A gravimetric method for the quantitative assessment of the products of electrolysis of water is presented. In this approach, the electrolysis cell was directly powered by 9 V batteries. Prior to electrolysis, a known amount of potassium hydrogen phthalate (KHP) was added to the cathode compartment, and an excess amount of KHCO3 was added to the anode compartment electrolyte. During electrolysis, cathode and anode compartments produced OH−(aq) and H+(aq) ions, respectively. Electrolytically produced OH−(aq) neutralized the KHP, and the completion of this neutralization was detected by a visual indicator color change. Electrolytically produced H+(aq) reacted with HCO3−(aq) liberating CO2(g) from the anode compartment. Concurrent liberation of H2(g) and O2(g) at the cathode and anode, respectively, resulted in a decrease in the mass of the cell. Mass of the electrolysis cell was monitored. Liberation of CO2(g) resulted in a pronounced effect of a decrease in mass. Experimentally determined decrease in mass (53.7 g/Faraday) agreed with that predicted from Faraday’s laws of electrolysis (53.0 g/Faraday). The efficacy of the cell was tested to quantify the acid content in household vinegar samples. Accurate results were obtained for vinegar analysis with a precision better than 5% in most cases. The cell offers the advantages of coulometric method and additionally simplifies the circuitry by eliminating the use of a constant current power source or a coulometer.
The quantity of a substance produced at the electrode and the quantity of electric charge passed are linked with Faraday’s laws of electrolysis. Many reports present coulometric methods of analysis of a variety of reagents [
Alternatively, the quantification of a reagent produced in an electrolysis cell without direct monitoring of charge is also well known [
Simple measurement of the volumes of hydrogen and oxygen gases produced at the cathode and anode in an electrolysis cell offers advantages of coulometric method without actually using a coulometer or any electronic equipment. In fact, early studies indicated that hydrogen-oxygen coulometer was easy to assemble and capable of measuring the quantity of electric charge with an accuracy of ±0.1% or better [
As stated in (
Electrolysis of water in cathode and anode compartments (
When a known quantity of KHP (a primary standard substance) is added to the cathode compartment, the endpoint of the neutralization reaction between electrolytically produced OH−(aq) and KHP can be determined visually with use of phenolphthalein as an indicator. Electrolysis is promptly stopped at the endpoint. Prior to electrolysis, an excess amount of KHCO3 is added to the anode compartment.
Concurrent liberation of chemically produced CO2(g) in addition to H2(g) and O2(g) amplified the effect of decreasing in the mass. Mass of the cell is continuously monitored before, during, and after electrolysis. From the mass change experienced by the cell, the amount of KHP can be determined to assess the stoichiometric relations between the electrolytic products of water. In this study, the feasibility of monitoring the mass change experienced by an electrolysis cell due to the liberation of gases and quantifying the electrolytic products of water was assessed.
Reagent-grade potassium nitrate, KHP, potassium hydrogen carbonate, and agar were obtained from Fisher. Millipore deionized water was used for preparation of reagents. Alfa Aesar platinum wire (0.762 mm diameter) was used as purchased. 0.1% phenolphthalein was used as a visual indicator to detect the completion of neutralization. Ohaus Pioneer balance (Model: PA114, 0.0001 g readability, and ±0.0002 g linearity) was used for mass measurements. Vinegar samples were locally purchased from a grocery store.
Electrolysis cell was made of spectrophotometer plastic cuvettes. A 3 mm hole was drilled on a side near the base of each cuvette, and the cuvettes were glued together using an epoxy. Bottom portions of two additional cuvettes were cut and glued on top (Figure
Schematic diagram of the water electrolysis cell. KNO3(aq) was used as an electrolyte in both compartments. KHCO3(aq) was added to the anode compartment. Effective concentrations of KNO3 and KHCO3 were 1 M each. Desired number of moles of KHP(aq) and a drop of 0.1% phenolphthalein were added to the cathode compartment. The volumes of electrolyte in each compartment were approximately 3.2 mL. Platinum wires were immersed in the electrolytes and directly connected to two 9 V batteries connected in parallel. The cell was placed on a digital balance housed in a glass cabinet.
Mass correction factor was determined to quantify an error associated with the mass loss due to the evaporation of electrolytes. In this experiment, mass of an idle cell (no electrolysis) was monitored at an interval of 5 minutes. Mass of the cell was monitored for 60 minutes (Figure
Monitoring the mass of an idle cell (no electrolysis).
Desired moles of KHP were added to the cathode compartment. Mass of the cell was monitored at a time interval of one minute for the first five minutes before electrolysis. After this initial rest period, electrolysis was started. Electrolysis was paused at a two-minute time interval, stir bars were maneuvered in up and down position, and the mass was recorded after a two-minute rest period. Electrolysis was promptly stopped when the cathode compartment electrolyte turned light pink indicating an endpoint of the titration. Mass measurement was continued for additional eight minutes. The results of this experiment are presented in Figure
Monitoring the mass of the electrolysis cell before, during, and after electrolysis.
A plot of the mass change experienced by the cell versus number of moles of KHP neutralized in the cathode compartment.
In view of confirming the stoichiometric relations between the moles of OH−(aq) and the mass loss due to the liberation of electrolytically produced H2(g) and O2(g) and chemically produced CO2(g), we monitored the charge passing through the electrolysis cell on Faraday MP potentiostat. Electrolysis was paused at 10 coulombs interval, stir bars were maneuvered in up and down position, and the mass was recorded after one-minute rest period (Figure
A plot of the mass of the electrolysis cell versus the charge passing through the cell.
A known volume of commercial vinegar sample was added to the cathode compartment and titrated against electrolytically produced OH−(aq). The quantity of acid in a sample was determined from the mass change experienced by the cell.
The mass of an idle cell linearly decreased with time (Figure
Faraday’s laws of electrolysis and (
Figure
The electrolytic titration experiment presented in Figure
A linear response of the plot (Figure
The efficacy of the electrolytic titration method presented in this paper was tested for the analysis of acid content in household vinegar samples. A known volume (500 μL) of vinegar was added to the cathode compartment and titrated against electrolytically produced OH−(aq) as presented in the experimental section. The molarity and percent acid content in the sample were determined from the mass change experienced by the cell (Table
Results of titration of an acid in commercial vinegar samples by electrolytically produced OH−(aq).
Percent acid content stated on manufacturer’s label | Mass change in electrolysis cell (g) | Moles of acid in (mmol) | Average molarity | Percent acid content determined from this study (95% CI, |
Percent RSD |
---|---|---|---|---|---|
4 | 0.0165 | 0.311 | 0.623 | 3.74% (±1.32) | 14.3 |
5 | 0.0222 | 0.419 | 0.838 | 5.03% (±0.53) | 4.3 |
7 | 0.0305 | 0.575 | 1.15 | 6.91% (±0.52) | 3.0 |
In this work, the feasibility of using an electrolysis cell for quantification of the electrolytic products of water from gravimetric measurement was tested. The cell presented in this paper enables in situ production of reagents and their direct quantification and does not require standardization of reagents. The electrolysis cell directly powered by 9 V batteries eliminates the requirement of a constant current source or a coulometer, yet offers advantages of coulometric method of titration. The cell utilizes minimal volume of reagents (3.2 mL or less in each compartment). The electrolysis cell is simple, transparent, and easy to fabricate. The cell eliminates the requirement of an external salt bridge or a fritted glass membrane. Linear response of the decrease in the mass of the cell to the moles of KHP added and to the charge passing through the cell validates the applicability of the method. An agreement between the estimated drop in mass (determined from Faraday’s laws of electrolysis) and the experimentally determined drop in mass underlines the feasibility of using the electrolysis cell for quantification of the electrolytic products. With acceptable values of relative standard deviation (better than 5% in most cases), the mass drop experienced by the cell quantitatively responds to the acid content in household vinegar samples.
The authors declare that there are no conflicts of interest regarding the publication of this paper.
The authors thank James O. Schreck and reviewers of the manuscript for their helpful comments and suggestions.