Rechargeable Metal-Ion Energy Storage

8/14/2019 Rechargeable Metal-Ion Energy Storage

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Rechargeable Metal-Ion Batteries for Energy Storage

CSUN AIMS2ProgramAttract, Inspire, Mentor and Support Students

Department of Education GrantCalifornia State University, Northridge

Kevin MirandaAugust 2013

CSUN Students: John Grijalva & Travis Van LeeuwenCSUN Advisers: B. Bavarian & L. Reiner


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Kevin Miranda CSUN AIMS22013 2

1 AbstractLithium-ion batteries dominate the energy storage industry and research & developmentfunding. Because of that the potential upside of sodium-ion batteries as a replacement isexplored. Drawbacks and highlights of sodium-ion cells are briefly discussed. The current

state of lithium-ion technology is explored and briefly discussed.Experimentation on a simple magnesium-copper cell is executed with note of open-cellpotential, corrosion, electrolyte, and corrosion inhibitor effectiveness. Water, alum,magnesium sulfate and acetic acid are tested as potential electrolytes for the cell. Openpotential is tested using VersaStat and the Ecorr vs Time test on MS-DOS. Averagepotential is scaled against each electrolyte. Open potential is analyzed as a function of time.

2 IntroducionPresent day rechargeable batteries have many limitations and relatively high costs. Lithium-ion batteries are the most ubiquitous of the bunch with high energy density, high capacity,

low memory effect, and long cycling life. However, lithium is expensive and dangerous.With its component materials becoming harder to extract or downright impossible to satisfyincreasing demand, lithium needs a replacementor at least a companion to fill in the tasksill-suited for lithium.Sodium-ion batteries present the solution to the cost problem that plagues lithium, withoutmuch loss in energy density or life cycling. The drawbacks of sodium ion batteries and itslack of popularity have caused it to be almost completely ignored by the research anddevelopment community. Although in its current state sodium does not present an immediatesolution, its unequaled abundance offers a good chance that a sodium battery could changethe face of energy storage.

As such, a goal to build an inexpensive battery that would light up an LED bulb for anextended period of time was conceived. The battery was to be built using cheap materials thatwere safe, easy to acquire, and resulted in at least 2.2 volts of potential. Magnesium andcopper were chosen as electrodes; many different electrolytes were tested to measurecorrosion, open potential degradation over time, and effectiveness of antifreeze as a corrosioninhibitor. The effects of antifreeze on open-cell potential over time were also noted.

3 Sodium-ion Batteries3.1The case for Sodium-ion

Due to the high availability of sodium on the earths crust, sodium ($135-$165 per ton oftrona) is extremely cheap when compared to lithium ($5000 per ton of lithium carbonate)[Slater]. Trona (trisodium hydrogendicarbonate dihydrate) is processed and sodiumcarbonate can then be extracted, making trona a cheap source for sodium ions. Also,lithium reserves in the future might be limited or dwindling and thus creating unstableand highly fluctuating prices [Egbue]. When taking analyst-projected optimistic andpessimistic estimates of lithium availability and demand, lithium falls short, where as thehigh abundance of sodium, as trona or seawater, means that demand of any size can be


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satisfied. Sodium-sulfur batteries are characterized by their high coloumbic efficiency,low cost, and ease of recyclability [1].

Figure 1: Projected demand of Lithium compared to projected HEV and EV demand [3].

3.2The drawbacks of Sodium-ion batteriesSodium-ion batteries tend to require high temperatures to operate efficiently since the

melting point of sodium is around 98C[1]. Sodium-sulfur and ZEBRA (Zero-EmissionBattery Research Activities) cells, which are based on Na-NiCl2, both operate at

temperatures in the range of 300C due to the requirement that both the anode andsodium are molten in the cell [2]. This requires that sodium-based cells employ efficientand robust battery management systems in order to maintain peak efficiency.Also, like lithium, sodium is highly reactive and explosive in water, which leads tosafety concerns. Precautionary measures must be taken when choosing sodium-cellmaterials. Sodium creates harsh environments within cells, which can result in corrosionthat in turn may lead to short circuits, unwanted conductivity, and dendrite growth in thesolid electrolyte interphase (SEI) [1].Sodium ions are nearly three times as large as those of lithium, 23g mol-1and 6.9g mol-1

respectively [3]. A heavier atom is harder to push across the electrodes, which equates toa lower energy density. A lower energy density essentially removes sodium-basedbatteries from ever being applied to the transportation sector since a heavier batterywould be required to hold the equivalent charge of a lithium-based one. Furthermore,

sodiums redox potential is relatively high when compared to lithium, E= -2.71V and

E= -3.04V respectively [2], which makes it harder to oxidize.

3.3Why Sodium?Battery consumers, the general populace and industrial, are price sensitive. Given themany applications and roles of batteries in modern life, including: hybrid-electric

vehicles (HEV), full electric vehicles (EV), mobile computing, smart grid energystorage, and many others, finding a cheap solution to the expected increase in demand inthe near future creates a big opportunity for manufacturers and suppliers to gain afoothold in the multibillion dollar battery industry. Sodiums low price and highavailability make it markedly important to research and develop into a contender in thefight for battery high energy density, near-infinite life cycle, decreasing weight andvolume, and high recyclability and safety.


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3.4Where Sodium makes senseLooking at sodiums strengths and weaknesses it makes for quite an easy argument torule out the possibility of sodium playing a huge role in the mobile or automotive sector.The chemistry behind sodium denies it this possibility due to its weight and low energy

density. However, since very little research has been given to sodium when compared tolithium, it isnt impossible to say that with the right application and chemistry sodiumcan eventually become efficient and dense enough to be applied in the mobile andautomotive industry. But, in its current state sodium is well suited for the energy storageindustry. Be it the smart grid storage on a national scale or solar energy storage at themicro-consumer level, the low cost and current research dictates that sodium is one ofthe best players in stationary applications. Already sodium displays a higher voltagepotential than the widely used nickel-metal hydride rechargeable battery [1], a sign ofsodium ion batteries potential.

Figure 2: Table of reduction potential of the both the half-cells and the full cell, respectively. 25 C [1].

3.5The chemistry of SodiumOne particular system, the sodium-sulfur rechargeable battery, uses sodium and sulfur as

the cathode and anode, respectively. Ceramic alumina (-Al2O3) serves as both theelectrolyte and the separator in this application [1]. This cell operates at temperatures

between 300 and 350 C (572 and 662 F), where sodium and sulfur both remain in aliquid state [1].As the sodium-sulfur battery discharges a solid forms, sodium polysulfide, at the sodiumcathode, which increases resistance and then eventually prevents further discharge from

occurring. The design requirements of the separator and electrolyte require that it bemanufactured in a solid tubular form, a difficult processing technique, so that it mayhave high Na+conductivity, low discharge, high energy density, strong mechanicalstability and a high resistance to corrosion.


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Figure 3: Schematic representation of Na-S cell [2].

However, at the current operating temperature range Na-S cell batteries are prone tomaterials degradation and corrosion. Sodium creates a very volatile and harshenvironment. The use of an inhibitor can help increase the longevity of this cell and itssubcomponents. Currently a Fe-75Cr alloy inhibitor is showing promising results whensprayed on the inner wall of the electrode casing [1].The Na-S cell in particular is not ready for mass production or practical use due to itsextremely high operating temperature, difficult tubular design, high rate of capacityfade, and other issues. Advancement in an effective planar design could increase energyand power densities and ease manufacturing, but this would require the rightcomponents and materials to withstand the internal environment [1].

4 Lithium Ion Batteries4.1The case for Lithium-ion

Lithium is the lightest metal in the periodic table and is the most easily reduced. Such alow redox potential, -3.07V, allows for lithium-ion batteries high specific energydensity, 150-190 watt-hours per kilogram [4] and an extremely high capacity per gram at3829 mAh, compared to sodiums 1165 mAh per gram [3]. Combined with lithium-ionslack of a memory effect, or capacity fade, lithium-ion cells hold the most capacity forthe most cycles with little mass. This very combination is what has made lithium-ion themost popular battery in consumer electronics and increasingly in the automotiveindustry.

4.2The drawbacks of Lithium-ion batteriesThe most significant problem in the production and operation of lithium-ion batteries issafety, say Tsivadze, Kulova and Skundin [4]. Lithiums extremely negative redoxpotential makes it very reactive, especially in water, and thus makes cells prone toexplosions and leakage. High reactivity is the reason why nearly all lithium-ion batteriesare made using non-aqueous electrolytes [6].Cost is lithiums second biggest drawback. Not only the cost of the lithium metal itself,but also cost is highly dependent on the price of cobalt, lithium-ion batteries mostpopular electrode. Tin, sodium-ions most popular electrode material, on the other handis extremely cheap and thus keeps sodium-ion relatively inexpensive [3]. As previously


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discussed in section 1.1, lithiums high cost per kilogram, about $5000, makes it veryexpensive in large-scale applications such as vehicles and energy storage.

4.3Where Lithium makes senseLithium is perfectly suited for applications where its high cost is of little to noimportance, where its energy density and capacity make the only viable option, and insmall-scale storage for consumer electronics and portables. Since 1991, the first yearlithium-ion was commercialized by Sony, very little has changed in the electrochemicalcell. Instead focus has remained on increasing structural improvements in order toincrease specific energy density [4]. This means that energy density and price is slated toremain stagnant unless a new source of lithium, thereby increasing supply, or a novelway increasing capacity, thereby decreasing demand, are discovered.

Figure 4: The battery pack of the Tesla Model S contains hundreds of lithium-ion cells [teslamotors.com]

Lithium-ions application where its price matters much less than in others is evident inthe luxury sedan the Model S, produced by Tesla Motors. With the price per kilowatt-hour near $1000 [5], only the most affluent of consumers are able to afford it. Outside ofluxury cars serving a niche market, lithium-ion has been outdone by cheaper lead-acidand nickel-metal hydride batteries found in hybrids like the Toyota Prius.However, in fully electric vehicles like the Model S, Honda Fit EV, Fiat 500e, and theNissan Leaf, lithium-ion is the only choice available that allows these cars to operateefficiently without exceeding mass or weight restrictions. Any other type of readilyavailable battery would make the range too low and curb weight too high.For small consumer electronics the high density, long life cycles, fast charging, and high

capacities highly outweigh the price. For this reason lithium-ion is nearly the soleprovider of energy for smartphones, laptops, tablets and other portables worldwide.


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Theory:

5 Designing the inexpensive battery cell5.1Goals

The tentative goals set forth for this experiment and research were as follows: to developa cell or combination of cells using readily available equipment and materials, low-costand safe electrolytes, and a resulting voltage open-cell potential in the range of 2.2-2.6volts, enough to light an LED light bulb.

5.2Theory and calculationsThe electrodes chosen were a copper, as the anode, and magnesium, as the cathode. Thereasoning behind this is due to the relative position of each electrode on the emf-seriestable [Figure 5]. Theoretically the open-cell potential would have been expected to equal

the absolute difference in the potentials listed in Figure 5:

|+ +|=|2.363 0.340|= 2.703

However the above calculation takes into account a few assumptions, like 1 molar ionicsolution of magnesium and copper at standard temperature and pressure. These idealrequirements were forgone during experimentation, not to lack of will, but due topracticality. Instead focus was placed on electrolytes and their effects on resulting open-cell potential.

Figure 5: emf-series Table


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5.3MaterialsFor the experiments, available copper and magnesium samples were used [Figure 6].Readily available salts and liquids were used, as per research goals. All electrolytes usedcould easily be prepared without much cost, difficulty, or danger. For electrolyte holder

and electrode separator, cotton pads were used due to their high permeability, lowelectrolyte leakage, and extremely low cost. Ethylene glycol, or antifreeze, was used as alow-cost corrosion inhibitor.

Figure 6: Polished Copper and Magnesium electrodes prior to testing.

5.4ProcedureFirst, three magnesium and three copper samples are grinded and polished using 240 grit(or lower) sandpaper in order to remove passive layer. The following aqueous solutionsare then prepared in a beaker: 200 mL of approximately 1 molar glacial acetic acidconsisting of 100 mL of water and 6 mL of glacial acetic acid, 200 mL of saturated alumpotash consisting of 14g alum potash per 100 mL of distilled water, 200 mL of saturatedmagnesium sulfate (Epsom salt) consisting of 25.5g per 100 mL of distilled water, and200 mL of pure water. pH and temperature of each solution was recorded.

Cotton pads are then soaked in the electrolyte currently being tested. The cotton pad isplaced between one magnesium sample [if the current test requires antifreeze, themagnesium sample must be submerged entirely in antifreeze prior] and one coppersample and pressed together so as to make good contact; this is one cell. Voltage is thenimmediately measured using a multimeter; this is the 1-cell voltage. The previous stepswere repeated in order to build a second cell but before measuring and recording voltageboth cells were stacked, alternating anodes and cathodes. Voltage was recorded again asthe 2-cell voltage. A third cell was built and the 3-cell voltage recorded [Figure 7].


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Figure 7: 3-cell stack and measuring open-cell potential.

Finally the 3-cell stack is held together using rubber bands, and an LED bulb is attached

to the outermost copper (positive terminal) and outermost magnesium (negativeterminal) samples. In order to insure the LED glows careful attention must be paid tomatch the positive terminal in the LED to the positive side of the 3-cell stack, thenegative terminal vis-a-vis.

Figure 8: 3-cell stack in final setup with lit LED and attached working and reference probes.

Using the VersaStat cell, the Ecorr vs Time test is run with data collecting for 1800seconds (30 minutes) at a data collection rate of 1 point per second (1800 points total).The working electrode (green) is affixed to the LED on the positive terminal, theterminal that shares the wire affixed to the outermost copper sample. The reference

electrode (white) is affixed to the LED on the negative terminal, the terminal that sharesthe wire affixed to the outermost magnesium sample. All other values on the test are leftat their default values. The test is then allowed to run.Data is saved once the test has completed. The 3-cell stack is disassembled, the cottonpads discarded. Note was taken of any visible corrosion.The entire procedure was repeated twice for every solution without coating magnesiumin antifreeze, and twice again when using the antifreeze coating for a total of four runsfor any given solution.


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6 Results and Conclusions6.1Data

Figure 9: Bar Graph of average potentials of a single cell.

Figure 10: Bar graph of average potential of antifreeze-coated single cell.

1.47

1.71

1.98

1.83

0.00

0.20

0.40

0.60

0.80

1.00

1.20

1.40

1.60

1.80

2.00

Potential(V)

Average Potential of Single Cell

Water

Acetic Acid

Alum

MgSO4

1.50

1.87 1.87 1.85

0.00

0.50

1.00

1.50

2.00

Potential(V)

Average Potential of

Antifreeze-coated Single Cell

Water +AF

Acetic Acid +AF

Alum +AF

MgSO4 +AF

Units (Volts) Water Water +AF Acetic Acid Acetic Acid +AF Alum Alum +AF MgSO4MgSO4

+AF

Average 1.47 1.50 1.71 1.87 1.98 1.87 1.83 1.85

Run #1 1.54 1.34 1.66 1.88 1.99 1.84 1.87 1.83

Run #2 1.40 1.65 1.76 1.86 1.97 1.90 1.78 1.86

Table 1: 3-cell open potential and averages.


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Figure 11: Average 3-stack Open-cell Potential as a function of time.

Figure 12: Average 3-Stack Open-cell Potential as a function of time.

6.2ObservationsFrom the data gathered in Table 1, it is evident that the electrolyte which resulted in thehighest voltage was Alum, with the chemical formula of KAl(SO4)2, followed byMagnesium Sulfate (Epsom salt), MgSO4, and Acetic Acid, C2H4O2, and lastly distilled


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water. Alum is clearly above the rest in terms of open-cell potential, especially with noinhibitor present.For water and acetic acid, coating the magnesium electrode in ethylene glycol(antifreeze) resulted in a higher single-cell open potential. The opposite was true for thesulfates, alum and Epsom salt, which experienced a slight decrease in potential or

negligible change.Corrosion-wise, water as an electrolyte resulted in very little to no visible corrosion[Figure 13]. When antifreeze was used, corrosion was nearly non-existent on some ofthe electrodes [Figure 14]. For acetic acid, alum and magnesium sulfate, the corrosionpattern resembled a powdery white coating on the magnesium and a very thing oxidationlayer on the copper samples. [Figure 15].

Figure 13: Electrodes after 30 minutes with water electrolyte. Pitting corrosion is slightly visible.

Figure 14: Electrodes after 30 minutes with water electrolyte and antifreeze coating. Corrosion is less evident

than before.

Figure 15: Corrosion layers on magnesium electrodes after 30 minutes with acetic acid electrolyte. Corrosion

layer is powdery white. Corrosion is slightly visible on copper electrodes.


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Every electrolyte was capable of lighting the LED when stacked to a 3-cell, and kept itlit throughout the 30 minute test on the VersaStat machine. From the graphs a familiarsharp decrease in the first few minutes of the test was common to all electrolytes,regardless of antifreeze use or not. After the point of sharp decrease the open-cell

potential decreased slowly until reaching a final point for the remainder of the test[Figures 11-12]. The tangential behavior of the graph is probably due to the corrosion onthe magnesium electrodes to reach a point of passive equilibrium. Once the corrosionlayer has stabilized, so does the resulting voltage.The effect of the antifreeze coating on the magnesium was evident in almost everysample tested. The antifreeze, for the most part, negatively affected the open-cellpotential but it decreased the amount of visible corrosion on every single run [Figures13-15].

6.3LimitationsMany unaccounted-for variables were present throughout the experiments and research.These include, but are not limited to: exact alloy composition of our electrode samples,surface area of copper and magnesium electrodes, volume of electrodes, quantitativemeasurement of contents in electrolyte solutions, constantly changing materials, andinconsistent testing methods.

6.4ConclusionsOverall the goal of creating an inexpensive cell that would light up and LED and achievean open potential of at least 2.2 volts was clearly achieved. No components or materialstested were expensive, hard to find, or posed any immediate risk.From the limited data gathered it can be concluded that alum is the better electrolyte as itpresented a 33% increase in single-cell open potential over water and nearly a 10%increase over magnesium sulfate, the next best. In terms of corrosion, water is the clearwinner with very little visible corrosion present, especially when using antifreeze as acorrosion inhibitor.Antifreeze served as a well-suited corrosion inhibitor due to its commonality andeffectiveness with the harsher electrolytes. Although open potential was decreased attimes, or even remained unchanged, antifreeze might serve essential when consideringcell life and degradation rates.Although the stated goals did not set forth the need for prolonging cell life, furthertesting with many different inhibitors with the most corrosive of electrolytes might helpincrease the understanding of antifreeze as an effective corrosion inhibitor. With theamount of data gathered, merely two testes per electrolyte, it isnt enough to concludethat antifreeze in fact played any beneficial or detrimental role on cell life or open-cellpotential.In addition, a more solid conclusion as the effectiveness or detriment of each electrolytecan only be achieved with more top-down testing, which would more clearly outlinepatterns and behaviors that could not be concluded with two tests per.


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Works Cited

[1.]Cheng, Fangyi, Jing Liang, Zhanliang Tao, and Jun Chen. "Functional Materials forRechargeable Batteries."Advanced Materials(2011): 1695-715. Wiley. Web. 14 Aug.2013.

[2.]Ellis, Brian L., and Linda F. Nazar. "Sodium and Sodium-ion Energy Storage Batteries."Current

Opinion in Solid State and Materials Science16.4 (2012): 168-77. Science Direct.Web. 14 Aug. 2013.

[3.]Slater, Michael D., Donghan Kim, Eungie Lee, and Christopher S. Johnson. "Sodium-ionBatteries."Advanced Functional Materials23.8 (2013): 947-58. Wiley. Web. 14 Aug.2013.

[4.]Tzivadze, A. Y., T. L. Kulova, and A. M. Skundin. "Fundamental Problems of LithiumIon

Rechargeable Batteries."Protection of Metals and Physical Chemistry of Surfaces49.2 (2013): 149-54. Springer. Web. 14 Aug. 2013.

[5.]Vyrynen, Antti, and Justin Salminen. "Lithium Ion Battery Production." The Journal ofChemical Thermodynamics46 (2011): 80-85. Science Direct. Web. 14 Aug. 2013.

[6.]Yoshino, Akira. "The Birth of the Lithium-Ion Battery."Angewandte ChemieInternational Edition51.24 (2012): 5798-800. Wiley. Web. 14 Aug. 2013.

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