Batteries, overview

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Batteries, Overview

From "Encyclopedia of Energy"

© 2004, Elsevier Inc

1 How Batteries Work

2 History

3 Types of Batteries and Their Characteristics

4 Applications

5 Environmental Issues

6 Future Outlook

Glossary

battery

An assemblage of cells connected electrically in series and/or parallel to provide the desired voltage and current for a given application

cell

An electrochemical cell composed of a negative electrode, a positive electrode, an electrolyte, and a

container or housing

electrode

An electronically conductive structure that provides for an electrochemical reaction through the change of oxidation state of a substance; may contain or support the reactant or act as the site for the electrochemical reaction

electrolyte

A material that provides electrical conduction by the motion of ions only; may be a liquid, solid, solution, polymer, mixture, or pure substance

primary cell

A cell that cannot be recharged and is discarded after it is discharged

secondary cell

A rechargeable cell

Batteries are an important means of generating and storing electrical energy They are sold at a rate of several billions of dollars per year worldwide They can be found in nearly all motor vehicles (e.g.,

automobiles, ships, aircraft), all types of portable electronic equipment (e.g., cellular phones, computers, portable radios), buildings (as backup power supplies), cordless tools, flashlights, smoke alarms, heart pacemakers, biomedical instruments, wrist-watches, hearing aids, and the like Batteries are so useful and ubiquitous, it is difficult to imagine how life would be without them Batteries, strictly speaking, are composed

of more than one electrochemical cell The electrochemical cell is the basic unit from which batteries are

built A cell contains a negative electrode, a positive electrode, an electrolyte held between the electrodes, and a container or housing Cells may be electrically connected to one another to form the assembly called a battery In contrast to the alternating current available in our homes from the electric utility company,

batteries deliver a direct current that always flows in one direction There are a few different types of

batteries: Primary batteries can be discharged only once and then are discarded; they cannot be recharged Secondary batteries are rechargeable Forcing current through the cells in the reverse direction can reverse the electrochemical reactions that occur during discharge Both primary and secondary batteries can be categorized based on the type of electrolyte they use: aqueous, organic solvent, polymer, ceramic, molten salt, and so on

1 HOW BATTERIES WORK

The electrical energy produced by batteries is the result of spontaneous electrochemical reactions The driving force for the reactions is the Gibbs free energy of the reaction, which can be calculated easily from data in tables of thermodynamic properties The maximum voltage that a cell can produce is calculated from this simple relationship:

where E is the cell voltage, ΔG is the Gibbs free energy change for the cell reaction (joules/mole), n is the number of electrons involved in the reaction (equivalents/mole), and F is the Faraday constant (96487 coulombs/equivalent) It is clear from Eq (1) that a large negative value for the Gibbs free energy of the cell reaction is desirable if we wish to have a significant cell voltage

In practical terms, cell voltages of 1 to 2 V are achieved when using aqueous electrolytes, and cell voltages

up to approximately 4 V are achieved when using nonaqueous electrolytes When larger voltages are

required, it is necessary to connect a number of cells electrically in series For example, 12-volt automotive batteries are composed of six 2-volt cells connected in series

An important property of batteries is the amount of energy they can store per unit mass This is the specific energy of the battery, usually expressed in units of watt-hours of energy per kilogram of battery mass

(Wh/kg) The maximum value of the specific energy is that which can be obtained from a certain mass of reactant materials, assuming the case that any excess electrolyte and terminals have negligible mass This

is called the theoretical specific energy and is given by this expression:

where the denominator is the summation of the molecular weights of the reactants

It can be seen from Eq (2) that the theoretical specific energy is maximized by having a large negative value for ΔG and a small value for ΣMw The value of ΔG can be made large by selecting for the negative

electrode those reactant materials that give up electrons very readily The elements with such properties are located on the left-hand side of the periodic chart of the elements Correspondingly, the positive electrode reactant materials should readily accept electrons Elements of this type are located on the right-hand side

of the periodic chart Those elements with a low equivalent weight are located toward the top of the periodic chart These are useful guidelines for selecting electrode materials, and they help us to understand the wide interest in lithium-based batteries now used in portable electronics Equations (1) and (2) give us a useful framework for representing the theoretical specific energy and the cell voltage for a wide range of batteries,

as shown in Fig 1 The individual points on the plot were calculated from the thermodynamic data These points represent theoretical maximum values The practical values of specific energy that are available in commercial cells are in the range of one-fifth to one-third of the theoretical values As expected, the cells using high equivalent weight materials (e.g., lead, lead dioxide) have low specific energies, whereas those using low equivalent weight materials (e.g lithium, sulfur) have high specific energies The lines in Fig 1 represent the cell voltages and simply represent the relationships given by Eqs (1) and (2)

FIGURE 1 Theoretical specific energy for various cells as a function of the equivalent weights

of the reactants and the cell voltage

The details of the electrochemical reactions in cells vary, but the principles are common to all During the discharge process, an electrochemical oxidation reaction takes place at the negative electrode The negative electrode reactant (e.g., zinc) gives up electrons that flow into the electrical circuit where they do work The negative electrode reactant is then in its oxidized form (e.g., ZnO) Simultaneously, the positive electrode reactant undergoes a reduction reaction, taking on the electrons that have passed through the electrical circuit from the negative electrode (e.g., MnO2 is converted to MnOOH) If the electrical circuit is opened, the electrons cannot flow and the reactions stop

The electrode reactions discussed in the preceding can be written as follows:

and

The sum of these electrode reactions is the overall cell reaction:

Notice that there is no net production or consumption of electrons The electrode reactions balance exactly

in terms of the electrons released by the negative electrode being taken on by the positive electrode

2 HISTORY

Volta's invention of the Volta pile in 1800 represents the beginning of the field of battery science and

engineering The pile consisted of alternating layers of zinc, electrolyte soaked into cardboard or leather, and silver Following Volta's report, many investigators constructed electrochemical cells for producing and storing electrical energy For the first time, relatively large currents at high voltages were available for

significant periods of time Various versions of the pile were widely used Unfortunately, there was a lack of

understanding of how the cells functioned, but as we now know, the zinc was electrochemically oxidized and the native oxide layer on the silver was reduced The cells were “recharged” by disassembling them and exposing the silver electrodes to air, which reoxidized them Inevitably, many other electrochemical cells and batteries were developed

John F Daniell developed a two-fluid cell in 1836 The negative electrode was amalgamated zinc, and the positive electrode was copper The arrangement of the cell is shown in Fig 2 The copper electrodes were placed in (porous) porcelain jars, which were surrounded by cylindrical zinc electrodes and placed in a larger container A copper sulfate solution was placed in the copper electrode's compartment, and sulfuric acid was put in the zinc electrode compartment

The electrode reactions were as follows:

and

FIGURE 2 A set of three Daniell cells connected in series.

Sir William R Grove, a lawyer and inventor of the fuel cell, developed a two-electrolyte cell related to the Daniell cell in 1839 Grove used fuming nitric acid at a platinum electrode (the positive) and zinc in sulfuric acid (the negative) Variants on this formulation were popular for a number of years Of course, all of these cells were primary cells in that they could be discharged only once and then had to be reconstructed with fresh materials

Gaston Planté invented the first rechargeable battery in 1860 It was composed of lead sheet electrodes with

a porous separator between them, spirally wound into a cylindrical configuration The electrolyte was sulfuric acid These cells displayed a voltage of 2.0 V, an attractive value During the first charging cycles, the

positive electrode became coated with a layer of PbO2 The charging operation was carried out using

primary batteries—a laborious process Because of the low cost and the ruggedness of these batteries, they remain in widespread use today, with various evolutionary design refinements

The electrode reactions during discharge of the lead-acid cell are as follows:

and

Notice that sulfuric acid is consumed during discharge and that the electrolyte becomes more dilute (and less dense) This forms the basis for determining the state of charge of the battery by measuring the specific gravity of the electrolyte The theoretical specific energy for this cell is 175 Wh/kg, a rather low value

compared with those of other cells

Waldemar Jungner spent much of his adult life experimenting with various electrode materials in alkaline electrolytes He was particularly interested in alkaline electrolytes because there generally was no net

consumption of the electrolyte in the cell reactions This would allow for a minimum electrolyte content in the cell, minimizing its weight Jungner experimented with many metals and metal oxides as electrode materials, including zinc, cadmium, iron, copper oxide, silver oxide, and manganese oxide

In parallel with Jungner's efforts in Sweden, Thomas Edison in the United States worked on similar ideas using alkaline electrolytes and many of the same electrode materials Patents to these two inventors were issued at nearly the same time in 1901

During the period since the beginning of the 20th century, a wide variety of cells have been investigated, with many of them being developed into commercial products for a wide range of applications Representative cells are discussed in the next section

3 TYPES OF BATTERIES AND THEIR CHARACTERISTICS

Batteries are usually categorized by their ability to be recharged or not, by the type of electrolyte used, and

by the electrode type Examples of various types of electrolytes used in batteries are presented in Table I Primary (nonrechargeable) cells with aqueous electrolytes are usually the least expensive and have

reasonably long storage lives Historically, the most common primary cell is the zinc/manganese dioxide cell used for flashlights and portable electronics in sizes from a fraction of an amp-hour to a few amp-hours Various modifications of this cell have been introduced, and now the most common version uses an alkaline (potassium hydroxide) electrolyte The materials in this cell are relatively benign from an environmental point

of view The electrode reactions in an alkaline electrolyte may be represented as follows:

and

The cell potential is 1.5 V, and the theoretical specific energy is 312Wh/kg Practical cells commonly yield 40

to 50Wh/kg

Another common primary cell is the zinc/air cell It uses an aqueous potassium hydroxide electrolyte, a zinc negative electrode, and a porous catalyzed carbon electrode that reduces oxygen from the air The overall electrode reactions are as follows:

and

TABLE I Electrolyte Types and Examples Aqueous electrolyte H2SO4, KOH

Nonaqueous electrolyte

Organic solvent Solid Li salt in ethylene carbonate-diethyl carbonate

Crystalline Na2O x 11Al2O3

Polymeric Li salt in polyethylene

Molten salt (high temperature) LiCl-KCl

The sum of these reactions gives this overall cell reaction:

The cell voltage is 1.6 V, and the theoretical specific energy is 1200 Wh/kg Practical specific energy values

of 200 Wh/kg have been achieved for zinc/air cells This cell has a very high energy per unit volume,

225Wh/L, which is important in many applications The materials are inexpensive, so this cell is very

competitive economically This cell is quite acceptable from an environmental point of view

The zinc/mercuric oxide cell has the unique characteristic of a very stable and constant voltage of 1.35 V In applications where this is important, this is the cell of choice It uses an aqueous potassium hydroxide electrolyte and is usually used in small sizes The electrode reactions are as follows:

and

The high equivalent weight of the mercury results in the low theoretical specific energy of 258 Wh/kg Because mercury is toxic, there are significant environmental concerns with disposal or recycling

The zinc/silver oxide cell has a high specific energy of about 100 Wh/kg (theoretical specific energy =430 Wh/kg), and because of this, the cost of the silver is tolerated in applications that are not cost-sensitive Potassium hydroxide is the electrolyte used here The electrode reactions are as follows:

and

This gives the following overall cell reaction:

The cell voltage ranges from 1.8 to 1.6 V during discharge Because of the high density of the silver oxide

electrode, this cell is quite compact, giving about 600Wh/L.

Primary (nonrechargeable) cells with nonaqueous electrolytes have been under development since the 1960s, following the reports of Tobias and Harris on the use of propylene carbonate as an organic solvent for electrolytes to be used with alkali metal electrodes It has long been recognized that lithium has very low equivalent weight and electronegativity (Table II) These properties make it very attractive for use as the negative electrode in a cell Because lithium will react rapidly with water, it is necessary to use a completely anhydrous electrolyte There are many organic solvents that can be prepared free of water Unfortunately, organic solvents generally have low dielectric constants, making them poor solvents for the salts necessary

to provide electrolytic conduction In addition, organic solvents are thermodynamically unstable in contact with lithium, so solvents that react very slowly and form thin, conductive, protective films on lithium are selected

TABLE II Some Properties of Lithium and Zinc

Lithium Zinc

Equivalent weight (g/equiv.) 6.94 32.69

Reversible potential (V) -3.045 -0.763

Electronegativity 0.98 1.65

Density (g/cm3) 0.53 7.14

Perhaps the most common lithium primary cell is that of Li/MnO2 It has the advantage of high voltage (compared with aqueous electrolyte cells), 3.05 V, and a high specific energy (~170 Wh/kg) The electrode reactions are as follows:

and

The electrolyte is usually LiClO4 dissolved in a mixture of propylene carbonate and 1,2-dimethoxy-ethane In general, the specific power that can be delivered by these cells is less than that available from aqueous electrolyte cells because the ionic conductivity of the organic electrolyte is much lower than that of aqueous electrolytes

An interesting cell is the one using fluorinated carbon as the positive electrode material Because this

material is poorly conducting, carbon and titanium current collection systems are used in the positive

electrode The electrode reaction is as follows:

The electrolyte commonly is LiBF4 dissolved in gamma butyrolactone

An unusual primary cell is that of lithium/thionyl chloride (SOCl2) The thionyl chloride is a liquid and can act

as the solvent for the electrolyte salt (LiAlCl4) and as the reactant at the positive electrode This cell can function only because of the relatively stable thin protective film that forms on the lithium electrode,

protecting it from rapid spontaneous reaction with the thionyl chloride The cell reaction mechanism that is consistent with the observed products of reaction is the following:

and

where SO' is a radical intermediate that produces SO2 and sulfur This cell provides a high cell voltage of 3.6

V and a very high specific energy but can be unsafe under certain conditions Specific energies of up to 700 Wh/kg have been achieved compared with the theoretical value of 1460 Wh/kg

Table III summarizes the characteristics of several representative primary cells

Primary cells with molten salt electrolytes are commonly used in military applications that require a burst of power for a short time (a few seconds to a few minutes) These batteries are built with an integral heating mechanism relying on a chemical reaction to provide the necessary heat to melt the electrolyte (at 400-500°C) A typical cell uses lithium as the negative electrode and iron disulfide as the positive electrode The following are example electrode reactions:

and

The molten salt electrolyte (a mixture of alkali metal chlorides) has a very high conductivity, allowing the cell

to operate at very high specific power levels in excess of 1 kW/kg

Rechargeable cells with aqueous electrolytes have been available for more than 140 years, beginning with the Planté cell using lead and sulfuric acid as discussed previously Modern Pb/PbO2 cells and batteries have received the benefit of many incremental improvements in the design and optimization of the system Current versions are very reliable and inexpensive compared with competitors Depending on the design of the cell, lifetimes vary from a few years to more than 30 years Specific energy values of up to approximately

40 Wh/kg are available

Alkaline electrolyte systems are available with a variety of electrodes Perhaps the most common alkaline electrolyte cell is the Cd/NiOOH cell It offers very long cycle life (up to thousands of charge-discharge cycles) and good specific power (hundreds of W/kg), albeit with a modest specific energy (35-55 Wh/kg) The cell voltage, 1.2 V, is lower than most and can be somewhat of a disadvantage The electrode reactions are as follows:

and

Of course, the reverse of these reactions takes place on recharge

TABLE III Properties of Some Lithium Primary Cells

Open circuit voltage

(V)

Practical specific energy

(Wh/kg)

Theoretical specific energy

(Wh/kg)

A very robust rechargeable cell is the Edison cell, which uses iron as the negative electrode, nickel

oxyhydroxide as the positive electrode, and an aqueous solution of 30w/o potassium hydroxide as the electrolyte During the early years of the 20th century, Edison batteries were used in electric vehicles They proved themselves to be very rugged and durable, although they did not have high performance The iron electrode reaction is as follows:

The positive electrode reaction is the same as for the Cd/NiOOH cell described earlier

The NiOOH electrode has proven itself to be generally the best positive electrode for use in alkaline

electrolytes and has been paired with many different negative electrodes, including Cd, Fe, H2, MH, and Zn Recently, the metal hydride/nickel oxyhydroxide cell (MH/NiOOH) has been a significant commercial success and has captured a large fraction of the market for rechargeable cells It offers a sealed, maintenance-free system with no hazardous materials The performance of the MH/NiOOH cell is very good, providing up to several hundred watts/kilogram peak power, up to about 85 Wh/kg, and several hundred cycles The

negative electrode operates according to the following reaction:

The very high theoretical specific energy and low materials cost of the zinc/air cell make it attractive for consumer use There have been many attempts to develop a rechargeable zinc/air cell, but with limited success due to the difficulties of producing a high-performance rechargeable air electrode

The electrode reactions during discharge for this cell are as follows:

and

During recent years, the cycle life of the air electrode has been improved to the point where more than 300 cycles are now feasible Development efforts continue In the meantime, various versions of a “mechanically rechargeable” zinc/air cell have been tested These cells have provision for removing the discharged zinc and replacing it with fresh zinc This approach avoids the difficulties of operating the air electrode in the recharge mode but creates the need for recycling the spent zinc to produce new zinc electrode material The status of some rechargeable cells with aqueous electrolytes is shown in Table IV

Rechargeable cells with nonaqueous electrolytes have been under development for many years, although they have been available on the consumer market for only a decade or so All of the types of nonaqueous electrolytes shown in Table I have been used in a variety of systems

Organic solvent-based electrolytes are the most common of the nonaqueous electrolytes and are found in most of the rechargeable lithium cells available today A typical electrolyte consists of a mixture of ethylene carbonate and ethyl-methyl carbonate, with a lithium salt such as LiPF6 dissolved in it Various other

solvents, including propylene carbonate, dimethyl carbonate, and diethyl carbonate, have been used Other lithium salts that have been used include lithium perchlorate, lithium hexafluoro arsenate, and lithium

tetrafluoro borate All of these combinations of solvents and salts yield electrolytes that have much lower conductivities than do typical aqueous electrolytes As a result, the electrodes and electrode spacing in these lithium cells are made very thin to minimize the cell resistance and maximize the power capability

Lithium is difficult to deposit as a smooth compact layer in these organic electrolytes, so a host material, typically carbon, is provided to take up the Li on recharge and deliver it as Li ions to the electrolyte during discharge Many types of carbon, both graphitic and nongraphitic, have been used as the lithium host

material In addition, a variety of intermetallic compounds and metals have been used as Li host materials All of the commercial Li rechargeable cells today use a carbon host material as the negative electrode

TABLE IV Rechargeable Aqueous Battery Status

System voltage Cell

(V)

Theoretical specific energy (Wh/kg)

Specific energy (Wh/kg)

Specific power (W/kg)

Cycle life (dollars/kWh) Cost

Fc/NiOOH 1.30 267 40-62 70-150 500-2000 >100

Note est., estimate