matls 4n03 batteries and fuel cells

31
Chapter 3 – Batteries and Fuel Cells 1 Batteries and Fuel Cells “You can start me up” -The Rolling Stones 3.1 Introduction Unlike the previous two chapters where the alternative energy sources solar and nuclear were discussed, the topic of batteries and fuel cells does not meet the criterion of “alternative”. A battery can only be used for so long before it has to be recharged and the energy for recharging must come from somewhere. Similarly the “fuel” in a fuel cell is obtained through the consumption of energy. Nevertheless, batteries/fuel cells do play a very significant role in the fight against global climate change, particularly in the area of transportation. A nuclear powered automobile is a far fetched idea, a solar powered car is less so, but still is a long way from commercialization. Therefore, if alternative energy sources become viable and economically competitive with fossil fuels, we still have the problem of gasoline consumption in cars. A high energy density battery design can solve, or at least help alleviate, the thorny issue of CO 2 emissions from vehicles, provided of course the energy for recharging the automotive battery is derived from alternative energy sources. Although most everyone on the planet knows what a battery is, it is useful to define some terms before discussing in more detail the physics behind the operation of a battery. A battery is a device which converts energy from a chemical reaction into electrical energy. The battery consists of several electrochemical cells, connected in series to produce a desired output voltage. A so-called primary battery is not easily recharged and is intended for a single use. A secondary battery is equivalent to a rechargeable battery. Note, a battery in use is being discharged, whereas charging refers to the application of a current in the reverse direction to the battery to restore it to its original state. A reserve

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Page 1: Matls 4N03 Batteries and Fuel Cells

Chapter 3 – Batteries and Fuel Cells

1

Batteries and Fuel Cells

“You can start me up”

-The Rolling Stones

3.1 Introduction

Unlike the previous two chapters where the alternative energy sources solar and nuclear

were discussed, the topic of batteries and fuel cells does not meet the criterion of

“alternative”. A battery can only be used for so long before it has to be recharged and the

energy for recharging must come from somewhere. Similarly the “fuel” in a fuel cell is

obtained through the consumption of energy. Nevertheless, batteries/fuel cells do play a

very significant role in the fight against global climate change, particularly in the area of

transportation. A nuclear powered automobile is a far fetched idea, a solar powered car is

less so, but still is a long way from commercialization. Therefore, if alternative energy

sources become viable and economically competitive with fossil fuels, we still have the

problem of gasoline consumption in cars. A high energy density battery design can

solve, or at least help alleviate, the thorny issue of CO2 emissions from vehicles, provided

of course the energy for recharging the automotive battery is derived from alternative

energy sources.

Although most everyone on the planet knows what a battery is, it is useful to define some

terms before discussing in more detail the physics behind the operation of a battery. A

battery is a device which converts energy from a chemical reaction into electrical energy.

The battery consists of several electrochemical cells, connected in series to produce a

desired output voltage. A so-called primary battery is not easily recharged and is

intended for a single use. A secondary battery is equivalent to a rechargeable battery.

Note, a battery in use is being discharged, whereas charging refers to the application of a

current in the reverse direction to the battery to restore it to its original state. A reserve

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Chapter 3 – Batteries and Fuel Cells

2

battery is one in which the electrolyte (typically) is isolated from the rest of the battery

prior to use. A fuel cell is an electrochemical cell that does not use consumable

electrodes, but instead feeds fuel continuously into the anode and cathode sides.

3.2 The electrochemical cell

As mentioned above, a battery consists of several electrochemical cells, also called

galvanic cells, which are hooked together in series in order to increase the total power.

Therefore to understand how a battery operates it is sufficient to consider just the

electrochemical cell. Much like a solar cell, a battery supplies electrons to drive some

external load. However in the case of a galvanic cell the electrons are produced by a

chemical reaction. From thermodynamics we know that the work done by the

electrochemical cell comes at a cost, which in turn implies that the chemical reactions

taking place within the cell must lead to a decrease in free energy. In fact for a reversible

process at constant temperature and pressure the maximum work done by the system, W,

is equal the free energy change G∆− .

The work performed when transporting an electric charge e with units of coulombs

through a potential difference E in volts is simply the product of eE. Now let’s switch to

a per mole basis. The total charge carried by one mole of positively charged ions of

valency +1 is 96,487 coulombs and this number, denoted by F, is known as Faraday’s

constant. Thus, the work produced by the electrochemical cell is:

nFEGW =∆−= 3.1

where n is the valency of the ions produced in the chemical reaction (more on this

below). The electric potential difference across the cell electrodes, E, is also known as

the electromotive force, or EMF, of the cell. Equation 3.1 tells us how much “juice” we

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Chapter 3 – Batteries and Fuel Cells

3

can get from a battery. Clearly to get the most bang for the buck we would like to find

reactions with the highest driving force G∆− .

The electrochemical cell consists of three main components: the anode, the cathode and

the electrolyte. To understand the role of each it is best to refer to a specific example

and, following the lead of D. R. Gaskell in his text “Introduction to Metallurgical

Thermodynamics”, let us investigate the so-called Daniell cell. The anode of the Daniell

cell is an electrode of pure Zn immersed in a compartment containing an aqueous

solution of ZnSO4. The chemical reaction that takes place when Zn is introduced into

solution is known as a half cell reaction and is written as:

Zn=Zn2+(in aq.soln.)+2e-

Here the reaction is written such that reading from left to right corresponds to the

favorable direction, ie. the free energy decreases when dissolving a Zn2+ ion. Zn is

referred to as the anode because the reaction taking place there is an oxidation reaction,

meaning the Zn loses electrons. Notice the valency of the reaction is n=2 and the two

electrons produced flow through the Zn electrode and into the external circuit.

The cathode in a Daniell cell is also a pure metal electrode, in this case Cu. It too is

immersed in an aqueous solution, but now containing CuSO4. The half cell reaction is:

Cu2+(aq. Soln.) + 2e- = Cu

which is a reduction reaction since the Cu is gaining electrons. The cathode then refers to

the electrode where reduction takes place.

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4

Fig. 3.1 Schematic depiction of the Daniell cell. Electrons flow from the

pure Zn anode to the Cu cathode. At the same time Zn 2+ ions are

dissolved from the anode and Cu 2+ is plated is plated onto the cathode.

Ions diffuse from side of the cell to the other across a porous separator.

Figure from http://www.chem-is-try.org/wp-

content/migrated_images/pengantar/pengantarkimia-

terjemah_img_105.jpg

The electrochemical cell is now complete. Electrons produced at the anode flow through

our device and arrive at the Cu cathode. In the meantime positively charged ions are

moving in the opposite direction. Zn2+ diffuses away from the Zn electrode into solution

and Cu ions are plated onto the Cu electrode. The electrolyte in the process is the

medium through which ionic conduction is occurring. In the case of the Daniell cell the

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5

electrolyte consists of two separate compartments containing an anolyte (ZnSO4) and a

catholyte (CuSO4), but for most cells the electrolyte is one phase. From this qualitative

discussion of the electrochemical cell it is clear that two important features are required

for an effective electrolyte. It must be a good ionic conductor, yet a poor electrical

conductor. Aqueous solutions are a popular choice for electrolytes, but there are other

possibilities such as the polymer membranes and the solid oxide materials.

How many volts does the Daniell cell produce or equivalently what is the cell EMF?

Equation 3.1 tells us how to proceed. Clearly we need the free energy of the total cell

reaction and to determine this free energy change we must first find the free energy when

all the species involved are in their standard states at the standard temperature of 298K.

Both the cathode and the anode consist of pure metals and thus Zn and Cu occur in their

standard states, but what about the electrolyte? Consider the following procedure. Water

is placed in contact with pure ZnSO4 and Zn2+ ions and SO42- ions are allowed to dissolve

in the water. The dissolution process will continue until the water is saturated with ions

or equivalently when the chemical potential of the ions is equal to the chemical potential

of the ZnSO4 solid. Then take the saturated aqueous solution and use it as the electrolyte

(more specifically the anolyte) in the cell. The construction just described results in a

standard state. The half cell reaction of Zn=Zn2+ (in standard state)+2e- is, from a free

energy change point of view, the same as the reaction:

Zn(s) ⇔ ZnSO4(s)

where the s notation emphasizes the fact that the free energy change can be computed

based on the reaction in the solid state. By similar reasoning the cathode half cell

reaction can be written as CuSO4(s)=Cu(s). In any electrochemical cell the total reaction

is found by simply adding the two half cell reaction, that is:

Zn(s)+CuSO4(s)=ZnSO4(s)+Cu(s)

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The standard free energy change for the above reaction is oG∆ =-213,040 Joules and,

from Eq. 3.1, the standard EMF for the cell is:

oE =1.104 Volts

From the above discussion it should be clear that the cell EMF will be different if the

concentration of ions in the electrolyte is something other than the standard saturated

values. The Nernst equation provides the necessary correction. For the specific case of

the Daniell cell, it is given by:

−=

4

4ln2 CuSO

ZnSOo

a

a

F

RTEE

3.2

where 4ZnSOa and

4CuSOa are the activities of ZnSO4 and CuSO4 respectively in the two

electrolyte compartments.

When the two electrodes are introduced into the electrolyte the battery begins to

discharge. At the anode Zn ions increase in concentration and at the cathode Cu ions are

depleted. The imbalance leads to a flow of Zn from anode to cathode across the porous

medium and a reverse flow of sulfate ions. The decrease of SO4 and Cu in the catholyte

means the activity a will decrease and by similar reasoning the activity term will increase.

From the Nernst equation it is clear that the cell EMF decreases over time, which, of

course, makes sense from a thermodynamic point of view. Eventually the cell will need

to be recharged. The recharging process involves applying a reverse voltage of sufficient

strength – greater than oE – such that the cell reaction moves in the opposite direction.

Zn will then precipitate on the anode and Cu will dissolve from the cathode.

Equation 3.2 refers specifically to the Daniell cell, but it can be generalized quite easily.

Imagine for example if instead of pure Cu and pure Zn electrodes, alloys containing Zn

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7

and Cu were utilized. Then the activities of these reacting species are no longer unity and

the Nernst equation must be rewritten as:

For any cell reaction given by:

aA + bB = cC + dD

the cell EMF is

−=

bB

aA

dD

cCo

aa

aa

nF

RTEE ln

3.3

The Daniell cell is capable of supplying 1.1 Volts. If we wire up 6 such cells in series,

we’re up to 6.6 Volts. But can we realistically expect to generate the theoretical cell

EMF? The answer of course is no. As we saw in the case of solar cells there are

processes that will rob the system of available voltage and decrease the efficiency. Some

of these processes are illustrated in Fig. 3.2 which plots the cell EMF vs the current.

There are three main sources of voltage loss. The first, labeled IR, is the internal

resistance of the cell and it includes resistance from flow through the two electrodes,

through the electrolyte and any resistance associated with interfaces between phases.

From Ohm’s law the IR loss contribution increases linearly with the operating current.

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Fig. 3.2 The voltage delivered by a cell is reduced by IR losses, activation

overpotential and concentration overpotential.

The curve labeled activation polarization is a little trickier to understand. Consider once

again the Daniell cell. At the cathode Cu ions from solution will precipitate on the pure

Cu electrode. Although there is a free energy gain during this process there is also, in

general, an activation barrier for the atomic process. Thus there is an overpotential

associated with the half reactions at both the cathode and anode. The aη term reflects the

decrease in cell voltage due to the activation overpotential that must be supplied. The

mathematical relationship between aη and the current is known as the Tafel equation and

is given by:

)ln(ibaa +=η 3.4

where a and b are constants and b is known as the Tafel slope.

The third factor that decreases the voltage of an electrochemical cell is the concentration

polarization. We have seen above that the electrons flowing in one direction from the

battery is compensated by a net flow of ions within the electrolyte. At low currents

positive ions are produced at the anode (Zn2+ in the Daniell cell) and transported through

the electrolyte at rates such that the concentration is more or less uniform. But as the

current increases there is a build up of ions at the electrode and now the ionic conduction

will contain a contribution, which is slow, from diffusion in the aqueous solution. The

concentration overpotential reflects this diffusion mechanism.

3.3 The lead acid battery

The Daniell cell can deliver around 1 Volt – we can do better. Indeed there are many

superior batteries available, such as: Nickel-cadmium batteries, silver oxide batteries,

nickel-metal hydride batteries, nickel-zinc batteries, etc. Our interest here is to discuss

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Chapter 3 – Batteries and Fuel Cells

9

battery materials and designs that, eventually, can be used to power automobiles. In

order to understand the materials requirements for electric vehicle applications it is

perhaps best to start by describing the most popular battery type and ask if it is up to the

task. The most popular battery is the one that starts your car every morning, the lead acid

battery.

It should come as no surprise that the anode in a lead acid battery is lead. Pure lead is a

very soft material and therefore the anode usually consists of Pb plated on a grid, as

shown in Fig. 3.3, where the grid material provides mechanical support. In addition, the

pure Pb is typically alloyed with 4-6 weight percent Antimony to increase the strength via

solid solution strengthening. In order to increase the number of electrons produced per

unit area of anode grid, grooves are machined into the lead to greater increase the

effective surface area.

Fig 3.3. Examples of lead acid battery grid anode grid designs.

Conventional (left) and radial (right).

The half cell reaction at the anode is

Pb=Pb 2+ + 2e-

and is combined with the Pb ion reaction with the electrolyte:

Pb 2+ + SO4 2- = PbSO4

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In the lead acid battery the electrolyte, the source of sulfate ions, is sulfuric acid H2SO4 at

a concentration of roughly 6molar (37% by weight).

Lead oxide PbO2 is the cathode material in a lead acid battery. Two crystalline forms of

the oxide are present: a rhombohedral phase denoted alpha-PbO2 and a tetragonal (rutile)

form designated as the beta phase. In each case the oxide is not stoichiometric and a

more accurate description of the structure would be PbOx with x varying from

approximately 1.85-2.05. Although two crystalline oxides are present in the cathode, the

difference in half cell potential for each phase amounts to just .01V. The cathode is

fabricated in the form of a highly porous structure to increase the active surface area.

The half cell reaction at the cathode can be written as:

PbO2 +4H+ + 2e- = Pb2+ + 2H2O

Pb2+ + SO4 2- = PbSO4

The overall cell reaction then is:

Pb + PbO2 + 2H2SO4 = 2PbSO4 + 2H2O

The Nernst equation applied to the lead acid batteries reads:

−=

42

2ln047.2SOH

OH

a

a

F

RTE

Notice, during discharge the battery produces water and depletes the amount of sulfuric

acid. The standard cell EMF for the lead acid design is 2.047 Volts, which is a factor of

two larger than the Daniell cell.

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The large open circuit voltage is one of the main advantages of the lead acid battery, but

there are others. Lead acid batteries were discovered in 1860 and the nearly 150 year

head start in research and development over modern battery designs has resulted in

batteries that are low cost, very reliable, low maintenance and available in a variety of

designs and sizes. On the negative side, lead acid batteries are an environmental

nightmare. The sulfuric acid electrolyte, of course, requires care when handling, but

more important is the use of lead itself. Pb is a toxic heavy metal and even in very small

concentrations can cause health concerns in children. In most other industrial

applications the trend in recent years has been toward lead free materials, such as lead

free solders in electronics packaging. The battery industry has instead relied on

recycling. Although recycling rates are quite high, >95%, the remaining discarded

batteries is cause for concern.

Currently lead acid batteries power small vehicles such as golf carts and fork lifts. But it

is doubtful the lead acid design will be the battery of choice for automobiles in the future

and the main reason is known to anybody who has actually installed a car battery – they

are heavy. To scale up from a golf cart to a conventional road vehicle will require a

substantial number of lead acid cells in series and therefore a large increase in weight.

An important figure of merit for a battery is the specific energy given in units of

Watts.hours/kg. A closely related quantity is the energy density Watts.hours/liter. Figure

3.4 plots the energy density and specific energy for various battery types. The

information is plotted as boxes to illustrate the variation in performance under different

operating conditions. It is clear that significant weight and space savings can be realized

by moving away from the traditional lead acid battery. A factor of roughly two in

performance can be gained by utilizing batteries of the Zn/MnO2 design and a factor of

almost four can be achieved with batteries based on lithium. With respect to materials

selection there are two ways to increase the specific energy of battery: materials leading

to higher cell EMF can be chosen and lighter materials can be used. As we will see

below, in both instances, there is no better choice than Li.

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Fig. 3.4. The energy density and specific energy for various battery

materials. Notice battery designs involving Li offer the best performance

in terms of weight and space savings.

3.4 Li ion batteries

As illustrated above, in order to find the cell EMF we can add the free energy of the

oxidation half reaction at the anode and then subtract the oxidation reaction at the cathode

(or add the reduction reaction). Since we add and subtract reactions it is valid to

reference all oxidation reactions to some standard reaction and the standard reaction in

most common use is:

H2 = 2H+ + 2e-

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Table I below lists several half cell reactions with their corresponding EMF with the

standard reaction involving hydrogen appearing in the middle and assigned an EMF of

zero. In order to maximize the cell EMF, and hence the battery power, anode materials

near the top of the list should be combined with cathode materials near the bottom. And

look what element is found at the very top. Not only does Lithium produce a large

voltage of 3.045 when oxidizing to Li ++e- , but Li is the third lightest element in the

universe! Given these two properties of Li it is not surprising that Li based battery

designs are the most promising in terms of specific energy.

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14

Lithium has the electronic structure 1S22S1. The lone electron occupying the 2S state

means that Li really wants to lose an electron. It is this property which places Li at the

top of the chart in table I, but the extreme reactivity of Li also makes it difficult to work

with. Pure Li metal is so violently reactive with air that it must be stored in oil. The

extreme reactivity of Li severely limits the choice of electrolyte materials or non-aqueous

solutions in rechargeable batteries. For example in Lithium/Manganese Disulfide

(Li/MoS2) batteries the electrolyte is LiAsF6 salt dissolved in a 50:50 mixture of

propylene carbonate and ethylene carbonate. In the Lithium/Titanium Disulfide (Li/TiS2)

design the electrolyte is again LiAsF6 but now dissolved in 2-methyl tetrahydrofuran and

tetrahydrofuran. Other choices for organic electrolyte solutions are equally

unpronounceable: dioxolane, diethyl carbonate and ethylmethyl carbonate. The major

drawback of organic electrolytes is the fact that the ionic conductivity is roughly two

orders of magnitude lower than typical aqueous solutions and the low conductivity

severely limits the current (or rate of discharge). The use of polymer films, rather than

liquids, as the electrolyte material has thus far not resulted in any improvement in

conductivity.

Although disposable Li batteries are in common use in consumer electronics, for

rechargeable high power applications, such as electric vehicles, battery designs based on

Li metal do not look promising. There is a possible solution however and it is the Li ion

battery. A critical materials feature of the Li ion battery is the use of so-called

intercalation compounds for the electrodes.

Consider the crystal structure of graphite as shown in Fig. 3.5. The structure is a

hexagonal ABAB stacking of sheets, one layer is often referred to as graphene, and in

each layer the carbon atoms form very strong 2sp2 hybrid bonds with three neighboring

atoms. Bond angles are 120o within the layers. The distance between successive layers

of grapheme sheets is fairly large and the interaction between layers is characterized by

weak van der Waals attraction. Graphite is used as pencil “lead” because it is easy to

shear and fracture the material along the basal planes.

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Fig. 3.5. The hexagonal crystal structure of graphite. There is strong

interatomic within each of the “grapheme” layers, whereas the interlayer

bonding is characterized by weak van der Walls attraction.

Since the bonding between AB layers in graphite is so weak it is fairly easy, in the

thermodynamic sense, to insert other atoms in between the layers, forming what is called

an intercalation compound. In the case of graphite popular choices for the intercalant

species are Ca, K and of interest here of course Li. In a particular layer Li can be inserted

until the point where there are six carbon atoms for every Li and the compound is

denoted LiC6. The Li atoms occupy the center of the carbon hexagons, but each Li atom

has no neighbor within any of the adjacent six hexagons. As the concentration of 1/6 Li

is reached the stacking sequence of the top and bottom graphene sheets changes to AA.

But the behavior of Li in an individual layer does not tell the whole story. Fig. 3.6 shows

the voltage during the discharge of a LixC6 intercalation material as a function of the Li

content denoted by x. Rather than a smooth curve between voltage and composition as

would be predicted by the Nernst equation, a series of plateaus, or stages, are observed.

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Chapter 3 – Batteries and Fuel Cells

16

Each level portion of the curve represents a two phase equilibrium where each of the

phases can be labeled stage 1, stage 2, etc. The phases themselves are shown

schematically in Fig. 3.7. Stage 3 for example refers to a structure where every third

layer consists of the 1/6 Li arrangement. (In the 2Lto 2 stage, the L refers to liquid-like

or disordered arrangement of Li within the layer. That is, Li still occupies the center of 6

member carbon rings, but there is a random arrangement of filled vs unfilled rings.)

Fig. 3.6. The voltage vs composition of Li in a graphite intercalation

compound. Each of the plateaus in the curve represents a different phase,

or stage, as shown in Fig. 3.7.

Fig. 3.7. A schematic of the staging behavior of a graphite intercalation

compound.

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Upon charging of a Li ion battery the Li is incorporated into the anode via the staging

process, whereas during discharge the reverse atomic transport occurs. Therefore we are

now in a position to write down the half cell reaction at the anode:

Li xC = C + xLi+ + xe-

Now what about the cathode? In order to avoid the use of pure Li metal anywhere in the

battery design, an intercalation compound for the cathode is again a good choice. There

are a few cathode materials in use such as the compounds LiCoO2, LiNiO2 and LiMn2O4,

but let us examine in more detail the most popular choice: LiCoO2. For LiCoO2

synthesized at low temperatures (~400oC) the spinel crystal structure is formed. This low

temperature phase does not perform well in battery applications and exhibits large

hysteresis upon intercalation and de-intercalation. Therefore the phase formed at higher

temperatures (~800oC) is preferred and its crystal structure is shown in Fig. 3.8 below.

Fig. 3.8. The unit cell of LiCoO2.

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Notice the structure is similar to that of graphite in that it can be viewed as a stacking of

hexagonal planes of atoms. Here the Li layers are intercalated between the oxygen

layers. Much like the graphite anode the cathode material undergoes a series of phase

transitions as the Li is intercalated into the LiCoO2. If we again let x represent the Li

atom fraction within the intercalation compound then the useful range of compositions is

1>x>.45. For x in range .9 to 1 a single hexagonal phase denoted H1 is stable whereas for

.78<x<.51 another single hexagonal phase H2 is stable. In the range of concentration

between .9 and .78 a two phase, H1 plus H2, region is observed. Finally, for .51<x<.45 a

monoclinic distortion associated with Li ordering is the thermodynamically favored

structure.

The half cell reaction at the cathode can be written as:

Li 1-xCoO2 + x Li+ + xe- = LiCoO2

For an overall cell reaction of:

Li xC + Li1-xCoO2 = LiCoO2 + C

And now for the good news: the cell voltage for the LiCoO2 battery is a huge 3.88V.

A schematic depiction of the Li ion battery is shown in Fig. 3.9. Because the Li moves

back and forth between the two intercalation compounds – the anode and cathode – the

design is often referred to as a “rocking chair” battery.

We have succeeded in choosing an anode and cathode material such that pure Li metal

does not exist within the battery, but suspiciously absent front the schematic diagram

above is the electrolyte. Four types of electrolytes have been used. As in the case of Li

metal batteries, liquid electrolytes consist of some organic solvent, usually a carbonate,

plus a Li salt. A second option is a gel electrolyte which is a salt and a solvent dissolve

in a high molecular weight polymer. Fluorinated polymers, such as

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polyvinylidene+hexafluoropropene (PVDF-HFP) combined with LiPF6 salt and a

carbonate solvent are a typical choice. In addition to liquid and gel electrolytes, polymer

and ceramic films have also been considered for use in Li ion batteries.

Fig. 3.9. Schematic depiction of the Li ions in a Li ion “racoking chair” battery. From P.

J.Bouwman “Lithium Intercalation in preferentially oriented submicron LiCoO2 films”,

PhD thesis, Universiteit Twente, 2002.

3.5 The hydrogen fuel cell

As mentioned in the introduction a fuel cell, as opposed to a battery, does not use an

anode and cathode material that is consumed during discharge. Instead the “fuel” is

continuously fed into both the cathode and anode ends of the cell. In a hydrogen fuel cell

H2 gas is introduced at the anode whereas oxygen O2 gas is introduced at the cathode

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side. The reaction product in this electrochemical process is water and therein lies the

attraction of the hydrogen fuel cell for energy applications. Harmful greenhouse gases

are produced nowhere in the reaction cycle.

The most popular design choice of the hydrogen fuel cell is the polymer electrolyte

membrane or PEM cell. (Interestingly the cell is often called a proton exchange

membrane fuel cell, but the abbreviation, PEM, is the same!) A schematic diagram of

the PEM fuel cell is shown in Fig. 3.10. The H2 gas flowing through the anode end

undergoes the following oxidation reaction:

H2(g) = 2H+ +2e-

Notice this is simply the standard hydrogen electrode reaction. As will be discussed in

more detail below a catalyst is introduced into the PEM design to aid in breaking the H2

molecule into two protons (H+ ions). Air or more specifically O2 gas flows into the

cathode side where it then reacts according to:

1/2 O2 (g) +2H+ + 2e- = H2O(l)

The liquid water formed must be removed from the cell. As with the anode side a

catalyst is again required to assist the reduction reaction. The overall cell reaction is thus:

H2 (g) + 1/2 O2 = H2O (l)

The standard cell potential is 1.23V. As we now know 1.23V is not great, but the use of

gaseous fuel, without any actual electrode material, means reasonable energy densities

are possible.

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Fig. 3.10. A schematic of the polymer electrolyte membrane (PEM)

hydrogen fuel cell. Note he catalysts on both side of the electrolyte film.

From http://www.lss.supelec.fr/Internet_php/Fuel_Cell/FC_scheme

The PEM fuel cell uses nothing but oxygen and hydrogen gas, generates power and

produces nothing but water as a bi-product. So what’s the problem? Why aren’t we all

driving hydrogen PEM fuel cell automobiles? Well, there are several problems. First

and foremost is a very general concern: where do we get the hydrogen gas? Hydrogen is

not very plentiful in the atmosphere and thus needs to be generated from some chemical

process. If the production of H2 involves the burning of hydrocarbons, as is done

currently, then we have done nothing to alleviate the problem global greenhouse gas

emissions. There are ongoing research efforts to utilize the power generated from solar

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cells to produce hydrogen through the electrolysis of water. But, like other photovoltaic

cell applications, the present designs are too expensive for hydrogen to compete with

traditional fossil fuels. Biological routes to produce hydrogen gas, so-called bio-

hydrogen, have also been proposed. Through the photosynthesis of algae, CO2 gas and

H2O are used in the generation of H2. Unfortunately the rate of H2 production is low and

it is safe to say that the technique requires further development.

The hydrogen source is only one major hurdle to be overcome in creating a cost effective

hydrogen PEM fuel cell, there are several others. As a first point of discussion consider

the electrolyte itself. Recall, an effective electrolyte will exhibit high ionic conductivity,

yet be electrically insulating. In addition, desirable properties of the electrolyte include

good thermal and hydrolytic stability and the ability to form the material into thin films.

The current PEM electrolyte of choice in fuel cells has the trade name Nafion and is from

a class of polymers known as perfluoro-sulfanated acids (PSFA). The chemical structure

of a PSFA is shown in the schematic of Fig. 3.11 and for Nafion the repeat units are m=1,

y=1, n=2 and the long chain backbone is in the range x=5-13.5.

Fig. 3.11. The repeat unit of a perfluoro-sulfanated acid (PSFA). The

electrolyte of choice in current PEM designs is Nafion, which is

characterized by m=1, y=1, n=2 and x in the range 5 to 13.5.

The repeat unit does not tell the full story of why Nafion is a good ionic conductor. Fig.

3.12 shows the microstructure of Nafion. The key feature being the long narrow channels

formed when the sulfonic acid functional groups self organize into hydrophilic water

channels. The water channels created are fairly wide, on average 24 A, such that ions can

pass readily. At the same time, the sidewalls of the channels provide mechanical and

thermal stability for the film.

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Fig. 3.12. Schematic diagram of the molecular structure of Nafion

showing wide hydrophilic channels. From K. Schmidt-Rohr and Q. Chen,

Nature Materials, 7, 75 (2008)

Since Nafion seems to fit the bill, what is the problem? Again there are a few. The first

is cost. Nafion is expensive to produce and a considerable amount of research is now

being conducted to find a lower cost alternative. A more important problem, however,

has to do with water management. During the operation of a hydrogen fuel cell the water

being produced must be continually removed. However, the water in the pores of the

Nafion film give rise to the good ionic conductivity. Therefore, water must be wicked

away from the cathode, but not at a sufficiently high rate such that the electrolyte dries

and its resistance increases. The problem is compounded at high temperatures. The cell

reaction to produce water is quite exothermic and relatively high operating temperatures

are to be expected. If the temperature exceeds 100oC, then evaporation of water from the

Nafion will occur and a dramatic drop in the ionic conductivity will follow.

An effective fuel cell electrolyte membrane must conduct protons, yet it must keep

separate the two components of the fuel – ie. H2 and O2 gas. Nafion also suffers from a

problem known as hydrogen crossover, where hydrogen in molecular form can diffuse

through the electrolyte to the cathode side. The H2 can then directly react with O2 to

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form undesirable peroxides. A seemingly simple solution to the cross over problem is to

fabricate thicker electrolyte films to decrease the H2 flux. However, thicker Nafion films

will increase the overall film resistance and decrease the power of the cell.

As you may guessed by now, Nafion is not the only problem facing the development of

hydrogen PEM fuel cells. Notice in Fig. 3.10 that catalysts are included in the cell

design. Without a catalyst accelerating the two half reactions, the current densities of a

hydrogen fuel cell would be too low for any practical application. Unfortunately, the

most effective catalyst known is platinum. It is not obvious that Nafion is a relatively

expensive polymer material, but it is pretty clear platinum is expensive. Pt is a precious

metal, which at the time of this writing is priced at roughly $40 (US) per gram. It is more

expensive than gold or silver and is several orders of magnitude more costly than steel.

To put the catalyst problem in perspective, various estimates have suggested that, for

PEM fuel cell technology to be cost competitive in large scale automotive applications,

the amount of Pt currently used in the fuel cell stacks must be reduce by a factor of five!

The splitting of H2 into ions of hydrogen at the anode side of the cell is not the problem.

This reaction is fairly easy to catalyze. The more difficult reaction occurs at the cathode

side. The catalyst design in current use consists of a graphite support decorated with

small particles of Pt. A fairly obvious scheme to decrease the Pt content is to fabricate

particles as small as possible, thereby creating more surface catalytic sites per unit

volume of the precious metal. Typically the carbon surface is functionalized, meaning

carboxylic, anhydride, phenolic or carbonyl groups are absorbed on the carbon surface.

The functional groups act as anchoring sites for the catalyst particles and aid in the

deposition of Pt. In recent years, several research groups have investigated the use of

carbon nanotubes as a support material to further decrease the Pt particle size. Although

Pt catalysts have been improved over time by altering the carbon support, it is difficult to

imagine these incremental changes overcoming the five fold decrease in Pt required for

the widespread use of hydrogen vehicles. Instead a fundamental break through is

required and there is hope!

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Before proceeding however a few more problems associated with the catalyst are worth

discussing. As noted above the optimal catalyst consists of a very high surface to volume

ratio of Pt, a situation which is thermodynamically unfavorable. Over time at elevated

temperatures the collection of Pt particles will tend to decrease the overall surface area;

the larger particles in the collection will grow whereas the smaller particles in the

distribution will shrink and eventually disappear. Slowing or stopping this process,

which is known as “coarsening”, is not an easy task. Another troubling aspect of the Pt

catalyst is poisoning. Carbon monoxide molecules absorbed at the Pt surface can greatly

reduce the performance of the catalyst. The effect is quite dramatic with only ppm levels

of CO within the oxygen gas stream having a substantial effect on performance. The

problem can be alleviated to some extent through the use of Pt-Ru alloys where Ru at the

surface more strongly attracts a CO molecule. However, Pt-Ru is not exactly a low cost

alternative and, moreover, the total amount of Pt needed remains unchanged.

An encouraging development in the quest for improved and less expensive PEM catalyst

is the intermetallic compound Pt3Ni. The crystal structure of the compound can be

viewed as an FCC unit cell with the corner atoms being Ni and all the face atoms Pt. In a

bulk specimen each 111 plane consists of 25% Ni atoms, however, at a 111 exposed

surface the Pt, for energetic reasons, tends to be enriched in the topmost plane such that a

pure Pt layer is formed. Thus we can see immediately a benefit of Pt3Ni. The surface,

where the catalytic reaction takes place, looks like pure Pt, but the overall Pt content has

been reduced by 1/4. But the story gets better. The Pt skin structure also changes the

electronic band structure of the material in such a way that the catalyst performance is

dramatically improved. Catalysts can be evaluated based on their activity, which is some

measure of how much faster a reaction proceeds in the presence of the catalyst. Fig 3.13

shows the results for Pt3Ni vs conventional Pt catalysts. The activity here is a bit

complicated. It is expressed in terms of a current flux at constant voltage of the catalyst

contained in a .1M HClO4 solution. We need not describe in detail the reaction occurring

by suffice it to say the activity measured is proportional to the oxygen reduction reaction

taking in a PEM cell. It is pretty clear the results are striking. For the 111 surface the

activity increases by over an order of magnitude. The effect is less for the 100 and 110

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surfaces, but a substantial increase is observed nevertheless. Catalyst improvement

through alloy design appears to be the promising route to developing a low cost fuel cell

for widespread automotive use.

Fig. 1.13. The activity of Pt3Ni catalysts vs pure Pt catalyst for three low

index surfaces. The increase in activity is very large, especially for the

(111) surface. From V. R. Stamenkovic, B. Fowler, B. S. Mun, G. Wang,

P. N. Ross, C. A. Lucas and N. M. Markovic, Science, 315, 493 (2007).

3.6 3.6 3.6 3.6 Hydrogen storageHydrogen storageHydrogen storageHydrogen storage

Let’s assume we have developed some method to produce large quantities of H2 gas

cheaply, we have found an attractive alternative to Nafion for an electrolyte material in a

PEM fuel cell and we have stumbled upon a low cost, yet highly active, catalyst. Are we

then ready to roll a hydrogen vehicle off the assembly line and sell it to the public? The

answer is no. There is still the very sticky problem of how to deliver the H2 gas to the

fuel cell on-board the vehicle.

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Hydrogen has a very high energy density by weight but is much worse than hydrocarbons

(gasoline) on a per volume basis. Thus, the simple solution of using pressurized gas

cylinders to store the H2 gas on board will result in a tank heavier than gas tanks in

automobiles. In addition pressurizing the gas costs energy such that the net efficiency

gain from a hydrogen vehicle is diminished. A final problem with the cylinder design is

one of safety. If the nozzle to a highly pressurized canister is fractured during, say, a

traffic accident, then the cylinder becomes a very dangerous high energy projectile.

An alternative solution to the fuel delivery problem is a metal hydride material for safely

and economically storing hydrogen. The material of choice should be able to absorb

hydrogen quickly during “filling” and must desorb hydrogen rapidly during the operation

of the vehicle. Another very crucial concern is how much hydrogen the material, by

weight, can absorb. The US, Department of Energy, Office of Efficiency and Renewable

Energy, as part of the FreedomCAR program, has published a list of technical targets for

required for viable designs of future hydrogen vehicles. Among other factors, the report

lists a goal of 6% by weight for an on-board hydrogen storage system by the year 2010.

Since hydrogen is the lightest element in the universe, a 6wt% goal represents a daunting

challenge. To understand just how difficult this task is consider the case of the Pd-H

system.

Many metal (M) hydrides exist at a stoichiometry as high as MH, but their formation

often involves a very exothermic reaction and the strong free energy gain means the

reverse reaction of pulling hydrogen out of the storage material is very difficult. The Pd-

H system offers a seemingly attractive alternative. Shown in Fig. 3.14 is the Pd-H phase

diagram where the x axis is not the usual atom or mole fraction. It is the ratio of H atoms

to Pd atoms in the hydride. In other words if the ratio H:Pd is equal to one then the

stoichiometry of PdH is reached. The y axis in the diagram is also a departure from what

we’re use to. Here it the pressure of hydrogen in atmospheres (the conversion to mmHg

for .024 atm is shown in parenthesis). The isotherms on the diagram illustrate, for each

temperature listed, the path in composition space the sample will follow as the pressure

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increases. For example at a temperature of 160oC the concentration will increase from

less than .05 H per Pd atom up to roughly .7 as the pressure is increased to 20atm. As the

pressure is decreased and H is depleted from Pd the system will follow the same isotherm

in the opposite direction. However, the absorption and desorption steps are not

completely reversible. Notice a two phase region between a low volume alpha and a high

volume beta phase exists and the miscibility gap ends at a critical point of approximately

295oC and 20atm. During the flat portion of the isotherm the composition does not

involve an increase in pressure, instead amount of alpha relative to beta is changing.

Because the volume change between phases is fairly large (~10%) a hysteresis is

observed between the loading and unloading of H. Nevertheless, the Pd-H system

represents a hydrogen storage system in which hydrogen can be easily and rapidly

removed and recharged.

Fig. 3.14. The pressure-composition phase diagram of the Pd-H system.

The y axis is atmospheres. Note, the miscibility gap between the low

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volume α and high volume β phases. From

http://131.104.156.23/Lectures/CHEM_462/462_chapter_1.html

A quick glance at the periodic table will reveal the problem with Pd-H. The atomic

weight of Pd is 106.4 and H, of course, is just 1. Therefore if the pressure is sufficiently

high to reach a composition of .7 in the hydride, then a weight percent of only 0.66% is

achieved – a value which is an order of magnitude less than the DOE target of 6.0wt%.

The materials design criterion for a suitable hydrogen storage system is in some sense

contradictory. A large binding energy between H and the material is required to absorb

sufficient amounts, yet the binding can’t be too strong or we will not be able to pull H

from the material to run our vehicle. On top of this delicate balancing act, we need a very

light weight H host. There is a considerable amount of research now being conducted to

identify a suitable storage material and we will discuss a few promising candidates.

A class of compounds known as the alanates is perhaps the leading contender. The

crystal structure of the compound NaAlH4 is shown below (Fig. 3.15) where the green

spheres represent the Al atoms and red correspond to Na. Surrounding each of the Al

atoms is a tetrahedral arrangement of hydrogen atoms. The structure can be viewed as a

packing of layers each consisting of an arrangement of Na and Al-H tetrahedral. Since

four H atoms surround each Al and because Na and Al are low density elements the

weight percent H in sodium alanate as 7.4% (the reversible H content however, is only

5.6%). A close cousin of the sodium alanate is lithium alanate LiAl4, which boasts an

even higher H content of 10.5%. The main hurdle with the alanates is the slow

desorption kinetics. Several recent studies have shown that catalysts like Ti and Zr can

greatly speed up the dissolution rate in NaAlH4, but thus far the rate increase has not

been enough and the search for improved catalysts is hampered because the precise role

played by Ti in the desorption process is not fully understood.

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Fig. 3.15. The unit cell of sodium alanate, NaAlH4, a candidate H storage

material. From http://www.ojwang.com/research.html

Another class of materials with the potential for use as on-board hydrogen storage

devices is the so-called metal organic framework (MOF) compound. One example is the

compound with stoichiometry Zn4O(BDC3) where the unit BDC3 represents the organic

part of the MOF designation and is the functional group 1,4-benzenedicarboxylate. The

atomic structure of the MOF is shown in Fig. 3.16 below. The blue tetrahedra represent

the Zn atoms, with each one being surrounded by four oxygen atoms, colored in red. The

four Zn atoms share a common O atom at the center, but the other sides of the tetrahedra

are bridged, in six places, by the BDC3 group. The result is a very open cube, the large

yellow sphere in Fig. 3.16 shows the available open space in the structure. With the cube

repeated over all space a material is that offers many binding sites for H is formed.

Indeed the Zn4O(BDC3) MOF exhibits a fairly high H weight percent of 4.5% (17.2 H

atoms per formula unit) at a temperature of 78K. Unfortunately the equilibrium H

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content decreases significantly with increasing temperature and at room temperature is

only 1%.

Fig. 3.16. The structure of the metal organic framework (MOF) material Zn4O(BDC3). The yellow sphere represents the large open space in the structure. From N. L. Rossi, J. Eckert, M. Eddaoudi, D. T. Vodak, J. Kim, M. O’Keefe and O. M. Yaghi, Science, 300, 1127 (2003).

Other hydrogen storage materials currently being considered include carbon nanotubes,

intercalated graphite and even ammonia, but the two most challenging problems of

absorption/desorption kinetics and high H content by weight have not been overcome and

it seems safe to conclude that a viable on-board H storage system is many years away.