Conductivity through Polymer Electrolytes and Its Implications in

Laboratory Experiment pubs.acs.org/jchemeduc

Conductivity through Polymer Electrolytes and Its Implications in Lithium-Ion Batteries: Real-World Application of Periodic Trends Owen C. Compton,† Martin Egan,‡ Rupa Kanakaraj,§ Thomas B. Higgins,‡ and SonBinh T. Nguyen*,† †

Department of Chemistry, Northwestern University, Evanston, Illinois 60208, United States Physical Science Department, Harold Washington College, Chicago, Illinois 60601, United States § Prosser Career Academy, Chicago, Illinois 60639, United States ‡

S Supporting Information *

ABSTRACT: Periodic conductivity trends are placed in the scope of lithium-ion batteries, where increases in the ionic radii of salt components affect the conductivity of a poly(ethyleneoxide)-based polymer electrolyte. Numerous electrolytes containing varying concentrations and types of metal salts are prepared and evaluated in either one or two laboratory sessions, requiring cooperation between all students in the classroom. The experiment is suitable for either high school students with a general chemistry background or undergraduate students in introductory general chemistry courses, as a number of fundamental topics can be discussed with this simple, inexpensive, and real-world-oriented project.

KEYWORDS: High School/Introductory Chemistry, First-Year Undergraduate/General, Laboratory Instruction, Physical Chemistry, Hands-On Learning/Manipulatives, Inquiry-Based/Discovery Learning, Applications of Chemistry, Conductivity, Electrolytic/Galvanic Cells/Potentials, Periodicity/Periodic Table

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Although LIBs are used daily by almost every member of our society, the simple chemical processes involved in their operation are rarely discussed in the general science curriculum as connections to the underlying principles that govern battery function (i.e., oxidation−reduction reactions) are not easily made. Exploring LIBs in the laboratory setting introduces additional challenges given the hazards associated with the liquid electrolytes (flammable organic solvents) and electrodes (toxic transition-metal oxides) commonly employed in these remarkable energy-storage devices.3 Fortunately, such issues can be circumvented by investigating safe and stable aqueous solutions of polymer electrolytes, comprising small quantities of lithium salts in aqueous poly(ethylene oxide) (PEO) solutions, which can serve as a vehicle to introduce LIBs to high school and undergraduate students through the fundamental topic of ionic radii and periodic trends. The chemistry behind battery functionality is commonly introduced to general chemistry students via the oxidation and reduction of electrodes in a primary cell. Although such electrochemical reactions can be visualized in the laboratory (e.g., the plating of a zinc electrode), the concepts associated with the generation of electricity via these reactions can be abstract to beginning students. Here, LIBs are much more accessible as an introduction to energy storage in the classroom, as their means for electricity generation (and recharging) relies

ince the turn of the 21st century, lithium-ion batteries (LIBs) have become ubiquitous in modern society with nearly all students relying on these energy sources to power their laptop computers, cellular phones, and digital media players. The function of these rechargeable power sources relies on the back-and-forth migration of lithium ions between two electrodes through a conductive medium (Figure 1).1,2

Figure 1. A schematic of a lithium-ion battery, illustrating the critical role that the electrolyte plays during the migration of lithium ions from anode to cathode during discharge. A separator is present in the functioning battery, which serves a role similar to a salt bridge in a conventional primary cell. No separator is utilized in the experiment carried out in this manuscript, as only the conductivity and resistivity of the electrolyte are studied and a functioning battery is not present. © 2012 American Chemical Society and Division of Chemical Education, Inc.

Published: August 22, 2012 1442

dx.doi.org/10.1021/ed200478d | J. Chem. Educ. 2012, 89, 1442−1446

Journal of Chemical Education

Laboratory Experiment

upon the shuttling of ions and electrons between two intercalation electrodes. Such a concept would be easy to grasp by nonchemistry-major students as the movement of charge within the battery can be demonstrated in the classroom.4,5 Given the aforementioned hazards associated with many LIB components, the safest means to introduce these power sources in the laboratory is through the preparation of polymer electrolytes. Polymer electrolytes have been investigated as a component of LIBs since the 1970s to reduce use of flammable organic solvents, while still efficiently conducting charge.6,7 They also provide functional advantages in allowing for the fabrication of flexible, lightweight batteries with versatile shapes and designs. As such, the investigation of polymer electrolytes is an ideal venue to connect real-world applications to the chemistry laboratory. Herein, we describe an easily deployed experiment for introducing LIBs to general chemistry students as they individually, or in small groups, prepare a series of electrolytes containing varying types and concentrations of lithium salts in predissolved aqueous solutions containing PEO. Whereas the presence of water in a LIB would yield a nonfunctional battery (safety warning: the content of a LIB will react violently with water), the dissolution of the salts in the PEO solution is critical to this experiment and water is used as a substitute in light of the hazards and waste associated with organic solvents. The instructor should explain that this experiment focuses on a simulated electrolyte for a LIB and not an actual functioning battery. In addition, the instructor should make it clear to the students that they should not attempt to open or disassemble a LIB or any alkaline battery. The activities described are suitable for the laboratory section of general chemistry in high school classrooms and introductory undergraduate courses, requiring two 45 min periods (one session to introduce the material via discussion and a directed exercise and a second session entailing an inquiry-based exercise) that can be combined into one 90 min period. Connections between the periodic trend of ionic radius and conductivity can be utilized to explain variations in performance of the electrolytes containing different types and concentrations of salts. Students compile and perform statistical analysis on results from an introductory experiment prior to an inquiry-based exercise where they test a hypothesis of their own design.

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Figure 2. (A) Digital image of an experimental setup where the conductivity or resistivity of a pure PEO-based electrolyte is measured. The electrodes are separated by a Styrofoam spacer to limit variance in conductivity or resistivity measurements. (B) A close-up image of the electrodes submerged in the electrolyte solution. (Any discernible unevenness in the electrolyte surface, or bending of the electrodes, in this image is due to light refraction and reflection when the photograph was takenthe surface of the gel is perfectly flat in reality.) Any variety of containing vessel can be utilized, but vessels of similar dimensions should be utilized for the whole class to promote reproducibility between the measurements of different laboratory groups.

proximity of the electrodes. All students, or lab groups, monitor the conductivity and resistivity of their solutions to allow for statistical analysis of the results. Connections between the conductivity of each solution and the ionic radius of the anion can be drawn from these results. To expedite the experiment, the students can measure only the conductivity of the electrolyte, given that resistivity is reciprocal to conductivity; however, obtaining both conductivity and resistivity values independently and comparing the two sets of data could generate a fruitful discussion regarding unit conversions and the accuracy of related measurements. This experiment can also be performed with the resistivity-measuring functionality standard on all multimeters, as some multimeters do not have a conductivity-measuring functionality. In addition to monitoring the lithium-salt-containing electrolytes, as described above, and explaining the connection between conductivity and resistivity and the ionic radii of the anions in the dissolved salts, a complementary experiment can be performed to vary the size of the cations in a series of dissolved salts. Again, PEO-based electrolyte solutions (10 g in 15 mL water) are prepared containing NaCl and KCl (0.1 M), with LiCl having already been tested. The conductivity and resistivity of these solutions can be monitored in the same fashion as described above for those electrolytes containing lithium salts. Upon statistical analysis of these results, further discussion can be introduced to emphasize the importance of ionic radius on the performance of LIBs.

EXPERIMENT

Relating Conductivity to Periodic Trends in Ionic Radius

The primary focus of this experiment is to demonstrate that varying the ionic radii of dissolved salts can improve the conductivity of a polymer electrolyte solution. The conductivity of a control sample comprising only dissolved PEO (20 kD, 10 g in 15 mL water) is compared to that of three different PEObased electrolytes additionally containing either LiCl, LiBr, or LiI at a concentration of 0.1 M (see the Supporting Information for preparation details). Once all components are dissolved, the resulting electrolyte is poured into a series of containing vessels having uniform diameters (e.g., centrifuge tubes, Petri dishes, beakers, etc.). A multimeter is used to monitor the conductivity and resistance of these solutions by placing the positive and negative electrodes at opposite ends of the containing vessel (Figure 2). Electrodes should be pressed against the side of the containing vessel, or punctured through a piece of Styrofoam or cardboard (serving as a consistent spacer) to mitigate errors in conductivity and resistivity measurement due to variation in the

Inquiry-Based Optimization of the Electrolyte

The second half of this laboratory experiment entails an inquiry-based exercise, where students explore the electrical properties of polymer electrolytes by manipulating different variables. The experimental setup for the conductivity and resistivity measurements is identical to that described above (Figure 2). Students set a dependent variable and then 1443




manipulate a second, independent variable while monitoring the conductivity and resistivity values for all samples. A number of variables can be investigated, including the concentration of the dissolved salts, the concentration of PEO, or the volume of the electrolyte. If available, other salts with different combinations of cation or anion can be employed in the electrolyte, further emphasizing the connection between the conductive and resistive properties of the electrolyte with the ionic radii of its components. Students write their own procedure for the experiment, analyze their results, and provide a discussion explaining the motivation behind their chosen variable and an assessment of its effect on the polymer electrolyte. These results are shared with the other students, who then collaborate to determine the composition of an ideal electrolyte.

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HAZARDS Lithium salts and PEO are irritants upon inhalation, ingestion, or direct contact with the skin. Lithium chloride is mutagenic, possibly tetrogenic, and may be toxic to the human reproductive system. Lithium iodide may be toxic to the kidneys and central nervous system. Sodium chloride and potassium chloride are commonly used in the kitchen, with the latter being used by high-blood-pressure patients as a substitute for table salt, so should be quite safe. However, to avoid accidental ingestion or exposure, all students must wear protective gloves (the least expensive is disposable latex gloves that can be easily obtained at a drug store) and eye protection when handling chemicals. If additional alkali metal salts are utilized during the inquiry-based portion of the experiment, material safety data sheets for each chemical should be examined before handling by students is allowed.

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Figure 3. Plot of the (A) conductivity and (B) resistivity of a PEObased aqueous electrolyte containing 0.1 M of a series of lithium salts (LiCl, LiBr, and LiI) as a function of the ionic radii of the anions. Plot of the (C) conductivity and (D) resistivity of a PEO-based aqueous electrolyte containing 0.1 M of a series of chlorine salts (LiCl, NaCl, and KCl) as a function of the ionic radii of the cations. These data points represent those collected by one pair of students and are representative of those values collected by all other student pairs in the classroom.

decrease and increase, respectively, with increased cation radii, as has been observed for other solutions of alkali chlorides.10 This observation can again be attributed to the ion-pair dissociation constant, which in this case increases with decreasing cation radius.11 For the Cl− series, a smaller cation such as Li+, with a large solvated radius,12 would interact more weakly with the Cl− anions in solution in comparison to a bulkier cation such as K+, which has a small solvated radius.12 That is, “the salt with the smallest crystallographic and consequently the largest solvated radius for the cation shows the smallest tendency toward ion-pair formation”.11 In comparison to the opposite trend mentioned above for the halide series, this observed trend for the alkali metal series reinforces the lesson that the ability of cations to pair with anions is a complex interaction that depends on several variables, for example, the lattice energy constant of the salt, the solvation media, and the polarizability of the ion pairs. A thought-provoking discussion can thus be generated in the classroom regarding the great number of competing forces at play in chemistry, where experiments such as those described herein often must be carried out to fully evaluate a system. During the inquiry-based study, a number of different variables were examined by the students. The key to this exercise is the manipulation of conductivity and resistivity by changing a single variable while holding all other aspects of the electrolyte composition constant. Because a number of variables can be examined, some with questionable purpose, the hypothesis behind the investigation should be approved by the instructor before the experiment is performed. One pair of students in the testing groups increased the concentration of the salt (LiCl) dissolved in the PEO-based electrolyte from 0.1 to 0.5 M, thereby enhancing the conductivity of the solution from 2.36 ± 0.43 to 4.62 ± 0.68 S m−1 (Figure 4A), and

RESULTS

This experiment was carried out by several groups of students in both an urban high school environment (Prosser Career Academy, Chicago, IL) and a college setting (Harold Washington College, Chicago, IL), with the most recent being a group of ten Harold Washington students during a summer session. The students worked in groups of two and all data reported herein are an assembly of their work. In the absence of any dissolved salt, the PEO electrolyte solution had a conductivity of 3.06 ± 0.95 S m−1 and a resistivity of 0.35 ± 0.07 Ω m. When lithium salts (LiCl, LiBr, LiI) were dissolved into the electrolyte, the conductivity increased in proportion with the ionic radii of the anions of the dissolved salts (Figure 3A). A near-linear relation (R2 = 0.99) was found to exist between the conductivity of the salt-containing electrolyte solutions and the ionic radii of the dissolved anions, with a similar trend observed for decreases in independently measured resistivity values (R2 = 0.99). This result has been observed for lithium halides in other solvents8 and should be attributed to an increase in the ion-pair dissociation constant (i.e., weaker ion pairing), as the radius of the anion is increased.9 In this experiment, the highly dispersed charge of a large anion (for example, I−) will interact more weakly with the Li+ cations in solution. This allows for higher charge mobility in the electrolyte in comparison to the Li+−Cl− ion pair where the smaller anion with higher charge density can interact strongly with the Li+ cations and slow their movement. Varying the ionic radii of the cations yields an opposite trend, where conductivity (Figure 3C) and resistivity (Figure 3D) 1444




Although the execution of the exercises presented herein is relatively simple, a number of systematic and experimental errors need to be avoided. The presence of ionic impurities in the water source can contaminate the sample, artificially enhancing the conductivity of the electrolyte. Ideally, ultrapure deionized water (>18 Ω cm) can be utilized to dissolve the electrolyte components; however, if this is not available, commercially purchased distilled water, readily available at grocery stores, can be used. (In the lack of financial resources, township- or city-provided water can be used for this experiment but only one faucet should be used and the water must be run through the faucet for several seconds prior to collection to avoid contamination with rust.) The distance between the electrodes during measurement also introduces variance to the collected conductivity and resistivity values. Vessels with similar dimensions should be employed throughout the class to ensure reproducibility in the values collected by the different groups of students. If students can avoid these simple sources of error, a fruitful statistical analysis of the results can be performed to introduce students to the concepts of significant figures, averaging, standard deviation, and error analysis in a laboratory setting. Additionally, these data can be plotted with error bars to encourage graph reading. Conductivity and resistivity values can be plotted against ionic radii, molar masses, and atomic numbers to demonstrate the near-linear relation between the former and the loose correlations between the latter two. As all of these concepts are introduced with the ubiquitous LIB in mind, this exercise should provide a focal point to enhance the retention of these basic concepts for students in general chemistry courses.

Figure 4. Plot of the (A) conductivity and (B) resistivity of a PEObased aqueous electrolyte containing 0.1, 0.25, and 0.5 M of LiCl. Plot of the (C) conductivity and (D) resistivity of a PEO-based aqueous electrolyte containing no salt, where PEO concentration was varied at 0.33, 0.66, and 1 mg mL−1. The data in panels A and B were collected by one pair of students and those in panels C and D were collected by another pair.14

decreasing resistivity from 0.42 ± 0.06 to 0.22 ± 0.03 Ω m (Figure 4B). Another pair varied the quantity of dissolved PEO in the solution from 5 to 15 g in 15 mL of water;13 however, only minimal changes in conductivity (Figure 4C) or resistivity (Figure 4D) were observed in their experiment, which they then explained using error analysis. Different combinations of cations and anions were also investigated, with one pair of students monitoring the [KCl, KBr, and KI] series, whereas another tested the [LiI, NaI, and KI] series. All of these series followed the trend described in the previous paragraph, where changes in the ionic radii of the anions or cations affected conductivity and resistivity.

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DISCUSSION The correlation of conductivity and resistivity of a salt solution with ionic radius is an easy means to introduce periodicity into the classroom under the guise of real-world applications in LIBs. Through the observed data, students can be asked which alkali halide salts should be used in the electrolytes of ionic batteries to achieve the best conductivity (and thus the greatest charge and discharge rate). They should be able to select Li+ as the best of the alkaline cations (hence the “L” in LIBs) and I− as the best anion in the halides. At this point, the instructor can stress the decrease of ion-pairing dissociation constant as an important consideration in the design of ionic batteries. Although periodic trends may not be readily accessible or memorable to students, this laboratory exercise provides a realworld-motivation anchor point upon which students can focus when studying this basic chemistry concept. The instructor can additionally introduce the use of noncoordinating anions such as [BF4]−, [ClO4]−, [PF6]−, [AsF4]−, and [CF3SO3]−, which further diminish ion-pairing beyond the halidesin the design of the modern LIBs.

Notes

ASSOCIATED CONTENT

* Supporting Information S

Perspective for instructors, list of compounds and associated hazards, detailed photographs with associated laboratory directions, modifiable student handouts. This material is available via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

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*E-mail: [email protected]. The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by a summer RET fellowship to R.K. by the Nanoscale Science and Engineering Center at Northwestern University (NSF Award EEC-0647560 through the Nanoscale Science and Engineering Initiative), an American Competitiveness in Chemistry Fellowship to O.C.C. (NSF Award CHE-0936924), and a URC grant to T.B.H. (NSF Award CHE-0629174).

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REFERENCES

(1) Nagaura, K.; Tozawa, K. Progress in Batteries and Solar Cells; JEC Press, Inc.: Brunswick, OH, 1990; Vol. 9. (2) Tarascon, J. M.; Armand, M. Nature 2001, 414, 359−367. (3) Balakrishnan, P. G.; Ramesh, R.; Prem Kumar, T. J. Power Sources 2006, 155, 401−414. (4) Treptow, R. S. J. Chem. Educ. 2003, 80, 1015−1020. (5) Collins, S. J. Chem. Educ. 2010, 87, 1018−1018.

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(6) Fenton, D. E.; Parker, J. M.; Wright, P. V. Polymer 1973, 14, 589−589. (7) Meyer, W. H. Adv. Mater. 1998, 10, 439−448. (8) Savedoff, L. G. J. Am. Chem. Soc. 1966, 88, 664−667. (9) Weaver, W. M.; Hutchison, J. D. J. Am. Chem. Soc. 1964, 86, 261−265. (10) Edwards, J. D.; Taylor, C. S.; Russell, A. S.; Maranville, L. F. J. Electrochem. Soc. 1952, 99, 527−535. (11) Graham, J. R.; Kell, G. S.; Gordon, A. R. J. Am. Chem. Soc. 1957, 79, 2352−2355. (12) Heyrovská, R. Chem. Phys. Lett. 1989, 163, 207−211. (13) We note that PEO has a high solubility in water where gram quantities of it can be dissolved in the same weight of water (see Dormidontova, E. E. Macromolecules 2002, 35, 987−1001 and references therein). The resulting solutions are extremely viscous. (14) We note that the values for the 0.1 M LiCl points presented in Figure 3 do not agree with those in Figure 4A,B because they were measured by different groups of students in different laboratories using different types of water and different experimental setups. One of the strengths of this project is students can use readily available equipment and materials to complete the laboratory measurements, which increases the accessibility of the experiment. A drawback is that the individual measurements show a high degree of variability from student group to student group. Nevertheless, the trends of the data do not change.

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Conductivity through Polymer Electrolytes and Its Implications in

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