Feasibility of Performing Concurrent Coulometric Titrations Using a Multicompartment Electrolysis Cell.

Feasibility of performing multiple coulometric titrations in a single course of electrolysis is presented. In these titrations, three pairs of cathode and anode compartments were connected with a network of electrodes and salt bridges. Passage of current through the cell caused concurrent electrolysis in cathode and anode compartments. Electrogenerated reagents produced in these compartments were used as titrants for quantifying the analyte samples. Endpoints of the titrations were determined from the visual color change of an indicator. The charge passing through the cell was monitored and Faraday's laws of electrolysis were applied to assess the quantitative relation between the charge and analyte concentration. Experimentally determined coulombs required to titrate aqueous potassium hydrogen phthalate, MnO4 -, OH-, and S2O3 2- were 0.100, 0.466, 0.103, and 0.0934 C, respectively. These results matched with estimated values of 0.0965, 0.482, 0.0965, and 0.0965 C, respectively. Agreement between the coulombs determined from experimental results and reaction stoichiometry suggests a feasible application of concurrent coulometric titrations. Efficacy of the method was tested for determining the active ingredients in household vinegar and vitamin C dietary supplement tablets. Quantities of acetic acid and ascorbic acid in these products were 5.1% and 980 mg, respectively, agreeing with the quantities determined from volumetric titrations (5.1% and 990 mg) and manufacturer's label (5.0% and 1000 mg).

Introduction

Coulometric Titrations

In coulometric titrations, a reagent is produced in an electrolytic reaction at a cathode or an anode. In a secondary reaction, an electrolytically produced reagent chemically reacts with the analyte sample and the endpoint of the titration is determined by a suitable detection method. Charge passing through the cell is monitored and Faraday’s laws of electrolysis are applied to determine the amount of reagent produced at an electrode. Coulometric titrations offer a number of advantages over conventional titrations;[1] in particular, they eliminate preparation and standardization of reagents. A small amount of reagent can be produced without volumetric quantification. Reagents that are unstable to use because of volatility or reactivity can be employed as coulometric titrants and dilution effects do not occur during the titration. Over several decades, numerous reports presented coulometric titrations applied to a wide field of applications.[1−30] Reviews of coulometric titrations illustrate a broad scope of its analytical applications.[7,13,18] A variety of substances have been analyzed by coulometric titrations, including porphyrins,[24] mercury complexes,[25] thiols,[26] cations and anions,[2,27,29] complexes of nickel,[21] biodiesels,[12] and reference standards.[5,6] Coulometric Karl–Fischer analysis was applied to characterize microemulsions[19] and porphyrin assemblies.[9] Laboratory experiments of an academic interest present coulometric analysis of common household products.[10,11] Several innovative methods demonstrate coulometric devices for inexpensive, rapid, and sensitive analysis. These methods include double-pulse compensation,[14] titration in a liquid drop,[15] application of a paper-based device,[16] and electrode modification with ion-selective membranes.[17,23] A study on electrocatalyzed water oxidation combines coulometry with cyclic voltammetry and quartz crystal microbalance.[22] Salient features of coulometric titrations continue to attract researchers from various areas of chemistry. However, there are relatively few reports on the coulometric methods leading to electrolytic production of reagents in multiple compartments and enabling quantification of multiple analyte samples in a single course of electrolysis.[16,18]

Feasibility of Concurrent Coulometric Titrations

As current passes through the cell, an electrolytically produced reagent (or titrant) chemically reacts with the analyte sample added to one of the compartments. Three cascading electrolysis cells producing three electrogenerated reagents enable three independent titrations. Continuous monitoring of charge permits the quantification of reagent produced in each compartment. In view of this hypothesis, in this article, we test the feasibility of performing concurrent coulometric titrations using a multicompartment cell. An electrolysis of water (eq ) produces the reagent OH–(aq) at the cathode. Three electrolysis cells connected in series facilitate electrolytic production of OH–(aq) ions, concurrently in three compartments. Since the same current passes through the electrolysis cells, the charge passing through the cells facilitates the quantification of OH–(aq) ions produced in each cathode compartment. In the secondary reaction, the electrogenerated reagent chemically reacts with the analyte sample, the primary standard KHP (KC8H4O4H) (eq ), and the endpoint of the titration is determined using a suitable acid–base indicator.

Exploring Other Electrolytically Produced Reagents for Titrations

Fe2+(aq) is produced in the cathode compartment by electrolytic reduction of Fe3+(aq) (eq ).[1] Three electrolysis cells connected in series facilitate electrolytic production of Fe2+(aq) ions, concurrently in three compartments. In the secondary reaction, Fe2+(aq) chemically reacts with MnO4–(aq) added as an analyte (eq ) and the endpoint of the titration is determined using MnO4–(aq) as a self-indicator.

Electrolysis of water in an anode compartment produces H+(aq) (eq ). Similarly, oxidation of I–(aq) produces I2(aq) (in form of I3–(aq) in presence of excess iodide) (eq ). These reagents chemically react with OH–(aq) and S2O32–(aq) (eqs and 8), respectively. We test the applicability of the cell for the quantification of OH–(aq) and S2O32–(aq), as the analytes.

Highlights of Concurrent Coulometric Titrations in the Cathode Compartments

Analyte 1: KHP(aq) (KC8H4O4H or C8H4O4H– in the ionic form)Visual indicator: 0.1% phenolphthalein

Analyte 2: MnO4–(aq)Visual indicator: MnO4–(aq) as self-indicator

Highlights of Concurrent Coulometric Titrations in the Anode Compartments

Analyte 3: OH–(aq)Visual indicator: 0.1% methyl orange

Analyte 4: S2O32–(aq)Visual indicator: 1% aqueous starch

Results and Discussion

Cathodic charge passing through the cell linearly changed with the concentrations of KHP (Figure ) and MnO4–(aq) (Figure ). Similarly, the anodic charge linearly changed with the concentrations of OH–(aq) (Figure ) and S2O32–(aq) (Figure ). These linear relations indicate that the amount of charge responsible for producing reagents in concurrent titrations quantitatively relates to analyte concentration. Blank trials of titrations for each analyte indicated nearly 0 C of charge. Blank trials ensured minimal contributions of interfering ions competing with the titration reaction. Faraday’s laws of electrolysis applied to eqs and 2 account for 0.0965 C/μmol of KHP(aq). Slope of the plot (0.100 C/μmol) agreed with this estimate (Figure ).

Figure 1

Plot of the charge passing through the cathodes versus the amount of KHP placed in the cathode compartments.

Figure 2

Plot of the charge passing through the cathodes versus the amount of MnO4–(aq) placed in the cathode compartments.

Figure 3

Plot of the charge passing through the anodes versus the amount of OH–(aq) placed in the anode compartments.

Figure 4

Plot of the charge passing through the anodes versus the amount of S2O32–(aq) placed in the anode compartments.

Mole relations (eqs and 4) between MnO4–(aq), Fe2+(aq), and Fe3+(aq) indicate that 5 F or 482 426.5 C charge is equivalent to 1 mol of MnO4–(aq) (or 0.482 C) charge equivalent to 1 μmol of MnO4–(aq). This value agreed with the slope of 0.466 C/μmol (Figure ).

Equations and 6 indicate that the passage of 1 F (or 96 485.3 C) charge corresponds to 1 mol of OH–(aq). This value (0.0965 C/μmol) estimated from eqs and 6 agreed with the experimentally determined value of 0.103 C/μmol (Figure ). Similarly, the value estimated from eqs and 8 (0.0965 C/μmol) agreed with the experimentally determined slope of 0.0934 C/μmol (Figure ). Quantitative results of the concurrent titrations of KHP(aq), MnO4–(aq), OH–(aq), and S2O32–(aq) are presented in Table .

Table 1

Estimated and Experimentally Determined Charges for Concurrent Coulometric Titrations

analyte	estimated charge per micromole of analyte (in C/μmol) from stoichiometry and Faraday’s laws of electrolysis	experimentally determined charge per micromole of analyte (in C/μmol) from the slope	coefficient of variation (%) (n = 5)
KHP(aq)	0.0965	0.100	3.6
MnO4–(aq)	0.482	0.466	3.3
OH–(aq)	0.0965	0.103	2.1
S2O32–(aq)	0.0965	0.0934	1.5

The linear relations between the charge and the amount of reagent were consistent with the simultaneous titrations using reagents from both compartments. Equations and 9 indicate that 0.0965 C charge accounts for neutralizing 1 μmol of acetic acid. The experimentally determined value of 0.0987 C/μmol of acetic acid agreed with this estimate. Similarly, an estimate (from eqs and 10) of 0.193 C/μmol of ascorbic acid agreed with the experimentally determined value of 0.189 C/μmol.

Quantities of acetic acid and ascorbic acid determined from simultaneous coulometric titrations were 5.1% and 980 mg, respectively, in agreement with the quantities determined from volumetric titrations (5.1% and 990 mg).

Conclusions

The coulometric titrations offer a number of advantages over conventional titrations[1] and the proposed concurrent coulometric titrations offer added features (Supporting Information, Table S1). Easy to construct, unbreakable and transparent poly(dimethylsiloxane) (PDMS) cell, disposable salt bridges, and small reagent volumes (about 2 mL in each compartment) add to the salient features of coulometric titrations.

Proposed concurrent coulometric titration method offers flexibility of combining multiple independent titrations. Three concentrations of an analyte can be concurrently titrated in three cathode compartments (as presented in Figures and 2), or three anode compartments (as presented in Figures and 4), or a combination of two different analytes simultaneously titrated in respective compartments (Figure ). The simultaneous titrations helped achieve six titrations in a single-course electrolysis.

Figure 5

Plot of the charge passing through the cathodes and the anodes versus the amounts of acetic acid (blue circles) and ascorbic acid (orange circles) titrated simultaneously. Microliters and milligrams of vinegar and finely powdered supplement tablet, respectively, were added to the compartments to achieve desired concentration.

Quantities of acetic acid and ascorbic acid determined from simultaneous titrations (5.1% and 980 mg, respectively) agreed with the quantities from manufacturer’s label (5.0% and 1000 mg, respectively). Relatively small coefficients of variation (mean value ≤3.6% for titrations presented in Figures –4 and ≤2.7% for titrations presented in Figure ) validated the repeatability of titrations. Deviations in titration results are presented in the Supporting Information section (Tables S2 and S3). Monitoring concentrations of acetic acid and ascorbic acid are important in view of diet intake.[31,32] Simultaneous coulometric method presented in this study exemplifies this monitoring.

Experimental Section

Materials

All reagents were obtained from Fisher Scientific and used as received. Platinum wire (0.5 mm diameter) was purchased from Alfa Aesar. Multicompartment electrolysis cell was made of poly(dimethylsiloxane) (PDMS) based on the procedure described previously.[33] Three pairs of electrolysis cells were connected by platinum wires and the platinum wires in the terminal cathode and anode compartments were connected to the coulometer (Figure ). Adjacent pairs of cathode and anode compartments were connected with disposable salt bridges. Obbligato-Objectives Faraday-MP potentiostat was connected to a computer via USB connection, serving as a coulometer. The current passing through the cell for the titrations ranged from 4 to 10 mA. Desired amounts of analyte sample, electrolyte, and indicator were added to the relevant compartments. The contents of the cell for each analyte are summarized in Table S4 in the Supporting Information.

Figure 6

Schematic diagram of the electrolysis cell: C1 and A1 represent the cathode and the anode compartments of the first cell, respectively. SB represents a salt bridge. C2, A2, C3, and A3 represent the cathode and anode compartments of the second and the third cell, respectively. Platinum wires served as terminal cathode and anode (in C1 and A3 compartments, respectively) and interconnecting electrodes between three cells.

The concentration of KHP was directly determined from its mass. The concentrations of other analytes (MnO4–(aq), OH–(aq), and S2O32–(aq)) were individually determined from volumetric titrations.

Concurrent Titrations Using Cathodically Produced OH–(aq) as the Reagent

Desired volumes of KHP(aq) were added to three cathode compartments as presented in Table S4. As current was passed through the cell, OH–(aq) ions produced in the cathode compartments chemically reacted with the KHP(aq). Completion of this secondary chemical reaction was monitored by visual color change of the indicator. The cathodic charge passing through the cell was recorded as the titration endpoint (colorless to pink) was reached in each cathode compartment. The titrations were repeated for three successively higher concentrations of KHP(aq). Thus, the titration of six concentrations of KHP (2.59–15.5 μmol) were performed in two courses of electrolysis. Average coulombs for five independent sets of experiments are presented in Figure .

Concurrent Titrations Using Cathodically Produced Fe2+(aq) as the Reagent

Electrolytically produced Fe2+(aq) served as a titrant in an independent set of titrations of MnO4–(aq). Analyte MnO4–(aq) serving as a self-indicator prompted the endpoint of the titration (purple to colorless). Multicompartment titrations of MnO4–(aq) were performed in the similar way as the titrations of KHP(aq). Figure presents the response of the cathodic charge passing through the cell to the concentration of MnO4–(aq) in 0.48–2.9 μmol range, determined from five independent sets of experiments.

Concurrent Titrations Using Anodically Produced H+(aq) and I2(aq) Reagents

During the course of electrolysis, H+(aq) ions were concurrently produced in the anode compartments (eq ). Validity of the cell was tested for the titration of OH–(aq) ions with the electrolytically produced H+(aq) ions. NaOH(aq) as an analyte sample with increasing concentrations was placed in three anode compartments with an indicator. Electrolytically produced H+(aq) chemically reacted with OH–(aq) (eq ). As stated earlier, the charge passing through the cell was recorded as the titration endpoint (yellow to red) was reached in each anode compartment. The titrations were repeated for successively higher concentrations of OH–(aq). Figure presents the response of the anodic charge passing through the cell to the concentration of OH–(aq) in 9.7–58 μmol range, determined from five independent sets of experiments.

Similar titrations were performed for the quantification of S2O32–(aq) using electrolytically produced I2(aq) reagent. Charge was recorded as the titration endpoint (colorless to black) has been reached. Figure presents the response of the anodic charge passing through the cell to the concentration of S2O32–(aq) in 10–62 μmol range, determined from five independent sets of experiments.

Simultaneous Titrations Using Reagents Produced in Both Compartments

In a course of electrolysis, reduction in the cathode compartment and oxidation in the anode compartment results in producing two reagents. These reagents were used as titrants for simultaneous titrations of two different analyte samples. We tested the feasibility of using both reagents to simultaneously titrate two analyte samples. In this course of electrolysis, we quantified active ingredients (acetic acid and vitamin C) from two household products. Concurrently produced OH–(aq) and I2(aq) in the respective compartments served as titrants for acetic acid and ascorbic acid, respectively (Table ). Desired moles of analyte samples were placed in the alternate compartments. This was achieved by placing vinegar samples in C1, C2, and C3 compartments and ascorbic acid samples in A1, A2, and A3 compartments (Figure ). Figure presents the electrolysis charge required to produce the titrant and reach the endpoint for each analyte sample. Five independent sets of experiments are presented. Five data points for each acetic acid concentration are nearly overlaid. Vertical and horizontal scatter for the data points for ascorbic acid represent small variations in the recorded charge and measured mass of analyte sample, respectively.

Table 2

Summary of Simultaneous Coulometric Titrations of Acetic Acid and Ascorbic Acid in Household Vinegar and Vitamin C Dietary Supplement Tablets

anode compartment	cathode compartment
analyte: vitamin C dietary supplement tablets	analyte: household vinegar
content determined: ascorbic acid (AA)	content determined: acetic acid (CH3COOH)
primary electrolytic reaction: 2I–(aq) → I2(aq) + 2e (eq 7)	primary electrolytic reaction: 2H2O(l) + 2e → 2OH–(aq) + H2(g) (eq 1)
10	9
visual indicator: 1% aqueous starch	visual indicator: 50 μL 0.1% phenolphthalein