Use of a pH indicator in the preparation and monitoring of electrolyte

Abstract

The application belongs to the technical field of electrochemistry, and particularly relates to application of a pH indicator in electrolyte preparation and monitoring. The application of the pH indicator in electrolyte preparation and monitoring comprises the following steps: adding the pH indicator into electrolyte to obtain electrolyte with a pH visual response function, and using the electrolyte with the pH visual response function to perform in-situ visual monitoring on local pH changes in an electrochemical system. The application can realize real-time and in-situ visual monitoring on local pH change processes in the electrochemical system by introducing the pH indicator into the electrolyte, can significantly reduce a corrosion reaction rate, and has an inhibition effect. Meanwhile, by using the proton-coupled electron transfer characteristics of the pH indicator, local acid-base environment changes at an electrode interface which are difficult to directly observe are converted into identifiable optical signals, which is conducive to revealing electrode reaction mechanisms, interface evolution processes and electrolyte state change rules.

Description

Technical Field

[0001] This invention belongs to the field of electrochemical technology, and specifically relates to the application of a pH indicator in electrolyte preparation and monitoring. Background Technology

[0002] pH is an important parameter characterizing the activity of hydrogen ions in a solution, typically described using a range of 0-14. pH=7 represents a neutral environment, pH<7 represents an acidic environment, and pH>7 represents an alkaline environment. pH is widely present in various scenarios in nature and human production and daily life. For example, human gastric juice is usually strongly acidic, with a pH of approximately 1; lemon juice, vinegar, and tomato juice have pH values ​​of approximately 2, 3, and 4, respectively; the pH values ​​of coffee, milk, and purified water increase sequentially, with purified water approaching neutral; human blood maintains a slightly alkaline environment, with a pH of approximately 7.35-7.45. As alkalinity increases, the pH values ​​of ammonia water, soapy water, and bleach can reach approximately 11, 12, and 13, respectively. In extreme acidic or alkaline environments, the pH of the sulfuric acid electrolyte in lead-acid batteries is close to 0, while the pH of concentrated sodium hydroxide solution is close to 14. Therefore, pH values ​​are applied throughout the natural environment, biological systems, and industrial production processes, and are closely related to the properties of substances, chemical reactions, and system stability.

[0003] In modern industrial production, the pH of a solution system is a key parameter affecting production efficiency, product quality, and equipment operational stability. Due to factors such as changes in raw material composition, side reactions, environmental conditions, and the non-uniformity of mass and heat transfer processes, the pH value of a system, whether local or overall, typically undergoes dynamic changes. Abnormal pH changes not only exacerbate equipment corrosion and material aging, shortening the lifespan of production equipment, but may also further lead to problems such as decreased raw material utilization, increased byproducts, increased energy consumption, and unstable product quality. In actual chemical production, local pH shifts can lead to decreased reaction selectivity; in metal processing and electroplating, pH changes affect deposition quality and surface uniformity; and in aquatic and soil treatment processes, uncontrolled pH can reduce treatment efficiency and increase operating costs. Furthermore, in marine engineering, ship hull steel structures are exposed to seawater for extended periods, and changes in seawater pH significantly affect the corrosion behavior of the steel. Therefore, achieving rapid and accurate monitoring of pH changes in a system has significant industrial value.

[0004] pH indicators are widely used in various detection and analysis scenarios to monitor and characterize changes in acidic and alkaline environments. pH indicators are compounds with acid-base responsive properties. Their molecular structures typically contain functional groups capable of protonation or deprotonation reactions; common groups include phenolic hydroxyl, carboxyl, amino, and imino groups. When the ambient pH changes, the electronic structure and conjugated state of the indicator molecule change accordingly, resulting in different colors, absorption spectra, or fluorescence signals. Based on this characteristic, pH indicators can transform acid-base changes that are difficult to observe directly into visually identifiable information in real time. Compared with traditional electrode detection and instrumental analysis methods, pH indicators have advantages such as lower cost, faster response, simpler operation, and ease of in-situ visual monitoring, thus showing broad application prospects in the fields of acid-base environment detection and process monitoring. Currently, pH indicators are widely used in water quality monitoring, environmental analysis, food quality and testing, biomedical diagnostics, chemical analysis, and educational experiments.

[0005] With the rapid development of electrochemical energy storage technology, the issue of local pH changes in electrochemical systems such as batteries, electrolyzers, and electrocatalytic reactors has attracted widespread attention. During electrochemical reactions, protons (H+) often accumulate near the electrode interface. + ) or hydroxide ions (OH-) - The generation, migration, and consumption of electrolytes lead to significant local pH gradients at the electrode / electrolyte interface. Abnormal fluctuations in local pH can trigger a series of adverse effects, including accelerated corrosion and damage to electrode materials, altered electrochemical reaction pathways, disruption of electrode interface structure and stability, induction of side reactions and byproducts, distortion of electrochemical sensing and measurement results, and impact on ion transport and charge transfer kinetics. These effects not only degrade the performance of the electrochemical system but may also affect the safety and reliability of device operation.

[0006] Localized pH fluctuations, often overlooked, can have a more insidious and destructive impact on electrochemical systems than overall pH changes. For example, in aqueous batteries, localized alkalization can induce metal ions to react with OH-. - Reactions generate byproducts that deposit, leading to electrode passivation and capacity decay. Localized acidification exacerbates electrode corrosion and active material dissolution. Furthermore, pH changes affect the deposition behavior of metal ions on the negative electrode, inducing dendrite growth and side reactions. In electrocatalytic systems, interfacial pH changes alter reaction kinetics and product selectivity, reducing overall system efficiency. More importantly, the hydrogen evolution reaction (HER) and oxygen evolution reaction (OER), commonly found in aqueous electrolytes, are always accompanied by H+. + or OH - The generation and consumption of electrolytes mean that the evolution of local pH can serve as an important characteristic signal reflecting interfacial reactions and electrolyte states.

[0007] Current methods for monitoring pH changes in electrochemical systems primarily rely on electrode probes, sampling analysis, or spectroscopic characterization techniques. These methods suffer from drawbacks such as operational complexity, insufficient real-time performance, and difficulty in achieving in-situ visualization. Therefore, developing a novel method for real-time, in-situ, and visualized monitoring of local pH changes in electrochemical systems is of great significance for studying electrode interface reaction processes, elucidating energy storage mechanisms, and improving the reliability of device operation. Summary of the Invention

[0008] The present invention aims to solve one or more technical problems existing in the prior art, and at least provide a beneficial solution. Specifically, the present invention provides an application of a pH indicator in electrolyte preparation and monitoring, which can realize real-time, in-situ, and visual monitoring of local pH changes in electrochemical systems.

[0009] The inventive concept of this invention: The application of the pH indicator of this invention in electrolyte preparation and monitoring includes the following steps: adding the pH indicator to the electrolyte to obtain an electrolyte with pH visualization response function, and using the electrolyte with pH visualization response function to perform in-situ visualization monitoring of local pH changes in the electrochemical system.

[0010] In addition to its acid-base response characteristics, the pH indicator of this invention also exhibits proton-coupled electron transfer (PCET) behavior. In an electrochemical environment, the indicator molecules can participate in interfacial reaction processes through the synergistic effect of proton and electron transfer. Therefore, its function is not limited to pH indication but may also influence the electrolyte microenvironment and interfacial reactions. Based on this PCET characteristic, the pH indicator shows multiple application potentials in electrolyte preparation and monitoring. On the one hand, the proton transfer process endows the electrolyte with a sensitive response to changes in the acid-base environment. It can not only be used for rapid colorimetric calibration and state determination during electrolyte preparation but also buffer local acid-base fluctuations through reversible protonation / deprotonation processes, thereby effectively suppressing side reactions such as hydrogen evolution and corrosion. On the other hand, for pH indicators with redox activity, the electron transfer process significantly enhances the charge transfer kinetics at the electrode / electrolyte interface and provides additional Faraday capacity for constructing redox-enhanced energy storage devices.

[0011] Furthermore, the PCET process establishes a correlation between chemical state and electrochemical behavior, and provides visualization through color response. When hydrogen evolution, oxygen evolution, electrode corrosion, or other side reactions occur in an electrochemical system, local pH, interfacial reaction rate, and electrolyte state often change simultaneously. By coupling chemical state (pH) with electrochemical behavior (potential, current, and capacity) using pH indicators, it is possible to track electrolyte performance degradation and side reaction accumulation in situ during long-term cycling, and to provide early warning through optical signals. Simultaneously, its electrochromic properties provide a simple and effective technical path for self-sensing and state visualization feedback in intelligent energy storage systems.

[0012] Therefore, a first aspect of the present invention provides the application of a pH indicator in electrolyte preparation and monitoring.

[0013] Specifically, the application includes the following steps: adding the pH indicator to the electrolyte to obtain an electrolyte with pH visualization response function, and using the electrolyte with pH visualization response function to perform in-situ visualization monitoring of local pH changes in the electrochemical system.

[0014] Specifically, after the pH indicator is introduced into the electrolyte, it can respond to the local pH changes generated during the electrochemical reaction and undergo a structural transformation through a reversible protonation and / or deprotonation process, thereby producing a recognizable color change, realizing in-situ visual monitoring of the local pH change process of the electrochemical system.

[0015] During electrochemical reactions, the electrode / electrolyte interface is usually accompanied by H+. + or OH - The generation, consumption, and transport processes of these molecules lead to localized acidification or alkalization near the interface. When the local pH changes, the pH indicator molecules undergo changes in electronic structure, conjugated system, and protonation state through proton-coupled electron transfer, resulting in changes in absorption spectrum and color. By monitoring these color changes, the local pH evolution process at the electrode interface can be characterized in real time, further reflecting the electrochemical reaction state and changes in the interfacial environment.

[0016] Preferably, the pH indicator comprises a compound that undergoes reversible protonation and / or deprotonation reactions under different acidic or alkaline conditions and produces a color response.

[0017] Preferably, the pH indicator contains a functional group that undergoes a proton transfer reaction.

[0018] Preferably, the functional group that undergoes the proton transfer reaction includes at least one of hydroxyl, carboxyl, amino, imino, and sulfonic acid groups.

[0019] Preferably, the pH indicator comprises at least one of the following: a naturally sourced indicator, a synthetic organic molecule, a metal complex, a conjugated polymer, and a pH-responsive nanomaterial.

[0020] Preferably, the natural source indicator is derived from at least one of plant, animal, or microbial metabolites.

[0021] Preferably, the plant is anthocyanin-rich plant tissue, such as at least one of purple cabbage, butterfly pea flower, hibiscus flower, blueberry, black goji berry, red onion, and grape skin.

[0022] Preferably, the animal includes cochineal insects used to extract cochineal red.

[0023] Preferably, the natural source indicator further includes litmus derived from lichen extracts.

[0024] Preferably, the natural source indicator includes at least one of litmus, anthocyanins, curcumin, betaine, and shikonin.

[0025] Preferably, the artificially synthesized organic molecule includes at least one of methyl orange, methyl red, phenolphthalein, phenol red, Congo red, bromocresol green, bromothymol blue, thymol blue, and alizarin red.

[0026] Specifically, alizarin red, a synthetic organic molecule, can also be used as an indicator of natural origin.

[0027] Preferably, the metal complex includes at least one of Chrome Black T and Xylenol Orange.

[0028] Preferably, the high molecular conjugated polymer includes polyphenylacetylene-pyridine conjugated polyelectrolytes.

[0029] Preferably, the pH-responsive nanomaterial comprises polyfluorene-ethylene conjugated polymer fluorescent nanoparticles.

[0030] According to some embodiments of the present invention, the pH indicator includes at least one of litmus, anthocyanin, methyl orange, methyl red, phenolphthalein, phenol red, Congo red, bromocresol green, bromothymol blue, thymol blue, alizarin red, chrome black T, xylenol orange, polyphenylene pyridine conjugated polyelectrolytes, and polyfluorene ethylene conjugated polymer fluorescent nanoparticles.

[0031] Specifically, the pH indicator of this invention is a compound that can respond to changes in environmental acidity or alkalinity and produce a recognizable color change. Its molecular structure typically contains functional groups such as hydroxyl, carboxyl, amino, imino, and sulfonic acid groups that can participate in proton transfer reactions. Different indicators have different color change ranges, color response characteristics, and applicable monitoring areas, and can be selected according to the working environment, interfacial reaction characteristics, and monitoring requirements of the target electrochemical system.

[0032] Preferably, the pH indicator is dispersed in the electrolyte in a dissolved state, or loaded on the surface of a carrier in contact with the electrolyte in a solid state.

[0033] Preferably, the carrier includes at least one of an electrode, a diaphragm, a current collector, a liquid reservoir, and an observation window.

[0034] Preferably, the concentration of the pH indicator in the electrolyte is 0.1 μmol·L⁻¹. -1 -5mol·L -1 More preferably, the concentration of the pH indicator in the electrolyte is 10 μmol·L⁻¹. -1 -500 μmol·L -1 .

[0035] Preferably, the electrolyte is an aqueous electrolyte, which includes at least one soluble salt of Zn, Mg, Al, Mn, Fe, Li, Na, and K; that is, at least one soluble salt of Zn, Mg, Al, Mn, Fe, Li, Na, and K.

[0036] Preferably, the electrochemical system includes at least one of a battery, an electrolyzer, an electrocatalytic reactor, an electrochemical sensor, and a supercapacitor.

[0037] Specifically, the electrochemical system is an electrochemical system in which local pH changes occur.

[0038] Preferably, the color response is used to characterize at least one of the following: local pH change at the electrode interface, hydrogen evolution reaction, oxygen evolution reaction, electrode corrosion, degree of side reaction, electrolyte state change, electrode activity change, battery health status, and failure process.

[0039] Specifically, by monitoring the color change of the pH indicator, the local pH change process near the electrode interface can be characterized in real time, and further reflected the hydrogen evolution reaction, oxygen evolution reaction, electrode corrosion, the degree of side reaction, electrolyte state change, electrode activity change, and the operating status of the electrochemical system.

[0040] Different pH indicators have different color change ranges, color change characteristics, water solubility, and applicable monitoring areas. Depending on the operating environment of different electrochemical systems and the local pH change range, pH indicators with different color change ranges can be selected for monitoring.

[0041] The typical pH indicators applicable to this invention, their color change range, color change characteristics, applicable monitoring areas, and water solubility are shown in Table 1.

[0042] Table 1: Typical pH indicators, their color change range, color change characteristics, applicable monitoring areas, and water solubility.

[0043] Table 1 lists only some typical pH indicators as examples and does not constitute a limitation on the scope of protection of this invention. Those skilled in the art can select one or more synthetic pH indicators, naturally derived pH indicators, or combinations thereof for the visual monitoring of local pH changes, based on the working environment, color resolution requirements, color change range, and electrochemical stability requirements of the target system.

[0044] According to some embodiments of the present invention, a single indicator whose color change range matches the working pH range of the target system can be selected to achieve visual monitoring of a specific range.

[0045] According to some embodiments of the present invention, two or more indicators with different color change ranges can be used in combination to achieve a wider pH range or a more advanced pH change process for visualization.

[0046] Compared with the prior art, the beneficial effects of the technical solution provided by the present invention are as follows: (1) By introducing a pH indicator into the electrolyte, the present invention can respond to the local pH changes generated during the electrochemical reaction and undergo structural transformation through reversible protonation and / or deprotonation processes, thereby generating identifiable color changes, and realizing real-time, in-situ visualization monitoring of the local pH change process of the electrochemical system.

[0047] (2) This invention utilizes the proton-coupled electron transfer characteristics of pH indicators to transform the local acid-base environment changes that are difficult to observe directly at the electrode interface into identifiable optical signals, which is beneficial for revealing the electrode reaction mechanism, interface evolution process and electrolyte state change law.

[0048] (3) The present invention can intuitively reflect the occurrence process of hydrogen evolution reaction, oxygen evolution reaction and other side reactions; at the same time, it can be used for auxiliary analysis and diagnosis of the operating status, failure process and health status of electrochemical system.

[0049] (4) The present invention can obtain information on changes in the acid-base environment of the electrode interface without complicated instruments and equipment, and has the advantages of fast response, simple operation, low cost and wide applicability. Attached Figure Description

[0050] Figure 1 This is a schematic diagram of the pH scale and typical medium acidity / alkalinity distribution of the present invention; Figure 2 This is a schematic diagram showing the electron-proton transfer behavior and structural state distribution of the pH indicator molecule of the present invention; Figure 3This is a schematic diagram of the pH-response color change mechanism of anthocyanins, a natural pH indicator of the present invention; Figure 4 This is a schematic diagram of the pH-responsive color-changing mechanism of alizarin red, a natural pH indicator of the present invention; Figure 5 This is a schematic diagram illustrating the pH-responsive color change mechanism of beetroot red, a natural pH indicator of this invention. Figure 6 This is a schematic diagram illustrating the pH-response color-changing mechanism of curcumin, a natural pH indicator of the present invention. Figure 7 This is a schematic diagram of the pH response color change mechanism of shikonin, a natural pH indicator of the present invention; Figure 8 The image shows a 0.2 g / L litmus aqueous solution with pH = 1-10 and its UV-Vis spectrum in Example 1 of this invention. Figure 9 This is a diagram showing the local color change during the monitoring of water electrolysis using litmus indicator in Example 2 of the present invention. Figure 10 This is a color change diagram of alizarin red solution at different pH values ​​in Example 3 of the present invention; Figure 11 This is a physical image of the two-electrode H-type electrolytic cell used for monitoring the health status of a Zn-C hybrid capacitor using alizarin red, as shown in Embodiment 4 of the present invention. Figure 12 This is a visualization of the alizarin red monitoring results in Example 4 of the present invention; Figure 13 This is a Tafel diagram of a Mg metal corrosion battery with and without litmus in MgCl2 electrolyte, as shown in Example 5 of the present invention. Detailed Implementation

[0051] To enable those skilled in the art to more clearly understand the technical solutions described in this invention, the following embodiments are provided for illustration. It should be noted that the following embodiments do not constitute a limitation on the scope of protection claimed by this invention.

[0052] Unless otherwise specified, the raw materials, reagents or devices used in the following examples are available from conventional commercial sources or can be obtained by existing known methods.

[0053] A schematic diagram of pH scale and acidity / alkalinity distribution in typical media is shown below. Figure 1 As shown. Figure 1 The pH levels in the range of 0-14 and their corresponding typical substances are given.

[0054] A schematic diagram of the electron-proton transfer behavior and structural state distribution of pH indicator molecules is shown below. Figure 2 As shown.

[0055] in, Figure 2Figure A in the diagram is a rectangular loop mechanism diagram of proton-coupled electron transfer, which includes the stepwise and cooperative transfer behavior of electrons and protons on the indicator molecule (Ind), and describes the interconversion between its different protonation states and redox states.

[0056] Figure 2 Figure B in the diagram shows the reaction network in the homogeneous electrolyte system, which describes the electron transfer (redox reaction) between the indicator molecule and the electrode and the proton transfer (color change) behavior of the indicator molecule in the solution. In addition, it also shows the hydration reaction of cation M in the solution and the free anion X.

[0057] Figure 2 Figure C in the diagram is the potential-pH diagram (Pourbaix phase diagram) of Ind, which shows the stable existence regions of different protonated and redox states of the indicator molecule in the potential-pH coordinate, as well as the transition boundaries between the structural states. E 0 ETPT This represents the standard potential corresponding to the first electron transfer step in the electron transfer-proton transfer (ETPT) pathway; E 0 PTET This represents the standard potential corresponding to the second electron transfer step in the proton-electron transfer (PTET) pathway.

[0058] Figure 3 A schematic diagram illustrating the pH-responsive color change mechanism of anthocyanins, a natural pH indicator; Figure 4 A schematic diagram of the pH-responsive color change mechanism of alizarin red, a natural pH indicator; Figure 5 A schematic diagram of the pH-responsive color change mechanism of beetroot red, a natural pH indicator; Figure 6 A schematic diagram illustrating the pH-responsive color-changing mechanism of curcumin, a natural pH indicator; Figure 7 This is a schematic diagram of the pH-response color change mechanism of shikonin, a natural pH indicator.

[0059] Example 1 Example 1 illustrates the color change of litmus indicator in aqueous solutions at different pH values. The specific process is as follows: (1) Prepare 0.05M H2SO4 solution (pH=1): Measure 0.543mL of 98% concentrated sulfuric acid into a beaker, add 160mL of deionized water and mix well, transfer to a 200mL volumetric flask, and make up to 200mL to obtain 0.05M H2SO4 solution, and put it into a reagent bottle for later use. (2) Preparation of 0.1M KOH solution (pH=13): Weigh 1.12g KOH solid into a beaker, add 160mL of deionized water and stir to dissolve, transfer to a 200mL volumetric flask, and make up to 200mL to obtain 0.1M KOH solution, which is then stored in a reagent bottle for later use. (3) Preparation of pH 1-6 series H2SO4 solutions: Take the 0.05 M H2SO4 solution prepared in step (1) as the pH=1 acidic standard solution; measure 10.0 mL, 1.0 mL, 100 μL, 10.0 μL and 1.0 μL of pH=1 acidic standard solution into reagent bottles, and then add 0.0 mL, 9.0 mL, 9.9 mL, 9.99 mL and 9.999 mL of deionized water respectively. After mixing evenly, H2SO4 solutions with a volume of 10 mL and pH values ​​of 1, 2, 3, 4 and 5 respectively are obtained; then measure 1.0 mL of pH=5 H2SO4 solution into reagent bottle, add 9.0 mL of deionized water, mix evenly and obtain 10 mL of pH=6 H2SO4 solution. (4) Preparation of pH 10-8 series KOH solutions: Take the 0.1M KOH solution prepared in step (2) as the pH=13 alkaline standard solution; measure 10.0μL of pH=13 alkaline standard solution into a reagent bottle, add 9.99mL of deionized water, mix well to obtain 10mL of KOH solution with pH=10; then measure 1.0mL and 100μL of pH=10 KOH solution into reagent bottles, and add 9.0mL and 9.9mL of deionized water respectively, mix well to obtain KOH solutions with a volume of 10mL and pH values ​​of 9 and 8 respectively; (5) Prepare litmus aqueous solution with pH 1-10: Using the solutions with pH=1, 2, 3, 4, 5, 6, 8, 9, 10 prepared in steps (3) and (4), and 10 mL of deionized water with pH=7 as solvents, add 2.0 mg of litmus powder to each solution and mix thoroughly until the litmus is completely dissolved to obtain litmus aqueous solution with pH=1-10.

[0060] The litmus aqueous solution with pH=1-10 prepared in step (5) was tested using a UV-Vis spectrophotometer. Deionized water was used as a reference. The absorption spectrum was collected in the wavelength range of 250-700nm, and the position of the characteristic absorption peak and the change of absorbance of the litmus solution under different pH conditions were recorded.

[0061] 0.2 g / L litmus aqueous solutions with pH = 1-10 and their UV-Vis spectra are shown below. Figure 8 As shown. Among them, Figure 8 The image at the top shows a physical sample of a 0.2 g / L litmus solution with a pH range of 1 to 10. Figure 8 The lower figure shows the UV-Vis spectra of litmus solutions at different pH values. Figure 8 It can be seen that the maximum absorption wavelength changes significantly between pH 3 and 4, corresponding to the color change from red to blue.

[0062] Depend on Figure 8It can be seen that litmus indicators exhibit significant color responses and absorption spectrum changes under different pH conditions. 7-Hydroxy-2-phenoxyoxanone, the most prominent chromophore in litmus, combines with a proton to form a red cation in acidic conditions, while in alkaline conditions it loses a proton to form a blue-violet anion with a conjugated structure. The reaction equation for its protonation-deprotonation reaction is as follows: .

[0063] Depend on Figure 8 As shown in the photograph of the litmus solution, the solution exhibits a color gradient with pH. Under strongly acidic conditions (pH=1-2), it is red or orange-red; as the pH increases to 3, the solution color gradually transitions from red to purple; within the pH range of 4-7, it is mainly purple; and as the pH continues to increase to 8-10, the solution gradually turns from blue-purple to blue. The color change is most pronounced in the pH range of 3-4, indicating that this range is the main color-changing region where litmus transitions from an acidic to a basic structure.

[0064] Depend on Figure 8 The UV-Vis spectrum shows that litmus exhibits strong absorption in the 250-350 nm region under different pH conditions, corresponding to its aromatic conjugated structure π→π Electron transitions occur; as pH increases, a distinct absorption band appears in the 550-650 nm visible light region, forming an absorption peak around 584 nm. This absorption peak is strongest at pH 4-5, indicating the coexistence of acidic and basic structures in the solution, resulting in strong visible light absorption and thus the typical purple color. Notably, the spectrum of litmus solution at pH 3 differs significantly from other samples, exhibiting strong absorption in the 511 nm region and relatively weaker absorption around 584 nm, consistent with the magenta characteristic observed in the physical image. This phenomenon suggests that the area around pH 3 is a region where the protonation / deprotonation equilibrium of litmus molecules changes rapidly, corresponding to the vicinity of its color change point.

[0065] The test results above show that litmus aqueous solution can produce obvious color changes in different pH ranges, especially when the pH is between 3 and 4, its maximum absorption wavelength shows the most significant change, which can theoretically be used to monitor electrochemical reactions in acidic electrolytes.

[0066] Example 2 Example 2 illustrates the application of litmus indicator in monitoring pH changes caused by hydrogen and oxygen evolution processes during water electrolysis. The specific process is as follows: (1) Dissolve 4.0 mg of litmus powder in 20 mL of dilute H2SO4 aqueous solution with pH 4, mix to obtain 0.2 g / L litmus aqueous solution, then add it to 20 mL of deionized water and stir until completely dissolved to obtain litmus indicator aqueous solution with pH value of about 4.3 and concentration of 0.1 g / L. (2) The litmus indicator aqueous solution with a concentration of 0.1 g / L obtained in step (1) is transferred to the electrolytic cell as the electrolyte for the water electrolysis reaction. A 1.0 cm × 1.0 cm carbon paper is used as the working electrode, a 1.0 cm × 1.0 cm platinum sheet (Pt) is used as the counter electrode, and an Ag / AgCl electrode is used as the reference electrode to assemble a three-electrode electrolysis system. (3) Turn on the power supply and set the working electrode potential to -2.0V (vs.Ag / AgCl). Electrolyze water in constant potential mode and continue the reaction for 10 minutes. During the reaction, the current continues to increase. Record the changes in the color of the electrolyte and the generation of bubbles on the electrode surface.

[0067] The chronoamperometry method was used, and the current was collected every 0.05s with a current range of 250mA. The electrolyte was an aqueous solution of litmus indicator with a pH of approximately 4.3 and a concentration of 0.1g / L, in which litmus was used as a pH indicator at a concentration of 0.1g / L.

[0068] Litmus indicator for monitoring local color changes during water electrolysis, such as... Figure 9 As shown. Among them, Figure 9 Figure ① shows the local color change during the chronoamperometry test at 0s. Figure 9 Figure ② shows the local color change during the chronoamperometry test over 60 seconds; Figure 9 Figure ③ shows the local color change during the chronoamperometry test at 120s. Figure 9 Figure ④ shows the local color change during the chronoamperometry test over 180 seconds.

[0069] Depend on Figure 9 It can be seen that as the current increases, after 60 seconds using the chronoamperometry method, bubbles gradually form on the working electrode at a potential of -2.0V. Figure 9 (See Figure ② in the text). Furthermore, under litmus visualization, the electrolyte color in the localized area of ​​the electrode changed from the initial magenta to bluish-purple. This phenomenon indicates that the pH value locally increased in the region near the working electrode, and that hydroxide ions (OH-) were generated at the working electrode. - The hydrogen evolution reaction in water electrolysis, which produces hydrogen gas, leads to a localized increase in pH near the electrode interface. This induces a structural transition and color change in the litmus indicator, causing it to change from an acidic to a basic state. The litmus indicator can track the increase in pH. The reaction equation for the hydrogen evolution reaction in water electrolysis is as follows: 2H₂O + 2e -→ 2OH - + H2↑.

[0070] As the reaction time progresses, the thickness of the discolored region gradually increases. When the thickness of the diffusion layer (discolored region) reaches approximately 2 mm, the discolored region gradually expands into the bulk electrolyte. Its spatial distribution gradually evolves from a local diffusion layer near the electrode surface into a non-steady-state mass transfer region controlled by both diffusion and natural convection. Figure 9 Figure ④ shows that the dynamic equilibrium of steady-state diffusion of ions and protons in the color-changing layer is gradually broken as the hydrogen evolution reaction continues.

[0071] The above results indicate that litmus in aqueous electrolytes has a sensitive visual response to pH changes during electrochemical processes and can be used as an indicator of the electrolytic reaction in this system.

[0072] Example 3 Example 3 illustrates the color change of alizarin red indicator in aqueous solutions at different pH values. The specific process is as follows: (1) Preparation of 1mM alizarin red standard solution: Weigh 6.85mg alizarin red powder, add 20mL of deionized water, dissolve it in the reagent bottle, mix well to obtain 1mM alizarin red aqueous solution; (2) Preparation of solutions with different pH values: The H₂SO₄ solution with pH=1 and the KOH solution with pH=13 prepared in Example 1 were used as standard solutions; 2.0 mL of H₂SO₄ solution with pH=1 and 2.0 mL of potassium hydroxide solution with pH=13 were measured into reagent bottles to serve as pH=1 and pH=13 solutions, respectively; 20 μL of H₂SO₄ solution with pH=1 was measured into a reagent bottle, 1.98 mL of deionized water was added, and the mixture was thoroughly mixed to obtain 2.0 mL of pH=3 solution; 20 μL of pH=3 solution was measured into a reagent bottle, 1.98 mL of deionized water was added, and the mixture was thoroughly mixed to obtain 2.0 mL of pH=5 solution; 20 μL of pH=13 KOH solution was measured into a reagent bottle, 1.98 mL of deionized water was added, and the mixture was thoroughly mixed to obtain 2.0 mL of pH=7 solution; pH=11 solution; then measure 20μL of pH=11 solution into a reagent bottle, add 1.98mL of deionized water, mix well to obtain 2.0mL of pH=9 solution; (3) Preparation of Alizarin Red solutions with different pH values: Add 200 μL of 1 mM Alizarin Red standard solution to the above 2.0 mL solutions with pH = 1, 3, 5, 7, 9, 11 and 13 respectively, mix well to obtain a series of solutions with different pH values ​​with a total volume of 2.200 mL and an Alizarin Red concentration of about 0.1 mM.

[0073] Color response test: The obtained samples were placed in transparent glass bottles of the same size, and photographed and recorded under the same lighting conditions and a white background. The color change patterns of alizarin red solution under different pH conditions were compared to evaluate its acid-base response characteristics and color gradient effect.

[0074] The color change graph of alizarin red solution at different pH values ​​is shown below. Figure 10 As shown. Figure 10 As shown, alizarin red exhibits a distinct and continuous color change from yellow, orange, pink, to purple as the pH value increases within the pH range of 1-13, demonstrating a good responsiveness to solution acidity and alkalinity. The color change is most pronounced in the pH range of 5-11, exhibiting high visual distinctiveness. The color differences between pH 1 and 3, and between pH 11 and 13, are relatively small, indicating that alizarin red exists in relatively stable acidic and basic structures under extremely acidic and strongly alkaline conditions, respectively.

[0075] Under strongly acidic conditions (pH=1-3), alizarin red solutions are pale yellow to yellow, with the colors being quite similar at pH=1 and pH=3, indicating that the molecules mainly exist in a protonated structure. As the pH increases to 5, the solution gradually changes from yellow to pale red; near neutral, it appears pale orange or flesh-colored, indicating that the molecules begin to deprotonate and the conjugated structure changes. Further increasing the pH to 9, the solution color deepens significantly to pink; at pH=11 and 13, it further transforms into purplish-red and purple, indicating that the phenolic hydroxyl groups in the molecule have essentially completed deprotonation, forming a stable basic structure.

[0076] The reaction equation for the protonation-deprotonation reaction of alizarin red is: .

[0077] The test results above show that alizarin red indicator can produce obvious color changes in different pH ranges and has good water solubility, and theoretically can be used for visual monitoring of electrochemical reactions.

[0078] Example 4 Example 4 illustrates the application of pH indicators in in-situ visual monitoring of the health status of energy storage devices (hybrid capacitors Zn-C). The specific process is as follows: (1) Preparation of zinc sulfate solution: Weigh 57.51g of ZnSO4 powder into a beaker, add 160mL of deionized water, mix well, then transfer to a 200mL volumetric flask and measure to 200mL to obtain 1M ZnSO4 solution, which is then stored in a reagent bottle for later use; repeat three times and combine to obtain 600mL of ZnSO4 solution. (2) Weigh 10.27 mg of alizarin red powder and add it to a beaker containing 600 mL of 1 M ZnSO4 solution. Mix well to obtain an alizarin red electrolyte with a concentration of 50 μmol / L in 1 M ZnSO4, which is called the alizarin red-ZnSO4 mixed solution. Since the 1 M ZnSO4 solution is acidic, the alizarin red-ZnSO4 mixed solution shows a magenta color. Take out 20 mL of the solution for later use (as a control sample before the reaction). (3) Transfer the remaining alizarin red-ZnSO4 solution into the H-type electrolytic cell, and install 1.0cm×1.0cm carbon paper as the working electrode and 1.0cm×1.0cm zinc sheet as the counter electrode at both ends of the H-type electrolytic cell in parallel and opposite directions. (4) Turn on the power supply and perform OCP test first. The initial voltage of the system is measured to be approximately -0.475V. Use this voltage as the starting point for cyclic voltammetry (CV). (5) Set the initial voltage to -0.475V, the voltage scan window to 0-1.8V, and the scan rate to 1mV / s. Perform cyclic voltammetry scans in the above alizarin red-ZnSO4 electrolyte. The number of scans is 11, and the reaction time is 11h. (6) After the CV scan is completed, take out 20 mL of the electrolyte after the reaction (as the post-reaction sample) and store it for later use.

[0079] The pH values ​​of the control sample reserved in step (2) before the reaction and the sample after the reaction in step (6) were measured using a pH meter, and the changes before and after were recorded. The pH value of the alizarin red-ZnSO4 solution before the reaction was 5.17, and the pH value of the electrolyte after the reaction was 5.32. This indicates that after 11 hours of redox cycle reaction, the pH value of the electrolyte showed signs of rising.

[0080] A physical image of a two-electrode H-type electrolytic cell used to monitor the health status of a Zn-C hybrid capacitor using alizarin red is shown below. Figure 11 As shown.

[0081] The visualization results of alizarin red monitoring in Example 4 are as follows: Figure 12 As shown.

[0082] like Figure 12 As shown, under the visualization effect of alizarin red, a horizontal dark band persistently appears in the two-electrode energy storage system as the reaction proceeds. This dark band appears only on the carbon electrode side, pointing perpendicularly from the carbon electrode surface to the Zn electrode. Its color is dark purple, deeper than the surrounding magenta, indicating a pH gradient between this banded environment and the overall environment. This is essentially closely related to the direction of the electric field in the energy storage system and to the Zn on the carbon electrode surface. 2+ H + There is an intrinsic link between ion desorption and migration.

[0083] The above results indicate that alizarin red has a good visual response to pH changes during electrochemical processes in ZnSO4 electrolyte and can be used as an indicator of the electrolytic reaction in this system.

[0084] Example 5 Application of pH indicators in improving the corrosion resistance of metal electrodes and detecting the operating status of electrodes.

[0085] Metallic magnesium (Mg) exhibits a low electrode potential and high chemical activity, while phosphorus (Pt), as an inert electrode, possesses excellent electrical conductivity and chemical stability. The Mg-Pt corrosion cell system exhibits a large potential difference, generating significant corrosion current and polarization behavior, making it an ideal model system for studying metal corrosion mechanisms.

[0086] By using Tafel polarization testing, key parameters such as self-corrosion potential (Ecorr) and self-corrosion current density (Icorr) can be obtained, thereby quantitatively evaluating the impact of additives on corrosion kinetics. The specific process is as follows: (1) Preparation of magnesium chloride solution: Weigh 19.04g of anhydrous MgCl2 powder into a beaker, add 160mL of deionized water, mix well, and then transfer to a 200mL volumetric flask. Quantify to 200mL to obtain 1M MgCl2 solution, and store in a reagent bottle for later use. (2) Take 50 mL of 1 M MgCl2 solution without litmus as the electrolyte of the control group, use a 1.0 cm × 1.0 cm Mg sheet as the working electrode, a 1.0 cm × 1.0 cm Pt sheet as the counter electrode, and Ag / AgCl as the reference electrode to assemble a three-electrode system; (3) Turn on the power and perform the OCP test first. The voltage of the system in the initial equilibrium state is approximately -1.544V (vs.Ag / AgCl). Use this voltage (±0.3V) to set the start and end points of the Tafel test. (4) Set the initial potential to -1.9V, the final potential to -1.3V, and the scan rate to 0.5mV / s. During the reaction, observe and record the phenomena in the electrolytic cell. After the scan is completed, use the built-in analysis function in the Shanghai Chenhua Electrochemical Workstation CHI440C program to calculate the Tafel slope of the cathode and anode, as well as the corrosion potential and corrosion current density. (5) Weigh 5.0 mg of litmus powder and add it to a beaker containing 50 mL of 1.0 M MgCl2 solution. Mix well to obtain a 1 M MgCl2 electrolyte with a litmus concentration of 0.1 g / L, which is called litmus-MgCl2 solution. (6) Take 50 mL of the above 0.1 g / L litmus-MgCl2 solution as the electrolyte and put it into the electrolytic cell; use a 1.0 cm × 1.0 cm Mg sheet as the working electrode, a 1.0 cm × 1.0 cm Pt sheet as the counter electrode, and Ag / AgCl as the reference electrode to assemble a three-electrode system. (7) First, perform an OCP test and measure the voltage of the system in the initial equilibrium state as approximately -1.008V (vs. Ag / AgCl). Use this voltage (±0.3V) as a reference to set the start and end points of the Tafel test. (8) Set the initial potential to -1.368V, the final potential to -0.768V, and the scan rate to 0.5mV / s. During the reaction, observe and record the phenomena in the electrolytic cell. After the scan is completed, calculate the Tafel slope of the cathode and anode, as well as the corrosion potential and corrosion current density in the software.

[0087] Tafel diagrams of Mg metal corrosion cells with and without litmus MgCl2 electrolytes are shown below. Figure 13 As shown.

[0088] Depend on Figure 13 It can be seen that without litmus, the Tafel slopes of the cathode and anode of the Mg-Pt corrosion cell are 0.278 and 0.124, respectively. After adding litmus, the Tafel slope of the Pt metal cathode decreases to 0.167, indicating that the introduction of litmus alters the cathode charge transfer process or the hydrogen evolution reaction control step. Lime may adsorb on the electrode surface, partially covering the active sites of the Pt cathode, thus affecting the hydrogen evolution kinetics. Simultaneously, the Tafel slope of the Mg metal anode increases from 0.124 to 0.252, indicating that litmus increases the activation energy or charge transfer resistance of Mg anode dissolution, making it more difficult for Mg to undergo electron-loss corrosion.

[0089] Further data analysis revealed that the intersection point obtained by extending the linear portions of the cathode and anode represents the logarithm of the self-corrosion current density and the self-corrosion potential, respectively. Without litmus, the logarithm of the cathodic corrosion current was -2.412, and the logarithm of the anodic corrosion current was -2.378. After adding litmus, the logarithm of the cathodic corrosion current decreased to -3.773, and the logarithm of the anodic corrosion current decreased to -3.774.

[0090] Compared to the blank control group without litmus, the introduction of litmus reduced the self-corrosion current from 10 -2.4 Reduced to 10 -3.8 The corrosion rate was reduced by approximately 1.4 orders of magnitude, indicating a significant decrease in corrosion rate, while simultaneously inhibiting Mg dissolution and the hydrogen evolution reaction at the Pt cathode. These test results demonstrate that litmus can effectively regulate the charge transfer process at the Mg and Pt electrode interface, significantly reducing the corrosion reaction rate and exhibiting excellent mixed-type corrosion inhibition.

[0091] In summary, this invention, by introducing a pH indicator into the electrolyte, can respond to local pH changes generated during electrochemical reactions. Through reversible protonation and / or deprotonation processes, a structural transformation occurs, resulting in a recognizable color change. This enables real-time, in-situ visual monitoring of local pH changes in the electrochemical system and effectively regulates charge transfer processes at the electrode interface, significantly reducing corrosion rates and exhibiting corrosion inhibition. Simultaneously, utilizing the proton-coupled electron transfer characteristics of the pH indicator, local acid-base environment changes at the electrode interface, which are difficult to observe directly, are transformed into recognizable optical signals, facilitating the revelation of electrode reaction mechanisms, interface evolution processes, and electrolyte state changes. Furthermore, the pH indicator of this invention can directly reflect the occurrence of hydrogen evolution reaction, oxygen evolution reaction, and other side reactions, and can be used for auxiliary analysis and diagnosis of the operating status, failure processes, and health status of electrochemical systems.

[0092] The above embodiments are only used to illustrate the technical solutions of the present invention and are not intended to limit the scope of protection of the present invention. Although the present invention has been described in detail with reference to preferred embodiments, those skilled in the art should understand that modifications or equivalent substitutions can be made to the technical solutions of the present invention without departing from the essence and scope of the technical solutions of the present invention.

Claims

1. The application of pH indicators in electrolyte preparation and monitoring, characterized in that, The application includes the following steps: adding the pH indicator to the electrolyte to obtain an electrolyte with pH visualization response function, and using the electrolyte with pH visualization response function to perform in-situ visualization monitoring of local pH changes in the electrochemical system.

2. The application according to claim 1, characterized in that, The pH indicator comprises compounds that undergo reversible protonation and / or deprotonation reactions under different acidic and alkaline conditions and produce a color response.

3. The application according to claim 1, characterized in that, The pH indicator contains functional groups that undergo proton transfer reactions.

4. The application according to claim 3, characterized in that, The functional group that undergoes the proton transfer reaction includes at least one of hydroxyl, carboxyl, amino, imino, and sulfonic acid groups.

5. The application according to claim 1, characterized in that, The pH indicator includes at least one of the following: natural source indicator, artificially synthesized organic molecule, metal complex, high molecular conjugated polymer, and pH-responsive nanomaterial.

6. The application according to claim 5, characterized in that, The natural source indicator includes at least one of litmus, anthocyanins, curcumin, betaine, and shikonin; And / or, the artificially synthesized organic molecule includes at least one of methyl orange, methyl red, phenolphthalein, phenol red, Congo red, bromocresol green, bromothymol blue, thymol blue, and alizarin red; And / or, the metal complex includes at least one of Chrome Black T and Xylenol Orange; And / or, the polymeric conjugated polymer includes polyphenylacetylene-pyridine conjugated polyelectrolytes; And / or, the pH-responsive nanomaterials include polyfluorene-vinyl chloride conjugated polymer fluorescent nanoparticles.

7. The application according to claim 1, characterized in that, The pH indicator is dispersed in the electrolyte in a dissolved state, or loaded on the surface of a carrier in contact with the electrolyte in a solid state.

8. The application according to claim 7, characterized in that, The carrier includes at least one of the following: electrode, diaphragm, current collector, liquid reservoir, and observation window.

9. The application according to claim 1, characterized in that, The pH indicator has a concentration of 0.1 μmol·L in the electrolyte. -1 -5mol·L -1 ; And / or, the electrolyte is an aqueous electrolyte, which includes at least one of the soluble salts of Zn, Mg, Al, Mn, Fe, Li, Na, and K; And / or, the electrochemical system includes at least one of a battery, an electrolyzer, an electrocatalytic reactor, an electrochemical sensor, and a supercapacitor.

10. The application according to claim 2, characterized in that, The color response is used to characterize at least one of the following: local pH change at the electrode interface, hydrogen evolution reaction, oxygen evolution reaction, electrode corrosion, degree of side reaction, electrolyte state change, electrode activity change, battery health status, and failure process.

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