In the Laboratory
Inexpensive and Disposable pH Electrodes Michael J. Goldcamp,* Alfred Conklin, Kimberly Nelson, Jessica Marchetti, Ryan Brashear, and Emily Epure Department of Chemistry, Wilmington College, Wilmington, Ohio 45177 *
[email protected] Electrodes for the measurement of pH are used in a wide variety of laboratory, field, and classroom environments. Commercially available pH electrodes are available from a variety of manufacturers at costs typically from about $30 up to about $500. These electrodes generally have long lifetimes but are subject to varying degrees of damage or improper storage. This sometimes makes their frequent use at all levels across the curriculum impractical. A number of different “homemade” electrodes are reported in the literature (1-3), and these often provide simple, effective means of allowing the use of electrodes without concern for the cost of electrode destruction. Additionally, inexpensive and disposable electrodes are ideally suited for field measurements. Ion-selective electrodes, including pH electrodes, operate by the generation of an electrical potential across a barrier, often a membrane, by diffusion of ions into or across the barrier (4, 5). This potential is described by the Nernst equation: E ¼ E0 þ ð2:303RT =nF Þ log aHþ where R is the ideal gas constant, F is the Faraday constant, T is the temperature, n is the number of electrons transferred (n = 1 for hydrogen ion), and aHþ is the activity of hydrogen ions. At 25 °C, the equation becomes E ¼ E0 þ ð59:1 mVÞpH such that the slope of the plot of potential versus pH is ideally 59.1 mV/pH unit. For selectivity, the membrane must allow absorption or passage of only the ion that is to be measured (Hþ for pH electrodes). Often this is achieved by incorporation of an ion-specific ligand, or ionophore. For a pH electrode, monodentate tertiary amines, especially those with substantial steric hindrance, should have high affinity for binding protons while having low affinity for coordination with metal ions due to steric factors and the inability to form chelate rings. Many tertiary amines have been shown to function well as ionophores for Hþ (6-9). A facile procedure for the construction of pH electrodes is described. Both traditional liquid-membrane electrodes (with plastic bodies and internal filling solutions) and coated-wire electrodes (10) (CWEs) using both metal wires and graphite pencil-lead “wires” (11-13) have been constructed. These pH electrodes utilize poly(vinyl chloride), PVC, membranes containing tribenzylamine as the ionophore (9). From this procedure, many electrodes can be produced with costs on the order of dollars per electrode. These low-cost pH electrodes provide for cost-effective use of pH electrodes across the science curriculum, including fieldwork. Additionally, instructors could adapt the procedures to develop instructional methods to teach the theory, construction, and operation of ion-selective electrodes by having 1262
Journal of Chemical Education
_
_
students construct and use their own pH electrodes in the laboratory. Experimental Section Materials, Equipment, and Chemicals Two Orion (Thermo Scientific, Beverley, MA) model 710A pH/ISE meters, one equipped with an Orion combination pH electrode and the other with the constructed electrode, were used to measure pH values of solutions. Poly(tetrafluoroethylene), PTFE, tubing (1/16 in. i.d.) was obtained from Restek (Bellefonte, PA). Silver wire and platinum wire (0.125 μm diameter) and gold metal pin connectors were obtained from A-M Systems, Inc. (Carlsborg, WA). An Ag/AgCl reference electrode was obtained from Cypress Systems, Inc. (Chelmsford, MA) or one may be prepared according to a procedure in the literature (3). High molecular weight poly(vinyl chloride), PVC, o-nitrophenyl octyl ether (NPOE), potassium tetrakis(4-chlorophenyl)borate (KTpClPB), tribenzylamine, and tetrahydrofuran (THF) were obtained from Aldrich (Milwaukee, WI). Concentrated nitric acid was obtained from Pharmco (Dallas, TX). Sodium hydroxide was obtained from Fisher Scientific (Hampton, NH). Standard solutions of pH 4, 7, and 10 for pH meter calibrations were obtained from Orion (Thermo Scientific, Beverley, MA). Preparation of Membrane Solution A solution of the membrane components was prepared by dissolving a mixture of the following quantities (or larger quantities in the same ratios) in a minimal volume of THF solvent (typically 1-2 mL): 0.035 g of PVC; 0.070 g of o-nitrophenyl octyl ether; 0.010 g of potassium tetrakis(4-chlorophenyl)borate; and 0.010 g of tribenzylamine. These quantities typically made 20-30 electrodes. Construction of the Liquid-Membrane Electrode PTFE tubing (∼1.5 mm i.d., 4-5 cm long) makes an excellent electrode body as it is relatively inert and flexible. To form the membrane in the tube, one end of the tube was dipped into the membrane solution until a plug of solution (2-4 mm thick) filled the end of the tube. The tube was allowed to dry for 24 h in an upright position. A silver wire was soldered to a metal pin on one end, and the other end of the wire was cut so that, when placed inside the electrode body, it comes near to but does not contact the membrane in the end of the electrode body. A piece of rubber tubing can be glued to the wire near the pin so that it snugly fits over the electrode body if desired; however, removal of the wire-pin assembly is necessary to replace the solution in the inner compartment periodically. The silver wire
_
Vol. 87 No. 11 November 2010 pubs.acs.org/jchemeduc r 2010 American Chemical Society and Division of Chemical Education, Inc. 10.1021/ed1005008 Published on Web 08/30/2010
In the Laboratory
Figure 1. Diagram and photograph of a constructed electrode.
was coated with silver chloride prior to usage (5). Finally, the electrode body was filled with a solution of a buffer of pH 4 or similar, and the wire was inserted into the body to complete the assembly (Figure 1). The electrode was soaked in the pH 4 buffer for 24 h before use and was stored in the same solution.
Figure 2. Comparison of pH values measured simultaneously using a commercial pH electrode (x axis) and the constructed pH electrode (y axis). The solid line represents the linear fit to the data, and the dashed line represents the theoretical x = y relationship for identical responses.
Construction of the Coated-Wire Electrodes (CWEs) Simple CWEs were constructed using the same membrane solution as described above. A suitable wire was chosen; platinum wires, graphite rods, or graphite pencil-lead refills worked well. To expose only a portion of the wire, it was cemented into an inert, nonconducting housing, such as PTFE tubing, with about 5 mm of the wire surface extending out of each end of the tubing. If a rigid “wire” such as graphite is used, it is convenient to cement electrical wiring to one end of the electrode using conducting epoxy for easier connection to the meter. One exposed end of the wire was dipped into the membrane solution so that it was coated with the membrane. After drying, a second dipping may be necessary to ensure complete coverage of the wire. The membrane was dried overnight. It was then stored in buffer (pH 4 or similar) in the same manner as the liquid-membrane electrode. Hazards Nitric acid is a strong acid that can cause severe burns. Sodium hydroxide is a strong base that is caustic and can cause severe burns. These acids and bases are hazardous and should be handled with care. When diluting concentrated acids, always remember to add the concentrated acid to water, not the reverse. Tetrahydrofuran is volatile, highly flammable, and may contain peroxides. Peroxides are explosive. They are easily formed from tetrahydrofuran and air upon exposure to ambient light and may accumulate up to dangerous concentration when tetrahydrofuran evaporated. Tribenzylamine, o-nitrophenyl octyl ether, and potassium tetrakis(4-chlorophenyl)borate may cause irritation to skin, eyes, and respiratory tract and may be harmful if swallowed or inhaled. Results and Conclusions Electrodes prepared by the above procedure, when used in conjunction with a separate reference electrode such as Ag/AgCl, are remarkably accurate when compared to a commercially available Orion pH electrode. After calibrating two pH meters (one with the Orion electrode and one with the constructed electrode) with pH 4, 7, and 10 standards, the electrodes were placed in the same starting pH 7 buffer, and the pH was varied by
r 2010 American Chemical Society and Division of Chemical Education, Inc.
_
Figure 3. The plot of pH (from commercial electrode) versus potential (from constructed electrode) shows linear response with a slope of -55 mV per pH unit.
adding aliquots of dilute HNO3 and dilute NaOH, sweeping several times from pH 2.5 to pH 10. As is typical for many measurements of pH, no effort was made to control ionic strength or activity of the test solution. For one typical data set, simultaneous pH measurements using the “homemade” electrode and the commercial Orion electrode (Figure 2) had an average difference of about 0.3 pH units (σ = 0.1; n = 35). After a series of measurements, each electrode was placed back into the pH 4, 7, and 10 calibration standards, and the pH values were measured using both the constructed and commercial electrode. It is notable that, of the two electrodes, the constructed electrode typically gave measured pH values closer to the known pH values of the buffers. The coated-wire pH electrodes performed similarly. For a typical data set from a graphite pencil-lead CWE, the average difference was 0.2 pH units between simultaneous measurements from the homemade and commercial electrodes (σ = 0.1; n = 35). When examining the change in electrochemical potential measured by the electrode as pH was varied, these constructed pH electrodes exhibited nearly Nernstian responses. The absolute
pubs.acs.org/jchemeduc
_
Vol. 87 No. 11 November 2010
_
Journal of Chemical Education
1263
In the Laboratory
values of the slopes of the plots of potential versus pH (Figure 3) ranged from about 50 to 55 mV/pH unit, which are close to the theoretical value of 59.1 mV/pH unit (at 25 °C). A number of common metal cations were tested for interference with the pH measurement by the electrodes. Aliquots of separate solutions containing Liþ, Naþ, Kþ, Mg2þ, Ca2þ, Ba2þ, Ni2þ, Cu2þ, Zn2þ, or Fe3þ were added to determine the effect on pH measurements. For the cations tested, there was a variation in the pH value of less than 0.1 pH units, even after addition to the test solution of concentrations of the interfering cations up to 0.1 M. These constructed pH electrodes have some limitations. The working range of the electrodes is about pH 3 to 10. Potentiometric response deteriorates outside those values. Under these conditions, these electrodes function for up to 2 months before their response degrades. Some electrodes still produced linear responses after 7 months, but with decreased slopes of only about 30-40% of the ideal Nernstian slope. Some electrodes exhibit decreased response if stored in direct sunlight, suggesting possible photodecomposition of membrane components or photoreduction of Agþ; storage in the dark prolongs lifetime. The electrode lifetime is also usually decreased if it is used in abrasive samples that may physically damage the membrane. Student-constructed electrodes have been successfully employed for the measurement of the pH of soil samples. Although the particulate matter of the soil when mixed with water significantly limited the lifetime of the electrodes, the pH values obtained for the soil samples were consistent with values obtained using traditional commercial electrodes.
1264
Journal of Chemical Education
_
Vol. 87 No. 11 November 2010
_
Acknowledgment Funding for this work was provided by the Instructional Development and Resources Committee of Wilmington College. Literature Cited 1. Scholz, F.; Steinhardt, T.; Kahlert, H.; Porksen, J. R.; Behnert, J. J. Chem. Educ. 2005, 82, 782–786. 2. Marafie, H. M.; Shoukry, A. F. J. Chem. Educ. 2007, 84, 793– 796. 3. East, G. A.; de Valle, M. A. J. Chem. Educ. 2000, 77, 97.4. 4. Wang, J. Analytical Electrochemistry, 2nd ed.; Wiley-VCH: New York, 2000. 5. Fry, C. H.; Langley, S. E. M. Ion-Selective Electrodes for Biological Systems; OPA/Harwood Academic Publishers: Amsterdam, 2001. 6. Schulthess, P.; Shijo, Y.; Pham, H. V.; Pretsch, E.; Ammann, D.; Simon, W. Anal. Chim. Acta 1981, 131, 111–116. 7. Erne, D.; Schenker, K. V.; Ammann, D.; Pretsch, E.; Simon, W. Chimica 1981, 35, 178. 8. Wakida, S.; Ohnishi, M.; Yamane, M.; Higashi, L.; Liu, J.; Wub, X.; Zhang, Z. Sens. Actuators, B 2000, 66, 153–155. 9. Han, W.-S.; Park, M.-Y.; Chung, K.-C.; Cho, D.-H.; Hong, T.-K. Electroanalysis 2001, 13, 955–959. 10. Martin, C. R.; Freiser, H. J. Chem. Educ. 1980, 57, 512–514. 11. Bond, A. M.; Mahon, P. J.; Schiewe, J.; Vicente-Beckett, V. Anal. Chim. Acta 1997, 345, 67–74. 12. Bendikov, T. A.; Harmon, T. C. J. Chem. Educ. 2005, 82, 439–441. 13. Goldcamp, M. J.; Underwood, M. N.; Cloud, J. L.; Harshman, S. J. Chem. Educ. 2008, 85, 976–979.
pubs.acs.org/jchemeduc
_
r 2010 American Chemical Society and Division of Chemical Education, Inc.
