Design, Development, and Characterization of an ... - ACS Publications

Jun 9, 2014 - The teaching of instrumental analysis for many small colleges and high schools continues to be stymied by high-cost, complicated mainten...
20 downloads 16 Views 2MB Size


Design, Development, and Characterization of an Inexpensive Portable Cyclic Voltammeter Jenna R. Mott, Paul J. Munson, Rodney A. Kreuter, Balwant S. Chohan, and Danny G. Sykes* Department of Chemistry and Forensic Science Program, Pennsylvania State University, University Park, Pennsylvania 16802 United States S Supporting Information *

ABSTRACT: The teaching of instrumental analysis for many small colleges and high schools continues to be stymied by high-cost, complicated maintenance, high power requirements, and often the sheer bulk of the instrumentation. Such issues have led us to develop inexpensive instruments as part of a SMILE initiative (small, mobile instruments for laboratory enhancement). This lab-based pedagogy, primarily aimed at engaging science and engineering students, has shown to significantly enhance the confidence and achievement of students in our technology-based courses. One instrument that has been designed, constructed, and characterized is the cyclic voltammeter (CV). CV is a versatile electroanalytical technique that monitors the redox behavior of chemical and biochemical species in solution. The CV instrument and technique readily lends itself to miniaturization and facilitates the practical application of CV analysis within standard undergraduate and advanced high school laboratory courses. The entire instrument was constructed for less than $50, thus allowing deployment of multiple apparatus in laboratories with a modest budget. Less noisy data was obtained when a commercial Pt working electrode was used, taking the total cost to less than $250. Together with details on how to construct the instrument, a series of experiments for the student-built CV have been described. KEYWORDS: First-Year Undergraduate/General, High School/Introductory Chemistry, Upper-Division Undergraduate, Analytical Chemistry, Laboratory Instruction, Hands-On Learning/Manipulative, Electrochemistry, Electrolytic/Galvanic Cells/Potentials, Laboratory Equipment/Apparatus, Oxidation/Reduction


construct, characterize, and troubleshoot small instruments. Underclassmen, including visiting school students, are then taught how to assemble the instruments from kits and use them in their chemistry laboratories. The instrument described herein facilitates the practical application of cyclic voltammetric (CV) analysis within standard undergraduate and high school laboratory courses. The instrument was designed for less than $50 (excluding Pt working electrode), clearly an inexpensive alternative to commercial CV instruments that can cost upward of $15,000. A recent report4 highlighted the cost issue by making use of a MicroLab sensor and interface as a potentiostat to control a Pine electrochemical cell that uses disposable electrodes for $1500/station; however, the instrument that we describe herein has been constructed by students for less than one-tenth of the cost of the Pine system. The CV remains under constant development, and the latest model has allowed us to obtain some remarkable data on a number of chemical systems. To encourage greater participation, we have also developed a series of lab experiments of varying complexity that could be easily geared toward a variety of chemistry and biochemistry courses. These laboratories also provided an opportunity to gauge the reliability, reproducibility, and limitations of the

nalytical instrumentation has changed significantly over the years, and equipment that was the mainstay of research institutes is now routinely found in undergraduate chemistry laboratories. The sensitivity and selectivity of analytical techniques has increased many-fold over this period. Likewise, the low cost and availability of such analysis has led to a broad increase in public awareness of health, safety, and environmental issues relating to pharmaceuticals, food, energy, and industrial processes. These advances demand that schools and colleges train students in the necessary STEM (science, technology, engineering, and mathematics) skills required to analyze a wide variety of samples with an equally wide array of instruments, thus enabling students to succeed in the modern workplace. Unfortunately, the teaching of analytical chemistry continues to be thwarted by the high-cost, complex maintenance, and bulk of many commercial instruments. For large institutions, providing each student access to an instrument requires either a sophisticated lab rotation program or the deployment of multiple instruments. These issues have led us to focus on developing a variety of small inexpensive instruments for undergraduate, high and middle school students, through a program that we have dubbed SMILE (small, mobile instruments for laboratory enhancement).1−3 At Penn State, students in the upper-level instrumental analysis course design, © XXXX American Chemical Society and Division of Chemical Education, Inc.

A | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


Figure 1. Circuit diagram for the CV instrument.

where, Ip is the current maximum (amps), n is number of electrons transferred, A is the electrode area (cm2), F is the Faraday Constant (C mol−1), D is the diffusion coefficient (cm2/s), C is the concentration (M), ν is the scan rate (V/s), R is the universal gas constant (J/mol·K), and T is temperature (K).

student-built CV in comparison to a commercial BASi Epsilon CV50W instrument.

THEORY The CV technique provides qualitative and quantitative information about inorganic, organic, and biochemical redox reactions within a potential range. The experiment is carried out in an electrolytic cell with three electrodes: the reference keeps the potential between itself and the working electrode constant, and the current is measured between the working and counter electrode. In a typical experiment, the potential of the working electrode is varied linearly from an initial value to a predetermined limit where the direction of the scan is reversed. A voltammogram is obtained by recording the current produced at the working electrode during the entire potential scan. Depending on the information sought, single or multiple cycles of varying scan rate and potential limits can be employed. Quantitative information on the analyte can be obtained using the Nernst and Randles−Sevcik equations. The Nernst equation (eq 1) describes the relationship between the potential applied to an electrode and the concentration of the redox species at the electrode surface E = E0 +

⎛ 0.0591 ⎞ ⎛ [Ox] ⎞ ⎜ ⎟ log ⎜ ⎟ ⎝ n ⎠ ⎝ [Red] ⎠


CV Instrument Design

The potentiostat circuit (Figure 1) makes use of a programmable interface controller (PIC). An alternative, breadboardbased version is described in the Supporting Information. The simpler potentiostat uses a commercial function generator which is tied to the 100k resistor on the inverting input of the LMC6484 op-amp, the output of which drives the counter electrode. Data obtained using either version were found to be identical. The CV consists of two parts (Figure 2): a commercial miniLAB-1008 data acquisition (DAQ) unit and a “studentbuilt” potentiostat. The miniLAB provides an interface for


where E is potential applied (V), E is formal reduction potential of the couple vs reference electrode (V), n is number of electrons, [Ox] is surface concentration of species Ox, and [Red] is surface concentration of species Red. The Randles−Sevcik equation (eq 2) specifies the peak current, Ip (Ipa, anodic or Ipc, cathodic), in terms of the analyte concentration, C. For simple redox events, Ip depends on the diffusional properties of the electroactive species and the scan rate Ip = 0.44463nFAC(nFvD/RT )1/2

Figure 2. (A) MiniLAB-1008 for data acquisition and (B) circuitry inside the student-built potentiostat.

(2) B | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


wire or a commercially available Pt disc (MF2013, BASi), and the counter (auxiliary) electrode, which was simply a 6-cm-long 2 mm Pt wire. These were housed in a 50 mL beaker or a shotglass (electrochemical cell) with a customized rubber or drilled Teflon top, shown in Figure 3.

analog to digital conversion. The potentiostat employs a Pt working electrode, a Pt counter electrode, and a simple studentbuilt Ag/AgCl reference electrode. A general list of components and costs are given in Table 1; a detailed list is provided in the Table 1. General List of Parts for the CV Part

Cost ($)

PIC-16F684 Programmable Chip LMC-6484 Operational Amplifier Plastic Housing Power Unit Resistors and Capacitors Ag wire Pt wire Pt working electrode Teflon stopper

2.00 3.00 5.00 5.00 4.00 2.00 25.00 207.00 3.00

Supporting Information. In order to build the instrument as shown, some machine work on the housing is required; this small additional cost is not reflected in the total. The compact nature of the circuit board lends itself to multiple packaging possibilities. The circuit diagram for the CV is shown in Figure 1, and detailed procedures for its construction are provided in the Supporting Information. The circuit consists of two main subcircuits, a wave generator and current amplifier. The power adapter (120 V/60 Hz to 12 V/300 mA) is not shown. The first circuit, comprising of the PIC-16F684, three variable resistors, and an output filter, is used to produce a triangle waveform in the ±2.5 V range. The amplitude, slope, and DC offset of this waveform are controlled by variable resistors. Earlier prototypes generated this using analogue circuitry; the PIC microcontroller design uses fewer parts. The triangular waveform is coupled to a second subcircuit, a potentiostat, which is built using the LMC-6484 op-amp. A phase-compensated op-amp drives the counter electrode. The loop is closed by the high input impedance voltage follower connected to the reference electrode which also provides a V applied signal. The current through the sample is detected by a transimpedance amplifier that has three selectable gains. The output of the amplifier is low-pass filtered to limit high frequency noise, which is then applied to the input of the A/D in the DAQ. Data processing and interpretation was performed with MathWorks-Matlab (7.14). If access to a commercial waveform generator is possible, then the output of the commercial unit can be tied to the 100k resistor on the inverting input of the LMC-6484 opamp, the output of which drives the counter electrode. The remaining circuit can then be built on a digital breadboard.

Figure 3. Setup for the CV cell, with electrodes and drilled Teflon stopper.

The Ag/AgCl reference electrode was readily constructed by students according to literature.5−8 One end of a silver wire was connected to the positive-terminal of a 9 V battery and the other end immersed in a 1 M solution of HCl for 50 s to ensure an even coat. A copper wire was connected to the negative terminal and also immersed in HCl. The data obtained using this electrode was comparable to that from a commercial electrode (MF2052, BASi).

The CV potentiostat was connected to the DAQ unit, which in turn was attached via USB to a computer. The acquisition program was written in Java and is programmed to plot voltage vs current. There are four switches on the potentiostat that allow adjustments to the receiver gain (low, medium, high), amplitude, frequency, and DC-offset. The amplitude and DCoffset were set at the start of each set of experiments based on the concentration of the test solutions. The frequency was adjusted at the start of each experiment so as to set the scan rate. Detailed procedures for the calibration are provided in the Supporting Information. Before each solution was analyzed, a clean electrochemical cell was filled to three-fourths of its capacity, and the three electrodes were carefully inserted through the top, submerged about 3 cm into the solution, and connected to the potentiostat. The solution was sparged with argon for 5 min while stirring. Following calibration, the scan rate was set using a stopwatch. After a preliminary run to determine the final settings for the reaction, the experimental data was saved, transferred to Excel, and plotted. Triplicate runs were performed to ensure reproducibility. For comparison, a BASi Epsilon CV50W electrochemical analyzer with a C3-stand and Faraday cage was employed for each experiment and the data allowed us to gauge the precision, accuracy, and quality of the student-built instruments. A Pt wire was used as the auxiliary, a student-built Ag/AgCl was the reference, and a Pt-wire or a Pt-disk electrode (BASi) was the working electrode. Slightly less data noise was observed with the commercial Pt-disk electrode. CV scans were made in three segments for the BASi and two segments for the student-built instrument: in general, an initial negative-going potential (cathodic) scan was followed by a reverse scan toward the anode and then finally allowed to return to the initial potential in the cathodic region. Data obtained from the BASi instrument was found to be similar to the student-built instrument.

General Procedures

Developed Lab Experiments

The CV experiment is carried out in an electrolytic cell with three electrodes: a student-built reference electrode, the working electrode, which was either a 6-cm-long 2 mm Pt

Experiment 1. Six solutions of K3Fe(CN)6 with concentrations ranging from 0.2 to 5.0 mM were prepared in 1 M KNO3, and used to generate a calibration curve of Ipa (mV) and

Reference Electrode

C | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


Table 2. Data from BASi and Student-Built Instrumenta,b [K3Fe(CN)6] (mM)

Ipa (mA)

Epa (mV)

Ipc (mA)

Epc (mV)

ΔEp (mV)

E° (mV)

−0.0235 −0.0510 −0.0782 −0.1235 −0.2450

190 240 247 247 278

0.0216 0.0431 0.0638 0.1009 0.2083

106 124 120 135 163

84 116 127 112 115

148 182 184 191 221

−0.0246 −0.0649 −0.1410 −0.2575

231 246 275 287

0.0224 0.0547 0.1202 0.2241

133 148 171 191

98 98 104 96

182 197 223 239

BASi 0.25 0.50 1.00 2.00 5.00 Student-built 0.33 0.75 1.50 3.00 a

BASi scan rate, 100 mV/s; student-built instrument scan rate, 95 mV/s. bPotentials are vs Ag/AgCl.

RESULTS AND DISCUSSION The SMILE program that we have developed for high school and undergraduate students provides an opportunity to reinforce the concepts of electrochemistry in a more engaging and realistic manner. The CV took less than 3 h to construct from kit form. The potentiostat employs a Pt working electrode, a Pt counter electrode, and a simple student-built or commercial Ag/AgCl reference electrode. The principle function of the potentiostat is to control and measure the current and potential of the working electrode. The potentiostat is interfaced with a DAQ unit, and an in-house JavaScript program allows signal data to be manipulated into a format that can be transferred to Excel for analysis. The potential of the working electrode is controlled with respect to the Ag/AgCl reference electrode. The reference electrode was constructed by students within 3 min; it is inexpensive, reliable, and gave consistently reproducible data even after four semesters of use. In order to calibrate the electrode potential, the Fe3+/Fe2+ redox couple was examined by a three-electrode BASi Epsilon instrument using a 10 mM K3Fe(CN)6 solution in 100 mM KNO3 as the test. The Ag/ AgCl electrode gave identical data to that obtained from a commercial electrode. The peak potential difference (ΔEp) and peak height (Ip) remained constant for several cycles, indicating that the fabricated electrode was stable and reliable. A minor correction factor9 was calculated and used in all subsequent runs. Each of the 3 h lab experiments was intentionally designed to be of varying technical complexity and sophistication and, thus, can be geared toward the aptitude and experience of students and the instructor. These CV experiments have been successfully conducted with high school, freshman, and senior college chemistry students. Detailed experimental protocols and instructor notes are provided as Supporting Information and offer helpful hints on how to prepare equipment, reagents, solutions, and students. Typical data as generated by our students is also included.

Ipc (mV) vs concentration (mM). From this plot, the concentration of an unknown K3Fe(CN)6 sample was calculated. The E° (mV) value, precision, and accuracy of the Fe(II)/Fe(III) redox couple was then compared to literature values. In a second series of experiments, the scan rate was varied from 100 to 300 mV/s, and its effect on Ip (mA) and Ep (mV) was observed. A plot of ΔEp (mV) vs square root of the scan rate (mV/s) allowed for discussion of the Randles−Sevcik equation. Experiment 2. Four known hydroquinone (H2Q)-buffer solutions between pH 1 and 6 were prepared and analyzed. A calibration curve of pH vs E° (mV) was generated, from which the pH of an unknown H2Q sample was calculated. In addition, the number of electrons transferred and the diffusion coefficient of the electrochemical system were calculated using the Nernst and the Randles−Sevcik equations. Experiment 3. The electrochemical mechanism of 4acetaminophenol (APAP) as it is oxidized to benzoquinone was examined. This transformation was observed by monitoring the shift in the redox potentials and the appearance/ disappearance of reaction intermediates while the pH was varied. A calibration curve of Ipa (mV) vs APAP concentration (mM) at a constant pH of 1.7 was prepared, from which the concentration of APAP in an unknown sample of Tylenol was calculated and compared to the manufacturers claim. Chemicals and CAS numbers. Potassium chloride (744740-7), potassium ferricyanide (13746-66-2), potassium nitrate (7757-79-1), hydroquinone (123-31-9), sodium monobasic phosphate (7558-80-7), sodium dibasic phosphate (7558-79-4), phosphoric acid (7664-38-2), acetaminophen (103-90-2), and Tylenol gelcaps (Extra Strength Rapid Release).


All solutions must be disposed of in designated containers. If skin or eye contact occurs, flush with water immediately. Never leave the working and auxiliary electrodes in buffer solution for long periods of time. The reference electrode should not run dry and can be stored in a saturated NaCl solution. Do not expose the Ag/AgCl electrodes to sunlight as decomposition will blacken the electrodes. Hydroquinone may be toxic by inhalation, skin absorption, and ingestion. Care must be taken when using phosphoric acid, it is extremely corrosive. Students should wear safety glasses and gloves. Acid spills should be neutralized with bicarbonate before disposal, and larger spills can be treated using Vermiculite.

Experiment 1

The ferro/ferricyanide couple was used to illustrate an electrochemically reversible inorganic redox system. It is a well-behaved couple and ideal for introduction of the basic concepts of electrochemistry and CV. Table 2 provides data from the BASi and the student-built instrument. Complete CV plots are provided in the Supporting Information. Calibration curves plotting Ipa and Ipc (mA) vs concentration (mM) for the instruments are shown in Figure 4. These plots D | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


in the Supporting Information, pertinent data is provided in Table 3. Plots of Ip (mA) vs square root of the scan rate (mV/s) are shown in Figure 5. The data illustrates that the redox potentials

Figure 4. Calibration curves for K3Fe(CN)6 solutions (solid lines, positive). Ipc (mA) vs concentration (mM): (solid red line) BASi instrument, R2 = 0.997; (solid blue line) student-built instrument, R2 = 0.999. Also plotted (dashed lines, negative) is Ipa (mA) vs Concentration (mM): (dashed red line) BASi instrument, R2 = 0.994; (dashed blue line) student-built instrument, R2 = 0.998.

Figure 5. Calibration curves for 2 mM K3Fe(CN)6 solutions (solid lines, positive). Ipc (mA) vs square root of the scan rate (mV/s): (solid red line) BASi instrument, R2 = 0.995; (solid blue line) student-built instrument, R2 = 0.988. Also plotted (dashed lines, negative) is Ipa (mA) vs square root of the scan rate (mV/s): (dashed red lines) BASi instrument, R2 = 0.999; (dashed blue line) student-built instrument, R2 = 0.987.

are consistent with the Randles−Sevcik equation, illustrating that the observed current is proportional to the concentration of the analyte. For the student-built instrument, the linear range of response falls between 0.3 and 10 mM. Ferricyanide samples of unknown concentration were calculated from these calibration curves. The formal potential (E°) is the mean of the anodic (Epa) and cathodic (Epc) peak potentials and was found to be 185 mV using the BASi and 210 mV using the student-built instrument vs Ag/AgCl, corresponding to 382 and 407 mV vs SHE, respectively. The lack of iR compensation contributed, in part, to the large peak-to-peak separation (ΔEp) away from the ideal Nernstian10−13 value of 59 mV. The ΔEp is also influenced by the poor quality and imperfections in the surface of the electrodes. The E° is highly dependent on the nature of electrolytes in solution. Literature14−16 provides values of 450 mV to 720 mV vs SHE at 25 °C for this system in a variety of common electrolytes; for a 1 M KNO3 solution, a value of 459 is reported. The E values using the student-built CV compare favorably with the values from the commercial instrument and with those reported in literature. The effect of scan rate on the ferro/ferricyanide redox potential and peak currents was also investigated. For the student-built instrument, the scan rate (mV/s) was adjusted manually using the frequency knob by timing the full voltage sweep and dividing the full cycle time by two. Complete CV plots when the scan rate is varied for both instruments are given

between the two instruments are comparable and that the formal redox potential value (E° or E1/2) is independent of scan rate. Scan rate plots are consistent with the Randles−Sevcik equation; the peak current (Ipa and Ipc) increases with square root of the scan rate and is proportional to concentration of the analyte. The peak-to-peak separations (ΔEp) increase slightly from 110 to 131 mV when the scan rate is changed from 68 to 200 mV/s for the student-built instrument; this is similar to what was observed for the BASi instrument. The value is slightly larger than the Nernstian value of 59 mV (at 25 °C) for a one-electron reversible redox10−13 couple, and predominantly arises from uncompensated Ohmic (iRu) drop and poor quality of the active surface of the electrodes. Overall, the data obtained from student-built instrument is comparable to that from the commercial instrument. Experiment 2

Quinones represent a class of compounds that are widely distributed in nature, and their basic structure is featured in many cofactors, coenzymes, and flavonoids, with numerous biological functions, including cellular respiration, blood

Table 3. Data from BASi and Student-Built Instrument as Scan Rate Is Varied for a 2 mM K3Fe(CN)6 Solutiona Scan Rate (mV/s)

(Scan Rate)1/2

Ipa (mA)

Epa (mV)

Ipc (mA)

Epc (mV)

E° (mV)

4.47 7.07 10.00 12.25

−0.0540 −0.0789 −0.1007 −0.1200

237 239 246 249

0.0609 0.0711 0.0902 0.1043

129 130 125 120

183 185 186 185

8.25 9.27 12.37 14.14

−0.093 −0.110 −0.134 −0.145

261 264 260 271

0.093 0.111 0.134 0.148

151 151 140 140

206 208 200 206

BASi 20 50 100 150 Student-built 68 86 153 200 a

Potentials are vs Ag/AgCl. E | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


coagulation, and photosynthesis.17 The biological action of quinones is linked to their electron transfer rates and redox potentials. The hydroquinone/p-benzoquinone (H2Q/Q) experiment18−27 provided an excellent illustration of the Nernst equation, its relationship to pH, and the features that make this redox couple of such importance to the biological world. Several known hydroquinone-buffer solutions at pH between 2 and 6 were analyzed with both a commercial and a studentbuilt CV. In both cases, a commercial Pt electrode was used as the working electrode. Complete voltammograms for the two instruments are given in the Supporting Information, and pertinent data is provided in Table 4. A typical voltammogram

redox species involved. Multiple scans at each pH show no change in the redox potentials or peak heights. A noticeable decrease in the Ipa as the pH increases was observed in the data from both instruments and is slightly more dramatic in the student-built instrument. We were careful in cleaning the working electrode between each run, so the decrease in Ipa is likely associated with a decrease in the concentration of the reduced species and/or of the diffusion coefficient associated with the species. The higher than expected ΔE values from both instruments is attributable to uncompensated Ohmic drop and imperfections in the student-built electrodes, indicative of the considerable resistance the electrons encounter when traveling between the redox species and electrode surface. We are currently investigating the anodic current drop phenomenon, along with the influence of various possible contaminants and interferents. Calibration curves of pH vs E° (mV) are presented in Figure 7. Plots from the BASi instrument and the student-built

Table 4. CV Data from BASi (2 mM Hydroquinone) and Student-Built Instrument (5 mM Hydroquinone), at Varying Solution pHa pH

Epa (mV)

Epc (mV)

E° (mV)

482 448 405 379 321

206 167 128 113 78

344 308 267 246 200

499 471 436 398

176 103 44 −12

338 287 240 193

BASi 2.0 3.0 4.0 5.0 6.0 Student-built 2.0 3.0 4.0 5.5 a

Potentials are given vs Ag/AgCl.

Figure 7. Calibration curves for hydroquinone-buffered solutions, E° vs pH: (red line) BASi instrument, 2 mM solution, R2 = 0.996, slope = 37 ± 4 mV/pH. (blue line) Student-built instrument, 5 mM solution, R2 = 0.993, slope = 43 ± 4 mV/pH. Potentials are given vs Ag/AgCl.

instrument provide slopes of 37 ± 4 and 43 ± 4 mV/pH, respectively. Although these values are not quite the ideal 59 mV/pH unit as expected10 for a 2-electron, 2-proton redox system, they do clearly show the pH dependence of the H2Q/Q redox potential, that the reaction involves about two electrons, and quality of data from the two instruments is comparable. Additional quantitative information regarding the H2Q/Q couple was obtained from the peak heights according to the Randles−Sevcik equation. Calculation of the surface area of the Pt working electrode (A = 0.071 cm2), together with the solution concentration and scan rate enabled calculation of the diffusion coefficient of the analyte (D). Using the voltammograms from the pH3 solutions (Figure 6), the Ip (mean average) for the BASi instrument was calculated to be 0.1358 mA (2 mM, 100 mV/s, 25 °C), and for the student-built instrument to be 0.3531 mA (5 mM, 76 mV/s, 25 °C). D for H2Q was determined to be 1.58 × 10−5 for the BASi, and 2.25 × 10−5 for the student-built instrument. Although we were unable to find a value in literature dealing with an identical system, these values are comparable and within range (10−4 to 10−6) of those obtained under different solvent and electrolyte conditions.22−29

Figure 6. Typical set of voltammograms. (Red line) BASi instrument, 2 mM hydroquinone at pH 3.0. Scanned from 0 mV to 1000 mV, then to −600 mV, and back to 0 mV, at scan rate of 100 (mV/s). (Blue line) Student-built instrument, 5 mM hydroquinone at pH 3.0. Scanned from 800 mV to −200 mV, and back to 800 mV, at scan rate of 74 (mV/s).

at pH 3 (Figure 6) shows one cathodic peak in the negativegoing scan attributed to the reduction of Q to H2Q and one coupled anodic peak in the positive-going scan corresponding to the oxidation of H2Q to Q. Data from both instruments are remarkably similar. As the pH of the buffer solution increases, the anodic and cathodic peak potentials shift to the more negative, indicating a decrease in the redox potential as predicted by the Nernst equation. The observed cathodic shift illustrates that the redox potential is dependent on the equilibrium concentration of the F | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


Experiment 3

In both cases, a commercial Pt working electrode was employed. Overall, the data is satisfyingly similar between the two instruments; although the Ipa values appear distorted, the key features are apparent and comparable. The voltammogram from the student-built instrument contains some additional features that arise from the wider potential window. The data illustrates that as the pH is varied, chemical reactions involving APAP can be mapped out. At high pH (>6) a reversible redox process is observed; the N-acetyl-pquinonimine (species 2 in Scheme 1) is stable, electroactive, and can be readily converted back to APAP by reversing the sweep back to negative potentials. The large separation is in part due to the sluggish electron transfer kinetics and is exacerbated by the poor quality of the electrodes. As the pH is raised toward 12, the oxidation process becomes less favorable. As pH is decreased, the concentration of 1 steadily declines, a change that is apparent in the cathodic current (Ipc). At acidic pH, the redox wave shifts to a more positive potential. As pH drops below 3, the reduction (cathodic) waves are reduced due to rapid protonation and hydration of 2 to yield the electrochemically inactive 4 (Scheme 1). The fast scan rates used in the experiment also allow small amounts of 2 to be observed at lower pH values; even at pH 2, some features that are attributable to the initial redox couple were observed. In extremely acidic medium, the acetamide moiety is chemically lost to form the final the benzoquinone product, 5. This mechanism is described as being “ec”, where an electron transfer step produces a species that then undergoes a chemical reaction. At such low pH, the redox couple associated with Q/ H2Q is observed, albeit poorly defined. A standard pure solution of benzoquinone/hydroquinone was used to verify these assignments. The second part of this experiment involved use of the Randles−Sevcik equation to construct a calibration curve of Ipa (mV) vs concentration (mM), at a constant pH of 1.7. This calibration was used to identify the concentration of an unknown APAP solution at this same pH. The Ipc is dependent on the solution pH; however, the Ipa value was found to be proportional to the concentration of the benzoquinone species (5) and so is ideal for the quantitative analysis of APAP (1) in unknown samples (Figure 9). From the calibration curves, the concentration of APAP in a sample of Tylenol Extra Strength was calculated to be 5.00 mM, thus verifying the claim on the

This experiment dealt with the chemical and electrochemical mechanism of 4-acetaminophenol (APAP, paracetamol, acetaminophen) as it is oxidized to benzoquinone (Scheme 1). Scheme 1. Proposed Mechanism for the Oxidation of APAP

APAP is selected not only on the basis of its redox properties but also because it is cheap and readily available. Many techniques have been developed for the quantitative determination of APAP in pharmaceutical and biological samples,30−32 and because the molecule is readily oxidized at low potentials, electroanalytical methods33,34 are particularly well suited. The experiment described herein is a modified version of that in literature.14,33 The oxidation of APAP to benzoquinone is pH dependent. The student-built CV monitored the shift in the redox potential, and the appearance/disappearance of intermediates as the pH was varied. The data generated at differing pH is displayed alongside that from the BASi instrument in Figure 8.

Figure 8. (A) Data from BASi instrument for varying pH of APAP solutions (5 mM). Scan segments were from 900 mV to 0 mV, and then back to 900 mV, at a scan rate of 100 (mV/s). (B) Data from the student-built instrument for varying pH of APAP solutions (5 mM). Scan segments were from 1500 mV to −500 mV, and then back to 1500 mV, at a scan rate of 102 (mV/s). Potentials are given vs Ag/ AgCl.

Figure 9. Calibration curves for APAP-buffered pH 1.7 solutions, Ipa vs concentration (mM): (red line) BASi instrument, R2 = 0.988. (blue line) Student-built instrument, R2 = 0.996. G | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


redox process, the relationship between current and potential using the CV technique, and the importance of calibration curves and statistical analysis of data far more clearly now than prior to the SMILE program.

label and the accuracy and reliability of the student-built instrument. Student Assessment and Feedback

We have conducted an informal assessment of the SMILE program within the instrumental analysis course. Over three semesters, half of the students built the CV and others built a conductivity detector. Although each group was assigned a specific instrument to build, every group performed experiments with both types of instruments. Students were then evaluated on their written responses to two questions on a final exam; the first question probed understanding of basic CV concepts discussed in lecture and reinforced in lab. Both groups performed equally well, receiving an average of 85%. The second question is not explicitly covered in class and probes understanding of how electrochemical signals are generated and measured. The students in the CV groups performed substantially better than the students who built the conductivity detector, 83% vs 71%, respectively. One factor that may contribute to such disparate results is whether students assigned the CV project inherently asked and researched the question on their own. Student feedback relating to the use of the CV instrument was also collected from our general chemistry students (Supporting Information), with students highly pleased with consistency of data and ease-of-use of the CV instrument. In addition, a more detailed postlab survey was conducted regarding use of such developmental instruments in the course, and suggestions were solicited on how the CV could be improved (Supporting Information). Overall, the students were very supportive of the SMILE program, finding it conducive to their studies and being collaborators with an ongoing research project. The instruments were described as being oversensitive, data as being noisy, and the transfer of data to a plotting program as being cumbersome. However, many constructive suggestions on how to overcome these issues and thus make the instrument better for future students was freely offered.


S Supporting Information *

Instructor notes and student handouts for the three experiments; student feedback. This material is available via the Internet at


Corresponding Author

*D. G. Sykes. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

ACKNOWLEDGMENTS The authors thank the staff of the Electronics Instrumentation Facility at Penn State. The SMILE project relies on student effort and feedback, and we are grateful for the time and patience that our undergraduate, high school, and middle school participants have given while developing the CV. We thank the Penn State Schreyer Institute for Teaching Excellence, and the Summer Experience program in the Eberly College of Science (SEECoS) for financial support. SEECoS is supported by the Upward-Bound Math and Science Center (UBMS) at Penn State, and a U. S. DOE TRIO grant.


(1) Dominguez, V. C.; McDonald, C. R.; Johnson, M.; Schunk, D.; Kreuter, R.; Sykes, D.; Wigton, B. T.; Chohan, B. S. The Characterization of a Custom-Built Coulometric Karl Fischer Titration Apparatus. J. Chem. Educ. 2010, 87, 987−991. (2) McDonald, C.; Johnson, M.; Schunk, D.; Kreuter, R.; Sykes, D.; Wigton, B.; Chohan, B. A Portable, Low-Cost, LED Fluorimeter for Middle School, High School, and Undergraduate Chemistry Labs. J. Chem. Educ. 2011, 88, 1182−1187. (3) Wigton, B.; Kreuter, R.; Sykes, D.; Chohan, B. The Characterization of an Easy-to-Operate Inexpensive Student-Built Fluorimeter. J. Chem. Educ. 2011, 88, 1188−1193. (4) Stewart, G.; Kuntzleman, T. S.; Amend, J. R.; Collins, M. J. Affordable Cyclic Voltammetry. J. Chem. Educ. 2009, 86, 1080−1081. (5) Ahn, M. K.; Reuland, D. J.; Chadd, K. D. Electrochemical Measurements in General Chemistry Lab Using a StudentConstructed Ag-AgCl Reference Electrode. J. Chem. Educ. 1992, 69, 74−75. (6) Sawyer, D. T.; Sobkowiak, A.; Roberts, J. L., Jr. Electrochemistry for Chemists, 2nd ed.; Wiley: New York, 1995; pp 189−190. (7) Kissinger, P. T.; Heineman, W. R. Laboratory Techniques in Electroanalytical Chemistry, 2nd ed.; Dekker: New York, 1996; pp 98− 100. (8) Inamdar, S. N.; Bhat, M. A.; Haram, S. K. Construction of Ag/ AgCl Reference Electrode from Used Felt-Tipped Pen Barrel for Undergraduate Laboratory. J. Chem. Educ. 2009, 86, 355−356. (9) Smith, T. J.; Stevenson, K. J. Handbook of Electrochemistry, 1st ed.; Zoski, C. G., Ed.; Elsevier Press: London, 2007. (10) Meites, L. Polarographic Techniques, 2nd ed.; Wiley: New York, 1965. (11) Bond, A. M. Modern Polarographic Methods in Analytical Chemistry; Marcel Dekker: New York, 1980; pp 27−45. (12) Bard, A. J.; Faulkner, L. R. Electrochemical Methods. Fundamentals and Applications, 2nd ed.; Harris, D., Swain, E., Robey, C., Aiello, E., Eds.; John Wiley and Sons: New York, 2001; p 241.

CONCLUSIONS The design and construction of a miniature scanning potentiostat that can generate reliable CV data at a fractional cost of a commercial instrument was described. The performance and accuracy of the student-built instrument have been successfully verified via a series of experiments of varying complexity. Excellent current response data corresponding to the change of analyte concentration was achieved for the ferricyanide redox couple, the hydroquinone couple, and the 4acetaminophenol experiments. Accordingly, linear calibration curves were obtained, and the precise concentration (or pH in case of the hydroquinone experiment) of unknown samples was obtained. Although the voltammograms obtained from the studentbuilt potentiostat were not as pretty as those of the commercial instrument, the data generated was reproducible. The ability to achieve precise amperometric detection of analytes of organic and inorganic systems using a three-electrode CV protocol was confirmed. Overall, the student-built potentiostat has the merits of accuracy, small size, low weight, portability, and very low cost. Student assessment and feedback suggests that the novel SMILE initiative is a pedagogical success in providing a rigorous, engaging, and memorable classroom and lab experience. Gathering data with this particular instrument has enabled our students to understand the concept behind the H | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education


(13) Nicholson, R. S.; Shain, I. Theory of stationary electrode polarography: single scan and cyclic methods applied to reversible, irreversible, and kinetic systems. Anal. Chem. 1964, 36, 706−723. (14) Van Benschoten, J. J.; Lewis, J. Y.; Heineman, W. R.; Roston, D. A.; Kissinger, P. T. Cyclic Voltammetry Experiment. J. Chem. Educ. 1983, 60, 772−776. (15) Sharpe, A. G. The Chemistry of Cyano Complexes of the Transition Metals; Academic Press: London, 1976. (16) Bott, A. W.; Jackson, B. P. Study of ferri-cyanide by cyclic voltammetry using the CV-50W. Curr. Sep. 1996, 15, 25−30. (17) Thomson, R. H. Naturally Occurring Quinones IV. Recent Advances; Blackie Academic: London, 1997. (18) Skoog, D. A.; West, D. M.; Holler, F. J. Fundamentals of Analytical Chemistry, 6th ed.; Saunders College Publishing: New York, 1992; Appendix 6. (19) Walczak, M. W.; Dryer, D. A.; Jacobson, D. D.; Foss, M. G.; Flynn, N. T. pH-Dependent Redox Couple: Illustrating the Nernst Equation using Cyclic Voltammetry. J. Chem. Educ. 1997, 74, 1195− 1197. (20) Wipf, D.; Wehmeyer, K. R.; Wightman, R. M. Disproportionation of quinone radical anions in protic solvents at high pH. J. Org. Chem. 1986, 51, 4760−4764. (21) Lammert, O. M.; Livingston, J.; Morgan, R. The Quinhydrone Electrode IV. J. Am. Chem. Soc. 1932, 54, 910−918. (22) Kolthoff, I. M.; Orlemann, E. F. The Use of Dropping Mercury Electrode as an Indicator Electrode in Poorly Poised Systems. J. Am. Chem. Soc. 1941, 63, 664−667. (23) Hemingway, A. A Direct-Reading pH Meter for Glass, Quinhydrone, and Hydrogen Electrodes. Ind. Eng. Chem., Anal. Ed. 1935, 7, 203−205. (24) Rosenthal, R.; Lorch, A. E.; Hammett, L. P. The Kinetics of the Quinhydrone Electrode Reaction. J. Am. Chem. Soc. 1937, 59, 1795− 1804. (25) Bates, R. G. Electrodes for pH measurement. J. Electroanal. Chem. 1961, 2, 93−109. (26) Aquino-Binag, C.; Pigram, P. J.; Lamb, R. N.; Alexander, P. W. Surface Studies of Quinhydrone pH Sensors. Anal. Chim. Acta 1994, 291, 65−73. (27) Eggins, B. R. Evidence for a one-electron intermediate in the anodic oxidation of hydroquinone in acetonitrile. Chem. Commun. 1972, 427. (28) Haimerl, A.; Merz, A. Catalysis of quinone-hydroguinone redox reactions at polypyrrole benzenesulphonate-coated platinum electrodes. J. Electroanal. Chem. 1987, 220, 55−65. (29) Peng, J.; Gao, Z.-N. Influence of micelles on the electrochemical behaviors of catechol and hydroquinone and their simultaneous determination. Anal. Bioanal. Chem. 2006, 384, 1525−1532. (30) Bosch, M. E.; Sanchez, A. J. R.; Rojas, F. S.; Ojeda, C. B. Determination of paracetamol: Historical evolution. J. Pharm. Biomed. Anal. 2006, 42, 291−321. (31) Fenk, C. J.; Hickman, N. M.; Fincke, M. A.; Motry, D. H.; Lavine, B. Identification and Quantitative Analysis of Acetaminophen, Acetylsalicylic Acid, and Caffeine in Commercial Analgesic Tablets by LC-MS. J. Chem. Educ. 2010, 87, 838−841. (32) Kagel, R. A.; Farwell, S. O. Analysis of Currently Available Analgesic Tablets by Modern Liquid Chromatography: An Undergraduate Laboratory Introduction to HPLC. J. Chem. Educ. 1983, 60, 163−165. (33) Miner, D. J.; Rice, J. R.; Riggin, R. M.; Kissinger, P. T. Voltammetry of Acetaminophen and Its Metabolites. Anal. Chem. 1981, 53, 2258−2263. (34) Wang, C.; Li, C.; Wang, F.; Wang, C. Covalent Modification of Glassy Carbon Electrode with L-Cysteine for the Determination of Acetaminophen. Microchim. Acta 2006, 155, 365−371.

I | J. Chem. Educ. XXXX, XXX, XXX−XXX