Abstract Development of a Teaching Coulometry

ABSTRACT

DEVELOPMENT OF A TEACHING COULOMETRY INSTRUMENT FOR THE DIRECT DETERMINATION OF SULFUR COMPOUNDS AND OF ZINC INDIRECTLY

by Jeralyne Beatriz Padilla Mercado

The development and characterization of a teaching coulometry instrument for on-line data acquisition is described. A constant current source connected to nonisolated Pt electrodes in a 150-mL beaker served as the cell where the iodine titrant is electrochemically generated and allowed to react with the analyte reducing agent. A photodiode monitored the darkening of the solution due to the starch-iodine complex permitting this titration curve to be stored by a multifunctional chemical analysis system used for teaching for subsequent graphical analysis. Characterization of the instrument for direct analyte determination is performed with ascorbic acid, thiols, thiosulfate, and bisulfite. The number of electrons per mole of thiol for iodine titration of glutathione and N-acetylcysteine varied as a function of pH, indicating different reaction pathways. Ascorbic acid and the thiols are determined in dietary supplements with a recovery of 90- 100%. The indirect determination of zinc after its complexation with cysteine was performed in alkaline media. The titration endpoint times of cysteine with zinc are proportionally longer as compared to cysteine itself. The determination of ascorbic acid and zinc in a supplement could be titrated sequentially without changing the sample. Recovery of zinc ranged from 96-112% with a RSD range of 6-11%.

DEVELOPMENT OF A TEACHING COULOMETRY INSTRUMENT FOR THE DIRECT DETERMINATION OF SULFUR COMPOUNDS AND OF ZINC INDIRECTLY

A Thesis

Submitted to the

Faculty of Miami University

in partial fulfillment of

the requirements for the degree of

Master of Science

Jeralyne Beatriz Padilla Mercado

Miami University

Oxford, Ohio

2017

Advisor: Neil D. Danielson

Reader: Jiangjiang Zhu

Reader: Dominik Konkolewicz

Reader: Richard Bretz

This Thesis titled

DEVELOPMENT OF A TEACHING COULOMETRY INSTRUMENT FOR THE DIRECT DETERMINATION OF SULFUR COMPOUNDS AND OF ZINC INDIRECTLY

Jeralyne Beatriz Padilla Mercado

has been approved for publication by

The College of Arts and Science

and

Department of Chemistry and Biochemistry

______Neil D. Danielson

______Jiangjiang Zhu

______Dominik Konkolewicz

______Richard Bretz

Table of Contents

Chapter 1. Introduction P. 1

Chapter 2. Iodine coulometry with on-line photocell detection for a multifunctional chemical analysis (MCA) system P. 21

Chapter 3. Indirect determination of zinc by thiol complexation and iodine coulometric titration with photodiode detection P. 59

Chapter 4. Conclusions and future directions P. 75

Appendix. Oxygen meter finger probe studies P. 77

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List of Tables Chapter 2 Table 2.1. Analytes determine in commercial products P. 27 Table 2.2. Average number of electrons per mole of analyte (% RSD). P. 28 Table 2.3. Standards and commercial products determined With the home-built coulometer. P. 30

Chapter 3 Table S1. Complete label listings of ingredients in the commercial samples. P. 74

Appendix Table A.1. Reproducibility studies with electric toothbrush. P. 81 Table A. 2. Attempts to make calibration curves of permanent marker paper chromatography separations. P. 82 Table A.3. Filter paper with permanent marker spots trials. P. 82 Table A.4. Trials of IR active compounds. P. 83 Table A.5. Teflon tape, nitrile glove, and salicylic acid sample in oximeter. P. 84 Table A.6. Blank trials with transfer pipette and Teflon tape. P. 85

List of Figures Chapter 1 Figure 1.1. A constant-current coulometer. P. 4 Figure 1.2. Coulometer with isolated electrodes and salt bridge versus coulometer with nonisolated electrodes. P. 4

Figure 1.3. Helical amylose units encasing 3I2 units. P. 7 Figure 1.4. Ascorbic acid structure. P. 15 Figure 1.5. Glutathione structure. P. 15 Figure 1.6. N-acetylcysteine structure. P. 15 Figure 1.7. Thiosulfate structure. P. 15 Figure 1.8. Bisulfite structure. P. 15 Figure 1.9. Cysteine structure. P. 17 2- Figure 1.10. Zncys2 complex at alkaline pH. P. 17

Chapter 2

Figure 2.1. Picture of constant-current coulometry instrument. From left to right: current source, electrodes in coulometric cell on top of magnetic photodiode and magnetic stirrer, and current-to-voltage converter circuit. MeasureNet station in the back. P. 24

Figure 2.2. Typical titration plot. P. 24 Figure S1. Proper alignment of electrodes, stir bar, and photocell. P. 33 Figure S2. Circuit diagram of the current-to-voltage converter. P. 34 Figure S3. Close-up picture of the breadboard used to construct the current-to-voltage converter circuit. P. 34 Figure S4. Titration plot of ascorbic acid titration using the 1 cm and 2 cm platinum electrodes as the anode and cathode, respectively. P. 35

Figure S5. Titration plot of ascorbic acid titration with 2 cm and 1 cm platinum electrodes as anode and cathode. P. 35 Figure S6. Ascorbic acid titration plot using the 1 cm Pt cylinder as the anode and the medium counter electrode (Platinum Inlay Cat. No. 476060 from Corning). P. 36 Figure S7. N-acetylcysteine titration plot using a dialysis membrane to cover the 1 cm cylinder cathode. P. 36

Chapter 3 Figure 3.1. Zinc-cysteine complex structure as described in the literature. P. 61 Figure 3.2. (a) The raw plot (voltage vs. time), (b) first derivative plot (first derivative versus average time), and (c) the normalized voltage plot (normalized voltage vs. time). P. 64 Figure 3.3. Titration of cysteine and cysteine with zinc standard showed the least variability in

endpoint measurements when using the buffer at its pKa, 9.2. P. 66 Figure 3.4. Effect of stir rate on titration endpoints. P. 67 Figure 3.5. Interference studies on cysteine titration endpoints. P. 69 Figure S1. Calibration curve of 1.24 x 10-1 mM cysteine with 0 - 1.74 x 10-5 mM zinc nitrate hexahydrate. The first four and last four points were taken 5 days apart using the same cysteine stock solution. P. 72 Figure S2. Cysteine and cysteine with zinc endpoints titrated in two different temperature baths. P. 72 Figure S3. Five-point calibration curve of titration of 1.3 x 10-1 mM cysteine with zinc standard ranging from 1.2 x 10-2 to 3.6 x 10-2 mM zinc. P. 73

Figure S4. Calibration curve of constant cysteine concentration with increasing zinc concentrations for ascorbic acid and zinc combined titrations. P. 73

Appendix Figure A.1. Heme group in hemoglobin. P. 78 Figure A.2. Picture of oximeter used in these studies. P. 79 Figure A.3. Green marker dot on electric toothbrush. P. 81

vii

Dedication

I want to dedicate this thesis to my parents, Ada and Rafael. I have worked tirelessly to reach my goals and make them proud. Thanks to their support and continuous words of encouragement I can say I am a master in chemistry.

viii

Acknowledgements

Words cannot adequately express how grateful I am to have had Dr. Neil D. Danielson as my thesis advisor. It has been an honor for me to work with him during these past two years. I greatly appreciate his time and effort in helping me with my research and writing. He always goes above and beyond and makes the best out of every mentoring opportunity. From my time in the Danielson Research Group I will take with me his love of teaching, creativity, and resourcefulness.

I also want to thank my committee members for taking the time to read my thesis and for providing helpful feedback on it. Dr. Jiangjiang Zhu for agreeing to be my graduate committee chair, Dr. Dominik Konkolewicz for his input during a crucial time of my zinc studies, and Dr. Richard Bretz for his insights.

I am indebted to Dr. Zuleika Medina Torres, Prof. Edgardo Ortíz Nieves, and Dr. Stacey López Rivera for encouraging me to pursue graduate studies. I hope I am making you proud.

Lastly, I want to thank the Danielson Research Group, the great friends I have met at Miami: Caitlyn; Andrew; Bryce, and my husband, Jean Pall, for their patience and support throughout this journey.

Chapter 1. Introduction The coulometry research described in this thesis is two-fold in nature. Specific aim 1 is the development and characterization of a teaching coulometry instrument compatible with a multi-chemical analysis (MCA) station network. The versatility of iodine coulometry for sulfur compounds as well as ascorbic acid is demonstrated. Specific aim 2 describes an analytical chemistry research project involving the indirect determination of zinc using cysteine and iodine coulometry. These aims are explained in more detail at the end of this chapter. What immediately follows is a summary of the fundamentals of coulometry, the variety of applications, and how coulometry has been presented in the teaching literature. I. Fundamentals of coulometry Coulometry is an electrochemical method used to quantitate inorganic and organic analytes by measuring the current and time needed to change their oxidation state. Coulometry has many advantages when compared to other analytical procedures. First, it is exceptionally good when small amounts of sample are measured. Second, unstable titration reagents that would not be used otherwise can be generated in situ, and third its quick analysis time makes automation possible. Constant-current coulometry requires straightforward data analysis because the measurement charge (Q) is directly proportional to the product of current (I) and time (t) as shown in the following equation: Q = It (1.1) Furthermore, Faraday’s law can be used to relate the measurement charge to the amount of analyte in solution:

molesanalyte = Q/nF (1.2) where n equals the moles of electrons participating in the redox reaction and F is Faraday’s constant, 96485.31 C/mol.1 Faraday’s constant represents the charge corresponding to one mole of electrons. This theory-derived quantity is the proportionality constant between the moles of analyte and the measurement charge which enables the use of coulometric determinations without calibration. a. Instrument set-up A typical constant-current coulometer is shown in Figure 1.1. Coulometers include a switch that starts a timer and current flow simultaneously. Coulometric cells are mounted on magnetic stirrers because stirring is needed to allow the species of interest to undergo a change in

1 oxidation state near the electrode where the titrant is generated. The movement of the analyte towards the electrode is called mass transfer. Aside from mechanical stirring, mass transfer is facilitated by ion movement in the electrolyte solution and by diffusion of species due to electrostatic attractions and concentration gradients. The leftmost electrode in Figure 1.1 is called the generator or working electrode and it is where the titrant is produced from the mediator, permitting the titration reaction of interest to take place. The counter or auxiliary electrode at the right is the source of species needed to complete the redox reaction for generation of the titrant. In order to have reliable results from coulometric analyses 100% current efficiency must be ensured, meaning that all current used corresponds to the production of the titrant which reacts to change the oxidation state of the target analyte. Counter electrodes are usually positioned in a separate compartment inside the cell to prevent the movement of interfering species towards the working electrode which would decrease current efficiency. This is a more important concern if the reaction kinetics are slow. A fritted-glass disk, as shown in Figure 1.1, allows the flow of redox species to the bulk of the solution. Coulometers can be used without this barrier as long as the reaction of the interfering species is slower than the redox reaction of interest. This is the case for determination of milligram amounts of cyclohexene by electrogenerated bromine forming 1,2-dibromocyclohexene with amperometric endpoint detection.2 The coulometric cell was constructed by the researchers by modifying a 100 mL round bottom flask and by using nonisolated platinum electrodes as the cathode and the anode. Nonisolated electrodes worked in this setup because 1,2-dibromocyclohexane cannot be reduced at the cathode and because the reaction of the cathodic product with species in the bulk solution is slow. These assumptions were tested by repeating the titration with isolated electrodes and the same results were obtained in both cases. Furthermore, the authors state that positive error could occur as a result of using nonisolated electrodes if the reaction between bromine and the analyte is slow. This would cause bromine to be reduced at the cathode leading to higher, yet erroneous, endpoints. The results of these titrations showed a precision of ± 1 ppt. Constant-current coulometry has also been used to determine ethylenediaminetetraacetate (EDTA) by electrogenerated copper(II).3 Spectrophotometric detection of the endpoint was possible due to the high molar absorptivity of the Cu-EDTA2- complex in the ultraviolet region. In this experimental setup, the electrodes are nonisolated because the copper(II) ion is readily

2 complexed with EDTA upon its electrolytic formation. Endpoint detection was done by eye and by graphical analysis of absorbance versus titration time plots. The use of nonisolated electrodes in coulometric determinations has also been reported by Kuntzleman and collaborators.4 Figure 1.2 shows a coulometer with isolated electrodes and salt bridge to the left, and a coulometer with nonisolated electrodes to the right. The salt bridge is needed to permit ion flow between both cells. Kuntzleman attributes shorter analysis time and the ability to use smaller instruments due to the absence of a barrier through which species would have to migrate. As shown in Figure 1.2, the use of nonisolated electrodes allows the use of one cell instead of two, making the titration instrument more compact. This homebuilt instrument with nonisolated electrodes was used to carry out neutralization reactions by applying a constant potential.4 In order to eliminate mass transfer of interfering species, an anion was added to precipitate the anode reaction product. Microliter aliquots of household cleaner and vinegar were titrated against electrogenerated hydroxide ions. The titration data were obtained in real time using a MicroLab data acquisition instrument. The endpoint was determined visually by use of phenolphthalein indicator. Endpoints were also determined graphically by finding the peak of the first derivative plot of current versus time plots because this is the procedure required to determine charge when constant potential coulometry is used, as shown in Equation 1.3. 푡 푞 = 퐼 푑푡 (1.3) ∫0

Figure 1.1. A constant-current coulometer.5

Figure 1.2. Coulometer with isolated electrodes and salt bridge versus coulometer with nonisolated electrodes.4

4 b. Coulometry detection methods Coulometric determinations can be coupled to different types of endpoint detection methods as it has been shown in the two examples above. This section will describe various endpoint detection methods. For example, amperometry is an electrochemical technique that is commonly used for this purpose. Amperometry consists of ion detection by monitoring the change of current in solution due to the electrolysis of the ion of interest.5 The reaction endpoint in this setup is considered to be the point in which the constant current output is disturbed. The setup for amperometric detection, also referred to as biamperometry, requires the use of two electrodes set at a constant voltage. The current between the twin electrodes is measured during the analysis. Graphical analysis is another method used in coulometric endpoint determinations. The electrical signal is plotted against time. These graphs can take many different forms but the overarching idea is that the change in signal is correlated to the endpoint time where the reaction of interest is finalized. Photometric endpoint determination is another method used in coulometry. Photometry relies on a color change in solution that can be correlated to the point in which all the analyte has undergone a change in oxidation state. The endpoint is typically considered to be the time in which the preselected color change is uniform in the titration cell. Visual indicators are chosen based on the redox reaction that takes place during analyte determination. For example, phenolphthalein can be used for acid and base determinations; the protonated form is colorless and the deprotonated form is pink. This change in color occurs at a pH of 8 – 9.6.5 Starch can be used as a visual indicator in titrations where iodine is electrogenerated. The chemistry behind starch-iodine detection is explained below. Iodine is known to be a mild oxidizing agent heavily used in coulometric titrations. As it was briefly explained above, such analyses use starch as the indicator of the titration endpoint. Starch is a carbohydrate composed of two types of molecules known as amylopectin and amylose.6 Amylose is the linear portion of starch which is the part involved in the molecule’s role as an indicator; it is a polymer composed of the unit α-D-glucose.5 In an aqueous solution where there 7 is an excess of iodine present, I2 molecules enter the amylose helix as shown in Figure 1.3. This structure is the chromophore responsible for the dark blue color of the starch-iodine complex. Calabrese and Khan studied the solution composition of the starch-iodine complex and found

5 that three molecular iodine species in the form of I2 are positioned within twelve amylose units, which is the structure shown in Figure 1.3. The effect of iodide concentration in coulometric determinations of arsenite(III) has been studied by D. A. Bell.8 High concentrations of iodide ions in solution populate the helical amylose structure instead of the 3I2 species disrupting the formation of the starch-iodine complex. Therefore, iodide concentrations of 0.025 M to 2 M were used for constant-current analyses of arsenite(III). The different current magnitudes applied were 5, 10, and 25 mA. The results of these tests showed that potassium iodide concentrations in the range of 0.1 to 2 M resulted in less precise measurements. Micro amounts of arsenite(III) were successfully determined in eighteen trials using 0.025 M KI with a standard deviation of 1%. II. Applications of coulometry a. Coulometry of organic reducing agents Several research articles that make use of the advantages of coulometry have been published. Ascorbic acid was determined using electrogenerated iodine by controlled-potential coulometry with isolated electrodes.9 The instrument had an electronic integrator that automatically displays the ascorbic acid concentration. Standard solutions of ascorbic acid were prepared to ensure method accuracy. Acetic acid was used as a titration medium, to make the standard solutions, and to extract the analyte from pharmaceutical samples. The titration times of this method ranged between two and three minutes and the relative standard deviation (RSD) of their measurements was within ± 0.1%. The construction of a coulometric cell that uses large electrode areas, small cell volumes, and a detector that provides rapid response times has been reported.10 This constant-current setup with isolated electrodes was used to titrate microgram amounts of thioglycolic acid and cyanide using electrogenerated iodide. The thiosulfate left in solution after these titrations was used to determine microgram amounts of iron(II), iodine, and copper(II). An added advantage with this setup is that the electrolyte solution can be used for up to fifty different determinations with reaction times ranging from ten to one hundred seconds. The detectors incorporated in this electrolytic cell include photometric and biamperometric flow-through detectors.

7 Figure 1.3. Helical amylose units encasing 3I2 units.

Biologically active thiols have also been titrated using electrogenerated iodine in methanol and using bromine in acetic acid.11 Methanol was used as the titration medium for iodine analyses because the reagent presented greater stability in an alcohol solvent. The use of acetic acid as a titration medium for bromine reactions is justified by the lower volatility of the titrant in this solvent. Cysteine, 2-thiouracil, 6-mercaptopurine, and 6-thioguanine were determined using both titrants but more accurate results were obtained when electrogenerated iodine was used. Isolated electrodes and biamperometric endpoint detection methods were used for these determinations. The method that was developed in this investigation was used to determine 6- mercaptopurine in pharmaceuticals and achieved recoveries greater than one hundred percent. The coulometric titration of carbimazole, a pharmaceutical drug that is activated after being metabolized, has been reported.12 Electrogenerated iodine was used in alkaline media to determine milligram amounts of carbimazole in standards and pharmaceutical tablets. Biamperometric endpoint detection was used in these determinations with a commercial coulometer of isolated electrodes. The results obtained include RSDs ranging from 0.01 to 0.23%. The coulometric determination of glutathione using amperometric endpoint detection has been previously reported.13 In this setup with isolated electrodes standard glutathione solutions were titrated with electrogenerated iodine yielding results with RSDs ranging from 1 – 6% and a limit of detection of 1.9 x 10-5 M. Human blood samples were analyzed using the same method with RSDs of 8 and 9%. Cysteine has been coulometrically determined using electrogenerated iodine with an amperometric endpoint detection method.14 Standard cysteine solutions were titrated with electrogenerated iodine yielding results with RSDs ranging from 1 – 3%. Cysteine tablets were analyzed using electrogenerated chlorine (1 – 3% RSD) and bromine (1 – 2% RSD). The analysis of starch and metabisulfite content in corn syrup using flow injection coulometry has been presented.15 To determine the starch content, diluted corn syrup is inserted in the flow injection system along with an electrolyte solution containing iodide. Iodine is electrogenerated in the chamber and its reaction with the starch is followed spectrophotometrically. Metabisulfite is determined in the flow injection chamber by titration with electrogenerated iodine. The computer-regulated analyses resulted in RSDs of less than 1.5% for sodium metabisulfite concentrations over the range of 3.5x10-4 to 2.9x10-2 M. Similar

8 results were obtained during the analysis of starch content: RSD of less than 1.4% and correlation coefficients of 0.997 of calibration curves of starch standard solutions. The determination of bromine, specifically the bromine number, using flow injection coulometry has been introduced as a suitable alternative to the methods used in industry.16 The bromine number has been defined as the grams of bromine that react with 100 grams of substance under a specific set of conditions. In this study, bromine was electrogenerated and reacted with olefins in solution. The titration endpoint was used to calculate the bromine number and olefin concentrations in petroleum samples. The results showed RSDs of 2%. b. Coulometry methods for cation detection Heavy metal detection methods are extremely important given the adverse health effects that these elements have in the human body when the accumulated amounts reach toxic levels.17 To take advantage of the inherent sensitivity of coulometric determinations, water samples have been analyzed for arsenic(III) content.18 In this investigation arsenic(V) was reduced to its electrochemically active form, arsenic(III), by microwave assistance in a closed vessel to reduce analyte loss. Microgram amounts per liter of contaminated sample (tap, sea, or waste water) were determined using iodine coulometry. The development of a coulometric method that can determine several heavy metals by electrodeposition followed by anodic stripping was reported.18 Anodic stripping is a two-step method in which the analyte in a dilute solution is reduced at the electrode. Electrodeposition is followed by a change in voltage direction which re-oxidizes the species. The resulting peak in the current versus applied potential plot is directly related to the amount of analyte present in the solution.5 Total concentrations of cadmium, zinc, copper, and lead were successfully determined. The method’s selectivity is dependent in part to the initial and final potentials that are used per trial. The supporting electrolytes, pH of solution and applied potentials for successful determinations of these four cations were determined. Mercury has been coulometrically determined by taking advantage of its catalytic role in the reaction between ferrocyanide and 1,10-phenanthroline as shown in Equation 1.4.19 Electrogenerated iodine is used to titrate the cyanide ion and mercury concentration is determined by a standard addition method. The concentration of cyanide ions is correlated to the concentration of mercury that catalyzes the reaction through stoichiometry. Mercury determination was performed in blood serum, urine, and fish samples. These were digested and

9 underwent liquid-liquid extraction before being titrated with iodine. Micro amounts of mercury were successfully determined using this method. 4- 2+ - Fe(CN)6 + 3(1,10-phenanthroline) ↔ ferroin + CN (1.4) Chromium has also been determined by direct electrolysis using a controlled-potential coulometric method.20 To eliminate possible interferences from other elements, chromium metal is converted to chromium(II) followed by its conversion to chromium(III). Micro amounts of chromium were determined by monitoring the current of a mercury working electrode versus a saturated calomel electrode during analyte redox. A standard chromium(III) solution was determined with precision of 1% and milligram amounts were determined with 0.1% precision. Mercury plated platinum electrodes have been used in the coulometric determination of sub- microgram amounts of cadmium and zinc standard solutions.21 These analytes were determined by anodic stripping, where a change in voltage is used after metal electrodeposition to cause their dissolution. Voltage changes for cadmium analyses ranged from -1.0 to -0.3 V versus saturated calomel electrode and from -1.3 to -0.7 V for zinc analyses. Separate and co-determinations of these heavy metals were successfully performed though the method proved to be more sensitive towards cadmium. A twelve-minute analysis of a 10-5 M cadmium solution resulted in a recovery of 10-5 grams of cadmium compared to a separate zinc determination of the same analysis time which yielded 10-10 grams of zinc from a 10-4 M zinc solution. The coulometric determination of zinc using ferrocyanide as a precipitating agent has been published.22 In this procedure ferrocyanide was electrolytically generated from ferricyanide in acidic pH. Upon its electrogeneration, the ferrocyanide precipitated with zinc as

K2Zn3[Fe(CN)6]2. This procedure permitted the determination of milligrams of zinc with average errors ranging from -1.1 to 2.0%. III. Summary of coulometry teaching publications There are many education research articles focused on teaching electrochemical analysis. The coulometric determination of weak acids using electrogenerated hydroxide has been shown.23 A commercial, photometric titration assembly with isolated electrodes was modified by addition of a 250 ml beaker to be used as a coulometric cell. The results of the titration of potassium hydrogen phthalate using thymol blue as an indicator show 0.19% RSD when using the potentiometric endpoint detection method. The photometric detection method showed 0.57% RSD.

A miniature coulometric cell made by Dabke and co-authors was used for neutralization reactions in a teaching setting.24 The polydimethylsiloxane (a silicone polymer) cell with separate anode and cathode compartments was used for the titration of base in household ammonia and acid in Lysol. Electrogenerated hydrogen ions were used to analyze ammonia content and methyl orange was used as the indicator. Electrogenerated hydroxide ions were used to determine acid and phenolphthalein was used to indicate the endpoint. Other experiments performed with the miniature cell include the analysis of iron supplement tablets and povidone- iodine. Iron was indirectly determined by using cerium(III) ions with ferroin as the indicator. Povidone-iodine was back-titrated using sodium thiosulfate excess and starch as the indicator. These coulometric titrations were compared to volumetric titrations of all the analytes and the results show that both methods are in good agreement and close to the label values. The use of electrogenerated reagents in a homebuilt instrument for volumetric analyses has been reported.25 The setup has separate anodic and cathodic compartments. The analytes are oxidized by electrogenerated reagents in the anode while water is reduced at the cathode. The hydroxide ions that are generated as a result of water reduction in the cathode are titrated with potassium hydrogen phthalate. The endpoint of this volumetric titration is signaled by phenolphthalein indicator and it is stoichiometrically related to the analyte concentration in the anode. This principle was used to determine ascorbic acid in supplement tablets with electrogenerated iodine using starch as the endpoint indicator. Iron in supplement tablets was determined with electrogenerated cerium(III) using ferroin indicator. The results of ascorbic acid and iron determinations showed standard deviations in the order of 10-3. Electrogenerated iodine has been used to determine total sulfite content in white wine samples.26 The free form sulfite content and the form that is bound to saturated compounds were both quantified using a commercial coulometer. The coulometric data was validated using an official distillation/titration method for sulfite analysis in foods from the Association of Analytical Chemistry (AOAC) International, a United States Department of Agriculture non- profit group. Using electrogenerated iodine 3.7 ± 0.2 mM of sulfite were found compared to 3.6 ± 0.2 mM found using the AOAC method. Constant-current coulometry was used to determine arsenite using electrogenerated iodine.27 A commercial coulometer with isolated electrodes was used and three different endpoint detection methods were compared. Volumetric pipets were calibrated before being used to

11 measure 1 mL aliquots of arsenite samples. Starch was used as the indicator for eye detection which found an average endpoint of 199.16 ± 0.71 seconds. Potentiometric endpoint detection was performed using a pH meter resulting in an average endpoint of 199.08 ± 0.23 seconds, while the amperometric method turned out to be the most precise one with an average endpoint of 199.12 ± 0.12 seconds. Determination of ascorbic acid in natural orange juice was performed using electrogenerated iodine and bromine with constant current.28 A commercial coulometer was used to determine ascorbic acid concentration in standards and in the juice. A gauze cylinder platinum electrode was used as the anode and was kept in a separate compartment from the wire spiral platinum cathode. Iodine was used as the indicator for the iodine method while methyl orange served as the visual indicator with bromine. Analyte recoveries from triplicate measurements of the standards were 101.5% for the iodine method and 102.1% for the bromine method. The higher oxidizing power of bromine compared to iodine was the explanation given to this greater recovery. These high yields were attributed to impurities that would impede a current efficiency of 100% during analyses. The building and characterization of a constant current coulometer that uses simple circuits has been described.29 This instrument is powered by two nine volt batteries and uses operational amplifiers as the electronic components. The setup is also composed of separate platinum electrodes in a 150-mL beaker with a stir bar. This coulometer is intended to be used with electrogenerated bromine as the titrant. Phenols and olefins were tested but hydrazine was found to be the best analyte to be titrated in this setup. Phenol had a slow reaction with bromine and olefins created too much cell resistance that impeded the flow of current needed to generate the titrant. Graphical endpoint determination methods of potentiometric measurements were used and student data showed that the standard deviation of 12 measurements was ± 1.5% hydrazine. IV. Specific Aims a. Specific Aim 1 Multifunctional chemical analysis (MCA) systems provide a viable alternative for large scale instruction while supporting a hands-on approach to more advanced instrumentation. These systems are robust and typically use student stations connected to a remote central computer for data collection, minimizing the need for computers at every student workspace. MCA networks offer multiple measurement capabilities, including temperature, potential, drop counting titration,

12 and spectrophotometry.30 However, constant current coulometry is still an uncommon MCA option. We have designed and characterized an inexpensive coulometry with photodiode detection instrument that can be used in conjunction with a MCA system. This instrument is composed of a selectable constant current source connected to twin platinum electrodes contained in a beaker mounted on top of a photodiode sensor connected to a homemade current- to-voltage circuit. The titrant is electrogenerated iodine and is used to determine analytes in a variety of commercial products. The initial iodide concentration from which the titrant is generated follows D. A. Bell’s research and is 0.03 M KI.8 Iodine acts as the oxidizing agent of the analyte in solution as shown below: - - Titrant generation at the anode: 2I → I2 + 2e (1.5) + - Reaction at the cathode: 2H + 2e → H2(g) (1.6) - In solution: I2 + Analytered → Analyteox + 2I (1.7) where the anode is the working electrode and the cathode is the counter electrode. We have chosen starch as the endpoint indicator because the presence of molecular iodine in solution occurs after all the analyte has been oxidized. At this point the I2 species enter the amylose units and the starch-iodine complex forms, signaling the titration endpoint. The darkening of the solution due to the formation of the blue iodine-starch complex is monitored by a photodiode mounted at the bottom of the beaker and the circuit voltage output is recorded as a function of time using the MeasureNet MCA system. The overall output profile looks similar to a titration curve; the first definite uptick (3 times the standard deviation of the baseline noise) in voltage is considered to be the endpoint time. The analytes to be coulometrically analyzed using the scheme presented in Equations 1.5 – 1.7. are ascorbic acid, glutathione, N-acetylcysteine, thiosulfate, and bisulfite. Their chemistry is explained below. Ascorbic acid, otherwise known as vitamin C, is an organic compound necessary for humans. Appropriate levels of ascorbic acid in the body help maintain healthy bones, skin, and cartilage. This vitamin cannot be synthesized by the body which makes its intake through food very important. Ascorbic acid deficiency may be treated by using dietary supplements or high ascorbic acid-containing foods such as oranges. Figure 1.4 shows the structure of ascorbic acid. Equation 1.8 shows the oxidation of ascorbic acid (AA) by iodine into dehydroascorbic acid (DHA). + - AA + I2 → DHA + 2H + 2I (1.8)

Glutathione is an antioxidant present in human bodies. The role of antioxidants in the body is to eliminate oxidizing species, such as free radicals, which can harm cells. In extreme cases, the existence of oxidizing species can lead to the formation of malignant tumors. The structure of glutathione is shown in Figure 1.5. As shown in Equation 1.9, the reduced form of glutathione (GSH) can be oxidized by iodine into GSSG, in a 2 to 1 ratio of the thiol to iodine. The disulfide bond is considered to be the oxidized form since the sulfur atoms lose a bond to hydrogen and gain a bond to one another. + - 2GSH + I2 → GSSG + 2H + 2I (1.9) Another important reducing agent present in the human body is N-acetylcysteine whose structure is shown in Figure 1.6. Just like glutathione, N-acetylcysteine is in charge of getting rid of reactive oxygen species that are harmful to the cells. Equation 1.10 shows the proposed 2 to 1 redox reaction of N-acetylcysteine with iodine, where R represents the organic portion of the thiol. + - 2RSH + I2 → RSSR + 2H + 2I (1.10) Thiosulfate is a compound that occurs naturally in nature and it is used to dechlorinate water. The structure of the thiosulfate ion is shown in Figure 1.7. It’s redox reaction with iodine is shown in Equation 1.11. 2- 2- - 2S2O3 + I2 → 2S4O6 + 2I (1.11) Bisulfites are commonly used as food preservatives to increase shelf life. Sun-dried fruits are one example of foods that contain bisulfites. The structure for this ion is shown in Figure 1.8. Its reaction with iodine in the presence of water is shown in Equation 1.12.15 - + + Na2S2O5 + 2I2 + 3H2O → 2H2SO4 + 4I + 2Na + 2H (1.12) The endpoints of standard solutions of these analytes can be predicted using Faraday’s law. Agreement between the predicted and resulting endpoints of standard solutions shows that the setup is suitable for the determination of the analytes. The analysis of commercial products containing these analytes can be performed subsequently and the resulting concentrations can be compared to the label values, when available. Two endpoint determination methods are to be to be compared in terms of precision in this research: eye detection/hand-timing and graphical analysis.

Figure 1.4. Ascorbic acid structure. Figure 1.5. Glutathione structure.

Figure 1.6. N-acetylcysteine structure. Figure 1.7. Thiosulfate structure.

Figure 1.8. Bisulfite structure.

15 b. Specific Aim 2

To extend the capabilities of the constant-current coulometry instrument, an approach for the indirect determination of zinc, an analyte that cannot be oxidized and quantified using electrogenerated iodine, has been formulated and tested. To the best of our knowledge, this represents a new analytical chemistry contribution to the mature field of coulometry. Among the inorganic, organic, and biomaterials available to aid in electrochemical determinations we have chosen cysteine as a chelating agent to indirectly determine zinc at an alkaline pH. Cysteine is an amino acid that contains a thiol as a side chain, as shown in Figure 1.9. The logarithmic acid dissociation constants of the carboxylate group, the sulfur, and the nitrogen are 1.71, 8.13, and 10.11 respectively.31 Biologically cysteine is the precursor of glutathione and it participates in protein structure and in metal ion chelation. Zinc was our choice of heavy metal due to its biological importance. It is a trace element present in about three thousand proteins of the human proteome. It is the second most abundant transition metal in the body after iron and has structural, catalytic, and regulatory roles.32 Zinc consumption is mainly achieved through foods such as meats, seafood, nuts, dairy products, legumes. The average daily intake necessary for healthy adults is 11 mg for males and 8 mg for females.33 Zinc deficiency can cause many health issues such as growth problems, skin problems, and fetal malformations. The maximum levels of zinc intake for both males and females is 40 mg per day to avoid zinc toxicity.33 Zinc determination in dietary supplements has been chosen given the importance of this metal in the human body. Aside from the foods known to provide a source of zinc, dietary supplements are used to increase total zinc intake. Dietary supplements are sold as tablets, caplets, lozenges, and in liquid form. These are regulated by the Food and Drug Administration (FDA) under the Dietary Supplement Health and Education Act (DSHEA) of 1994.34 The DSHEA sets dietary supplements apart from food and drug products while setting requirements for manufacturers and distributors. The FDA requires their labels to be standardized in terms of claims, ingredients, intended use, and safety information. The FDA is also in charge of handling reports of complaints and/or side effects after the dietary supplements are on the market.35

Figure 1.9. Cysteine structure.

2- Figure 1.10. Zncys2 complex at alkaline pH.

Complex formation between zinc and cysteine can be explained by Pearson’s theory of hard and soft acids and bases (HSAB).36 The theory uses the definition of the Lewis acid base theory where a Lewis base donates an electron pair and is accepted by a Lewis acid. The categorization of species as either hard or soft depends on their degree of polarizability which denotes the ability to form dipoles as a result of their attraction and/or repulsion by external charges. Hard acids and bases have a low degree of polarizability; soft acids and bases are more polarizable. The HSAB theory states that in general soft acids form bonds with soft bases, while hard acids bond preferably to hard bases. Cysteine is considered a soft Lewis base because it has electron pairs readily available to donate through π bonding. The chelation between cysteine and zinc is possible because zinc is a borderline Lewis acid and can accept electrons from the ligand. Zinc determination through its complexation with cysteine is possible due to the stability of this chelate. The formation constant of this complex was experimentally determined using the Bjerrum method and has been reported in the literature.37 This method consists of the titration of a mixture of cysteine hydrochloride and zinc nitrate with sodium hydroxide while monitoring the pH. Subsequent mathematical analyses yielded concentration step-wise formation constants of 7.2x109 and 6.9x108 for the 2:1 complex of cysteine to zinc ion. The results of this investigation also show that the binding sites in the amino acid are the -SH and -NH2 groups, as shown in Figure 1.10. This structure in which the zinc ion has a distorted tetrahedral when forming two 5- membered rings is in accordance with infrared spectra and X-ray crystallography studies of the compound.31,38 By use of a pH titration, the complex formation of cysteine and zinc was studied.39 The stability constants calculated in this investigation were used to calculate the concentrations of 2+ - 2- - 2- different complex species [Zn , Zn(HL)2, Zn3L3(HL) , Zn3L4 , ZnL(HL) , and ZnL2 ] over a pH 2- range of 5 through 8. The predominant species in alkaline pH is the Zncys2 complex with a calculated stability constant 1.6x1018.39 Using a sodium tetraborate decahydrate buffer with pH 9.2 in which cysteine is fully deprotonated the following titration setup will be used to determine zinc:40 - - Titrant generation at the anode: 2I → I2 + 2e (1.13) - - Reaction at the cathode: 2H2O + 2e → 2OH + H2(g) (1.14) - - - - In solution: 2I2 + RS + 4OH → RSO2 + 4I + 2H2O (1.15)

18 where the number of electrons transferred between iodine and cysteine is equal to four as established by studies of iodine reactions with thiols in alkaline medium.41,42 On the basis of a 2- stable Zncys2 complex in the bulk solution a retardation in titration endpoint will be observed. Calibration curves can be used to quantitate zinc in solution and the method can be extended to zinc determination in dietary supplements. V. References (1) Skoog, D. A.; Holler, F. J.; Nieman, T. A. Principles of Instrumental Analysis, 5th ed.; Brooks/Cole, 1997. (2) Evans, D. H. J. Chem. Educ. 1968, 45 (2), 88–90. (3) Williams, K. R.; Young, V. Y.; Killian, B. J. J. Chem. Educ. 2011, 88 (3), 315–316. (4) Kuntzleman, T. S.; Kenney, J. B.; Hasbrouck, S.; Collins, M. J.; Amend, J. R. J. Chem. Educ. 2011, 88 (11), 1565–1568. (5) Harris, D. C. Quantitative Chemical Analysis, 9th ed.; W. H. Freeman and Company: New York, NY, 2016. (6) Teitelbaum, R. C.; Ruby, S. L.; Marks, T. J. J. Am. Chem. Soc. 1980, 102 (10), 3322– 3328. (7) Calabrese, V. T.; Khan, A. J. Polym. Sci. Part A Polym. Chem. 1999, 37 (15), 2711–2717. (8) Bell, D. A. J. Chem. Educ. 1978, 55 (12), 815. (9) Karlsson, R. Talanta 1975, 22, 989–993. (10) Rüttinger, H. H.; Spohn, U. Anal. Chim. Acta 1987, 202, 75–84. (11) Pastor, T. J.; Barek, J. Mikrochim. Acta 1989, 97 (5–6), 407–413. (12) Ciesielski, W.; Krenc, A. Anal. Lett. 2000, 33 (8), 1545–1554. (13) Budnikov, G. K.; Ziyatdinova, G. K.; Valitova, Y. R. J. Anal. Chem. Transl. from Zhurnal Anal. Khimii Orig. Russ. Text 2004, 59 (6), 573–576. (14) Ziyatdinova, G. K.; Grigor ’eva, L. V; Budnikov, G. K. J. Anal. Chem. Orig. Russ. Text © 2007, 62 (12), 1176–1179. (15) Taylor, R. H.; Rotermund, J.; Christian, G. D.; Ruzicka, J. Talanta 1994, 41 (1), 31–38. (16) Taylor, R. H.; Winbo, C.; Christian, G. D.; Ruzicka, J. Talanta 1992, 39 (7), 789–794. (17) National Organization for Rare Disorders. Heavy Metal Poisoning https://rarediseases.org/rare-diseases/heavy-metal-poisoning/. (18) Jurica, L.; Manova, A.; Dzurov, J.; Beinrhor, E. Fresenius J. Anal. Chem. 2000, 366, 260–

266. (19) Rohm, T. J.; Purdy, W. C. Anal. Chim. Acta 1974, 72 (1), 177–182. (20) Meites, L. Anal. Chim. Acta 1958, 18, 364–372. (21) Gardiner, K. W.; Rogers, L. B. Anal. Chem. 1953, 25 (9), 1393–1397. (22) Lingane, J. J.; Hartley, A. M. Anal. Chim. Acta 1954, 11, 475–481. (23) Beilby, A. L.; Landowski, C. A. J. Chem. Educ. 1970, 47 (3), 238–239. (24) Dabke, R. B.; Gebeyehu, Z.; Thor, R. J. Chem. Educ. 2011, 88 (12), 1707–1710. (25) Scanlon, C.; Gebeyehu, Z.; Griffin, K.; Dabke, R. B. J. Chem. Educ. 2014, 91 (6), 898– 901. (26) Lowinsohn, D.; Bertotti, M. J. Chem. Educ. 2002, 79 (1), 103–105. (27) Tackett, S. L. J. Chem. Educ. 1972, 49 (1), 52–54. (28) Bertotti, M.; Moreira Vaz, J.; Telles, R. J. Chem. Educ. 1995, 72 (5), 445–447. (29) Grimsrud, E.; Amend, J. J. Chem. Educ. 1979, 56 (2), 131–133. (30) MeasureNet Technology, Ltd. http://www.measurenet-tech.com/index.html. (31) Bell, P.; Sheldrick, W. S. Zeitschrift fur Naturforsch. - Sect. B J. Chem. Sci. 1984, 39 (12), 1732–1737. (32) Pace, N.; Weerapana, E. Biomolecules 2014, 4 (2), 419–434. (33) National Institutes of Health. Zinc Fact Sheet for Health Professionals https://ods.od.nih.gov/factsheets/Zinc-HealthProfessional/. (34) U.S. Food and Drug Administration. Dietary Supplements http://www.fda.gov/Food/DietarySupplements/default.htm. (35) Levinson, D. R. Dietary Supplements: Structure/Function Claims Fail to Meet Federal Requirements https://oig.hhs.gov/oei/reports/oei-01-11-00210.pdf. (36) Pearson, R. G. J. Am. Chem. Soc. 1963, 85 (22), 3533–3539. (37) Li, N. C.; Manning, R. A. J. Am. Chem. Soc. 1955, 77 (20), 5225–5228. (38) Shindo, H.; Brown, T. L. J. Am. Chem. Soc. 1965, 87 (9), 1904–1909. (39) Perrin, D. D.; Sayce, I. G. J. Chem. Soc. A Inorganic, Phys. Theor. 1968, 53–57. (40) Sigma Aldrich. Safety Data Sheet Sodium tetraborate decahydrate; 2016. (41) Ciesielski, W.; Krenc, A.; Zlobinska, U. Chem. Analityczna 2005, 50, 397–405. (42) Ciesielski, W.; Zakrzewski, R. Chem. Analityczna 2006, 51, 653–677.

Chapter 2. Iodine coulometry with on-line photocell detection for a multifunctional chemical analysis (MCA) system Jeralyne B. Padilla Mercado, Eri M. Coombs, Jenny P. DeJesus, Stacey Lowery Bretz, Neil D. Danielson* Department of Chemistry & Biochemistry, Miami University, 651 E. High Street, Oxford, OH 45056, USA *Corresponding author e-mail address: [email protected]

Abstract Multifunctional chemical analysis (MCA) systems provide a viable alternative for large scale instruction while supporting a hands-on approach to more advanced instrumentation. These systems are robust and typically use student stations connected to a remote central computer for data collection, minimizing the need for computers at every student workspace. MCA networks offer multiple measurement capabilities however, constant current coulometry is still an uncommon MCA option. We have designed and characterized a coulometry instrument with photocell detection that can be used in conjunction with a MCA system. Electrogenerated iodine is used to determine more unusual analytes such as thiols (glutathione and N-acetylcysteine), thiosulfate, and bisulfite. The determination of standard solutions and commercial products by the authors and by undergraduate students are shown. Analyte recoveries range from 90 to 100% and relative standard deviations of triplicate measurements are in the 0.7-5%.

Keywords: First-Year Undergraduate/General, Upper-Division Undergraduate, Analytical Chemistry, Coulometry, Laboratory Equipment/Apparatus, Constant-Current, Iodine, Ascorbic Acid, Thiol Compounds

I. Introduction

The substantial body of teaching literature that covers coulometric determinations shows the continued importance of this electrochemical technique. This analytical method uses in situ- generated titrants for analyte determination. Different titrants can be electrogenerated to determine electroactive species, acids and bases, and even complexing reactants. Iodine, bromine, hydroxide, and hydrogen ions have been reported in the literature.1–13 These home-built instruments recommend either isolating the anodic and cathodic electrodes or instrument simplification using nonisolated electrodes.11,13,14 Coulometric analyses can be coupled to various endpoint detection methods to find the time in which titrations are finished. For example, colorimetric endpoint determinations require the use of chemical indicators that signal the completion of the analyte reactions. Starch is used as the visual indicator when iodine is the electrogenerated titrant.1,3,4,8,9 The use of twin polarized electrodes for amperometric measurements have been used for cyclohexene, hydrazine, biologically active thiols, and carbimazole determinations.6,7,10,11 Some authors have used two or more different endpoint determination methods to compare their precision. Comparison of eye detection vs graphical analysis, potentiometry versus eye detection, potentiometry versus amperometry and eye detection, and spectrophotometry vs amperometry have all been published in this Journal and elsewhere.2,5,12–14 The multifunctional chemical analysis (MCA) system from MeasureNet (www.measurenet- tech.com) consists of student workstations which can be used to work independently.15 Workstations are connected to a controller that connects to a computer and can send the information to be printed or to the Internet. The probes available for the workstations include temperature, pH, conductivity electrodes, pressure sensor, multi-function drop counter, among others. Constant-current coulometry is not an option for this system. Just as previous research in which flow injection analysis and liquid chromatography have been adapted to be used with a MCA system, the present research sets out to incorporate constant-current coulometry as an option.16 The goals of this research are to develop and characterize a constant-current coulometry instrument that provides on-line detection that is compatible with a MCA system. A simple detection system based on a photocell mounted at the bottom of the reaction beaker monitors the darkening of the solution due to starch-iodine complex. The change in resistance of the photocell

22 is monitored by an operational amplifier circuit and the voltage output is sent to the MCA workstation to provide a graphical output to determine the titration endpoint. Electrogenerated iodine is the titrant used to determine ascorbic acid, glutathione, N-acetylcysteine, thiosulfate, and bisulfite in standards and real samples. II. Apparatus a. Instrument Design

The instrument that we have designed is shown in Figure 2.1. The selectable constant-current source, LakeShore Cryotronics Model 121, shown in the picture to the left, was set to 10 mA and connected using wired with alligator clips to the generating and counter electrodes.17 The electrodes, 1 and 2 cm concentric platinum wire and gauze cylinders, are contained in a 150 mL beaker that serves as the titration cell and is mounted on top of a magnetic stirrer. As shown in Figure 2.1, there is a Styrofoam cup lid between the beaker and the magnetic stirrer. The plastic lid is used to mount the photocell sensor, a cadmium-sulphide photoresistor (Radioshack), which monitors the darkening of the solution due to the formation of the blue starch-iodine complex. The beaker is covered with aluminum foil to minimize sources of error in photometric readings due to the surroundings. Figure S1 in the Supporting Information shows the proper placement of the stir bar, the photocell, and the electrodes. This alignment ensures that the photocell is not blocked by the other two components of the cell, minimizing sources of signal noise. The leads from the photocell sensor are connected to a home-built operational amplifier circuit shown in Figure 2.1 to the right, on the blue tray.

Figure 2.1. Picture of constant-current coulometry instrument. From left to right: current source, electrodes in coulometric cell on top of magnetic photocell and magnetic stirrer, and current-to- voltage converter circuit. MeasureNet station in the back.

Figure 2.2. Typical titration plot.

A diagram of the operational amplifier circuit is shown in Figure S2, included in the Supporting Information. The circuit, which was constructed using a Global Specialties breadboard, is essentially a standard voltage operational amplifier circuit.18 A close-up picture of the breadboard is shown in the Supporting Information in Figure S3. One lead of the photocell is connected to the variable output of the trimmer resistor (Radioshack) and the other lead is connected to the inverting input of the 741 operational amplifier (Radioshack). The non- inverting input of the 741 chis is grounded. The photoresistor detects the change in ambient light during the titration; its resistance increases as a function of solution darkening due to the starch indicator. The ratio of the circuit feedback resistor to the resistance of the photocell times the input voltage from the trimmer resistor is the output voltage.19

The iodine titrant is electrogenerated at the anode from I- and the reduction of hydrogen ion to hydrogen gas takes place as the cathode. During the titration, the electrogenerated iodine oxidizes the analyte in solution and the photocell senses ambient light giving a stable baseline reading. The darkening of the solution due to the formation of the blue starch-iodine complex signals the endpoint of the titration.20 The change in color of the indicator blocks ambient light from the photocell, causing an increase in its resistance (R) and a subsequent increase in the voltage output (smaller negative number).

The voltage output of the amplifier circuit is recorded as a function of time using the MeasureNet MCA system. It looks similar to a titration curve as shown in Figure 2.2. The initial portion at around -1600 mV is the baseline reading when the analyte titration is taking place. At around 300 s the solution turns a uniform blue color signaling the titration endpoint. The final stable portion near 0 v, after 350 s, is the result of the intense color of the starch-iodine complex. The endpoint is defined as the first definite uptick in voltage. This is determined by selecting a range of 20 data points before the uptick from baseline. The standard deviation of the voltage signal of this range is multiplied by 3 and added to the voltage value of the last data point from this range, as shown in Equation 2.1. The time corresponding to the calculated voltage is considered the endpoint of the titration.

Endpoint voltage (mV) = 3σ (mV) + voltage of last point(mV) (2.1)

25 b. Nonisolated electrode combinations

To determine the efficiency of the simple setup using nonisolated electrodes, an ascorbic acid standard solution was titrated using different electrode combinations in a 0.03 M H2SO4 medium. The ascorbic acid structure and reaction with iodine are shown in Table 2.1. The 1 cm cylinder and 2 cm concentric gauze platinum electrodes were initially used as the anode and cathode. Figure S4 shows that this setup resulted in a straight baseline; measurements yielded 1.57% relative standard deviation (RSD). Although the RSD value for titrations when the leads were inverted was 0.752%, Figure S5 shows that the baseline had a gradual increase, making it difficult to determine the area before the endpoint. Subsequent electrode combinations used the 1 cm platinum cylinder as the working electrode. A medium counter electrode (Platinum Inlay Cat. No. 476060 from Corning) yielded plots with a lot of noise in the uptick area (Figure S6). The RSD for these measurements was 13.9%. When using a small Pt counter electrode (CHI 102 2mm diameter working electrode from CH Instruments, Inc.) the starch indicator did not turn blue, so no defined curve was observed and no endpoint was determined. A dialysis membrane (Fisherbrand® Dialysis Tubing 10 cm long and 3.2 cm wide) was used to cover the cathode during the titration of a N-acetylcysteine standard in a 0.03 M H2SO4 medium. The 1 cm cylinder and 2 cm concentric gauze platinum electrodes were used as the cathode and anode respectively. A sample titration plot is shown in Figure S7 and demonstrates that the curves have gradual increases in the baseline. It was concluded that the most convenient electrode setup was the 1 cm cylinder as the anode and the non-isolated 2 cm concentric gauze as the cathode. III. Procedure and results a. General procedure

Standards of known concentration are prepared and titrated in triplicate prior to analyzing the commercial product samples. The amount of standard in aliquot needed to reach a predetermined endpoint is calculated using Faraday’s electrolysis equation. Equation 2.2 shows a sample calculation when the target endpoint is 300 s, where I is the current, t is the time, n is the number of electrons transferred per mole of analyte, and F is Faraday’s constant. The mean grams found of the standard triplicates is compared to the theoretical value. A small percent error between the theoretical and experimental amounts can be used as a check to ensure that the titration setup works at essentially 100% current efficiency.

I×t 0.0100 A × 300 s grams analyte = ×molar mass = C ×molar mass (2.2) n×F n × 96485.31 mol e−

The titration of various analytes by this instrument with electrogenerated iodine follows a basic setup. A clean 150 mL beaker is covered on the outside with aluminum foil. One mL volume of 3% starch is added to the beaker along with 0.5 g of KI salt, giving an [I-] of 0.03 M as suggested by the literature.1 The analyte aliquot to be titrated is added, followed by addition of 50 mL of buffer. Water is added to reach a total volume of 130 mL. Ascorbic acid, N- acetylcysteine, bisulfite, glutathione, and thiosulfate were analyzed following this general procedure. The chemistry and sample preparation of each analyte is explained below. b. Commercial product analyses

Table 2.1 shows the structures of the analytes that have been determined in this experiment along with their reaction with molecular iodine, the titrant. The reaction chemistry as indicated in Table 2.1 is as expected except for N-acetylcysteine (NAC). The reason for not identifying the reaction product will be explained in the pH study of thiol compounds

Structure Name Reaction with titrant + − Ascorbic acid AA + I2 → DHA + 2H + 2I

− N-acetylcysteine 2RSH + I2 → ? +2I

Bisulfite Na2S2O5 + 2I2 + 3H2O − → 2H2SO4 + 4I + 2Na+ + 2H+

+ − Glutathione 2GSH + I2 → GSSG + 2H + 2I

2− 2− − Thiosulfate 2S2O3 + I2 → S4O6 + 2I

Table 2.1. Analytes determined in commercial products Ascorbic acid, also known as vitamin C, is an organic acid found in foods and is also used as a food additive because it is a biologically important compound. Its structure is shown in Table 2.1. Since it is a reducing agent it can undergo a redox reaction with iodine as the oxidizing agent. Iodine is electrogenerated and its reaction with ascorbic acid (AA) to form dehydroascorbic acid (DHA) and iodide is shown in Table 2.1. 2.78 mg of AA standard were

-4 taken and titrated in 3.18 x 10 M HNO3 at pH 3.14; a quantity of 2.73 ± 0.02 mg was found (98.2% recovery, 0.772% RSD, n=3). The number of electrons transferred in a redox reaction is affected by the pH of the titration medium. This experimental variable is rarely considered in most coulometry studies. We have performed coulometric titrations of cysteine, N-acetylcysteine, and glutathione with electrogenerated iodine in a pH range of 2.0 to 9.0. The results of these studies are shown in Table 2.2. The unexpected conclusion is that cysteine and N-acetylcysteine generally involve a 2 electron transfer per mole while glutathione reacts with iodine as expected with a one electron transfer per mole. The glutathione oxidized dimer can be easily explained as the product for the one electron reaction. However, the product of cysteine or N-acetylcysteine is more difficult to formulate. The average number of electrons transferred per moles of analyte that lie between whole numbers, such as 3.5 and 2.6, can possibly be explained by multiple pathways in which these thiols (either RSH or RS-) are oxidized to a radical which can convert by different pathways to a RSSR or RSOH form.21 Based on these results it can be concluded that an acidic pH is best when determining these analytes. Care to minimize the presence of dissolved O2 is - 22 important which under acidic conditions can react with I to form I2. Because of the greater n value, the sensitivity of the titration would be improved at higher pH but Faraday’s law may not hold and a calibration curve would be required. In addition, at a pH greater than 9, I2 will react with 2OH- to form I- and IO- plus water; the hypoiodite can react with a thiolate anion.22

pH Cysteine N-acetylcysteine Glutathione

2.0 2.4 (1.1%) 1.9 (2.1%) 1.0 (3.4%)

4.6 2.2 (0.86%) 2.6 (0.60%) 1.3 (3.0%)

7.0 1.8 (1.8%) 2.7 (0.89%) 1.2 (5.1%)

9.0 3.5 (2.3%) 2.0 (1.2%) 2.1 (4.1%)

Table 2.2. Average number of electrons per mole of analyte (% RSD, n=3).

N-acetylcysteine (NAC) is a thiol used to support liver and lung function. Its structure is shown in Table 2.1 along with its reaction with iodine, where -SH is highlighted as the moiety

28 that participates in the redox reaction and R- represents the rest of the molecule. A 5.07 mg portion of the NAC standard was taken and titrated in 0.03 M H2SO4. A value of 4.80 ± 0.10 mg was found (94.7% recovery, 2.14% RSD). The N-Acetyl-L-Cysteine dietary supplement with a label value of 500 mg NAC per capsule by Jarrow Formulas® was the commercial product analyzed and an amount of 507 ± 19 mg were found (101% recovery, 3.76% RSD, n=3). Bisulfite is a food additive to dried fruits and vegetables and is intended to increase the shelf life of those products. The structure of the metabisulfite ion and its reaction with iodine in the presence of water are shown in Table 2.1.23 A 1.48 mg amount of a metabisulfite standard was taken and titrated in 1 M acetic acid/ 1 M sodium acetate, pH 4.7. A value of 1.42 ± 0.02 mg was found (95.9% recovery, 1.21% RSD, n=3). The commercial product samples selected were dried apricots, Mediterranean Apricots by Mariani®, and dried golden raisins, California Golden Raisins by Sun Maid®. To determine any sample matrix effect, 3.5 g portions of dried fruit were blended with 140 mL of 0.2 M NaOH, acidified to pH 1-2 and sonicated to eliminate any sulfite as SO2. Portions of that solution were then brought to pH 7 before coulometric titration with endpoint times averaging 200 s. The same procedure was done but the acidification step was omitted to determine sulfite plus the matrix; titration times were about 350 s. After subtracting for the matrix, the raisins and apricots were found to range from 715-750 mg bisulfite/kg. The Codex Alimentarius Commission from the Food and Agriculture Organization of the United Nations and the World Health Organization have determined that the maximum level of sulfites in apricots are 2000 mg/kg in the final product and 1500 mg/kg for bleached raisins.24,25 Possibly this procedure could also be adopted for the determination of sulfite in wine.9 c. Commercial product analyses by undergraduate student

An undergraduate student and co-author determined a variety of analyte standards and commercial products. The results of titrations of ascorbic acid, N-acetylcysteine, and glutathione are shown in Table 2.3. Glutathione is an important thiol that has antioxidant properties. The structure of this organic compound and its reaction with iodine are shown in Table 2.1. Based on the one electron transfer found in acidic and neutral pH (Table 2.2), glutathione follows a more standard reaction pathway to the oxidized dimer as compared to the cysteine compounds. The recoveries as compared to the label values for glutathione and NAC (both 500 mg capsule) are good; the high result as compared to the likely label value (90 mg ascorbic acid in 200 mL apple juice) is not unexpected for a natural fruit containing product. Thiosulfate is a sulfur-containing

29 compound that is commonly used to dechlorinate water. Its structure and its redox reaction with iodine are shown in Table 2.1. The results of thiosulfate titrations are shown in Table 2.3. Both commercial products lacked label values for the analyte (although the patent of the Ultra Swim Shampoo indicated 1% thiosulfate by weight), so a standard addition method was used to analyze these commercial products. Standard Commercial Commercial Name Titration media Recovery product product (% RSD, n=3) recovery 99.7% Ascorbic acid H SO , pH 2.0 Apple juice 121% 2 4 (0.885% RSD) N-Acetyl-L- Cysteine dietary 107% N-acetylcysteine H2SO4, pH 2.0 supplement by 89% (4.34% RSD) Jarrow Formulas®

Glutathione Reduced K HPO /KH PO , 99.3% dietary Glutathione 2 4 2 4 100% pH 7.0 (1.54% RSD) supplement by Jarrow Formulas® BettaSafeTM Water 4.27% by Conditioner for weight K HPO /KH PO , 100% Betta Fish Thiosulfate 2 4 2 4 pH 7.0 (0.693% RSD) Ultra Swim Chlorine 0.718% by Removal weight Shampoo Table 2.3. Standards and commercial products determined with the home-built coulometer. IV. Discussion

The coulometry student experiment (see Supporting Information) has been part of our instrumental analysis lab course every semester since Fall 2015. Initially more traditional samples containing ascorbic acid such as fruit drinks and fruit flavored snacks were analyzed. To provide more experiment variation on a semester basis, the determination of sulfur compounds has and will be continued to be implemented. About 60-80 students, mainly junior level

30 chemistry and biochemistry majors, have taken this required one hour lecture, two credit course per year. Student rotation through the experiments was required due to limited instrumentation. The Thursday “lecture” was taught as a guided self study in which the students would work in groups of four to answer key questions about the technique using an on-line textbook as the primary reference. This would give them the necessary fundamentals to do the corresponding experiment the following Tuesday. Students also work in groups of 3 to 4 students to complete the work in a 3.5 hour lab period. Coulometry is taught in conjunction with cyclic voltammetry because both experiments are amperometric in nature (see Student Experiment handout in the Supporting Information). This electrochemistry experiment is one of four long laboratory reports that are written up in a journal format; the other three are fluorescence, atomic emission spectroscopy, and liquid chromatography/flow injection. The other experiments conducted in this course are Beers Law and its limitations, Raman spectroscopy, atomic absorption, mass spectrometry, and gas chromatography. In addition, students work on a two period project which is intended to extend the scope of samples that can be analyzed for a particular instrument. For coulometry, this could be analysis of various dried fruits for bisulfite or the indirect detection of zinc using thiol complexation.26 V. Acknowledgements

This material is based upon work supported by the National Science Foundation under Grant Number 1432466, Principal Investigator S.L. Bretz. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the author(s) and do not necessarily reflect the views of the National Science Foundation. VI. References

(1) Bell, D. A. J. Chem. Educ. 1978, 55 (12), 815.

(2) Tackett, S. L. J. Chem. Educ. 1972, 49 (1), 52–54.

(3) Bertotti, M.; Moreira Vaz, J.; Telles, R. J. Chem. Educ. 1995, 72 (5), 445–447.

(4) Scanlon, C.; Gebeyehu, Z.; Griffin, K.; Dabke, R. B. J. Chem. Educ. 2014, 91 (6), 898– 901.

(5) Rüttinger, H. H.; Spohn, U. Anal. Chim. Acta 1987, 202, 75–84.

(6) Pastor, T. J.; Barek, J. Mikrochim. Acta 1989, 97 (5–6), 407–413.

(7) Ciesielski, W.; Krenc, A. Anal. Lett. 2000, 33 (8), 1545–1554.

(8) Dabke, R. B.; Gebeyehu, Z.; Thor, R. J. Chem. Educ. 2011, 88 (12), 1707–1710.

(9) Lowinsohn, D.; Bertotti, M. J. Chem. Educ. 2002, 79 (1), 103–105.

(10) Grimsrud, E.; Amend, J. J. Chem. Educ. 1979, 56 (2), 131–133.

(11) Evans, D. H. J. Chem. Educ. 1968, 45 (2), 88–90.

(12) Beilby, A. L.; Landowski, C. A. J. Chem. Educ. 1970, 47 (3), 238–239.

(13) Kuntzleman, T. S.; Kenney, J. B.; Hasbrouck, S.; Collins, M. J.; Amend, J. R. J. Chem. Educ. 2011, 88 (11), 1565–1568.

(14) Williams, K. R.; Young, V. Y.; Killian, B. J. J. Chem. Educ. 2011, 88 (3), 315–316.

(15) MeasureNet Technology, Ltd. http://www.measurenet-tech.com/index.html.

(16) Mayo, A. V.; Loegel, T. N.; Bretz, S. L.; Danielson, N. D. J. Chem. Educ. 2013, 90 (4), 500–505.

(17) Lake Shore Cryotronics. Model 121 Programmable Dc Current Source http://www.lakeshore.com/products/DC-Current-Sources/Model-121-DC-Current- Source/Pages/Overview.aspx.

(18) Global Specialties. PB-10 Externally Powered 840 Tie-Point Breadboard http://www.globalspecialties.com/solderless-breadboards/breadboards-mounted/item/82- pb-10.html.

(19) Diefenderfer, A. J. Principles of Electronic Instrumentation, 2nd ed.; W. B. Saunders: Philadelphia, PA, 1979.

(20) Calabrese, V. T.; Khan, A. J. Polym. Sci. Part A Polym. Chem. 1999, 37 (15), 2711–2717.

(21) Uchiyama, S.; Sekioka, N. Electroanalysis 2005, 17 (22), 2052–2056.

(22) Ciesielski, W.; Zakrzewski, R. Chem. Analityczna 2006, 51, 653–677.

(23) Taylor, R. H.; Rotermund, J.; Christian, G. D.; Ruzicka, J. Talanta 1994, 41 (1), 31–38. 32

(24) Codex Alimentarius Commission. CODEX STAN 130-1981 Standard for Dried Apricots; 1981.

(25) Codex Alimentarius Commission. CODEX STAN 67-1981 Standard for Raisins; 1981.

(26) Padilla Mercado, J. B.; Konkolewicz, D.; Bretz, S. L.; Danielson, N. D. Indirect Determination of Zinc by Thiol Complexation and Iodine Coulometric Titration with Photocell Detection. Microchem. J. 2017, 134 DOI: 10.106/j.microc.2017.05.013.

VII. Supporting information

Figure S1. Proper alignment of electrodes, stir bar, and photocell.

Figure S2. Circuit diagram of the current-to-voltage converter.

Figure S3. Close-up picture of the breadboard used to construct the current-to-voltage converter circuit.

Figure S4. Titration plot of ascorbic acid titration using the 1 cm and 2 cm platinum electrodes as the anode and cathode, respectively.

Figure S5. Titration plot of ascorbic acid titration with 2 cm and 1 cm platinum electrodes as anode and cathode.

Figure S6. Ascorbic acid titration plot using the 1 cm Pt cylinder as the anode and the medium counter electrode (Platinum Inlay Cat. No. 476060 from Corning).

Figure S7. N-acetylcysteine titration plot using a dialysis membrane to cover the 1 cm cylinder cathode.

Chapter 2. Supporting Information continued

Iodine coulometry with on-line photocell detection for a multifunctional chemical analysis (MCA) system

Instructor Notes for Coulometry Materials needed to build instrument  Current source  Leads with alligator clips (6)  Twin platinum electrodes (1 and 2 cm concentric)  150 mL beaker  Aluminum foil  Magnetic stir bar  Magnetic stirrer  Photodiode  Circuit o Global Specialties PB-10 externally powered 840 tie-point breadboard; ±12 V o +5 V variable resistor set at +2.15 V o 741 operational amplifier from RadioShack o Resistor 3.3 kΩ o Connecting wires o Power supply cord o MeasureNet connector Commercial samples to be analyzed  Apple juice  N-acetyl-cysteine dietary supplement: 500 mg NAC per capsule by Jarrow Formulas ®  Glutathione Reduced dietary supplement: 500 mg per capsule by Jarrow Formulas ®  BettaSafeTM water conditioner  Ultra Swim Chlorine Removal Shampoo Chemicals and reagents  Sulfuric acid (CAS 7664-93-9)  Nitric acid (CAS 7697-37-2)

 Glacial acetic acid (CAS 64-19-7)  Sodium acetate (CAS 127-09-3)  Sodium phosphate (CAS 7558-79-4)  3% starch solution  Potassium iodide  L-ascorbic acid (CAS 50-81-7)  N-acetyl-cysteine (CAS 616-91-1)  Sodium metabisulfite (CAS 7681-57-4)  Glutathione (CAS 70-18-8)  Sodium thiosulfate pentahydrate (CAS 10102-17-7) Lab ware  Balance  Deionized water  150 mL beakers  100 mL volumetric flasks  100 mL graduated cylinder  Pipets of various sizes: 1.00 mL, 5.00 mL Building the instrument  Figure 1 in the manuscript shows a possible set up for all the instrument parts: current source, electrodes, beaker, photodiode, magnetic stirrer, and circuit.  Figure S2 shows the circuit diagram of the current-to-voltage converter.  Figure S3 provides a close-up picture of how the breadboard was used to create the circuit. MeasureNet parameters The following are the MeasureNet parameters that work best for all coulometric titrations described in this experiment. y-maximum = 0 y-minimum = -2500 x-maximum = 3000 x-minimum = 0

s rate = 2 Sample preparation  The following solutions have been used as buffers for titrations: o 0.03 M sulfuric acid for ascorbic acid and N-acetyl-cysteine titrations o pH 7 phosphate buffer for glutathione and thiosulfate titrations o Borate buffer for bisulfite titrations  Standard solutions can be prepared in 100 mL volumetric flasks.  Aliquots of juice sample can be taken directly from the container.  N-acetylcysteine and glutathione capsules are prepared the same way. Capsules are emptied in a beaker with 50 mL of buffer solution and sonicated for 15 minutes. This solution is quantitatively transferred to a 100-mL volumetric flask and becomes the stock. In case the dietary supplements are tablets, these should be pulverized in a mortar and pestle before being sonicated in buffer.  Aliquots of BettaSafe and Ultra Swim shampoo can be used to make stock solutions in 100 mL volumetric flasks. The Ultra Swim shampoo stock should be dilute to avoid interferences from surfactants.

Titration procedure A clean 150 mL beaker is covered on the outside with aluminum foil. The following are added to the beaker:  Magnetic stir bar  1 mL of 3% starch  0.5 g of KI  The analyte aliquot to be titrated  50 mL of buffer  Water is added to reach a total volume of 130 mL

The beaker is placed on top of the photodiode and the stir rate is set to 240 rpm. The electrodes are lowered into the beaker; the leads are connected and the MeasureNet is started simultaneously. Students should avoid to walk near the instrument during the titration to avoid blocking ambient light from the photodiode. Classroom lights should be kept at the same

39 intensity through all measurements. The titration is stopped when the solution turns dark blue; at this point the MeasureNet shows an S-curve plot. Data analysis The MeasureNet file can be transferred to an Excel sheet. The area in which the uptick in baseline occurs should be noted. The endpoint is within these points. To find the actual endpoint the 20 voltage data points before this visual uptick should be highlighted. The following equation is used to calculate the endpoint based on these voltage values. Endpoint voltage (mV) = 3σ (mV) + voltage of last point(mV) The standard deviation of these values is multiplied by three; the resulting value is added to the last point of the selected range. The time corresponding to this calculated endpoint voltage is the endpoint of the titration. A more involved experiment is the determination of bisulfite in dried fruits such as Mediterranean Apricots by Mariani® and California golden raisins (Sun Maid®).

1. Fruit sample matrix only (bisulfite decomposed at low pH): -Blend 3.5 grams fruit in 140 mL 0.2 M NaOH. For 3 minutes at speed #9. -Add acid to drop pH to 1-2. -Sonicate for 15 minutes to evolve SO2. -Take an aliquot of 20 mL, adjust the pH using a NaOH pellet, add 120 mL phosphate buffer (pH 7, 0.1M/0.1M), 0.5 g KI, 1 mL 3% starch and titrate.

2. Fruit sample matrix plus bisulfite: -Blend 3.5 grams fruit in 140 mL 0.2 M NaOH. For 3 minutes at speed #9. -Sonicate for 15 minutes. -Take an aliquot of 10 mL, spike one with a known amount of metabisulfite and do another one without a spike, add 120 mL phosphate buffer (pH 7, 0.1M/0.1M), 0.5 g KI, 1 mL 3% starch and titrate using dry ice and Saran wrap.

The difference between titration endpoint times (Part 2 – Part 1), corrected for the different aliquots taken, corresponds to the time required to titrate the bisulfite.

Chapter 2. Supporting Information continued

Answers to Student Coulometry Experiment Questions 1. What causes the formation of gas bubbles from the cathode during the titration of ascorbic

acid with I2 and give the reaction? + - Reduction of 2H + 2e ↔ H2(g) 2. To ensure better selectivity and accuracy of a coulometric titration, the use of dry ice might be justified to displace dissolved oxygen from the solution before starting the titration. Calculate the overall Eo and whether it is positive or negative for the chemical

reaction involving ascorbic acid and O2. See Introduction, p. 1. Also, write the balanced

reaction involving ascorbic acid and O2.

+ - o O2 + 2H +2e → H2O2 E = 0.682 V. + O2 + ASCORBIC ACID ↔ H2O2 + DEHYROASCORBIC ACID Eo = 0.682 – 0.39 = +0.29 V. Reaction proceeds to the right o 3. Explain whether arsenic(III) could be titrated by coulometry with I2 (E = 0.622 V) indicating the chemical reaction involving the analyte and titrant. What is the overall Eo for the reaction? The standard reduction reaction and potential:

As5+ + 2e- → As3+ Eo = 0.56 V. 3+ 5+ - o As + I2 ↔ As + 2I E = O.622-0.56= +0.06 V. Reaction proceeds to the right 2+ 4. Could the determination of Fe be done by coulometry using I2? The standard reduction reaction and potential: Fe3+ + 1e- → Fe2+ Eo = 0.770 V.

2+ 3+ - o Fe + I2 ↔ Fe + 2I E = 0.622-0.770 = -0.15 V. Reaction proceeds to the left. o 5. If electrogenerated Br2 was used as the titrant instead of I2 (E = 0.622 V), the results for ascorbic acid in various real samples would probably be high as compared to the expected values? Why? Hint: Is a positive Eo for the overall reaction involving some interferent - - o and titrant more likely for I2 or Br2? Br2(aq) + 2e → 2Br E = 1.098 V.

Interfering reducing agents with Eo for the reduction half reaction below 1.098 V will o all react with Br2, potentially more than those with E values below 0.622 V for I2 as the titrant.

6. What is the Karl Fischer titration and describe the chemistry (do a google search)?

The Karl Fischer method is used to determine water using iodine coulometry as shown below. B = organic base such as pyridine. + - B-I2 +B-SO2 + B + H2O ↔ 2BH I + BSO3 + - BSO3 + ROH ↔ BH ROSO3

Chapter 2. Supporting Information continued

CHM 375 Spring 2017 Experiment 10. Voltammetry and Coulometry

Pre-lab question: In Part II Coulometry, what oxidizing agent is reacting with what sample? Part I. Cyclic Voltammetry Background Reading: Harvey, Chapter 11; Section D. Cyclic voltammetry is a technique used to characterize various electrochemical parameters for a redox process. It can be used to determine concentration, the standard potential, the number of electrons involved, the diffusion coefficient, the reversibility, the extent of diffusion control and the kinetics of the reaction. Thus, it is a diagnostic tool used to study reactions in addition to being a semi-quantitative method. In cyclic voltammetry, a linearly increasing potential is applied to the electrodes in a cell as a function of time. When a potential sufficient to cause an oxidation or reduction of the analyte to occur is reached, a current arises. Diffusion of the redox active species to the electrode surface limits the current. The diffusion rate is in turn determined by the concentration of the species. Thus, the current is proportional to concentration. The resulting current vs. voltage plot is called a voltammogram. The following values can be derived from the voltammogram (see figures below).

When the voltage is scanned from negative to positive values, Epa = anodic peak potential and ipa = anodic peak current (negative by convention) from the profile. Epa is always greater than Epc. When the voltage is scanned from positive to negative values, Epc = cathodic peak potential and ipc = cathodic peak current (positive by convention) from the profile. Epa is always greater than Epc.

The fundamental equation relating the current measured to other parameters is the 1/2 Randles-Sevcik equation: ip = 0.4463nFAC (nFvD/RT)

ip = peak current in amperes or coul/s n = no. of electrons in the redox process F=Faraday constant in coul/mol A = electrode area, cm2 C = concentration, mol/cm3 v = scan rate, V/s D = diffusion coefficient, cm2/s R= gas constant in JK-1mol-1 T= temperature in K o 5 3/2 1/2 1/2 At 25 C, this equation simplifies to ip = 2.69 x 10 n AD Cv . This equation can be used in several ways. Note that for a given process n and D are constants. For a given electrode A is fixed. Thus, the variables are i, C, and v. If the scan rate is held constant for a series of experiments, the equation can be rewritten I=K C where K = 2.69 x 105n3/2AD1/2v1/2 For quantitative analysis, the scan rate is held constant and the current measured for a series of standard solutions can be used to construct a calibration curve. From the equation of the experimentally determined best-fit line and the current measured for an unknown the concentration of the unknown can be calculated. If the current resulting from measurements on one solution of known concentration is recorded at several scan rates, the Randles-Sevcik equation can be rewritten:

I=K v12 where K = 2.69 x 105n3/2AD1/2C If a plot of i vs v1/2 is linear, this shows that the redox process is diffusion controlled. For a reversible redox couple, both the oxidation and reduction steps can be observed and the average potential E = formal potential for the reversible couple. The formal potential E’ can be considered to be equal to the standard potential Eo if the activity coefficients for the reduced and oxidized species are approximated as 1 (dilute solution). E +E E= pa pc 2 It can be shown that

0.059 ΔE =E ̶ E= p pa pc n

Rearranging this equation, ΔEp, the peak potential separation, can be used to estimate the number of electrons involved in the redox process. Note: Surface films on the working electrode often give very large ΔEp values resulting in inaccurate (low) values for n. For a simple well-behaved reaction (no side reactions or adsorption to the electrode), the current for the reverse scan peak should be approximately the same as for the peak on the forward scan. This means the ratio value can be rounded to 1. i pa =1 ipc A. Ferricyanide – ferrocyanide redox couple In this experiment, you will investigate the following reaction: 3ˉ 4ˉ Fe(CN)6 + eˉ  Fe(CN)6 The formal potential and the peak current ratio will be determined, and the effect of diffusion control will be investigated. Procedure (be sure all four students are involved)

The 0.5 M KNO3 (MW 101.1) supporting electrolyte (250 mL) needs to be prepared by and for two pairs of students. Prepare a 50-mL solution of 10 mM ferricyanide (MW 329.26) and dilute to the mark with 0.5 M KNO3. The supporting electrolyte ensures that the analyte migrates to the electrode surface due to diffusion and is not attracted or repelled as a function of the applied potential on the electrode. The high concentration of KNO3 masks the analyte charge. Be sure to clean the Pt working electrode before use with Emory cloth and then assemble the cell with a platinum working electrode, an Ag/AgCl reference electrode, and a Pt auxiliary electrode. Fill with about 10 mL 0.5 M KNO3, the background electrolyte. Deoxygenate by bubbling with N2 for about a minute. Cover the solution to prevent oxygen from re-entering the cell during the experiment. Connect the electrodes and set the scan parameters. Set the initial potential at 0.8 V. Set the limiting potentials at 0.8V and -0.2V and the final potential equal to the initial potential. Set the scan rate at 100 mV/sec, the initial scan direction as negative, and the current sensitivity as recommended by the instructor. The initial scan direction is from positive to negative potential because ferricyanide, our reactant, can only undergo reduction first.

Record a background voltammogram of the supporting electrolyte to check that it is close to zero. NOTE: Remember that cyclic voltammetry is a diffusion controlled process and is done on a quiescent solution. But the solution must be stirred between scans in order to restore initial conditions at the electrode surface.

Rinse and refill the cell with 10 mL of 10 mM K3Fe(CN)6 already prepared in 0.5 M

KNO3. Degas for 3 minutes. Record a CV using the same scan parameters. Use the same solution to investigate the effect of scan rate on peak height. Record CVs (scanning negatively first) at the following scan rates: 20, 50, 100 and 200 mV/sec. Check to make sure the current is increasing as a function of scan rate. Save these voltammograms on the computer. You will need to use the peak current values. Repeat the 100 mV/sec scan but now let the solution be stirred during the voltage sweep. Save this voltammogram. Calculations (Results) 1. Determine the standard potential Eo for the iron redox couple from E’ (see p. 2). You o may use any one CV to get this data. E’ = E – Ereference. The reference electrode potential is 0.226 v. Compare to the tabulated Eo value of 0.430 for the ferricyanide/ferrocyanide reduction at pH =7.

2. The instrument may not give reliable ipa and ipc current values in the output table. To

check, find the maximum current ipa and ipc values at a constant voltage from your

voltammograms using your data files. Plot ipc (the peak current for the cathodic process) vs. v0.5 (the square root of scan rate). Is the reduction process diffusion controlled?

ipa 3ˉ 3. Calculate for the 10 mM Fe(CN)6 solution scanned at 100 mV/S. Does the reaction ipc appear to be well-behaved (no electrode adsorption or side reactions)? See previous page for background. 4. Given the following data from a voltammogram, calculate D, the diffusion coefficient for 3ˉ 4 2 Fe(CN)6 . Be sure to include units. ip = 1.086 x 10ˉ amperes, A = 0.0707 cm , C = 1.00 x 10ˉ5 moles/cm3, and v = 0.100 V/s. A reasonable value is on the order of 10-5 to 10-6 cm2/sec.

Questions (as part of Discussion) 3 1. If the reduction of Fe(CN)6 ˉ was completely irreversible, what would a typical cyclic voltammogram look like? Draw a labeled sketch and explain its shape. 2. Explain why the stirred cyclic voltammogram looks different than the unstirred one. Why is the signal noisy and consider what happens to the product when formed at the electrode? Part B. Determination of ascorbic acid in fruit juice Ascorbic acid can be easily oxidized to dehydroascorbic acid as shown below and this peak current is proportional to the concentration of ascorbic acid. The corresponding reduction back to ascorbic acid is likely not favorable due to the instability of dehydroascorbic acid or slow kinetics. Ascorbic acid → Dehydroascorbic acid + 2 H+ + 2e- To determine ascorbic acid in fruit juice, the standard addition method should be used because the proportionality constant (K’) relating current (I) and concentration (C) will likely be different for the standard ascorbic acid solution as compared to a juice aliquot (portion) due to the sample matrix. See the background information given in Experiment 5 (atomic absorption) for standard addition.

I. Procedure (each pair of students has their own unknown) 1. For two pairs of students, weigh out accurately about 0.11xx g ascorbic acid (MW = 176.12) to prepare a 25 mM standard solution in a 25-mL volumetric flask. Dilute with 0.5 M KNO3. 2. Clean and sonicate the Pt electrode for a few min before each sample measurement. Pipet 5 mL of the unknown and 5 mL of the 0.5 M KNO3 into the assembled cell and purge with N2 as before. Scan the voltammograms as oxidation from initial -0.2 v to 0.8 v final. 3. Run voltammograms for the following solutions 1-4 as described in the table. 4. Check that the oxidation current (at about 0.6 v) is increasing proportional to ascorbic acid concentration. Save all voltammograms.

Solution # Volume unknown Volume Standard Volume KNO3 (mL) (drink) mL (ascorbic acid) mL 1 5 1 4 2 5 2 3 3 5 3 2 4 5 4 1

Calculations

1. The instrument may not tabulate reliable ipa and ipc current values. Find the maximum ipa current at a constant voltage (likely near 0.6 v) from your voltammograms using your data files. Plot the |oxidation current| (y axis) versus mg/L ascorbic acid. Determine the trendline and correlation coefficient. Set y =0 and determine the x intercept as before in Experiment 5. Determine the mg/L in the original drink sample (remember there is a dilution factor of 2). 2. Compare to the expected label value of 100% recommended daily allowance (RDA) ascorbic acid (75 mg women, 90 mg men) per serving (container size). The container size of your unknown is probably 177 mL (use the actual value). Discuss the fact that there is real juice in your drink sample but the manufacturer likely spikes the juice with ascorbic acid to ensure there is 100% RDA. 3. From an estimate of the noise fluctuation from the background voltammogram taken from part A involving ferricyanide and the slope of the standard addition plot, determine the detection limit of ascorbic acid and give units. Because this is standard addition, do not include the intercept in this calculation.

Questions (as part of Discussion) 1. Determine whether the ascorbic acid voltammogram is reversible and explain your reasoning. 2. How does your detection limit for ascorbic acid compare to that of 9 x 10-5 M reported by Pisoschi AM, Danet AF, Kalinowski, S J Automated Methods Management Chemistry (2008) Article ID 937651?

Introduction When explaining cyclic voltammetry, be sure to describe Figure 11.47a and 11.47b (p. 745) in Chapter 11 of Harvey. Google cyclic voltammetry images and include an analogous figure to Figure 11.47b as part of your report. In your report, in addition to the required plots, you should have a representative voltammogram for 1. Unstirred ferricyanide/ferrocyanide, 2. Stirred ferricyanide/ferrocyanide, and 3. Ascorbic acid standard addition. Be sure to check the long report guidelines under Resources on Canvas.

Part II. Coulometric Generation of Iodine for the Titration of Vitamin C or a Thiol Vitamin C, or ascorbic acid, is found in many fruits and vegetables, particularly in citrus fruit juices. It is also one of the more popular food additives of our modern food-processing technology. Ascorbic acid, which is a moderately strong organic acid, is a rather good reducing agent, so it is not particularly stable in solution, being capable of undergoing air oxidation.

Thiols (RSH) such as glutathione (a tripeptide of glutamic acid, cysteine, and glycine) and N-acetylcysteine are often prescribed as supplements and are touted as antioxidants. Both can cause neutralization of free radicals and reactive oxygen compounds as well as maintaining Vitamin C in its reduced form. Glutathione is specifically involved in iron metabolism. N- acetylcysteine is the antidote to acetaminophen poisoning by raising the dropped glutathione level in the liver and aiding in acetaminophen poisoning.

Iodine is a mild oxidizing agent that can determine quantities of ascorbic acid. However, because of the volatility of the iodine molecule, it is difficult to work with standard solutions of the reagent. One means of getting around this problem is to generate the iodine electrically by the oxidation of iodide ion. The electrical generation is carried out in a quantitative way by using a constant current and timing the period during which the current flow takes place. Since current (in amperes) is a measure of coulombs per second, we can then calculate the number of coulombs passed. When applied to a titration, the technique is called a coulometric titration, and it is performed in the following way. An electronic constant current circuit is set up to generate a constant current flow between two platinum electrodes in a solution. At the anode, in the presence of iodide, we get the half-reaction 49

- - I  1/2I2 + e For each electron (1.60 x 10-19 coulombs) passed, we get one iodide ion oxidized. For each faraday (F = 96,487 coulombs), one-half mole of iodine is produced. At the cathode in

+ - aqueous solution, typically the reaction will be H + e  1/2H2(g). The production of hydrogen gas is not important for our consideration because it is not involved in the rest of the chemistry in the cell.

The standard reduction reactions and potentials for the titration reaction are: - - o I2 + 2e  2I E = 0.622 V Dehydroascorbic acid + 2H+ + 2e-  ascorbic acid Eo = 0.39 V The iodine in solution will react spontaneously with a reducing agent such as ascorbic acid according to the equation below to give dehydroascorbic acid and iodide because the Eo for this reaction (0.622 – 0.39 = + 0.23 V) is positive in sign.

HO-C == C-OH O= C ____C =O | | | |

- + HOCH2- CH(OH)-CH C =O + I2  HOCH2-CH(OH)CH C =O + 2I + 2H \O/ \ O/

The iodine in solution can alternatively react spontaneously with a reducing agent such as a thiol (N-acetylcysteine or glutathione) according to the equation below because the Eo for this reaction (0.622 – 0.22 = + 0.40 V) is also positive in sign. - RSH + I2  RS-SR + 2I This titration reaction must be done at low pH to ensure there is sufficient H+ to be reduced at the cathode. The end point of the titration can be determined by adding a quantity of starch solution to the cell. Starch combines with iodine to form a complex that has an intense blue color in aqueous solution. Since free iodine will not be present until all the ascorbic acid or bisulfite is oxidized, the solution will not turn blue until the end point. By timing how long (t) the current I (known from the constant current source) must be generated to get the solution to turn blue, we can calculate I(t) which is the number of coulombs passed, and from this value, we can calculate the grams of ascorbic acid or bisulfite present in solution.

W = [(I) (t) (Molar mass)] / nF where W = grams substance oxidized, I = current in amperes (coulomb/sec), t = time in sec, and F = Faraday (96,487 coulombs/equivalent). The molar mass is 176 for ascorbic acid and the number of electrons lost or gained per molecule of interest during the electrochemical process (n) is 2. The molar mass for glutathione is 307.3 and that for N-acetylcysteine is 163.2; the number of electrons transferred is 2.

II. Procedure (be sure all four students are involved)

A diagram of a constant current coulometry instrument using electrogenerated I2 for the determination of a selected reducing agent is shown below. In our case, the instrument is a constant current source connected to Pt electrodes immersed in a beaker of iodide and the ascorbic acid sample solution at low pH (dilute acid), on top of a photodiode sensor on a magnetic stirrer. The signal due to the decrease in ambient light hitting the photodiode positioned below the beaker is measured using MeasureNet.

For lemonade:

3.4 g of lemonade mix containing 0.009 g ascorbic acid in 25 mL flask, filled to volume with dilute acid. 7.5 mL of this lemonade solution (= 0.0027 g ascorbic acid = 300 s equivalence

51 point) was transferred to the foil-covered beaker, along with 0.5 g KI and roughly 3 mL starch, and was then filled with dilute acid to cover the electrodes. The titration was then run.

For fruit snack:

One pack of 12 fruit snacks is 90 mg ascorbic acid. Two gummies were sonicated in dilute acid, and this concentrated gummy solution was transferred to a 100 mL volumetric and filled to volume with acid, so we have a stock dilute gummy solution containing 0.015 g ascorbic acid. 18 mL of this stock (= 0.0027 g ascorbic acid = 300 s equivalence point) was transferred to the foil- covered beaker along with 0.5 g KI and roughly 3 mL starch, and was then filled with dilute acid to cover the electrodes. The titration was then run.

For glutathione or N-acetylcysteine tablet: Weigh tablet. Crush to a powder in a mortar and pestle and take a portion for weighing. Add this to the coulometry beaker with dilute sulfuric acid and sonicate before titrating.

1. To a 100 mL volumetric flask, add about 0.5 g KI, and about 100 mL of 0.05 M sulfuric acid. A drink or fruit snack unknown will be provided; add an appropriate volume of sample solution for a titration time of 5 min assuming a current of 0.010 amps. Add about 1 mL of 3% starch indicator and dilute the solution close to the 140-mL mark with RO water. 2. Rinse the electrodes with DI water. Immerse electrodes completely and without touching in the solution and turn on the magnetic stirrer (slow speed). Make sure the electrode leads are not touching. Make sure the alligator clips are firmly touching the electrodes. Attach the photodiode operational amplifier circuit to the electrodes and to the MeasureNet station; see the instructor for help. The settings for MeasureNet should be 0 for maximum and -1500 for minimum y values; x value is 3000 s. Start the electrolysis current (record the setting, probably 10 mA or 0.010 amps) and take data using MeasureNet until a uniform blue color of the solution has persisted. Be save the file on MeasureNet and turn off the current. Repeat the measurement with at least two additional portions of prepared solution. The midpoint time between the plateau regions can be considered the endpoint time.

Calculations for ascorbic acid 1. Calculate the amount of ascorbic acid in the juice sample (in mg) per serving size. Make sure to include appropriate dilution factors in your calculation to get the actual amount of ascorbic acid in the mL or g of juice sample taken. Then calculate the g ascorbic acid in a serving size. Recommended daily allowance (RDA) (100%) of ascorbic acid is 75 mg women or 90 mg men. 2. Compare to the expected label value by calculating the relative error. A low recovery may be due to the age of the sample; how do you expect ascorbic acid degraded? See question 2 below.

Calculations for thiol 1. Calculate the amount of thiol (glutathione or N-acetylcysteine) in mg in the tablet. Make sure to include appropriate dilution factors in your calculation to get the actual amount of thiol in the original tablet. 2. Compare to the expected label value by calculating the relative error. A low recovery may be due to the age of the sample; how do you expect thiol degraded? See question 2 below.

Part of the Introduction (the sentence below and Questions 1 and 2) Describe briefly what coulometry is and indicate at least one advantage of a coulometric titration over a regular buret titration. 1. What causes the formation of gas bubbles from the cathode during the titration of ascorbic acid with I2 and give the reaction? 2. To ensure better selectivity and accuracy of a coulometric titration, the use of dry ice might be justified to displace dissolved oxygen from the solution before starting the titration. Calculate the overall Eo and whether it is positive or negative for the chemical reaction involving ascorbic acid and O2. See Introduction, p. 1. Also, write the balanced reaction involving ascorbic acid and O2. + - o O2 + 2H +2e → H2O2 E = 0.682 v.

Part of the Discussion (Questions 3-6) o 3. Explain whether arsenic(III) could be titrated by coulometry with I2 (E = 0.622v) indicating the chemical reaction involving the analyte and titrant. What is the overall Eo for the reaction? The standard reduction reaction and potential: As5+ + 2e- → As3+ Eo = 0.56 v.

2+ 4. Could the determination of Fe be done by coulometry using I2? The standard reduction reaction and potential: Fe3+ + 1e- → Fe2+ Eo = 0.770 v. o 5. If electrogenerated Br2 was used as the titrant instead of I2 (E = 0.622 v), the results for ascorbic acid in various real samples would probably be high as compared to the expected values? Why? Hint: Is a positive Eo for the overall reaction involving some interferent and - - o titrant more likely for I2 or Br2? Br2(aq) + 2e → 2Br E = 1.098 v. 6. What is the Karl Fischer titration and describe the chemistry (do a google search)?

Long Report Guidelines Abstract. Approximately a 250 word abstract which summarizes the objectives, methods, pertinent results, and main conclusions. Introduction. See previous notes in this experiment handout for specific content. See Long Report Guidelines in your lab manual or Canvas. I will provide at least one related literature reference (ascorbic acid) on Canvas that you need to describe in the Introduction using at least three sentences. You need to find one analogous recent (preferably last 5 years) research paper (not textbook chapter or review paper) pertaining to analytical chemistry involving cyclic voltammetry or coulometry using Science Citation Index or Google Scholar to summarize in the Introduction. This paper should be summarized using at least three sentences in your Introduction. Copy the first page of the paper and attach this to your report. Experimental. See Long Report Guidelines in your lab manual or Canvas. Results and Discussion. See previous notes in this experimental handout for specific content. See Long Report Guidelines in your lab manual or Canvas. Conclusion. See Long Report Guidelines in your lab manual or Canvas. References. See Long Report Guidelines in your lab manual or Canvas. Attach: 1. Duplicate lab notebook data pages and calculations of each type. 2. First page of related reference paper described in the Introduction. Additional help with formal lab reports (especially your first one in this course) can be gained from http://labwrite.ncsu.edu (this website is a project funded by the National Science Foundation that provides extensive assistance for writing formal laboratory reports). LabWrite

54 contains a great deal of useful information and guidance, but you cannot absorb it all in one night. Start working on your lab report as soon as practical after the data have been collected and the calculations have been performed. If you try to write the report in one night, it will very likely be poorly organized and not well thought-out.

The PreLab section of LabWrite is a useful tool to help you organize your thoughts about the lab report. It asks questions about concepts, objectives, purpose and hypotheses that you should answer before moving on to the formal report. The answers to these questions will guide you in writing the components of a formal report.

The PostLab section of LabWrite should be the most useful part to you in organizing your report. You can use the site in tutor mode, or in self-guide mode. In tutor mode, you write the report within LabWrite and then paste it into your own word processing program. In self-guide mode, you write the report directly into your own word processing program with help from LabWrite.

The lab report does not need to be written in the actual order: Abstract, Introduction, Methods, Results, Discussion, Conclusions, and References. A recommended order follows.

When writing formal reports, start with the methods section. Write up the method that describes the experiment you performed. Your notebook entries will provide the basis for writing this section, but there are formatting differences between what you write in a notebook, and what you would write for the methods section of a formal report or a journal article. The methods section must be written in third person passive voice (“The sample was dried in the oven”, not “I (or we) dried the sample in the oven”). This is also the preferred format for the other sections of the report. First or second person (“I” or “we” or “you”) may not be used anywhere in the report except the conclusion section. No results are presented in the methods section.

The results section can be written next. In this section the main findings of your experiments are reported with tables, graphs or other figures (visuals) that summarize those findings. Many beginning students make the mistake of cramming the results section with tables and graphs with no explanation of what the tables and graphs actually show. Each table, graph, or other figure should be captioned and accompanied by a paragraph that describes and summarizes the information presented in the visual.

Continue on to write an introduction, discussion, and conclusion section. Note that the abstract is written last, as is the title (you can’t summarize what you haven’t written). Keep in mind that the work you turn in is the report, not the entries in LabWrite (this site helps you with the report, but you are ultimately responsible for writing an accurate and readable report). Avoid redundant statements and edit the report to make a document that is easy to read. The important components of each of the sections of a formal report are summarized below:

 Title: The title should capture the most important component(s) of the report. It should be short, typically less than ten words. It should not be written as a complete sentence, and it should not be identical to the title of experiment in the lab manual. It should be descriptive.

 Abstract: Concisely (<300 words) summarize the objectives, the methods, the pertinent results (at least one quantitative example), and main conclusions of the experiment. The reader should be able to understand the major points of the experiment from reading the abstract. This is not a short introduction. The abstract may be written in the present tense.

 Introduction: The structure of the Introduction can be thought of as an inverted triangle - the broadest part at the top representing the most general information and focusing down to the specific problem you studied. Organize the information to present the more general aspects of the topic early in the Introduction (such as textbook fundamentals of the technique), then narrow toward the more specific topical information that provides context (two related literature papers), finally arriving at your statement of purpose and rationale. You may cite your textbook Harvey and the references I have provided on Canvas for that experiment. I will provide at least one related literature reference on Canvas that you need to describe (using at least three specific sentences) in the Introduction. You need to find one analogous related reference using Science Citation Index or Google Scholar to summarize in (using at least three specific sentences) in the Introduction.

 Methods: This section, written in third person, passive voice, is based on your notebook and describes the experiment as performed by the author. It is not written as a set of bullet points or recipe format (no sentence subject), and it should not be identical to the

procedure summaries written in this manual for the experiments. You should avoid repeatedly using a single sentence to relate a single action; this results in very lengthy, wordy passages. A related sequence of actions can be combined into one sentence to improve clarity and readability. Indicate exactly how the experiment was done (it may be slightly different than what was in the lab handout). Be sure instruments employed and their relevant parameters are described. Results of experiments are not presented in this section.

 Results: This section reports both qualitative and quantitative data derived from the experiment presented in a concise, readable and understandable format including appropriate use of tables and graphs. This section should be predominately text, but tables, graphs, and other figures that provide an effective presentation of data should be used. When it is stated to plot some parameter versus another parameter, this always means y axis versus x axis. Do not mix this up. Be sure the plots are labelled correctly. Rules for significant figures, use of proper units for all numbers, etc. must be followed. Any numerical error analyses (standard deviations, Q-tests, t-tests, etc.) should be presented here. Calculations that you may have performed to generate your results should be in your notebook pages (or Excel® spreadsheets), but the equations that you used to generate any reported numbers that are derived from your raw data should be provided in this section.

Present and discuss your results in order in prose form, as in a journal article. Use subheadings. Number any accompanying tables or figures (graphs) consecutively and refer to them by number within the body of this section. Each table should have a title and each figure should have a legend (short description). All numbers (in the text, tables, or graphs) should have the correct number of significant figures and SI units. Significant figures are usually dependent on the precision of the glassware, no better than 1 part per 500. Therefore, three sig. figs. expressed in scientific notation (ie. 2.54 x 10-3 M) is usually appropriate.

 Discussion: This section summarizes and interprets your results and presents your main conclusions with justifications for the conclusions. If the experiment was driven by a hypothesis you should restate that hypothesis at the beginning of this section, and your

discussion should be organized around whether your results support or do not support the hypothesis. If the experiment was driven by some other purpose, the relationship between that purpose and your results should be discussed. A discussion of the error analysis should be part of this section. Experimental difficulties that may have affected results and that are not covered by the numerical error analysis should be brought up. Comparison of your results to literature results may be a part of this section, depending on the experiment. Possible improvements in experimental design may be appropriate for this section also.

 Conclusion: This should be a short (one or two paragraph) section that summarizes what you have learned as a result of performing this experiment. Describe the significance of your results in terms of what you learned. Describe the significance of the experiment to the field of study. What future specific work could be done to extend the scope of the experiment?

 References: This section contains citations for all journal articles, books, etc. used in putting the report together. To as great an extent as possible, primary sources should be cited. This means that websites and encyclopedias (including Wikipedia) should not be cited. A consistent citation style should be used for all sources: authors, title, journal name, volume, (year), inclusive pages. Example: J. S. Fritz, Z. Yan, P. R. Haddad, Modification of ion chromatographic separations by ionic and nonionic surfactants, J Chromatogr. A 997 (2003) 21-31.

For more complete presentations on all of these sections, refer to the LabWrite website. Additional guidance is provided in this manual in the chapters that describe the experiments requiring partial or complete formal reports. The Howe Writing Center (http://www.units.muohio.edu/writingcenter/) can assist you with your writing skills.

Chapter 3. Indirect determination of zinc by thiol complexation and iodine coulometric titration with photocell detection Jeralyne B. Padilla Mercado, Dominik Konkolewicz, Stacey Lowery Bretz, Neil D. Danielson* Department of Chemistry & Biochemistry, Miami University, 651 E. High Street, Oxford, OH 45056, USA *Corresponding author. E-mail address: [email protected] https://doi.org/10.1016/j.microc.2017.05.013 Abstract We have extended the capabilities of iodine coulometry for the determination of ionic zinc through its complexation with cysteine and subsequent titration of the thiol. The titration endpoint times of cysteine with zinc are proportionally longer as compared to cysteine itself. A homebuilt constant-current coulometer is employed for this method development. The instrument is composed of a constant-current source connected to nonisolated platinum electrodes contained in a 150-mL beaker (the coulometric cell) that sits on a photocell mounted in a Styrofoam lid on top of a magnetic stirrer. Photometric endpoint detection is facilitated by connecting the photocell detector to an operational amplifier circuit; its output, recorded on a multifunctional chemical analysis system, resembled a sigmoidal curve. The titration conditions are optimized in terms of pH, temperature, stirring rate, and cysteine-to-zinc concentration ratios. Calibration curves generated as time versus increasing zinc concentrations in the presence of a constant cysteine concentration are used to determine zinc content in dietary supplements. The limit of detection is 5.5 × 10-5 M Zn and the limit of quantification is 6.2 × 10-5 M Zn. Percent relative standard deviations are at 4%. The recoveries of zinc from liquid zinc, enhanced zinc lozenges, and zinc tablets generally range from 80 to 120%. Titration of ascorbic acid and then zinc in the same solution is successfully performed with recoveries greater than 95% for both analytes. A comparison of the coulometry Zn data to atomic absorption results is also made for most samples. Keywords: Iodine; Coulometry; Cysteine; Zinc; Chelation 1. Introduction

Zinc is a biologically important metal ion given its presence in proteins with structural, catalytic, and regulatory roles [1]. It is necessary to maintain healthy levels of zinc in the human body by proper intake through foods and dietary supplements. These supplements are regulated

59 by the Food and Drug Administration (FDA) under the Dietary Supplement Health and Education Act of 1994 [2]. Under this act the FDA requires dietary supplement manufacturers to include specific information such as claims, ingredients, intended use, and safety information. All complaints and side effects as a result of dietary supplement use can be reported to the FDA only by customers [3]. Coulometry can be used to determine analytes directly and indirectly. For example, ascorbic acid is often determined by using electrogenerated iodine. When the two-electron transfer redox reaction between the iodine and the ascorbic acid is finished, the excess of iodine in solution complexes with starch and a deep blue color signals the endpoint [4]. In this direct determination, iodine is the titrant, ascorbic acid is the analyte of interest, and starch is the indicator [5]. An example of metal cation determination by coulometry has been shown by the titration of zinc with electrogenerated ferrocyanide at acidic pH [6]. Potentiometry is used as the endpoint detection method of the zinc precipitate K2Zn3[Fe(CN)6]2. The titrant was generated at the isolated anode and was used to determine 0.3 – 30 mg of zinc standards with average percent errors around - 0.3%. The determination of mercury by has been performed indirectly by coulometric titration [7]. Mercury is the catalyst of the reaction between ferrocyanide and phenanthroline. In this coulometric application electrogenerated iodine is used to titrate cyanide ions, a product of the mercury-catalyzed reaction. The concentration of cyanide ions found is correlated to the concentration of mercury(II) in solution by the stoichiometry of the catalysis reaction. In general, indirect coulometric methods are not common in the literature. In this work, we were able to take advantage of the complexation between zinc and cysteine to determine ionic zinc indirectly using electrogenerated iodine. The formation of this stable complex can be explained. Pearson’s theory of hard and soft acids and bases predicts cysteine, classified as a soft Lewis base due to the electron pairs it has available to donate through π bonding, and zinc ion, classified as an intermediate Lewis acid, should interact favorably. Because the Zn2+ electron configuration is d10 and this metal ion when complexed has no geometric preference based on ligand field stabilization energy, the lowest energy state for many zinc complexes is tetrahedral [8], reasonable considering the octet rule. The predominant Zn- cysteine species in solution have been studied as a function of pH showing 100% of the Zn2+ is 2- complexed as Zn-cys2 above pH 8 [9]. The step-wise formation constants for this 1:2 zinc to

60 cysteine chelate have been reported as 7.2 × 109 and 6.9 × 108 giving an overall stability constant of about 1.6 × 1018 [9,10]. Figure 3.1 shows the structure of the Zn-cysteine complex [11,12].

Fig. 3.1. Zinc-cysteine complex structure as described in the literature [9–12]. Iodine is known to be a mild oxidizing agent which makes it a suitable titrant to determine reducing agents such as cysteine. Iodine is generated from iodide at the anode and water is reduced to hydroxide and hydrogen gas at the cathode. Previous studies of the titration of thiols - - - with iodine in alkaline media have concluded that the redox reaction, 2I2 + RS + 4OH → RSO2 - + 4I + 2H2O , involves a four-electron transfer [13,14]. The primary objective of this research is to determine zinc using constant-current coulometry through its complexation with cysteine and its subsequent titration using iodine. The developed method will be tested by analyzing commercial product samples containing zinc, specifically dietary supplements. The primary advantage of this approach is that ascorbic acid and zinc, often prescribed together as cold prevention supplements, can be determined sequentially in the sample using the same coulometric method. A series titration of two different analytes by iodine coulometry is an unusual approach. 2. Material and methods

Details of the instrument used to carry out this research will be published in a separate article [16]. The teaching constant-current coulometer was constructed by using a LakeShore Cryotronics Model 121 programmable DC current source (www.lakeshore.com) set at 10 mA for all measurements. Components attached to the current source were concentric platinum electrodes, a 1-cm cylinder and a 2-cm wire gauze serving as the anode and cathode, respectively. The generating and counter electrodes were nonisolated in a 150-mL beaker that served as the coulometric cell. Aluminum foil was used to cover the sides of the beaker to prevent secondary light pathways from entering the beaker and reaching the sensor. A magnetic stir bar (2 cm × 0.5 cm) was placed inside the cell and the setup was mounted on a magnetic

61 stirrer set at 240 rpm. Photometric detection was made possible by placing a cadmium-sulphide photoresistor (Radioshack) mounted on a Styrofoam cup plastic lid under the coulometric cell. The leads from the sensor were connected to a home-built operational amplifier circuit. A Global Specialties breadboard (PB-10 externally powered 840 tie-point) (www.globalspecialties.com) was used to build the circuit. The circuit consisted of a trimmer resistor with +5 V applied connected to one lead of the photoresistor with the other lead connected to the inverting input of an operational amplifier. A negative feedback was incorporated across the output and inverting input by use of a 3.3 kΩ resistor. The non-inverting input of the amplifier was grounded. The photoresistor sensed ambient light during the titration; its resistance increased as a function of solution darkening due to the color indicator. The ratio of the circuit feedback resistor to the resistance of the photocell times the inverting input voltage indicated that a decrease in voltage output signaled the titration endpoint [15]. The voltage output of the operational amplifier was adapted to be used with a multifunctional chemical analysis (MCA) system from MeasureNet Technology Ltd. (www.measurenet-tech.com), enabling the performance of coulometric measurements with this teaching instrumentation. The MCA parameters were set to x = [0 – 3000 s], y = [0 – -2500 mV], and rate = 2 s/sample. The electrodes were connected and the MCA system was set to collect data simultaneously. The contents of the coulometric cell for titration include 0.5 g of KI, 1 mL of 3% starch solution, L-cysteine stock solution [3.3 × 10-3], zinc nitrate hexahydrate stock solution [8.0 × 10- 4], and sodium tetraborate decahydrate (0.09 M) for a final volume of 130 mL. The cysteine and sodium tetraborate were obtained from Sigma-Aldrich (www.sigmaaldrich.com). Sodium tetraborate decahydrate (pKa = 9.2) was used as the buffer at pH 9.2. This pH would ensure that the cysteine is deprotonated and can readily complex to zinc. In addition, borate does not interact with Zn2+, a problem with many other buffer choices such as carbonate. Four brands of zinc supplements (all obtained from Amazon) were initially tested for Zn using the method described in the previous paragraph. These were: Finest Nutrition 50 mg Zn caplets, Thorne Research 15 mg Zn vegetarian capsules, Nature Made® 30 mg Zn tablets, and Solgar® 22 mg Zn tablets. All tablets were prepared by crushing to a fine powder, ashed in a furnace rising to a temperature of 800°C for 30 min, and diluted using 12 mL of aqua regia and

62 water to a 100 mL final volume. Two additional types of commercial samples (all obtained from Amazon) were prepared to determine their zinc content using the method described in the previous paragraph. A 5.00 to 100 mL dilution was used to make a stock solution from Kirkman® Zinc Liquid (20 mg Zn in 5 mL). The second type of dietary supplements was lozenges: Life Extension® Enhanced Zinc Lozenges (18.75 mg Zn per lozenge). One lozenge was dissolved in 250 mL of deionized water. Aliquots were filtered using 0.2 μm Phenex syringe filters from Phenomenex (www.phenomenex.com) prior to titration. The complete listing of ingredients on the label is shown in Table S1 (Supplementary information). The following procedure was used to titrate supplements with both ascorbic acid and Zn. The titration medium used to determine ascorbic acid content was a dilute nitric acid solution (3.18 × 10-4 M) with the same quantities of KI and starch previously indicated. Then, 4.45 g of sodium tetraborate decahydrate were added to the same beaker, creating a 9.2 pH buffer solution in situ and making the blue starch-iodine color disappear. Cysteine but no additional KI or starch was added to perform the subsequent titrimetric determination for Zn. The Bluebonnet Earthsweet Zinc Lozenges, and Sunkist Zinc Throat Lozenges (5 mg Zn per lozenge); one vegetarian The Bluebonnet Earthsweet Zinc Lozenges (15 mg Zn per lozenge) and Sunkist Zinc Throat Lozenges (5 mg Zn per lozenge) were chosen because they have 100 mg and 50 mg of ascorbic acid each. The complete listing of ingredients on the label is shown in Table S1. One lozenge was dissolved in 250 mL deionized water. Aliquots were filtered using 0.2 μm Phenex syringe filters from Phenomenex (www.phenomenex.com) prior to titration. Atomic absorption data were taken using a Perkin-Elmer AAnalyst 200 instrument (www.perkin-elmer.com). 3. Results and discussion The constant current teaching instrument had been characterized by successfully determining ascorbic acid, N-acetyl-cysteine, glutathione, thiosulfate, and bisulfite. Standard solutions were determined first to show that the setup worked for each. Determinations of these analytes in commercial samples such as apple juice, dietary supplements, and dried fruits were successfully performed. The relative standard deviations of these measurements range between 0.67% and 5.49%. Recoveries were calculated from 89% to 121% [16].

Initial studies of indirect determination of zinc were performed using 3- mercaptopropionic acid (MPA) as the chelator. A change in the titration endpoint as a function of

63 zinc concentration was observed. The step-wise formation constants of the 2:1 MPA-Zn complex are 5.6 × 106 and 1.1 × 106, resulting in a stability constant of 6.3 × 1012 [17]. The zinc-cysteine complex has a higher stability constant, which led to its use as the preferred chelating agent in this method. 3.1 Titration plots

The MeasureNet system plotted titration curves of voltage (mV) versus time (s) in Excel. As shown in Figure 3.2a, from 0 to ~300 seconds the photocell sensed ambient light and gave a baseline reading. At around 300 s the solution started to turn blue, blocking the light reaching the photocell, and increasing its resistance. The increase in resistance is translated into a decrease in the absolute value voltage as predicted for the operational amplifier circuit. The resulting plot resembled a sigmoidal curve.

Fig. 3.2. (a) The raw plot (voltage vs. time), (b) first derivative plot (first derivative versus average time), and (c) the normalized voltage plot (normalized voltage vs. time). 3.2 Endpoint determination method: plot analysis

The coulometric titrations were followed by eye by removing the aluminum foil from the 150-mL beaker. The authors were looking for a uniform blue color from the starch-iodine complex. It was concluded that this uniform blue color corresponds to the point of inflection of the S-type titration plots (Figure 3.2a). The first derivative (Δvoltage/Δtime) values were plotted against time with the goal of selecting the peak of the curve as the endpoint of the titration. This

64 method was unsuccessful because no definitive peak was observed possibly due to noise, as shown in Figure 3.2b. Thus, the voltage values of the steep slopes were normalized using this equation, ynorm = y – ymax / ymin – ymax, where ynorm was the normalized voltage value, y was the raw voltage value, and ymin and ymax were the minimum and maximum voltage values from the raw data. A plot of ynorm (range 0.8 – 0.2) versus average time was created and the resulting regression was used to find the midpoint of the normalized values (Figure 3.2c). The midpoint was found by substituting 0.5 for the y and solving for x in the regression equation, yielding an endpoint of 409 s in the example shown in Figure 3.2c. This method provided a more objective calculation of the endpoint time. 3.3 Endpoint retardation as a function of zinc concentration

The four-electron transfer reaction between cysteine and iodine in alkaline media was confirmed by solving the Faraday electrolysis equation. Substituting the average endpoint time of a cysteine titration into the equation, the number of electrons was found to be 3.9, as expected [14]. During initial trials of the coulometric method it was noted that titration endpoints were longer as a function of zinc concentration in solution, permitting the creation of calibration curves. Different cysteine to zinc concentration ratios were tested to find the range in which the calibration plots showed linearity. A constant cysteine concentration of 1.27 × 10-1 mM was titrated with 7.71 × 10-4 – 5.10 × 10-3 mM Zn and a negative slope was obtained. Then, 2.51 × 10-1 mM cysteine was titrated with 0 – 4.79 × 10-1 mM Zn. The steepest slope occurred where the cysteine concentration was at least three times higher than the zinc concentration. The endpoints for some of these trials were greater than 700 s which led to the use of more diluted aliquots. As shown in Figure S1 in the Supplementary Information, from 0 – 1.74 × 10-5 mM taken in triplicate, there is an increase in titration endpoint proportional to the concentration of zinc when added to 1.24 × 10-1 mM cysteine. 3.4 Cysteine solution stability As shown in Figure S1 (Supplementary information), the creation of a calibration curve on different days showed an error of about 10% for the endpoint measurements of the 4th zinc standard (10.0 × 10-3 Mm) when titrated 5 days after standard #3. It was concluded that cysteine was subject to air oxidation. The cysteine solution stability was tested over a period of 5.5 h on the same day it was made. During the first 2.5 h, a 0.09 M sodium tetraborate buffer with pH 8.6

65 was used as the medium to titrate cysteine aliquots. The average endpoint (n=3) of these titrations was 525 ± 20.6 s, 3.92% RSD. The subsequent titrations were made with the buffer at pH 9.0. The average endpoint (n=3) for these titrations was 639 ± 29.6 s, 4.62% RSD. It was concluded that the cysteine stock solution was stable during a period of at least 5.5 h. 3.5 pH influence on endpoints The influence of pH on titration endpoints was studied after it was noted that the cysteine stability studies yielded different average endpoints at different pH values. Figure 3.3 shows titration endpoints taken in triplicate at three different pH values, the buffer’s pKa of 9.2 and pKa ± 1. In these series of experiments, a solution of 2.4 × 10-1 mM of cysteine was titrated followed by titration of a solution with 8.4 × 10-2 mM zinc. No significant difference in time was noted between cysteine only and cysteine plus zinc at a pH corresponding to 8.2. The time difference between the solutions is more significant and reproducible at the buffer pKa 9.2 than at the pH of 10.2. The buffer pH was set at 9.2 for subsequent measurements.

Fig. 3.3. Titration of cysteine and cysteine with zinc standard showed the least variability in endpoint measurements when using the buffer at its pKa, 9.2. 3.6 Endpoint dependence of temperature and stirring rate

The %RSD of the average standard endpoints measured at room temperature ranged from 1 to 10%. The titration cell was placed in a warm temperature bath (33 °C) as an attempt to maintain a constant temperature and compare measurement variability at room temperature and 33 °C. Figure S2 shows the titration of 1.3 × 10-1 mM cysteine and titration of the same concentration with 2.1 × 10-2 mM zinc. Titration endpoint variabilities (n=3) were lower at room

66 temperature. An increase in endpoint time was observed when the coulometric cell was placed in a 33 °C water bath. However, subsequent titrations were performed by placing the coulometric cell in a room temperature water bath. A stirring rate optimization study was also performed. Four trials of the same stock solution of a zinc gluconate sample were performed at different stir rates: 60 rpm, 125 rpm, 350 rpm, and 700 rpm and the endpoints were found to be 439 s, 585 s, 593 s, and 744 s, as shown in Figure 3.4. Faster stir rates resulted in increased titration times due possibly to some change in the cysteine-zinc complex stability. No appreciable change in endpoint with variation of the stirring rate (125 and 700 rpm) was noted with iodine coulometry of ascorbic acid.

Fig. 3.4. Effect of stir rate on titration endpoints. 3.7 Optimized calibration curves

Once the titration media pH, cysteine to zinc concentration ratios, and temperature bath settings were optimized the calibration curves showed significant improvements. The resulting calibration curve using the optimized parameters when titrating 1.3 × 10-1 mM cysteine with 1.2 × 10-2 to 3.6 × 10-2 mM zinc is shown in Figure S3. The linear regression for this curve was found to be y = 7732.4x + 401.53 and the correlation coefficient was 0.9966. The limit of detection and limit of quantification are 5.5 × 10-5 M and 6.2 × 10-5 M, respectively. 3.8 Dietary supplements 3.8.1 Liquid zinc The Zinc Liquid from Kirkman® has a label value of 20 mg Zn per 5 mL. A calibration curve was made and triplicate measurements of 1 mL aliquots from the 100 mL stock solution

67 were titrated. The resulting endpoints were interpolated and 23.81 ± 1.65 mg of zinc per 5 mL of liquid zinc were found using this method. The RSD was 6.93% and zinc recovery was calculated as 119%. The Zn atomic absorption result showed a found value of 20 mg. 3.8.2 Zinc lozenges The Enhanced Zinc Lozenges from Life Extension® have a label value of 18.75 mg zinc per lozenge. The endpoints of 3 mL aliquots take in triplicate from a 250 mL stock solution were interpolated in the linear regression equation of a cysteine-zinc calibration curve. The amount of zinc found in the lozenge was found to be 15.65 ± 2.18 mg Zn; an 83.5% recovery and 13.9% RSD. 3.8.3 Zinc tablets

Recoveries from dietary tablets (Nature Made and Solgar), caplets (Finest Nutrition), and capsules (Thorne Smith) having more complicated samples matrices (see Table S1) ranged from 60 to 80%. The reason for the low recoveries of the tablets is not clear. Possibly ZnO formation has occurred during the ashing step but aqua regia should convert this to zinc ion. A comparison to zinc atomic absorption data showed only the Solgar tablet had a similar low recovery. Therefore, in general, the method of sample preparation should not be a factor. Interference studies were performed because cysteine can chelate other metals present in these supplements such as calcium and magnesium ions (although more weakly with stability constants of 102 and 104 respectively. Figure 3.5 shows the average endpoints of cysteine (1.3 × 10-1 mM), cysteine with calcium nitrate tetrahydrate (1.2 × 10-2 mM), cysteine with magnesium nitrate hexahydrate (1.2 × 10-2 mM), and cysteine with 1.2 × 10-2 mM Zn and 8.2 × 10-3 mM phosphate. Based on the average endpoints of these trials it can be stated that the presence of calcium cations does not appear to influence cysteine titration endpoints. Magnesium cations caused cysteine endpoint retardation showing that its presence can possibly interfere with zinc determinations. Lastly, zinc cations in the presence of phosphate caused shorter cysteine titration endpoints, indicating phosphate may be a potential interferent. However, an equimolar combination of zinc and calcium phosphate titrated similar to that as zinc alone, indicating the interference was possibly alleviated (result not shown).

Fig. 3.5. Interference studies on cysteine titration endpoints. 3.8.4 Ascorbic acid and zinc combination titrations

Combined titrations of ascorbic acid followed by zinc determination were performed. The Bluebonnet Zinc Lozenge contains 15 mg Zn and 100 mg ascorbic acid per lozenge. The Sunkist supplement contains 5 mg Zn and 50 mg ascorbic acid. A three-point calibration curve, shown in Figure S4, was created using this method to ensure linearity. The linear regression was found to be y = 1.2 × 107x + 356.5 and the correlation coefficient was 0.9959. Ascorbic acid titration endpoints were used to calculate the analyte amounts by use of the Faraday electrolysis equation. The cysteine-zinc titration endpoints were interpolated in the linear regression of the calibration curve. Determination of 3 mL aliquots taken in triplicate from a 250 mL solution of the dissolved Bluebonnet lozenge yielded 95.8 ± 5.5 mg ascorbic acid (95.8% recovery and 5.72% RSD) and 16.8 ± 1.9 mg Zn (112% recovery and 11.1% RSD). The closeness of the determined ascorbic acid to the label value was unexpected considering the list of natural ingredients (Table S1). The Zn atomic absorption result showed a found value of 15 mg. Volumes of 9 mL aliquots taken in triplicate from 250 mL stock solution were used to determine the same two analytes in the Sunkist lozenge. This supplement was found to have 95.0 ± 2.1 mg ascorbic acid (190% recovery and 2.22% RSD) and 4.80 ± 0.31 mg zinc (96% recovery and 6.43% RSD). The high recovery of ascorbic acid was expected because of the likely natural citrus based extracts in the lozenge (Table S1). The Zn atomic absorption result showed a lower recovery, a found value of 3.9 mg. 69

4. Conclusions We have shown that iodine coulometry can be extended to permit the determination of zinc ion. We have done this by taking advantage of the high stability constant of the cysteine- zinc complex. It was discovered that the coulometric titration of cysteine when complexed with zinc at pH 9.2 is slower as compared to uncomplexed cysteine. This endpoint retardation was proportional to the concentration of zinc with excess cysteine in solution permitting generation of a linear calibration curve to quantify zinc in dietary supplements. One drawback of this method is that standards and samples should be run on the same day because the L-cysteine solution is subject to air oxidation. A minor issue is the titration must be run at a controlled alkaline pH to ensure reproducible zinc-cysteine complexation and because variation of the number of electrons has been shown for iodine coulometry of thiols as a function of hydroxide concentration [14]. Low analyte recoveries for zinc obtained with this coulometry method when compared to atomic absorption results were indicated for certain tablets that contain Ca/Mg ions and various anions as well. Further study concerning potential interferences and how to mask them is required. Advantages of this approach have been demonstrated. A low cost coulometry instrument has been shown to be useful for the indirect determination of zinc using well-established chemistry involving iodine and thiol reducing agent. The same chemistry and instrument for the titration of ascorbic acid and Zn2+ in series for the analysis of cold prevention supplements has been shown to be viable. The method described herein could also potentially be used to determine other transition metal cations that have high stability constants with cysteine or an alternative thiol such as MPA [18].

Funding This material is based upon work supported by the National Science Foundation under Grant Number 1432466, Principal Investigator S.L. Bretz. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the author(s) and do not necessarily reflect the views of the National Science Foundation. References [1] N. Pace, E. Weerapana, Zinc-Binding Cysteines: Diverse Functions and Structural Motifs, Biomolecules. 4 (2014) 419–434. [2] U.S. Food and Drug Administration, Dietary Supplements, (n.d.). http://www.fda.gov/Food/DietarySupplements/default.htm.

[3] D.R. Levinson, Dietary Supplements: Structure/Function Claims Fail to Meet Federal Requirements, (2012). https://oig.hhs.gov/oei/reports/oei-01-11-00210.pdf. [4] V.T. Calabrese, A. Khan, Amylose-iodine complex formation without KI: Evidence for absence of iodide ions within the complex, J. Polym. Sci. Part A Polym. Chem. 37 (1999) 2711–2717. [5] R. Karlsson, Iodometric determination of ascorbic acid by controlled potential coulometry, Talanta. 22 (1975) 989–993. [6] J.J. Lingane, A.M. Hartley, Coulometric titration of zinc with ferrocyanide, Anal. Chim. Acta. 11 (1954) 475–481. [7] T.J. Rohm, W.C. Purdy, Kinetic-coulometric determination of mercury in biological samples, Anal. Chim. Acta. 72 (1974) 177–182. [8] T. Dudev, C. Lim, Tetrahedral vs Octahedral Zinc Complexes with Ligands of Biological Interest: A DFT/CDM Study, 122 (2000) 11146–11153. [9] D.D. Perrin, I.G. Sayce, Complex formation by nickel and zinc with penicillamine and cysteine, J. Chem. Soc. A Inorganic, Phys. Theor. (1968) 53–57. [10] N.C. Li, R.A. Manning, Some Metal Complexes of Sulfur-containing Amino Acids, J. Am. Chem. Soc. 77 (1955) 5225–5228. [11] H. Shindo, T.L. Brown, Infrared Spectra of Complexes of L-Cysteine and Related Compounds with Zinc(II), Cadmium(II), Mercury(II), and Lead(II), J. Am. Chem. Soc. 87 (1965) 1904–1909. [12] P. Bell, W.S. Sheldrick, Preparation and Structure of Zinc Complexes of Cysteine Derivatives, Zeitschrift Fur Naturforsch. - Sect. B J. Chem. Sci. 39 (1984) 1732–1737. [13] W. Ciesielski, A. Krenc, U. Zlobinska, Potentiometric Titration of Thioamides and Mercaptoacids With Iodine in Alkaline Medium, Chem. Analityczna. 50 (2005) 397–405. [14] W. Ciesielski, R. Zakrzewski, Iodimetric Titration of Sulfur Compounds in Alkaline Medium, Chem. Analityczna. 51 (2006) 653–677. [15] A.J. Diefenderfer, Principles of Electronic Instrumentation, 2nd ed., W. B. Saunders, Philadelphia, PA, 1979. [16] J.B. Padilla Mercado, E.M. Coombs, J.P. DeJesus, S.L. Bretz, N.D. Danielson, Iodine coulometry with on-line photocell detection for a multi-functional chemical analysis (MCA) system, J. Chem. Ed. (2017) (to be submitted) [17] Q. Fernando, H. Freiser, Chelating Properties of B-Mercaptopropionic Acid, J. Am. Chem. Soc. 80 (1958) 4928–4931. [18] C.G. Halliday, M.A. Leonard, Selective Photometric Titration of Calcium or Magnesium with EDTA Using Various Thiols as Masking Agents, Analyst. 112 (1987) 83–86.

Supplementary Information

Figure S1. Calibration curve of 1.24 × 10-1 mM cysteine with 0 - 1.74 × 10-5 mM zinc nitrate hexahydrate. The first four and last four points were taken 5 days apart using the same cysteine stock solution.

Fig. S2. Cysteine and cysteine with zinc endpoints titrated in two different temperature baths.

Fig. S3. Five-point calibration curve of titration of 1.3 × 10-1 mM cysteine with zinc standard ranging from 1.2 × 10-2 to 3.6 × 10-2 mM zinc.

Fig. S4. Calibration curve of constant cysteine concentration with increasing zinc concentrations for ascorbic acid and zinc combined titrations.

Brand Ingredients Zinc gluconate, cellulose gel, calcium Nature Made® tablets carbonate, stearic acid, croscarmellose sodium, hydroxypropyl methylcellulose, magnesium stearate, polyethylene glycol. Zinc gluconate, cellulose (plant origin), Finest Nutrition caplets dicalcium phosphate, croscarmellose, silica, vegetable magnesium stearate, vegetable stearic acid. Zinc picolinate, dicalcium phosphate, Solgar® tablets microcrystalline cellulose, vegetable stearic acid, vegetable cellulose, vegetable magnesium stearate. Zinc picolinate, microcrystalline cellulose, Thorne Research capsules hypromellose (derived from cellulose) capsule, leucine, silicon dioxide. Purified water, sorbitol, glycerine, citric acid, Kirkman® Zinc Liquid zinc alpha ketoglutarate, natural raspberry flavor, potassium sorbate. Life Extension® Enhanced Zinc Lozenges Zinc acetate, dextrose, peppermint flavor, stearic acid, vegetable stearate, silica. Vitamin C (L-ascorbic acid, sodium ascorbate), Zinc (zinc gluconate, zinc citrate), natural orange flavor, EarthSweet® [juice Bluebonnet Earth Sweet Chewables zinc concentrates (wild berry, cranberry, prune, lozenges cherry, strawberry, grape, raspberry and bilberry fruits, grape seed, raspberry extracts), cane crystals], orange crystals, stearic acid, vegetable magnesium stearate. Vitamin C (sodium ascorbate), zinc (zinc citrate, zinc gluconate), sorbitol, sucrose, Sunkist lozenges mannitol, flavors (orange, honey lemon, cherry), colors (red carmine, annatto, turmeric), sucralose. Table S1. Complete label listings of ingredients in the commercial samples.

Chapter 4. Conclusions and future directions I. Specific Aim 1

The use of constant-current coulometry has a substantial presence in the teaching literature. Much of the published work uses homebuilt instruments to teach this analytical technique. We have shown the versatility of iodine coulometry to provide facile analyte variation for this experiment on an on-going semester basis. In this work, we have demonstrated that iodine titration can be used in our constant-current coulometry instrument to determine ascorbic acid, N-acetylcysteine, bisulfite, glutathione, and thiosulfate. The successful determination of these analytes in acidic media (except bisulfite which is more stable at neutral or alkaline pH) via their oxidation by electrogenerated iodine shows that our instrument can excel the capabilities of previously published work. Our simple design provides another advantage over commercial coulometry instruments. The exposed fashion of the cell, photodiode, and current-to-voltage converter circuit allows a rich pedagogical experience with this analytical method. Students can see the parts of the instrument and have an opportunity to relate them to their functions. Furthermore, the ability to determine titration endpoints by eye or by graphical analysis provides a comparison of the ease and objectivity of the methods. II. Specific Aim 2

To the best of our knowledge, iodine coulometry has not been used to determine zinc ions. We have extended the capabilities of this chemistry to determine the concentration of zinc cation indirectly by taking advantage of its complexation with cysteine at alkaline pH and subsequent titration of cysteine by iodine. The titration time was found to be proportional to the zinc concentration. The method has been developed and optimized in terms of cysteine-to-zinc concentration ratios, titration media pH, stirring rate, and reaction temperature. The optimized parameters were used to successfully determine zinc content in a variety of dietary supplements by use of a calibration curve. The series titration of ascorbic acid followed by zinc determination via cysteine complexation was also performed successfully. L-ascorbic acid was quantified by use of Faraday’s law of electrolysis and zinc content was found by interpolation of the linear regression equation of a calibration curve.

III. Potential Future Work

The teaching constant-current coulometry instrument described herein can potentially be used to determine other analytes. There is a wide variety of dried foods or beverages such as wine that contain bisulfite. In addition, the cysteine complexation method should be adaptable for the determination of lead in paint chips. The formation constant of the lead-cysteine complex has been reported as 1.58 x 1012.1 References (1) Li, N. C.; Manning, R. A. Some Metal Complexes of Sulfur-Containing Amino Acids. J. Am. Chem. Soc. 1955, 77 (20), 5225–5228 DOI: 10.1021/ja01625a006.

Appendix: Oxygen meter finger probe studies

I. Introduction a. Plethysmographs: Pulse oximeters A plethysmograph is a device used to measure volume changes in an organ due to air or blood flow. Several different types of plethysmographs exist; these differ in the mechanism used to measure changes in volume. Water, air, strain gauge, impedance and photoelectricity are common types of plethysmographs, each with unique uses.1 In light of the experiments carried out here it is important to define photoplethysmographs. These devices utilize light sources with photodetectors to monitor blood flow. Pulse oximeters are a type of plethysmograph that are placed on earlobes or fingertips to measure oxygen saturation in the blood.2 These are suitable locations for pulse oximeter placement because the device operates under the assumption that arterial blood is anything that pulses and absorbs red and infrared light. Pulsatile tissue is the first requirement for a measurement of oxygen saturation, or the fraction of hemoglobin (HbO2) to reduced hemoglobin (Hb) as shown in Equation A.1. By monitoring pulsatile signals, the device provides information about oxygenation of arterial blood.

HbO2 SpO2% = ×100% (A.1) HbO2+Hb To explain the second half of the assumption by which pulse oximeters work (anything that absorbs red and infrared light) the structure of these two compounds will be explained. Hemoglobin is the molecule in charge of oxygen transportation in the body. It is a protein composed of four subunits, each containing a heme group. A heme group is a cyclic structure, also called porphyrin, containing an iron atom in the center as shown in Figure A.1. Every iron atom can bond to one molecular oxygen atom; when the four hemes are bound to oxygen the whole protein is named oxygenated hemoglobin. The loss of molecular oxygen due to pH changes in the protein’s environment, also known as the Bohr effect, results in a reduced hemoglobin structure.

Figure A.1. Heme group in hemoglobin. The iron(II) in reduced hemoglobin is 70 pm out of the plane of the porphyrin. It is in a high spin d6 configuration and absorbs red light. Bonding with oxygen causes the iron(II) atom to sit on the plane of the porphyrin. In this structure the iron is low spin, diamagnetic, and absorbs blue light.3 These differences in electronic arrangements and absorption allow the use of pulse oximeters to monitor blood oxygenation. The Beer-Lambert Absorption Law states that the amount of light absorbed by a compound is proportional to its concentration as shown in Equation A.2, where A is the absorbance, ε is the molar absorptivity, l is the light path length, and c is the concentration of the compound of interest. The light absorption of hemoglobin and reduced hemoglobin have been studied and quantitatively described in the wavelength range from 450 to 1000 nm.4 The millimolar absorptivity of hemoglobin is higher than that of reduced hemoglobin at 660 nm, 0.83 to 0.07 L/mmol×cm respectively. In the near infrared region, specifically at 940 nm, the trend is reversed and the hemoglobin ε value is 0.17 L/mmol×cm compared to 0.28 L/mmol×cm for reduced hemoglobin. This difference in absorbance drives blood oxygenation monitoring. A = εlc (A.2) b. Parts of the oximeter The oximeter used for these studies was the Oxi Go Quick Check Pro by Oximeter Plus, 5 shown in Figure A.2. The display shows SpO2, pulse rate, and a pulse bar. Its SpO2% range is 35 – 99% and the pulse rate range is 30 – 235 beats per minute. The light sources, a red LED at 660 nm and an IR LED at 940 nm, are positioned on top the finger. The bottom part of the finger clip contains the photodetector, making it a transmittance meter since the attenuation of light by blood absorption is the basis of the measurement.

Figure A.2. Picture of oximeter used in these studies. c. Measurement Pulse oximeter measurements are time-dependent because they perform background noise subtraction in real time. This is done by alternating the LED sources.2 The red LED is turned on, then it is turned off. Next, the IR LED is turned on and off, and then both are kept turned off. Alternating the light sources and then keeping both off is done hundreds of times per second and the idea is to detect any scattered light that is constant through all three phases and subtract it from the measurement. The light source alternation produces four signals, two per LED: AC at 660nm, DC at 660nm, AC at 940 nm, and DC at 940 nm, where AC is alternating current and DC is direct current. The alternating current arises from the blood volume increase due to the heartbeat. The unchanging, direct current comes from the steady state absorbance of the background which includes room light, body fluids, tissue, and bones. The circuit amplifies and filters the alternating current signals before converting them to a SpO2% value. d. Purpose The idea that drove the experiments described in this appendix was the desire to take advantage of the parts oximeter: light source, detector, and readout, to use as a detector for chromatographic separations. Specifically, the oximeter would become an affordable, portable, and quantitative component for paper chromatographic separations.

By looking at the SpO2% equation it is understood that a high percent (close to 99%) corresponds to a high concentration of oxygenated hemoglobin. In this case the blood looks red, it is absorbing blue, and the highest absorption occurs in the infrared region (950 nm). The opposite is true for reduced hemoglobin. Based on this information the following experiments were performed.

II. Experiments and Results a. Filter paper studies Filter paper was the first surface tested to be used to introduce samples into the finger oximeter. A green permanent marker was used to make a dot on the tip of a filter paper cut as a rectangle. Green was chosen because green dyes absorb in the red region of the visible spectrum. Whatman filter papers of the following numbers were tested: 1. 1PS, 2, 40, 41, 42, 54, and 541. Every filter paper was introduced into the oximeter opening and waved. All pulse rate readings were between 161 – 163 rpm, showing that the “pulsatile” signal was kept constant. The SpO2% values found were from 77 to 97%. Thus, the Whatman no. 1 filter paper was selected to be used for subsequent studies. Whatman no. 1 filter papers were used to test different green colored writing utensils. Crayons, finger paint and permanent markers with different shades of green were tried. The pulse rate readings for these experiments were between 160 and 163 rpm. SpO2% readings ranged from 81 to 97%. It was expected that the darker green colors would show high SpO2% values, just like oxygenated hemoglobin. No correlation was seen between the darkness of the green color and the resulting SpO2%. Malachite green is an organic compound used as a dye. It has an absorption band at 620 nm. At low pH values, it turns yellow. Two stock solutions of malachite green were prepared, a green and a yellow one. At pulse rates of 162 rpm the green solution resulted in SpO2 of 98% and the yellow yielded 86%. An additional dye stock solution was made using a Fast Green FCF dye. This dark solution produced a darker stain on the Whatman filter paper but resulted in a lower SpO2% value than that of the malachite green, 80%. These studies did not correlate with our hypothesis of darker green colors producing higher SpO2% values. An electric toothbrush (Kroger brand, Vibraclean®) was employed to hold the colored filter papers for the SpO2% reproducibility tests. Different permanent markers were used to draw a dot on Whatman filter paper no. 1: lime green, yellow, orange, brown, and blue. The setup for these studies is shown in Figure A.3. Table A.1. shows the results of these tests. The variability of SpO2% values is relatively low. No apparent trend was seen between marker darkness and

SpO2%.

Figure A.3. Green marker dot on electric toothbrush.

Permanent marker color Average SpO2% (N=4) Relative standard deviation Lime green 91% 3% Yellow 87% 7% Orange 87% 9% Brown 93% 8% Blue 81% 4% Table A.1. Reproducibility studies with electric toothbrush. Paper chromatographic separations of permanent markers were performed on Whatman no 1 filter paper. The solvent used was 1:1 deionized water to isopropyl alcohol 90% (CVS brand). Separations of 12 different markers were used to stain filter papers and were run in duplicates. The colors were black, red, green, blue, yellow, orange, brown, light blue, aqua, green, and lime green. The results of the blue and yellow marker separations were blue and yellow color gradients, respectively. Three different blue spots (dark, medium, and light) and two different yellow spots (dark and light) were hole-punched and tested to create calibration curves. The results of these tests in which the electric toothbrush was used are shown in Table

A.2. No apparent trend can be seen between color hue and SpO2% values. A white paper was run as a blank and was found to have a SpO2% value of 86% (1 %RSD). Reproducibility of these trials is good.

Color (hue) Average SpO2% (N=4) Relative standard deviation Blue (light) 81% 15% Blue (medium) 76% 13% Blue (dark) 94% 6% Yellow (light) 91% 3% Yellow (dark) 89% 4% Table A.2. Attempts to make calibration curves of permanent marker paper chromatography separations. Filter papers with black, red, and green permanent marker spots were inserted in the oximeter while using a finger as the probe. The resulting SpO2% values are shown in Table A.3.

The permanent marker spots increased the SpO2% readings confirming the hypothesis. No conclusion could be drawn from the specific colors.

Trial Average SpO2% (N=8) Relative standard deviation Finger 97% 3% Finger and filter paper 98% 3% Black 99% 1% Red 99% 1% Green 99% 2% Table A.3. Filter paper with permanent marker spots trials. b. Targeting the IR region The dye and permanent marker studies were inconclusive which made us target the IR region exclusively. Several compounds were considered, including long-chain esters such as fruit flavor and perfumes, coumarin, warfarin, fluorescein, N-butyl acetate, N-propyl acetate, ethyl acetate, methyl acetate, aspirin, salicylate, naproxen, ibuprofen, and acetaminophen. N-butyl acetate and methyl acetate solutions were tested using Whatman filter paper no. 1. The results of these solutions and their respective blanks are shown in Table A.4. Readings were only obtained for blanks and dried filter papers. Blank readings were higher than filter papers with sample on them.

Name Average SpO2% (N=3) Relative standard deviation Blank 98% 1% 2 drops of N-butyl acetate 96% 2% (dried) 2 drops of N-butyl acetate No reading (“finger out”) - (wet) Blank 97% 4% 2 drops of methyl acetate 94% 7% (dried) 2 drops of methyl acetate No reading (“finger out”) - (wet) Table A.4. Trials of IR active compounds. Salicylic acid was chosen as another compound absorbing in the IR region. An attempt to obtain readings from the solid form of this compound was preceded by a search of a noninterfering material in which to insert the sample in the oximeter. The options were a non-CH polymer like Teflon tape, a CH-containing polymer like propylene, nylon filter, and nitrile gloves. Teflon tape was chosen first because it is a polymer with repeating units of (C2F4)n. In theory, it would not interfere with oximeter readings. Different tests were performed including wrapping Teflon tape a finger to test this hypothesis followed by nitrile glove trials and the measurement of 0.03 grams of salicylic acid. Table A.5 shows the results of these tests. Teflon tape could be wrapped up to three times around the finger and measurements could still be recorded. The %RSD increased as a function of amount of Teflon tape used. Trials in which the tape was wrapped four and five times around the finger, and where a purple nitrile glove was used gave no measurements. The insertion of salicylic acid caused no apparent change in SpO2% values.

Trial Number of Average Trial %RSD description trials SpO2% No tape 8 100% 0.5% Tape wrapped Teflon tape two times 8 98% 3% around finger No tape 3 99% 1% Tape wrapped Teflon tape three times 3 92% 10% around finger No glove 8 99% 0.4% Blue nitrile Tape wrapped glove twice around 8 99% 1% finger Teflon tape bundle under 6 99% 1% Teflon tape and finger salicylic acid Sample in tape bundle under 6 99% 0.4% finger Table A.5. Teflon tape, nitrile glove and salicylic acid sample in oximeter. c. Meat studies Using colored filter paper to exclusively target the red LED of the pulse oximeter and using an IR-active compound to target the IR LED did not give conclusive results. This led us to find a system like the one the oximeter was built for. Meat preservatives such as nitrites are added before the packaging stage. Nitrites can prevent growth of disease-causing microorganisms and they cause the red, cured meat color. Meat contains myoglobin which is an oxygen-storage protein in muscle tissue. Myoglobin is a heme-containing protein just like hemoglobin. Reduced myoglobin looks purple and nitrosomyoglobin looks red. By adding nitrites to raw meat, the meat industry ensures that nitric oxide reacts with myoglobin to produce the red color. It was hypothesized that oximeter readings of red, raw meat could be taken

84 followed by measurement of partially cooked meat and the SpO2% readings would decrease as a function of loss of red color. Rectangles with the size of the LED sources and the photodetector were cut out from two sides of a transfer pipette to avoid interferences from the plastic. Raw meat was placed inside the cutout bulb of the transfer pipette and Teflon tape was wrapped around it to keep the sample in place. The results of blank tests are shown in Table A.6. Subsequent meat trials were as follow: blank (pipette and Teflon tape) 52%, pipette with tape and raw meat 82%, and pipette with tape and heated meat 99%. These results were incongruent with our hypothesis. The variability of

SpO2% results for these trials was very high.

Trial Number of trials Average SpO2% %RSD Finger by itself 3 99% 0.6% Finger with pipette 3 98% 0.6% Finger with pipette 3 76% 6% and Teflon tape Table A.6. Blank trials with transfer pipette and Teflon tape.

III. Discussion and Possible Future Work The tests performed with the pulse oximeter show that the algorithm it uses to calculate

SpO2% values might be too specific. Attempts to target the red LED exclusively and the IR LED only were unsuccessful probably because of this specificity. Going forward with the attempt to use the finger pulse oximeter as a detector for separations two things must be kept in mind. First, both light sources should be targeted. Second, pathlength should be kept constant during measurements. This can be achieved by using a foam finger, for example, or any cylindrical object that causes the least interference at the 660 nm and 940 nm wavelengths. An idea for possible future work is to use tubing made of non-interfering material through which dyes can be pumped through. A peristaltic pump would probably send waves of the dye and mimic the pulsatile signal the oximeter is designed to detect. The oximeter can be clamped to the tubing and be used to measure absorbance of dyes. It is also possible that it could be used to detect the blue shift and absorbance peak enhancement of Nile Blue in the presence of 6 3.5 mM sodium dodecyl sulfate, 휆푚푎푥 635 nm to 590 nm.

IV. References

(1) Elgendi, M. On the Analysis of Fingertip Photoplethysmogram Signals. Curr. Cardiol. Rev. 2012, 8 (1), 14–25 DOI: 10.2174/157340312801215782. (2) Tremper, K. K. Pulse Oximetry. Chest 1989, 95 (4), 713–715. (3) Miessler, G. L.; Tarr, D. A. Inorganic Chemistry, 2nd ed.; Prentice Hall: New Jersey, 1999. (4) Zijlstra, W. G.; Buursma, A.; Meeuwsen-van der Roest, W. P. Absorption Spectra of Human Fetal and Adult Oxyhemoglobin, de-Oxyhemoglobin, Carboxyhemoglobin, and Methemoglobin. Clin. Chem. 1991, 37 (9). (5) Oximeter Plus. Oxi Go Quick Check Pro http://oximeterplus.com/pro-7. (6) Mitra, R. K.; Sinha, S. S.; Pal, S. K. Interactions of Nile Blue with Micelles, Reverse Micelles and a Genomic DNA. J. Fluoresc. 2008, 18 (2), 423–432 DOI: 10.1007/s10895- 007-0282-1.