Electrochemistry 15

7/30/2019 Electrochemistry 15

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Page Notifications OffChemWiki: The Dynamic Chemistry Textbook > Analytical Chemistry > Electrochemistry > Electrochemistry6: Electrochemical Energy Storage and Conversion

Electrochemistry 6: Electrochemical Energy Storage and

Conversion

One of the oldest and most important applications of electrochemistry is to the storage and conversion

of energy. You already know that a galvanic cell converts chemical energy to work; similarly, an

electrolytic cell converts electrical work into chemical free energy. Devices that carry out these

conversions are called batteries. In ordinary batteries the chemical components are contained within the

device itself. If the reactants are supplied from an external source as they are consumed, the device is

called a fuel cell.

1. Introduction2. Physical limitations of battery performance3. Primary and secondary batteries4. The lead-acid storage cell5. The LeClanch "dry cell"

6. Timeline of battery development6.1. Lithium batteries as incendiary devices6.2. Biological Batteries6.3. The fuel cell6.4. Microbobial Fuel Cells

7. Summary8. Outside Links9. Contributors

Introduction

The term battery derives from the older use of this word to describe physical attack or "beating";

Benjamin Franklin first applied the term to the electrical shocks that could be produced by an array of

charged glass plates. In common usage, the term "call" is often used in place of battery. For portable

and transportation applications especially, a battery or fuel cell should store (and be able to deliver) the

maximum amount of energy at the desired rate (power level) from a device that has the smallest

possible weight and volume. The following parameters are commonly used to express these attributes:

Storage capacity or charge density, coulombs/liter or coulombs/kg;Energy density, J/kg or watt-hour/lbPower density, watts/kgVoltage efficiency, ratio of output voltage to ELifetime: shelf-life (resistance to self-discharge) or charge/recharge cycles

A very nice compilation of these and other properties of many common battery types can be found at

this "What's the best battery?" site.

Physical limitations of battery performance

The most important of these are:

Effective surface area of the electrode. A 1-cm2 sheet of polished metal presents far less activsurface than does one that contains numerous surface projections or pores. All useful batteries anfuel cells employ highly porous electrodes. Recent advances in nanotechnology are likely to great
http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Contributorshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Contributorshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Outside_Linkshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Outside_Linkshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Microbobial_Fuel_Cellshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#The_fuel_cellhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Lithium_batteries_as_incendiary_deviceshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Timeline_of_battery_developmenthttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#The_LeClanch.c3.a9_.22dry_cell.22http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Introductionhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6:_Electrochemical_Energy_Storage_and_Conversionhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6:_Electrochemical_Energy_Storage_and_Conversionhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6:_Electrochemical_Energy_Storage_and_Conversionhttp://chemwiki.ucdavis.edu/index.php?title=Special:PageAlerts&id=5547http://www.batteryuniversity.com/partone-3.htmhttp://www.uh.edu/engines/epi510.htmhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Contributorshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Outside_Linkshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Summaryhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Microbobial_Fuel_Cellshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#The_fuel_cellhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Biological_Batterieshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Lithium_batteries_as_incendiary_deviceshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Timeline_of_battery_developmenthttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#The_LeClanch.c3.a9_.22dry_cell.22http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#The_lead-acid_storage_cellhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Primary_and_secondary_batterieshttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Physical_limitations_of_battery_performancehttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6%3a_Electrochemical_Energy_Storage_and_Conversion#Introductionhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_6:_Electrochemical_Energy_Storage_and_Conversionhttp://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistryhttp://chemwiki.ucdavis.edu/Analytical_Chemistryhttp://chemwiki.ucdavis.edu/http://chemwiki.ucdavis.edu/index.php?title=Special:PageAlerts&id=5547


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.

Current density of electrode surface. Expressed in amperes m2, this is essentially a measure of thcatalytic ability of the electrode, that is, its ability to reduce the activation energy of the electrotransfer process.Rate at which electroactive components can be delivered to or depart from the active electrodsurface. These processes are controlled by thermal diffusion and are inhibited by the very narropores that are needed to produce the large active surface area.Side reactions and irreversible processes. The products of the discharge reaction may tend to reawith the charge-storing components. Thermal diffusion can also cause self-discharge, limiting thshelf life of the battery. Recharging of some storage batteries may lead to formation of less activ

modifications of solid phases, thus reducing the number of charge/discharge cycles possible.

Clearly, these are all primarily kinetic and mechanistic factors which require a great deal of

experimentation to understand and optimize.

Primary and secondary batteries

A secondaryor storage battery is capable of being recharged; its electrode reactions can proceed in

either direction. During charging, electrical work is done on the cell to provide the free energy needed to

force the reaction in the non-spontaneous direction. A primary cell, as expemplified by an ordinary

flashlight battery, cannot be recharged with any efficiency, so the amount of energy it can deliver islimited to that obtainable from the reactants that were placed in it at the time of manufacture.

The lead-acid storage cell

The most well-known storage cell is the lead-acid cell, which was invented by Gaston Plant in 1859 and

is still the most widely used device of its type.

The cell is represented by

Pb(s) | PbSO4(s) | H2SO4(aq) || PbSO4(s), PbO2(s) | Pb(s)

and the net cell reaction is

Pb(s) + PbO2(s) + 2 H2SO4(aq) 2 PbSO4(s) + 2 H2O

The reaction proceeds to the right during discharge and to the left during charging. The state of charge

can be estimated by measuring the density of the electrolyte; sulfuric acid is about twice as dense as

water, so as the cell is discharged, the density of the electrolyte decreases.


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Lead-acid storage cell and detail of its

plate construction

The technology of lead-acid storage batteries has undergone remarkably little change since the late 19th

century. Their main drawback as power sources for electric vehicles is the weight of the lead; the

maximum energy density is only about 35 Ah/kg, and actual values may be only half as much. There are

also a few other problems:

The sulfuric acid electrolyte becomes quite viscous when the temperature is low, inhibiting the flow ions between the plates and reducing the current that can be delivered. This effect is well-known anyone who has had difficulty starting a car in cold weather.These batteries tend to slowly self-discharge, so a car left idle for several weeks might be unable t

start.Over time, PbSO4 that does not get converted to PbO2 due to lack of complete discharge gradua

changes to an inert form which limits the battery capacity. Also, "fast" charging causes rapid evolutioof hydrogen from the water in the electrolyte; the bubbles form on the lead surface and can tePbO2 off the positive plate. Eventually enough solid material accumulates at the bottom of th

electrolyte to short-circuit the battery, leading to its permanent demise.

The LeClanch "dry cell"

The most well-known primary battery has long been the common "dry cell" that is widely used to power

flashlights and similar devices. The modern dry cell is based on the one invented by Georges Leclanch in

1866. The electrode reactions are

Zn Zn2+ + 2e

2 MnO2 + 2H+ + 2e Mn2O3 + H2O

Despite its name, this cell is not really "dry"; the electrolyte is a wet paste

containing NH4Cl to supply the hydrogen ions. The chemistry of this cell

is more complicated than it would appear from these equations, and

there are many side reactions and these cells have limited shelf-lifes due

to self discharge. (In some of the older ones, attack by the acidic

ammonium ion on the zinc would release hydrogen gas, causing the

battery to swell and rupture, often ruining an unused flashlight or other

device.) A more modern version, introduced in 1949, is the alkaline cell

which employs a KOH electrolyte and a zinc-powder anode which permits the cell to deliver higher

currents and avoids the corrosive effects of the acidic ammonium ion on the zinc.

Timeline of battery development

Although the development practical batteries largely paralelled the expansion of electrical technology
http://www.geocities.com/bioelectrochemistry/leclanche.htm


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rom a out t e mi -19t century on, it is now t oug t t at a very primitive in o attery was

apparently in use more than 2000 years ago. The brief popularity ofelectrically powered automobiles in

the 1920's encouraged storage battery development. The widespread use of portable "personal"

electrical devices has kept the search for better batteries very much alive.

See here for some interesting pictures of very early electrochemistry cells.

"Baghdad Battery" - 1000 BCE?

Earthenware jars containing an iron rod

surrounded by a copper cylinder were

discovered near Baghdad in 1938. They are

believed to have been used by the Parthian

civilization that occupied the region about

2000 years ago as a source of electricity to

plate gold onto silver.

More on the Baghdad Battery

Allesandro Volta 1782

His "Voltaic pile", a stack of zinc and silver disksseparated by a wet cloth containing a salt or aweak acid solution, was the first battery known to

Western civilization.

Sir Humphry Davy 1813

Davy builds a 2000-plate battery that occupies889 square feet in the basement of Britain'sRoyal Society. His earlier batteries provided power
http://www.corrosion-doctors.org/Biographies/DavyBio.htmhttp://en.wikipedia.org/wiki/Alessandro_Voltahttp://www.unmuseum.org/bbattery.htmhttp://www.parthia.com/parthia_history.htm#Overviewhttp://physics.kenyon.edu/EarlyApparatus/Electricity/Electrochemical_Cell/Electrochemical_Cell.htmlhttp://www.ieee.org/organizations/pes/public/2004/may/peshistory.html


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lighting (carbon arc).

Michael Faraday, 1830'sFaraday discovered the fundamentals of galvaniccells and electrolysis that put electrochemistry ona firm scientific basis.

1836 - Daniell cell (also known as a Crow's

Foot or Gravity cell.)

John Daniell (English chemist and meterologist)developed the first modern storage cell basedon Faraday's principles. This consists of a largeglass jar with a copper star-shaped electrode inthe bottom and a zinc "crow's foot" shapedelectrode suspended near the top. The bottom ofthe jar was filled with a concentrated coppersulfate solution. On top of this was poured dilutesulfuric acid, whose lower density kept it on top.This was the first practical battery to find wideuse to power telegraphs and railway signalingsystems and home doorbells.

1839 - William Grove (Welsh)

Grove was best known in the 19th century for

his "nitric acid battery" which came into wide

use in early telegraphy.

Now, however, he is most famous for his "gas

voltaic battery" in which discovered "reverse

electrolysis": the recombination of H2 and O2

following electrolysis of water at platinum

electrodes. This was the first demonstration of

what we now know as the hydrogen-oxygen

fuel cell (see below.)

1859 -Gaston Plant (French)

Invents the first lead-acid storage cell whichconsisted of two sheets of lead separated by arubber sheet, rolled into a spiral and immersed indilute sulfuric acid.
http://www.geocities.com/bioelectrochemistry/plante.htmhttp://www.100welshheroes.com/en/biography/williamgrovehttp://www.cartage.org.lb/en/themes/Biographies/MainBiographies/D/Daniell/1.htmlhttp://www.chemheritage.org/EducationalServices/chemach/eei/mf.html


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1866 - Georges Leclanch (French)

By 1868 twenty thousand Leclanch cells werebeing used in telegraph systems. The originalLeclanch cells were built in porous pots whichwere heavy and subject to breakage. Withintwenty years other inventors had modified thedesign into what we now know as "dry cells"which became widely used in the first flashlights(1909) and in battery-powered radios of the1920s.

1881 - Faure and others

Development of the first practical lead-acidstorage cell. The major improvement overPlant's design was the addition of a paste ofPbSO4 to the positive plate.

1905 Nickel-iron cell

(Thomas Edison) Edison, who was as much achemist as an all-aroundinventor, thought that thelead in Plant-type cells
http://inventors.about.com/library/inventors/bledisonbiography.htmhttp://www.energizer.com/learning/historyofflashlights.asphttp://en.wikipedia.org/wiki/Georges_Leclanch%C3%A9


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ma e t em too eavy, anthat having acid in contactwith any metal was aninherently bad idea. Aftermuch experimentation, hedeveloped a successfulalkaline battery. TheEdison cell uses an ironanode, nickel oxide cathode, and KOH electrolyte.This cell is extremely rugged and is still used in

certain industrial applications, but it was neverable to displace the lead-acid cell as Edison hadhoped.

1950s

A similar cell, employing a nickel anode instead of

iron, was the first rechargeable cell that was smallenough to be used in portable consumer devices.Its main disadvantage is that it is ruined bycomplete discharge.

1949 - Alkaline dry cell - Lew Urry(Eveready Battery Co.)

First commercial alkaline dry cell. These substituteKOH for the corrosive NH4Cl used in the older dry

cells and last 5-8 times longer.

1947 - Mercury cell

(Ruben and Mallory, 1950's)

This was one of the first "button"-type cells

which were widely used in cameras and hearing

aids. The constancy of the 1.34 v output made

them popular for use in sensitive instruments

and cardiac pacemakers. The net cell reaction

is

Zn(s) + HgO(s) ZnO(s) + Hg(l)

Most countries have outlawed sales of these

cells in order to reduce mercury contamination

of the environment.

Nickel-Cadmium (NiCad) cells

The NiCad cell quickly become one of the mostpopular rechargeable batteries for small consumerdevices. They can deliver high current and

undergo hundreds of chargedischarge cycles.Because cadmium is an environmental toxin, theiruse is being discouraged.
http://www.duracell.com/procell/about/background.asphttp://www.energizer.com/company/companyhistory.asphttp://www.almostfabulous.com/canadians/name/u/urrylew.phphttp://www.worldwideschool.org/library/books/hst/biography/Edison/chap48.htmlhttp://www.ieee-virtual-museum.org/collection/tech.php?id=2345874&lid=1


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1959 - Fuel cell - Francis Bacon (UK)

The first practical fuel cell was developed byBritish engineer Francis Bacon (1904-1992). Thishydrogen-oxygen cell used an alkaline electrolyteand inexpensive nickel electrodes.

Late 1960's - Nickel-metal hydride cells

The hydride ion H

would be an ideal cathode

material except for the fact that its oxidation

product H2 is a gas. The discovery that certain

compounds such as LiNi5 and ZrNi2 can act as

"hydrogen sponges" made it practical to

employ metal hydrides as a cathode material.

One peculiarity of Ni-MH cells is that recharging

them is an exothermic process, so that proper

dissipation of heat must be allowed for. These

batteries are widely used in cell phones,

computers, and portable power tools. The

electrode reactions take place in a

concentrated KOH electrolyte:

Cathode (+): NiOOH + H2O + e Ni(OH)2 +

OH

Anode (-): (1/x) MHx + OH (1/x) M + H2O

+ e

1990s - Lithium cells (Sony Corp.)

Lithium is an ideal anode material owing to its

low density and high reduction potential,

making Li-based cells the most compact ways

of storing electrical energy. Lithium cells are

used in wristwatches, cardiac pacemakers and

digital cameras. Both primary (non-

rechargeble) and rechargeable types have been

available for some time. More recent

applications are in portable power tools and

perhaps most importantly, in electric-powered

or hybrid automobiles.

Modern lithium cells operate by transporting

Li+ ions between electrodes into which the ions

can be inserted or intercalated. Cathodes are

lithium transition-metal oxides such as LiCoO3,

while anodes are lithium-containing carbon,

LiC6. The species that undergoes oxidation-

reduction is not lithium but the transition
http://www.sony.net/Fun/SH/1-24/h2.htmlhttp://en.wikipedia.org/wiki/Francis_Thomas_Bacon


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metal, e.g. Co(III)-Co(IV).

See this Wikipedia article for a very

thorough description of current lithium battery

technology and for recommendation on how to

store and use these cells.

Comparison of Li cell types

Lithium batteries as incendiary devices

There have been numerous reports of fires and explosions associated with lithium batteries. In 2006,

the Dell Corporation had to recall 4.1 million Sony batteries that had been shipped with Dell's laptop

computers and were judged to be at risk owing to a manufacturing defect.

This illustrates the difficulty of concentrating a large amount of chemical energy into a small package,

which is of course the goal of all battery developers eager to meet commercial demands ranging from

consumer personal electronics to electrically-powered cars. The fully-charged Li+-deficient lithium cobalt

oxide cathodes are inherently unstable, held in check only by a thin insulating membrane which, if

accidentally breached, can lead to thermal runaway involving gaseous oxygen, carbon, organic solvents,

and (in some cases) lithium chlorate all the components necessary for a fierce fire.

Much research has gone into the development of fail-safe membranes. In one type, made by ExxonMobil

and targeted at the automotive market, the pores are designed to close up and thus inhibit the passage

of lithium ions when the temperature rises above a safe level.

Biological Batteries

Finally, we should mention the biological batteries that are found in a

number of electric fish. The "electric organs" of these fish are modified

muscle cells known as electrocytes which are arranged in long stacks. A

neural signal from the brain causes all the electrocytes in a stack to become

polarized simultaneously, in effect creating a battery made of series-connected cells. Most electric fish

produce only a small voltage which they use for navigation, much in the way that bats use sound for

echo-location of prey. The famous electric eel, however, is able to produce a 600-volt jolt that it employs

to stun nearby prey.

The fuel cell

Conventional batteries supply electrical energy from the chemical reactants stored within them; when

these reactants are consumed, the battery is "dead". An alternative approach would be to feed the

reactants into the cell as they are required, so as to permit the cell to operate continuously. In this case

the reactants can be thought of as "fuel" to drive the cell, hence the term fuel cell.
http://helium.vancouver.wsu.edu/~ingalls/eels/index.htmlhttp://whozoo.org/fish/teleosts/electric.htmhttp://www.powerstream.com/Pli.htmhttp://en.wikipedia.org/wiki/Lithium_ion_battery


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Although fuel cells were not employed for practical purposes until space exploration began in the 1960's,

the principle was first demonstrated in 1839 by Sir William Grove, a Welsh lawyer and amateur chemist.

At the time, it was already known that water could be decomposed into hydrogen and oxygen by

electrolysis; Grove tried recombining the two gases in a simple apparatus, and discovered what he called

"reverse electrolysis" that is, the recombination of H2 and O2 into water causing a potential

difference to be generated between the two electrodes:

anode: H2(g) 2 H+ + 2e E = 0 v

cathode: O2 + 2 H+ + 2e H2O(l) E = +1.23 v

net: H2(g) + O2(g) H2O(l) E = +1.23 v

It was not until 1959 that the first working hydrogen-oxygen fuel cell was developed by Francis Thomas

Bacon in England. Modern cells employ an alkaline electrolyte, so the electrode reactions differ from the

one shown above by the addition of OH to both sides of the equations (note that the net reaction is

the same):

anode: H2(g) + 2 OH 2 H2O + 2 e

E = 0 v

cathode: O2 (g) + 2 H2O + 2 e 2 OH E = +1.23 v

net: H2(g) + O2(g) H2O E = +1.23 v

Although hydrogen has the largest energy-to-mass ratio of any fuel, it cannot be compressed to a liquid

at ordinary temperatures. If it is stored as a gas, the very high pressures require heavy storage

containers, greatly reducing its effective energy density. Some solid materials capable of absorbing largeamount of H2 can reduce the required pressure. Other fuels such as alcohols, hydrocarbon liquids, and

even coal slurries have been used; methanol appears to be an especially promising fuel.

Schematic diagram of a modernhydrogen-oxygen fuel cell.

Commonly used electrolytes areNaOH solution, phosphoric acid,

or solid oxides. A major limitationof any oxygen-consuming fuel

cell is the slow rate of thereduction of this element at acathode. The best cathode

surfaces are usually made ofplatinum, which is a major cost

factor in fuel cell design.
http://en.wikipedia.org/wiki/Francis_Thomas_Baconhttp://chem.ch.huji.ac.il/~http://www.100welshheroes.com/en/biography/williamgrove


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One reason for the interest in fuel cells is that they offer a far more efficient way of utilizing chemical

energy than does conventional thermal conversion. The work obtainable in the limit of reversible

operation of a fuel cell is 229 kJ per mole of H2O formed. If the hydrogen were simply burned in oxygen,

the heat obtainable would be H = 242 kJ mol1, but no more than about half of this heat can be

converted into work so the output would not exceed 121 kJ mol1. This limit is a consequence of the

Second Law of Thermodynamics. The fraction of heat that can be converted into work is a function of

how far (in temperature) the heat falls as it flows through the engine and into the surroundings; this

fraction is given by (1 - Thigh)/Tlow. At normal environmental temperatures of around 300K, Thigh

would have to be at least 600 K for 50% thermal efficiency.

The major limitation of present fuel cells is that the rates of the electrode reactions, especially the one in

which oxygen is reduced, tend to be very small, and thus so is the output current per unit of electrodesurface. Coating the electrode with a suitable catalytic material is almost always necessary to obtain

usable output currents, but good catalysts are mostly very expensive substances such as platinum, so

that the resulting cells are too costly for most practical uses. There is no doubt that if an efficient, low-

cost catalytic electrode surface is ever developed, the fuel cell would become a mainstay of the energy

economy.
http://www.chem1.com/acad/webtext/thermeq/TE3.html#ENGINEhttp://www.chem1.com/acad/webtext/thermeq/TE3.html


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A portable fuel cell developed by

NASA/JPL

[link]

This "carbon-air" fuel cell was built by William

Jaques in the mid-1800s. It was based on an

1855 design by E. Becquerel and put out 1.5

kw. The electrolyte was molten KNO3 (hence

the furnace) and the anode reaction was

C + 4OH CO2 + 2H2O + e

Microbobial Fuel Cells

Certain types of bacteria are able to oxidize organic compounds to carbon dioxide while directly

transferring electrons to electrodes. These so-called electricigen organisms may make it possible to

convert renewable biomass and organic waste directly into electricity without the wasted energy and

pollution produced by direct combustion. In one experiment, a graphite electrode immersed in ordinary

mud (containing humic materials) was able to produce measurable amounts of electricity.

Summary

Make sure you thoroughly understand the following essential ideas which have been presented above. It

is especially imortant that you know the precise meanings of all the highlighted terms in the context of

this topic.

A battery is a galvanic cell in which some of the free energy change associated with a spontaneouselectron-transfer reaction is captured in the form of electrical energy.

A secondary or storage battery is one in which the electron-transfer reaction can be reversed byapplying a charging current from an external source.A fuel cell is a special type of battery in which the reactants are supplied from an external source as

power is produced. In most practical fuel cells, H+ ions are produced at the anode (either from H2 or

hydrocarbon) and oxygen from the air is reduced to H2O at the cathode.

The cathodic reduction of O2 is kinetically limited, necessitating the use of electrode surfaces having

high catalytic activity.The electrodes in batteries must have very high effective surface areas, and thus be highly porous.This requirement may conflict with the other important one of efficient diffusion of reactants andproducts in the narrow channels within the pores.Batteries and fuel cells designed to power vehicles and portable devices need to have high charge-to

weight and charge-tovolume ratios.

Outside Links

Another very good electric eel siteFor more information, see this article by Derek Loveley of U. Mass Amherst.Smithsonian fuel cell project A well-designed site with information on many types of fuel cells and nice section (with pictures) of early ones.Other major fuel cell types A more compact site with good illustrations

Contributors
http://www.nfcrc.uci.edu/fcresources/FCexplained/FC_Types.htmhttp://fuelcells.si.edu/http://www.asm.org/microbe/index.asp?bid=43711http://www.chm.bris.ac.uk/webprojects2001/riis/electriceels.htmhttp://nobelprize.org/nobel_prizes/physics/laureates/1903/becquerel-bio.htmlhttp://www.jpl.nasa.gov/releases/2002/release_2002_94.html


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Electrochemistry 15

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Transcript of Electrochemistry 15