollutants was also almost completely blocked, making the membrane anti-fouling (Supplementary\nFig. 27).”\nAdsorption of heavy metal ions and positively charged pollutants\nThe adsorption of heavy metal ions by TpPa-SO H membranes was carried out using a 50 mL mixed\n3\nsolution containing 0.001 M of each metal ion: Al³⁺, Mn²⁺, Fe³⁺, Co²⁺, Ni²⁺, Cu²⁺, and Zn²⁺. A known\nmass of the TpPa-SO H membrane was immersed into the initial metal ion solution. The system was\n3\nmaintained under magnetic stirring at 600 rpm for 12 hours to facilitate the adsorption process.\nFollowing this, the concentration of metal ions in the solution was measured to determine the extent\nof metal ion adsorption onto the membrane. The amount of metal ions adsorbed was calculated using\nthe following formula:\n𝑐0 −𝑐𝑡\nwhere Q represents the adsorption capacity (mg/g), V is the volume of the solution (mL), c is the\n𝑄 = ×𝑉 0\n𝑚\ninitial concentration of metal ions (mg/L), c is the concentration after adsorption (mg/L), and m is the\nt\nmass of the TpPa-SO H membrane (g).\nusing 50 ppm solutions of Rhodamine B and Crystal Violet, respectively, as model pollutants. The\nadsorption process was conducted under the same experimental conditions as described for the heavy\nmetals, with a stirring speed of 600 rpm for 12 hours. After adsorption, the concentrations of\nRhodamine B and Crystal Violet in the solution were quantified by UV-Vis spectrophotometer,\nallowing for the determination of adsorption capacity of the TpPa-SO H membrane.\n3\nmembrane.\na, Photographs showing the solutions before and after Rhodamine B (RB) and Crystal Violet (CV)\nadsorption test. b, Optical images of the TpPa-SO H membrane before and after exposure to the dyes,\n3\ndemonstrating no significant changes. c, Adsorption capacity (Q, mg/g) of the membrane for both\npollutants. Error bars represent the standard deviation from at least three independent tests.\nSupplementary Fig. 27. Energy conversion performance of the TpPa-SO H membrane after 12\n3\nhours of exposure to positively charged pollutants.\na, Comparison of the open-circuit voltage of the pristine membrane and membranes soaked in Crystal\nViolet and Rhodamine B, showing negligible changes. b, Comparison of the maximum power density\nof the pristine and polluted membranes, showing stable energy output. c, Power density tested with\nexternal load for the pristine and soaked membranes. Error bars represent the standard deviation from at least three independent tests.\nQ8:The system design confines H₂ and Cl generation to the terminal electrodes, which may constrain\n2\nscalability. The authors should comment on potential design adaptations to improve scalability for\nindustrial applications.\nA8: Thank you very much for the helpful suggestion! We agree that the confinement of H and Cl\n2 2\ngeneration to the terminal electrodes (Pt cathode for HER and graphite anode for CER, as shown in\nFig. 1b) could appear limiting at first glance. However, this design choice is primarily driven by safety\nconsiderations: by isolating the HER and CER at opposite ends of the serially connected units, we can\nprevent the mixing of H and Cl gases, reducing risks that are prevalent in conventional chlor-alkali\n2 2\nprocesses where gases are produced in shared compartments.\nOur system can be adapted for higher scalability. As illustrated in Supplementary Fig. 42, the\ndesign allows for straightforward adaptations through serial and parallel connection to adjust\nproduction rates.\nIncreasing the number of serially connected units (e.g., from 18 to more, as depicted in\nSupplementary Fig. 42a) raises the overall electrochemical reaction voltage. This enhancement in\nvoltage accelerates the chlorine evolution reaction (CER) kinetics, thereby boosting Cl production\n2\nrates. For instance, our experiments with varying unit numbers (Supplementary Fig. 29) showed that\nstacking more units (up to 18 in the integrated device, Supplementary Fig. 30) linearly increases the\nopen-circuit voltage (Fig. 5e), enabling higher throughput while keeping the terminal electrodes fixed\nfor gas isolation.Caution is needed to avoid excessively high voltages, which may trigger competing\noxygen evolution reactions and lower Faraday efficiency.\nTo amplify the reaction current and thus scale up the overall Cl yield, multiple series stacks can\n2\nbe connected in parallel (Supplementary Fig. 42b). This configuration leads to proportionally higher\ncurrent densities and gas production volumes.\nBeyond these electrical configurations, further optimizations can be implemented to adapt the\nsystem for industrial-scale deployment. For example, refining the electrode switching protocol and\noptimizing Ag/AgCl electrode size (e.g., enlarge the electrode size, automating the periodic reversal\nof Ag/AgCl electrodes to sustain Cl⁻ transport) can minimize downtime and enhance continuous\noperation (Science 2024, 385, 1438-1444). Reducing reservoir volume and electrode spacing within\nunits could lower internal resistance and improve efficiency (International Journal of Hydrogen Energy\n2003, 28, 35-41; Journal of Cleaner Production 2021, 319).\nWe have incorporated this discussion into the revised manuscript, including a reference to\nSupplementary Fig. 42 for visual clarity, to better highlight the system's scalability potential. We\nrevised the Discussion section,\n“The modular design of our system allows scalable applications by connecting multiple units in\nseries or parallel (Supplementary Fig. 42). Further improvements could be achieved by reducing\nelectrode spacing and automating the electrode switching.”\nSupplementary Fig. 42. Modular scaling of chlorine production.\na, Increasing the number of series-connected units raises the electrochemical reaction voltage, thereby\naccelerating chlorine production. b, Increasing the number of parallel-connected units raises the\nelectrochemical reaction current, thereby enhancing chlorine production.\nQ9: The rationale for using 18 cells (achieving ~2.95 V) to drive spontaneous Cl₂ evolution should be\nmore clearly linked to the electrochemical requirements and electrode polarization behaviour.\nA9: Thank you very much for the valuable suggestion!\nThe stack voltage V supplied by the series-connected units must exceed the sum of (i) the\nstack\nthermodynamic gap for CER + HER (~1.36 V between Cl /Cl⁻ and H+/H in acid), (ii) the kinetic\n2 2\noverpotentials at the graphite (CER) and Pt (HER) electrodes at target current densities, and (iii) the\nohmic voltage drop across the electrolytes and membrane (iR ). As illustrated in Supplementary Fig.\nin\n29a, when the series number is ≤8, the effective reaction voltage V = V - iR remains below\nreaction stack in\nthe onset of the CER/HER, resulting in negligible Faradaic current (Supplementary Fig. 29b, c; inset\nfor 8 units). Increasing the unit number raises V and consequently V , moving the operating\nstack reaction\npoint into the kinetically accessible regime and unlocking substantial current. For instance, 12, 15, and\n18 units deliver progressively higher operating currents (Supplementary Fig. 29c-d). With 18 units, the\nopen-circuit voltage reaches ~2.95 V, and after subtracting iR , the reaction provides sufficient\nin\nheadroom above the CER/HER onset to achieve the highest Cl production rate among the tested\n2\nconfigurations (Supplementary Fig. 29d, e).\nWe revised the third paragraph of Spontaneous chlorine production section inour manuscript,\n“Equation 1 tells us that each osmotic power generation unit can produce a maximum of 0.184 V\nwhen 5 M/0.5 M HCl was used for chlorine production. This means that 8 units are required to initiate\nthe spontaneous production of chlorine. Considering the electrochemical overpotential, internal\nresistance voltage drops, we tested serially connected units from 8 to 18 (Fig. 5d and Supplementary\nFig. 29). Then the terminal electrodes of the serial osmotic generators were replaced with a graphite\nelectrode and a Pt electrode. The Ag/AgCl electrodes of the feed and permeate reservoirs were\nswitched every 600 s for continuous operation. The hydrogen generated on the Pt electrode was\ncollected using the water drainage method, while chlorine production rate was calculated from the\nionic current data. Obviously, The V of the system increased linearly with the number of units in\noc\nseries, reaching 2.95 V when the unit number was 18 (Fig. 5e), resulting in an enhanced chlorine\nproduction rate. When the number of units were lower, the reaction was either slower or could not be\ninitiated.”\nSupplementary Fig. 29. Performance of the system with varying numbers of serially connected\nunits.\na, Schematic diagram illustrating the relationship between voltage and current for different numbers\nof units, showing the increasing electrochemical reaction voltage and current density as the number of\nunits increases. b, Diagram illustrating the difference in current behavior when the number of units is\n≤8 versus >8. c, Open-circuit voltage and current density versus time for different numbers of units\nduring operation. d, Average current density and corresponding chlorine production rate for systems\nwith different numbers of units. The inset shows the current density for the 8-unit system. e, Hydrogen\nproduction rate for systems with varying numbers of serially connected units. Error bars represent the\nstandard deviation from at least three independent tests.\nQ10: Generating Cl₂ in an open system raises safety concerns. While HER and CER are spatially\nseparated, the authors should discuss gas handling and containment strategies to support practical\nimplementation.\nA10: Thank you very much for reminding us of this safety consideration! This problem is also\nencountered by the conventional chlor-alkali process. The Cl₂ can be sealed to ensure safety. In our\nsystem, the Cl₂ generation compartment can also be sealed, and we can maintain a negative pressure\nwithin the sealed compartment. Sealed compartments would allow direct piping of Cl to storage or\n2\nprocessing units, while negative pressure (e.g., via vacuum-assisted collection) would prevent outward\nleakage by drawing gases inward (J. Modern Chlor‐Alkali Technology 2001; Handbook of Chlor-\nAlkali Technology 2005). These adaptations are compatible with our system's modular configuration.\nHere, we revised the Discussion section in our manuscript to,\n“...Regarding chlorine handling, we can seal the anode compartment and maintain slight negative\npressure to prevent chlorine leakage, ensuring both safety and directed transport to storage units.”"}, {"round": 2, "reviewer_comments": {"reviewer_1": "All my concerns have been well addressed in this revision.", "reviewer_2": "I am satified with the revision and the response of reviewers' comments. so I suggest a positive response for publication.", "reviewer_3": "The authors have addressed most questions well. The revised manuscript is suitable for publication."}, "author_response": "Response to Reviewer 1\n(Q=question/comment; A=answer)\nComments: All my concerns have been well addressed in this revision.\nResponse: Thank you for participating in the review process and for the valuable effort contributed.\nResponse to Reviewer 2\n(Q=question/comment; A=answer)\nComments: I am satisfied with the revision and the response of reviewers' comments. so I suggest a positive response for publication.\nResponse:Thanks a lot! We highly appreciate your effort in helping us improve this manuscript!\nResponse to Reviewer 3\n(Q=question/comment)\nComments: The authors have addressed most questions well. The revised manuscript is suitable for\npublication.\nResponse: We highly appreciate your effort in helping us strengthen this work!"}]

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