Here we discuss different of the battery designs currently used, some of the chemistry
involved, and advantages and disadvantages of each design. We have also included some useful definitions and a list of
parameters to guide you in matching your battery requirements to a specific battery design.
- More Technical Resources on the PowerStream Web Site
Table of Contents
- Anode: The electrode where oxidation (loss of electrons) takes place. While discharging, it is the
negative electrode; while charging it becomes the positive electrode.
- Amps: Also known as Amperes. This is the rate at which electrons flow in a wire. The units are
coulombs per second, or since an electron has a charge of 1.602 x 10-19 coulombs, an amp is 6.24 x 10+18
electrons per second. Think of marbles rolling through a tube. If 6.24 x 10+18 pass by in 1 second you wuld have an
amp of marbles.
- Amp hours: Also known as ampere hours. This a measure of the amount of charge stored or used.
For example if you had an amp of marbles flowing out of your tube into a bucket for an hour, you would have one amp-hour of
marbles in the bucket ( 6.24 x 10+18 times 3600 seconds = 2.2 x 10+22 marbles. A 1 amp hour battery
contains enough charge to supply 1 amp for 1 hour, if you discharge at the same rate. Usually if you discharge faster than the
rate at which the the amp hours were specified you will get fewer amp-hours out. You may notice that amp-hours and coulombs
measure the same quantity-charge. One amp-hour is 3600 coulombs, but amp hours are easier to use in battery design. So
remember, amps are flow ( "this motor requires 2 amps to run at 1800 rpm.") Amp hours measure capacity, quantity, or amount of
charge ("this 100 amp-hour battery will supply 2 amps for 50 hours before recharge." Amp-hours are amps times hours, not
amps divided by hours.
So Amp-Hours, (AH), or milliamp-Hours (mAH) is a measure of the size of the battery a 10 mAH battery has half
the capacity of a 20 mAH battery, even though they may be in the same physical package.
- Batteries: Two or more electrochemical cells, electrically interconnected, each of which contains two
electrodes and an electrolyte. The redox (oxidation-reduction) reactions that occur at these electrodes convert electrochemical
energy into electrical energy. In everyday usage, 'battery' is also used to refer to a single cell.
- C:C represents the capacity of a battery divided by 1 hour, its units are amps. It represents a 1
hour discharge rate using the nominal capacity of the battery. So a discharge rate of 10C for a 5AH battery would be 50 amps.
The concept of "C" is also used for charge currents, since both charge and discharge properties are proportional to the
capacity of the battery, so a 5C charge rate for a 5 AH battery would be 25 amps.
- Capacity: The total quantity of electricity or total ampere-hours available from a fully charged cell
or battery.
- Cathode: The electrode where reduction (gain of electrons) takes place. When discharging, it is the
positive electrode, when charging, it becomes the negative electrode.
- Charge: The conversion of electrical energy, provided in the form of current from an external source,
into chemical energy stored at the electrodes of a cell or battery.
- Discharge: The conversion of the chemical energy of a cell into electrical energy, which can then be
used to supply power to a system.
- Discharge curve: A plot of cell voltage over time into the discharge, at a constant temperature and
constant current discharge rate.
Each curve in this graph represents cell performance at a different discharge rate. The farther right the
curve ends, the lower the discharge rate (Crompton 31.4).
- Dry cell: A Leclanché cell, so called because of its non-fluid electrolyte (to prevent spillage). This is achieved by
adding an inert metal oxide so that the electrolyte forms a gel or paste.
- Efficiency: For a secondary cell, the ratio of the output on discharge to the input required to
restore it to its initial state of charge under specified conditions. Can be measured in ampere-hour, voltage, and watt-hour
efficiency.
- Electrolyte: The chemistry of a battery requires a medium that provides the ion transport mechanism
between the positive and negative electrodes of a cell.
- Energy density (specific energy): These two terms are often used interchangeably. Energy
density refers mainly to the ratio of a battery's available energy to its volume (watt hour/liter). Specific energy
refers to the ratio of energy to mass (watt hour/kg). The energy is determined by the charge that can be stored and the cell
voltage (E=qV).
- Fuel cell: A cell in which one or both of the reactants are not permanently contained in the cell,
but are continuously supplied from a source external to the cell and the reaction products continuously removed. Unlike the
metal anodes typically used in batteries, the fuels in a fuel cell are usually gas or liquid, with oxygen as the oxidant. The
hydrogen/oxygen fuel cell is the most common. In this fuel cell, hydrogen is oxidized at the anode:
half-reaction |
V vs SHE |
2H2 > 4H+ + 4e- |
0 |
4H+ + O2 + 4e- > 2H2O |
1.2 |
Hydrogen/oxygen fuel cell systems work well in space travel applications because of their
high efficiency, high power-to-weight and volume ratios, and usable reaction product (water). They can function for many months
as long as fuel is supplied and therefore the energy density cannot be measured.
- Half-reaction: Refers to the chemical processes occurring at each electrode. The potential of the two
half-reactions add to give us the overall cell potential. We can see this in the zinc mercury cell, for example:
Location |
Reaction |
Potential |
Anode |
Zn + 2OH- > Zn(OH)2 + 2e- |
1.25 V |
Cathode |
HgO +H2O + 2e- > Hg + 2OH- |
0.098 V |
Overall |
Zn + HgO + H2O > Zn(OH)2 + Hg |
1.35 V |
- Polarization: The voltage drop in a cell during discharge due to the flow of an electrical current.
The cell's internal resistance increases with the buildup of a product of oxidation or a reduction of an electrode, preventing
further reaction.
- Power: Defined by voltage (V) and current (I), P=VI.
Since V=IR, P=I2R and P=V2/R
Power also can be described by energy emitted per unit of time: P=E/t.
Thus E=VIt=qV.
- Power density (specific power): Power density is the ratio of the power available from a battery to
its volume (watt/liter). Specific power generally refers to the ratio of power to mass (watt/kg). Comparison of power to cell
mass is more common.
- Primary cells: A cell that is not designed for recharging and is discarded once it has produced all
its electrical energy.
- Prismatic: Just a word to say that the cells are not cylindrical, as nature intended battery cells to
be, but fit nicely into a parallelepiped or any other such flattened shape.
- Reserve cell: A cell that may be kept inactive and which is activated by adding an electrolyte or
electrode, or melting an electrolyte in a solid state.
- Secondary cells: A cell capable of repeated use. Its charge may be fully restored by passing an
electric current through the cell in the opposite direction to that of discharge, thus reversing the redox reactions.
No one battery design is perfect for every application. Choosing one requires compromise. That's why it's
important to prioritize your list of requirements. Decide which ones you absolutely must have and which you can compromise on.
Here are some of the parameters to consider:
- Voltage: Normal voltage during discharge, maximum and minimum permissible voltages, discharge curve
profile
- Duty cycle: Conditions the battery experiences during use. Type of discharge and current drain, e.g.,
continuous, intermittent, continuous with pulses, etc.
- Temperature: In storage and in use. Temperatures that are too high or too low can greatly reduce
battery capacity.
- Shelf life: How rapidly the cell loses potential while unused.
- Service life: Defined either in calendar time or, for secondary cells, possible number of
discharge/charge cycles, depending on the battery application. Service life depends on battery design and operational
conditions, i.e., the stress put on a battery. For stationary and motive power application, the end of service life is defined
as the point at which a battery's capacity drops to 80% of its original capacity. Exceptions would include car batteries where
the service life ends when the capacity falls below 60%.
- Physical restrictions: These include dimensions, weight, terminals, etc.
- Maintenance and resupply: Ease of battery acquisition, replacement, charging facilities,
disposal.
- Safety and reliability: Failure rates, freedom from outgassing or leakage; use of toxic components;
operation under hazardous conditions; environmentally safe
- Cost: Initial cost, operating cost, use of expensive materials
- Internal resistance: Batteries capable of a high-rate discharge must have a low internal
resistance.
- Specific energy: As discussed in the definition section, this is a measurement of possible stored
energy per kilogram of mass. This number is purely theoretical as it does not take into account the mass of inactive materials,
nor the variation in chemical reactions.
- Specific power: Also defined in the definitions section, a P=E/t, so the specific power is discussed
at a specific discharge rate. It is possible for batteries with a high specific energy to have a low power density if they
experience large voltage drops at high discharge rates. Specific power and specific energy can be compared in a Ragone
plot.
- Unusual requirements: Very long-term or extreme-temperature storage; very low failure rate; no
voltage delay, etc.
Of course the ideal battery would perform well in all these areas with a long shelf and service life, high
specific energy and specific power, low initial and maintenance costs, low environmental impact, and good performance in a
variety of conditions (temperatures, duty cycles, etc.). When you find one that meets all these requirements, let us know! In
the meantime, we have to make do with batteries that work very well in specific applications.
Primary Batteries
Leclanché Cells(zinc carbon or dry cell)
Anode: Zinc
Cathode: Manganese Dioxide (MnO2)
Electrolyte: Ammonium chloride or zinc chloride dissolved in water
Applications: Flashlights, toys, moderate drain use
The basic design of the Leclanché cell has been around since the 1860s, and until World War II, was the
only one in wide use. It is still the most commonly used of all primary battery designs because of its low cost, availability,
and applicability in various situations. However, because the Leclanché cell must be discharged intermittently for best
capacity, much of battery research in the last three decades has focused on zinc-chloride cell systems, which have been found
to perform better than the Leclanché under heavier drain.
This figure shows typical discharge curves for general-purpose Leclanché zinc chloride D-size cells
discharge 2 h/day at 20º C. Solid linezinc chloride; broken lineLeclanché (Linden 8.18). The
zinc-chloride cell has a higher service life and voltage than the Leclanché (at both higher and lower discharge
rates).
In an ordinary Leclanché cell the electrolyte consists (in percent of atomic weight) of 26%
NH4Cl (ammonium chloride), 8.8% ZnCl2 (zinc chloride), and 65.2% water. The overall cell reaction can be
expressed:
Zn + 2MnO2 +2NH4Cl > 2MnOOH + Zn(NH3)2Cl2
E=1.26
The electrolyte in a typical zinc chloride cell consists of 15-40% ZnCl2 and 60-85% water,
sometimes with a small amount of NH4Cl for optimal performance. The overall cell reaction of the zinc chloride as
the electrolyte can be expressed:
Zn + 2MnO2 + 2H2O + ZnCl2 > 2MnOOH + 2Zn(OH)Cl
MnO2, is only slightly conductive, so graphite is added to improve conductivity. The cell voltage
increases by using synthetically produced manganese dioxide instead of that found naturally (called pyrolusite). This does
drive the cost up a bit, but it is still inexpensive and environmentally friendly, making it a popular cathode.
These cells are the cheapest ones in wide use, but they also have the lowest energy density and perform poorly
under high-current applications. Still, the zinc carbon design is reliable and more than adequate for many everyday
applications.
Anode: Zinc powder
Cathode: Manganese dioxide (MnO2) powder
Electrolyte: Potassium hydroxide (KOH)
Applications: Radios, toys, photo-flash applications, watches, high-drain applications
This cell design gets its name from its use of alkaline aqueous solutions as electrolytes. Alkaline battery
chemistry was first introduced in the early 60s. The alkaline cell has grown in popularity, becoming the zinc-carbon
cell's greatest competitor. Alkaline cells have many acknowledged advantages over zinc-carbon, including a higher energy
density, longer shelf life, superior leakage resistance, better performance in both continuous and intermittent duty cycles,
and lower internal resistance, which allows it to operate at high discharge rates over a wider temperature range.
Zinc in a powdered form increases the surface area of the anode, allowing more particle interaction. This
lowers the internal resistance and increases the power density. The cathode, MnO2, is synthetically produced because
of its superiority to naturally occurring MnO2. This increases the energy density. Just as in the zinc carbon cell,
graphite is added to the cathode to increase conductivity. The electrolyte, KOH, allows high ionic conductivity. Zinc oxide is
often added to slow down corrosion of the zinc anode. A cellulose derivative is thrown in as well as a gelling agent. These
materials make the alkaline cell more expensive than the zinc-carbon, but its improved performance makes it more cost
effective, especially in high drain situations where the alkaline cell's energy density is much higher.
The half-reactions are:
Zn + 2 OH- > ZnO + H2O + 2 e-
2 MnO2 + H2O + 2 e- >Mn2O3 + 2
OH-
The overall reaction is:
Zn + 2MnO2 > ZnO + Mn2O3 E=1.5 V
There are other cell designs that fit into the alkaline cell category, including the mercury oxide, silver
oxide, and zinc air cells. Mercury and silver give even higher energy densities, but cost a lot more and are being phased out
through government regulations because of their high toxicity as heavy metals. The mercury oxide, silver oxide, and zinc air
(which is being developed for electronic vehicles) are all discussed below.
Mercury
Oxide Cells
Anode: Zinc (or cadmium)
Cathode: Mercuric Oxide (HgO)
Electrolyte: Potassium hydroxide
Applications: Small electronic equipment, hearing aids, photography, alarm systems, emergency beacons,
detonators, radio microphones
This is an obsolete technology. Most if not all of the manufacture of these cells has been stopped by
government regulators. Mercury batteries come in two main varieties: zinc/mercuric oxide and cadmium/mercuric oxide. The
zinc/mercuric oxide system has high volumetric specific energy (400 Wh/L), long storage life, and stable voltage. The
cadmium/mercuric oxide system has good high temperature and good low temperature (-55 C to +80 C, some designs to +180 C) and
has very low gas evolution.
Basic Cell Reaction |
Voltage |
Electrochemical Efficiency |
Zn + HgO = ZnO + Hg |
1.35 V |
820 mAH/g(Zn), 250 mAH/g(Hg) |
Cd + HgO + H2O = Cd(OH2) + Hg |
0.91 V |
480 mAH/g(Cd) |
The electrolytes used in mercury cells are sodium and/or potassium hydroxide solutions,
making these alkaline cells. These cells are not rechargeable.
Zinc/Air Cells
Anode: Amalgamated zinc powder and electrolyte
Cathode: Oxygen (O2)
Electrolyte: Potassium hydroxide (KOH)
Applications: Hearing aids, pagers, electric vehicles
The zinc air cell fits into the alkaline cell category because of its electrolyte. It also acts as a partial
fuel cell because it uses the O2 from air as the cathode. This cell is interesting technology, even aside from the
question "how do you use air for an electrode?" Actually, oxygen is let in to the cathode through a hole in the battery and is
reduced on a carbon surface.
A number of battery chemistries involve a metal oxide and zinc. The metal oxide reduces, the zinc becomes
oxidized, and electric current results. A familiar example is the old mercury oxide/zinc batteries used for hearing aids. If
you leave out the metal oxide you could double the capacity per unit volume (roughly), but where would you get the oxygen?
Right!
First let's look at the electrochemical reactions and find that the open cell voltage should be 1.65
volts:
Location |
Half Cell reactions |
Voltage |
Anode |
Zn2+ + 2OH- > Zn(OH)2 |
1.25 |
Cathode |
1/2 O2 + H2O + 2e > 2 OH- |
0.4 |
Overall |
2Zn +O2 +2H2O > 2Zn(OH)2 |
1.65 |
The electrolyte is an alkali hydroxide in 20-40% weight solution with water. One disadvantage is that since
these hydroxides are hygroscopic, they will pick up or lose water from the air depending on the humidity. Both too little and
too much humidity reduces the life of the cell. Selective membranes can help. Oxygen from the air dissolves in the electrolyte
through a porous, hydrophobic electrodea carbon-polymer or metal-polymer composite.
Since there is no need to carry around the cathode, the energy density of these batteries can be quite high,
between 220300 Wh/kg (compared to 99123 Wh/kg with a HgO cathode), although the power density remains low. However,
the use of potassium or sodium hydroxides as the electrolyte is a problem, since these can react with carbon dioxide in the air
to form alkali carbonates. For this reason large zinc air batteries usually contain a higher volume of CO2 absorbing
material (calcium oxide flake) than battery components. This can cancel out the huge increase in energy density gained by using
the air electrode.
This cell has the additional benefits of being environmentally friendly at a relatively low cost.
These batteries can last indefinitely before they are activated by exposing them to air, after which they have
a short shelf life. For this reason (as well as the high energy density) most zinc-air batteries are used in hearing aids.
There is a company promoting them for use in electric vehicles also because they are environmentally friendly and cost
relatively little. The idea is to have refueling stations where the zinc oxide waste can be replaced by fresh zinc pellets.
Aluminum / Air Cells
Although, to our way of thinking, the metal/air batteries are strictly primary, cells have been designed to
have the metal replaceable. These are called mechanically rechargeable batteries. Aluminum/air is an example of such a
cell. Aluminum is attractive for such cells because it is highly reactive, the aluminum oxide protective layer is dissolved by
hydroxide electrolytes, and it has a nice, high voltage. The overall chemical reaction is:
Location |
Half Cell reactions |
Voltage |
Anode |
Al + 4 OH-> Al(OH)4- + 3e |
-2.35 |
Cathode |
3/4 O2 + 3/2 H2O + 3e> 3OH- |
0.40 |
Overall |
Al + 3/2 HO + 3/4 O2 > Al(OH)3 |
2.75 V |
As I mentioned above, alkali (chiefly potassium hydroxide) electrolytes are used, but so also are neutral salt
solutions. The alkali cell has some problem with the air electrode, because the hydroxide ion makes a gel in the porous
electrode, polarizing it. The typical aluminum hydroxide gel is a problem on either electrode because it sucks up a lot of
water. Using a concentrated caustic solution prevents this, but is very reactive with the aluminum electrode, producing
hydrogen gas. Another way to prevent the gel formation is to seed the electrolyte with aluminum trihydroxide crystals. These
act to convert the aluminum hydroxide to aluminum trihydroxide as the crystals grow. To prevent hydrogen gas evolution tin and
zinc have been used as corrosion inhibitors. A number of additives are used to control the reactions. A disadvantage of the
alkaline electrolyte is that it reacts with atmospheric carbon dioxide.
Aluminum / air cells have also been made for marine applications. These are "rechargeable" by replacing the
seawater electrolyte until the aluminum is exhausted, then replacing the aluminum. Some cells that are open to seawater have
also been researched. Since salt water solutions tend to passivate the aluminum, pumping the electrolyte back and forth along
the cell surface has been successful. For those cells that don't need to use ocean water, an electrolyte of KCL and KF
solutions is used.
Air electrodes of Teflon-bonded carbon are used without a catalyst.
v
Lithium
Cells
Applications: Pacemakers, defibrillators, watches, meters, cameras, calculators, portable, low-power
use
Lithium battery chemistry comprise a number of cell designs that use lithium as the anode. Lithium is gaining
a lot of popularity as an anode for a number of reasons. In this comparison of anode materials, we can see some reasons
why:
Anode |
Atomic mass (g) |
Standard potential (V) |
Density g/cm3 |
Melting point ºC |
Electrochemical Equivalence (Ah/g) |
Li |
6.94 |
3.05 |
0.54 |
180 |
3.86 |
Na |
23.0 |
2.7 |
0.97 |
97.8 |
1.16 |
Mg |
24.3 |
2.4 |
1.74 |
650 |
2.20 |
Al |
26.9 |
1.7 |
2.7 |
659 |
2.98 |
Ca |
40.1 |
2.87 |
1.54 |
851 |
1.34 |
Fe |
55.8 |
0.44 |
7.85 |
1528 |
0.96 |
Zn |
65.4 |
0.76 |
7.1 |
419 |
0.82 |
Cd |
112 |
0.40 |
8.65 |
321 |
0.48 |
Pb |
207 |
0.13 |
11.3 |
327 |
0.26 |
Note that lithium, the lightest of the metals, also has the highest standard potential of all the metals, at
over 3 V. Some of the lithium cell designs have a voltage of nearly 4 V. This means that lithium has the highest energy
density. Many different lithium cells exist because of its stability and low reactivity with a number of cathodes and
non-aqueous electrolytes. The most common electrolytes are organic liquids with the notable exceptions of SOCl2
(thionyl chloride) and SO2Cl2 (sulfuryl chloride). Solutes are added to the electrolytes to increase
conductivity.
Lithium cells have only recently become commercially viable because lithium reacts violently with water, as
well as nitrogen in air. This requires sealed cells. High-rate lithium cells can build up pressure if they short circuit and
cause the temperature and pressure to rise. Thus, the cell design needs to include weak points, or safety vents, which rupture
at a certain pressure to prevent explosion.
Lithium cells can be grouped into three general categories: liquid cathode, solid cathode, and solid
electrolyte. Let's look at some specific lithium cell designs within the context of these three categories.
v
Liquid cathode lithium cells:
These cells tend to offer higher discharge rates because the reactions occur at the cathode surface. In a
solid cathode, the reactions take longer because the lithium ions must enter into the cathode for discharge to occur. The
direct contact between the liquid cathode and the lithium forms a film over the lithium, called the solid electrolyte interface
(SEI). This prevents further chemical reaction when not in use, thus preserving the cell's shelf life. One drawback, though, is
that if the film is too thick, it causes an initial voltage delay. Usually, water contamination is the reason for the thicker
film, so quality control is important.
LiSO2 LithiumSulfur Dioxide
This cell performs very well in high current applications as well as in low temperatures. It has an open
voltage of almost 3 V and a typical energy density of 240280 Wh/kg. It uses a cathode of porous carbon with sulfur
dioxide taking part in the reaction at the cathode. The electrolyte consists of an acetonitrile solvent and a lithium bromide
solute. Polypropylene acts as a separator. Lithium and sulfur dioxide combine to form lithium dithionite:
2Li + 2SO2 > Li2S2O4
These cells are mainly used in military applications for communication because of high cost and safety
concerns in high-discharge situations, i.e., pressure buildup and overheating.
LiSOCl2 Lithium Thionyl Chloride
This cell consists of a high-surface area carbon cathode, a non-woven glass separator, and thionyl chloride,
which doubles as the electrolyte solvent and the active cathode material. Lithium aluminum chloride (LiAlCl4) acts
as the electrolyte salt.
The materials react as follows:
Location |
Reaction |
Anode |
Li > Li+ + e- |
Cathode |
4Li+ + 4e- + 2SOCl2 > 4LiCl + SO2 + S |
Overall |
4Li + 2SOCl2 > 4LiCl + SO2 + S |
During discharge the anode gives off lithium ions. On the carbon surface, the thionyl chloride reduces to
chloride ions, sulfur dioxide, and sulfur. The lithium and chloride ions then form lithium chloride. Once the lithium chloride
has deposited at a site on the carbon surface, that site is rendered inactive. The sulfur and sulfur dioxide dissolve in the
electrolyte, but at higher-rate discharges SO2 will increase the cell pressure.
This system has a very high energy density (about 500 Wh/kg) and an operating voltage of 3.33.5 V. The
cell is generally a low-pressure system
In high-rate discharge, the voltage delay is more pronounced and the pressure increases as mentioned before.
Low-rate cells are used commercially for small electronics and memory backup. High-rate cells are used mainly for military
applications.
Solid cathode lithium cells:
These cells cannot be used in high-drain applications and don't perform as well as the liquid cathode cells in
low temperatures. However, they don't have the same voltage delay and the cells don't require pressurization. They are used
generally for memory backup, watches, portable electronic devices, etc.
LiMnO2
These account for about 80% of all primary lithium cells, one reason being their low cost. The cathode used is
a heat-treated MnO2 and the electrolyte a mixture of propylene carbonate and 1,2-dimethoyethane. The half reactions
are
Anode |
Li > Li+ + e |
Cathode |
MnIVO2 + Li+ + e >
MnIIIO2(Li+) |
Overall |
Li + MnIVO2 > MnIIIO2(Li+) |
At lower temperatures and in high-rate discharge, the LiSO2 cell performs much better than the
LiMnO2 cell. At low-rate discharge and higher temperatures, the two cells perform equally well, but
LiMnO2 cell has the advantage because it doesn't require pressurization.
Li(CF)n Lithium polycarbon monofluoride
The cathode in this cell is carbon monofluoride, a compound formed through high-temperature intercalation.
This is the process where foreign atoms (in this case fluorine gas) incorporate themselves into some crystal lattice (graphite
powder), with the crystal lattice atoms retaining their positions relative to one another.
A typical electrolyte is lithium tetrafluorobate (LiBF4) salt in a solution of propylene carbonate
(PC) and dimethoxyethane (DME).
Anode |
Li > Li+ + e |
Cathode |
MnIVO2 + Li+ + e >
MnIIIO2(Li+) |
Overall |
Li + MnIVO2 > MnIIIO2(Li+) |
These cells also have a high voltage (about 3.0 V open voltage) and a high energy density (around 250 Wh/kg).
All this and a 7-year shelf life makes them very suitable for low- to moderate-drain use, e.g., watches, calculators, and
memory applications.
v
Solid electrolyte lithium cells:
All commercially manufactured cells that use a solid electrolyte have a lithium anode. They perform best in
low-current applications and have a very long service life. For this reason, they are used in pacemakers
LiI2Lithium iodine cells use solid LiI as their electrolyte and also produce LiI as the cell
discharges. The cathode is poly-2-vinylpyridine (P2VP) with the following reactions:
Anode |
2Li > 2Li+ + 2e |
Cathode |
2Li+ + 2e + P2VP· nI2 > P2VP· (n1)I2 + 2LiI |
Overall |
2Li + P2VP· nI2 > P2VP·
(n1)I2 +2LiI |
LiI is formed in situ by direct reaction of the electrodes.
Lithium-Iron Cells
The Lithium-Iron chemistry deserves a separate section because it is one of a handful of lithium metal systems
that have a 1.5 volt output (others are lithium/lead bismuthate, lithium/bismuth trioxide, lithium/copper oxide, and
lithium/copper sulfide). Recently consumer cells that use the Li/Fe have reached the market, including the Energizer. These
have advantage of having the same voltage as alkaline batteries with much more energy storage capacity, so they are called
"voltage compatible" lithiums. They are not rechargeable. They have about 2.5 times the capacity of an alkaline battery of the
same size, but only under high current discharge conditions (digital cameras, flashlights, motor driven toys, etc.). For small
currents they don't have any advantage. Another advantage is the low self-discharge rate10 years storage is quoted by the
manufacturer. The discharge reactions are:
Type |
Reaction |
Nominal Voltage |
Range |
FeS2 Version |
2 FeS2 + 4 Li > Fe + 2Li2S |
1.6 Volts |
1.6-1.4 V |
FeS Version |
FeS + 2Li > Fe + Li2S |
1.5 Volts |
1.5-1.2 V |
Both Iron sulfide and Iron disulfide are used, the FeS2 is used in the Energizer. Electrolytes are
organic materials such as propylene carbonate, dioxolane and dimethoxyelthane
Magnesium-Copper Chloride Reserve Cells
The magnesium-cuprous chloride system is a member of the reserve cell family. It can't be used as a primary
battery because of its high self-discharge rate, but it has a high discharge rate and power density, so it can be made "dry
charged" and sit forever ready, just add water. The added advantage of being light-weight has made these practical for portable
emergency batteries.
It works by depositing copper metal out onto the magnesium anode, just like the old copper-coated
nail experiment.
Variations of this battery use silver chloride, lead chloride, copper iodide, or copper thiocyanate to
react with the magnesium.
The water does not have to be pure, sea water, tap water, or even bio-derived waste fluids
have been used. The torpedo batteries force seawater through the battery to get up to 460 kW of power to drive the propeller.
Type |
Reaction |
Nominal Voltage |
Range |
Mg CuCl |
Mg + 2 CuCl > MgCl2+ 2 Cu |
1.6 Volts |
1.5-1.6V |
Secondary batteries
Anode: Sponge metallic lead
Cathode: Lead dioxide (PbO2)
Electrolyte: Dilute mixture of aqueous sulfuric acid
Applications: Motive power in cars, trucks, forklifts, construction equipment, recreational water
craft, standby/backup systems
Used mainly for engine batteries, these cells represent over half of all battery sales. Some advantages are
their low cost, long life cycle, and ability to withstand mistreatment. They also perform well in high and low temperatures and
in high-drain applications. The chemistry lead acid battery half-cell reactions are:
half-reaction |
V vs SHE |
Pb + SO42- > PbSO4 + 2e- |
.356 |
PbO2 + SO42- + 4H+ + 2e- >
PbSO4 + 2H2O |
1.685 |
There are a few problems with this design. If the cell voltages exceed 2.39 V, the water breaks down into
hydrogen and oxygen (this so-called gassing voltage is temperature dependent, for a chart of the temperature dependence click
here ). This requires replacing the cell's water. Also, as the hydrogen and oxygen vent from the cell, too high a concentration
of this mixture will cause an explosion. Another problem arising from this system is that fumes from the acid or hydroxide
solution may have a corrosive effect on the area surrounding the battery.
These problems are mostly solved by sealed cells, made commercially available in the 1970s. In the case of
lead acid cells, the term "valve-regulated cells" is more accurate, because they cannot be sealed completely. If they were, the
hydrogen gas would cause the pressure to build up beyond safe limits. Catalytic gas recombiners do a great deal to alleviate
this problem. They convert the hydrogen and oxygen back into water, achieving about 85% efficiency at best. Although this
doesn't entirely eliminate the hydrogen and oxygen gas, the water lost becomes so insignificant that no refill is needed for
the life of the battery. For this reason , these cells are often referred to as maintenance-free batteries. Also, this cell
design prevents corrosive fumes from escaping.
These cells have a low cycle life, a quick self discharge, and low energy densities (normally between 30 and
40 Wh/kg). However, with a nominal voltage of 2 V and power densities of up to 600 W/kg, the lead-acid cell is an adequate, if
not perfect, design for car batteries.
Anode: Cadmium
Cathode: Nickel oxyhydroxide Ni(OH)2
Electrolyte: Aqueous potassium hydroxide (KOH)
Applications: Calculators, digital cameras, pagers, lap tops, tape recorders, flashlights, medical
devices (e.g., defibrillators), electric vehicles, space applications
The cathode is nickel-plated, woven mesh, and the anode is a cadmium-plated net. Since the cadmium is just a
coating, this cell's negative environmental impact is often exaggerated. (Incidentally, cadmium is also used in TV tubes, some
semiconductors, and as an orange-yellow dye for plastics.) The electrolyte, KOH, acts only as an ion conductor and does not
contribute significantly to the cell's reaction. That's why not much electrolyte is needed, so this keeps the weight down.
(NaOH is sometimes used as an electrolyte, which doesn't conduct as well, but also doesn't tend to leak out of the seal as
much). Here are the cell reactions:
Reaction |
V vs SHE |
Cd + 2OH- > Cd(OH)2 + 2e- |
0.81 |
NiO2 + 2H2O + 2e- > Ni(OH)2 + 2OH- |
0.49 |
Cd +NiO2 + 2H2O > Cd(OH)2 + Ni(OH)2 |
1.30 |
Advantages include good performance in high-discharge and low-temperature applications. They also have long
shelf and use life. Disadvantages are that they cost more than the lead-acid battery and have lower power densities. Possibly
its most well-known limitation is a memory effect, where the cell retains the characteristics of the previous cycle.
This term refers to a temporary loss of cell capacity, which occurs when a cell is recharged without being
fully discharged. This can cause cadmium hydroxide to passivate the electrode, or the battery to wear out. In the former case,
a few cycles of discharging and charging the cell will help correct the problem, but may shorten the lifetime of the battery.
The true memory effect comes from experience with a certain style of NiCad in space use, which were cycled within a few percent
of discharge each time.
An important thing to know about "conditioning " a NiCd battery is that the deep discharge spoken of is not a
discharge to zero volts, but to about 1 volt per cell.
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Anode:Hydrogen Gas
Cathode: Nickel oxyhydroxide
Electrolyte: Potassium hydroxide
Applications:Space satellites that require long cycle life, over 40,000 cycles. Nickel/Hydrogen
batteries have a high self-discharge rate, something like 80% a month, which isn't a problem for satellite applications.
The NiH2 cell is a welded pressure vessel. It has a high specific energy, 60WH/kg, long life, can
tolerate overcharge and cell reversal, but has a low volumetric energy density, 50 WH/liter.
Here are the cell reactions:
Location |
Reactions |
Voltage |
Anode |
½H2 + OH- > H2O + e- |
0.83 |
Cathode |
NiOOH + H2O + e- > Ni(OH)2 + OH- |
0.52 |
Overall |
NiOOH + ½H2 > Ni(OH)2 |
1.35 |
In order to get the hydrogen gas into solution a Teflon-bonded platinum black catalyst is used, similar to
that used in fuel cells. This platinum electrode has the added advantage that it can recombine oxygen with hydrogen extremely
fast. Since the only bad chemical reaction during over charge is the creation of oxygen at the positive electrode this means
that the Nickel/Hydrogen battery is impossible to overcharge (though there may be a thermal runaway problem if the excess heat
isn't dissipated.) A similar reaction keeps any damage from being done if the cell is reverse-charged.
The battery weight for a 10kW satellite is about 350 kg, or 770 lbs.
Anode: Rare-earth or nickel alloys with many metals
Cathode: Nickel oxyhydroxide
Electrolyte: Potassium hydroxide
Applications: Cellular phones, camcorders, emergency backup lighting, power tools, laptops, portable,
electric vehicles
This sealed cell is a hybrid of the NiCd and NiH2 cells. Previously, this battery was not available
for commercial use because, although hydrogen has wonderful anodic qualities, it requires cell pressurization. Fortunately, in
the late 1960s scientists discovered that some metal alloys (hydrides such as LiNi5 or ZrNi2) could store
hydrogen atoms, which then could participate in reversible chemical reactions. In modern NiMH batteries, the anode consists of
many metals alloys, including V, Ti, Zr, Ni, Cr, Co, and Fe.
Except for the anode, the NiMH cell very closely resembles the NiCd cell in construction. Even the voltage is
virtually identical, at 1.2 volts, making the cells interchangeable in many applications. Here are the cell reactions:
Location |
Reactions |
Voltage |
Anode |
MH + OH- > M + H2O + e- |
0.83 |
Cathode |
NiOOH + H2O + e- > Ni(OH)2 + OH- |
0.52 |
Overall |
NiOOH + MH > Ni(OH)2 + M |
1.35 |
The anodes used in these cells are complex alloys containing many metals, such as an alloy of V, Ti, Zr, Ni,
Cr, Co, and (!) Fe. The underlying chemistry of these alloys and reasons for superior performance are not clearly
understood, and the compositions are determined by empirical testing methods.
A very interesting fact about these alloys is that some metals absorb heat when absorbiong hydrogen, and some
give off heat when absorbing hydrogen. Both of these are bad for a battery, since we would like the hydregen to move easily in
and out without any energy transfer. The successful alloys are all mixtures of exothermic and endothermic metals to achieve
this.
Hydrogen Storage Metals Comparison:
Material |
Density |
H2 Storage Capacity |
LaNi5 |
8.3 |
0.11 g/cc |
FeTi |
6.2 |
0.11 |
Mg2Ni |
4.1 |
0.15 |
Mg |
1.74 |
0.13 |
MgNi Eutectic |
2.54 |
0.16 |
liquid H2 |
0.07 |
0.07 |
Please notice that the density of hydrogen stored in a metal hydride is higher than that of pure liquid
hydrogen! Commercial NiMH batteries are mostly of the rare earth-nickel type, of which LaNi5 is a representative.
These alloys can store six hydrogen atoms per unit cell such as LaNi5H6. Even misch metal nickel alloys
are used to save the cost of separation.
The electrolyte of commercial NiMH batteries is typically 6 M KOH
The NiMH cell does cost more and has half the service life of the NiCd cell, but it also has 30% more
capacity, increased power density (theoretically 50% more, practically 25% more). The memory effect, which was at one time
thought to be absent from NiMH cells, is present if the cells are treated just right. To avoid the memory effect fully
discharge once every 30 or so cycles. There is no clear winner between the two. The better battery depends on what
characteristics are more crucial for a specific application.
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Anode: Molten sodium
Cathode: Molten sulfur
Electrolyte: Solid ceramic beta alumina (ß"-Al2O3)
Applications: Electric vehicles, aerospace (satellites)
This cell have been studied extensively for electric vehicles because of its inexpensive materials, high cycle
life, and high specific energy and power. Specific energies have reached levels of 150 W-h/kg and specific powers of 200 W/kg.
The half-reactions are:
half-reaction |
V vs SHE |
2Na > 2Na+ + 2e- |
|
3S + 2e- > S32- |
|
2Na + 3S > Na2S3 2.076 V
Despite these advantages there are couple of disadvantages serious enough that other alternatives, such as
lithium-ion, nickel-metal hydride, and lithium polymer, have emerged as the most promising solutions to electric vehicle power.
One is that the power output is very small at room temperature. The temperature must be kept at around 350 ºC to keep the
sulfur in liquid form and to be effective in motive power applications. This is achieved through insulation or heating through
the cells own power. This lowers the energy density.
The second problem has to do with electrolyte breakdown, which is one of the principal causes of sodium sulfur
cell failure. The electrolyte, ceramic beta"-alumina, has several attractive characteristics. It has all the benefits of a
solid electrolyte with the added qualities of a high ionic conductivity with a small electronic transfer, all with the added
benefit of being a solid. However, ceramic beta"-alumina also is brittle and develops microfissures. Thus the liquid sodium and
sulfur come in contactwith explosively violent results.
Recently, some research efforts have focussed on replacing the molten sulfur cathode with a poly(disulfide)
such as poly(ethylenedisulfide), (SSCH2CH2)n. These cells can be discharged just above the
melting temperature of Na (90 °C). The net cell reaction becomes:
2 Na + (SSR)n=Na2SSR
where the discharge reaction involves scission of the S-S disulfide linkage in the polymer backbone, and
charge involves repolymerization of the resulting dithiolate salt.
One of these is the sodium/metal chloride, which in addition to beta"-alumina has a secondary electrolyte
(NaAlCl4) to conduct ions from the first electrolyte to the cathode. This is necessary because the metal chloride is
a solid.
v Nickel/Sodium Cells These are specialty cells made by one manufacturer in England, Beta Research. They have advantages for electric
vehicles. The cell runs hot, about 300 degrees C, but this isn't a worry, since they heat themselves up during discharge. The
discharge reaction is:
Location |
Half Reaction |
Voltage |
Charge |
2 NaCl + Ni Z > 2Na +NiCl2 |
|
Discharge |
NiCl2 + 2 Na > Ni + 2NaCl |
2.58 V |
The electrolyte on the nickel side of the alumina separator is sodium
tetracloroaluminate.NaAlCl4, which melts at 151 degrees C.. Energy density is 100 to 150 Wh/kg. These use an
aluminum oxide ceramic as a separator, similar to that of the sodium-sulfur cell. They have the same danger of rupture of the
separator, but have a unique solution to the problem. The cell is encased in a two-wall steel thermally insulated package. If
the separator breaks the energy is confined within this package. A cell that is broken in this way has a low resistance, so it
can continue to reside in the battery pack without causing a vehicle break-down. This double-insulated case also prevents the
cell from spilling in car crashes.
There are no higher-voltage reactions or other side reactions, so the inventors
claim that up to the point of full charge the cell is 100% coulomb efficientmeaning that the amp-hours you put in is
exactly the same as the amp-hours you get out. Overcharging does not damage the cell, so the battery packs are easy to keep in
balancejust overcharge the whole pack.
It seems that the cell has no self-discharge if the batteries are cold,
(solid blocks of sodium don't migrate at room temperature) and that a pack requires about 24 hours to get to temperature with a
230 VAC input to the pack heater.
Anode: Carbon compound, graphite
Cathode: Lithium oxide
Electrolyte:
Applications: Laptops, cellular phones, electric vehicles
Lithium batteries that use lithium
metal have safety disadvantages when used as secondary (rechargeable) energy
sources. For this reason a series of cell chemistries have been developed using
lithium compounds instead of lithium metal. These are called generically
Lithium ion Batteries.
Cathodes consist of a a layered crystal (graphite) into which the lithium is intercalated. Experimental cells
have also used lithiated metal oxide such as LiCoO2, NiNi0.3Co0.7O2,
LiNiO2, LiV2O5, LiV6O13, LiMn4O9,
LiMn2O4, LiNiO0.2CoO2.
Electrolytes are usually LiPF6, although this has a problem with aluminum corrosion, and so
alternatives are being sought. One such is LiBF4. The electrolyte in current production batteries is liquid, and
uses an organic solvent.
Membranes are necessary to separate the electrons from the ions. Currently the batteries in wide use have
microporous polyethylene membranes.
Intercalation (rhymes with relationnot inter-cal, but in-tercal-ation) is a long-studied process which
has finally found a practical use. It has long been known that small ions (such as lithium, sodium, and the other alkali
metals) can fit in the interstitial spaces in a graphite crystal. Not only that, but these metallic atoms can go farther and
force the graphitic planes apart to fit two, three, or more layers of metallic atoms between the carbon sheets. You can imagine
what a great way this is to store lithium in a batterythe graphite is conductive, dilutes the lithium for safety, is
reasonably cheap, and does not allow dendrites or other unwanted crystal structures to form.
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Manganese-Titanium (Lithium) Cells
Anode: Lithium-Titanium Oxide
Cathode: Lithium intercalated Manganese Dioxide
Electrolyte:
Applications: Watches, other ultra-low discharge applications
This technology might be called Manganese-Titanium, but it is just another lithium coin cell. It has
"compatible" voltage 1.5 V to 1.2 Volts, like the Lithium-Iron cell, which makes it convenient for applications that
formerly used primary coin cells. It is unusual for a lithium based cell because it can withstand a continuous overcharge at
1.6 to 2.6 volts without damage. Although rated for 500 full discharge cycles, it only has a 10% a year self-discharge rate,
and so is used in solar charged watches with expected life of 15+ years with shallow discharging. The amp-hour capacity and
available current output of these cells is extremely meager. The range of capacities from Panasonic is 0.9 to 14 mAH (yes, 0.9
milliamp hours). The maximum continuous drain current is 0.1 to 0.5 mA.
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Rechargeable
Alkaline Manganese Cells
Anode: Zinc
Cathode:Manganese Dioxide
Electrolyte: Potassium Hydroxide Solution
Applications: Consumer devices
Yes, this is the familiar alkaline battery, but specially designed to be rechargeable, and with a hot new
acronymRAM (haven't I seen that acronym somewhere before?). In the charging process, direct-current electrical power is
used to reform the active chemicals of the battery system to their high-energy charge state. In the case of the RAM battery,
this involves oxidation of manganese oxyhydroxide (MnOOH) in the discharged positive electrode to manganese dioxide (MnO2), and
of zinc oxide (ZnO) in the negative electrode to metallic zinc.
Care must be taken not to overcharge to prevent electrolysis of the KOH solution electrolyte, or to charge at
voltages higher than 1.65 V (depending on temperature) to avoid the formation of higher oxides of manganese.
Nickel
Zinc Cells
Anode: Zinc
Cathode: Nickel oxide
Electrolyte: Potassium hydroxide
Applications:Electric vehicles, standby load service
The combination of nickel and zinc is very interesting because
of the low cost and low toxicity of the constituents. There have been many technical obstacles, but a string of recent patents
and a commercial start-up based on a KOH electrolyte holds great promise for applications where light weight is an issue.
The nickel/zinc battery uses zinc as the negative electrode and nickel hydroxide as the positive. The
discharge reactions are:
Location |
Half Reaction |
Voltage |
Anode |
Zn + 2OH- > Zn(OH)2+ 2e |
1.24 V |
Cathode |
2NiOOH + 2H2O > 2Ni(OH)2 + 2OH- |
0.49 V |
Overall |
2NiOOH + Zn + 2H2O > 2Ni(OH)2 + Zn(OH)2 |
1.73 |
These cells run between 1.55 and 1.65 V. Theoretical energy density is 334 Wh/kg, or about 1.3 kg of nickel
and 0.7 kg of zinc per kilowatt-hour. The internal resistance of nickel/zinc batteries is remarkably low, which makes this
system particularly attractive for high charge and discharge rates
Practical specific energy is around 60 Wh/kg. The technical problems that have plagued these batteries so far
are dissolution of the zinc in the electrolyte, and uneven redepositing of the zinc during charging. Progress in these
batteries has been mostly in the improvement of the zinc electrode. The charging is tricky because the termination voltage is a
strong function of temperature.
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Iron Nickel Cells
Anode: Iron
Cathode: Nickel oxyhydroxide
Electrolyte: Potassium hydroxide
Applications:
This battery was introduced by Thomas Edison. It is a very robust battery: it can withstand
overcharge, overdischarge, and remaining discharged for long periods of time without damage. It is good for high depths of
discharge and can have very long life even if so treated. It has low energy density, a high self-discharge rate, and evolves
hydrogen during both charge and discharge. It is often used in backup situations where it can be continuously charged and can
last for 20 years.
The chemistry involves the movement of oxygen from one electrode to the other: 3Fe + 8NiOOH + 4H2O=8 Ni(OH)2
+Fe3O4.
Half Reaction |
Voltage |
Fe + 2OH- > Fe(OH)2 +2e- |
|
3Fe(OH)2 + 2OH- > Fe3O4 + 4H2O +
2e- |
|
The open circuit voltage of this system is 1.4 V, and the discharge voltage is about 1.2 V. The
electrolyte is 30% KOH solution, with some additives.
The ability of this system to survive frequent cycling is due to the low solubility of the
reactants in the electrolyte. The formation of metallic iron on charge is slow because of the low solubility of the
Fe3O4, which is good and bad. It is good because the slow formation of iron crystals preserves the
electrode morphology. It is bad because it limits the high rate performance: these cells take a charge slowly, and give it up
slowly.
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Iron Air Cells
The Iron/Air is another of the air-electrode batteries. The electrochemistry is as follows:
Half Reaction |
Voltage |
O2 + 2Fe +2H2O=2Fe(OH)2 |
|
O2 +2H2O +2e=H2O2 +2(OH) |
|
These batteries require a high degree of support, since the CO2 must be taken out of
the air in order to prevent potassium carbonates forming in the KOH electrolyte. They have been built in large backup systems.
The air electrode consists of a catalyst on a support. For example a carbon particle substrate held together with Teflon,
coated with a silver complex catalyst. Support is provided by a silver-plated nickel screen.
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Iron Silver Cells
These have a very high energy density, and a good cycle life. It is an alkaline battery with a
KOH electrolyte, and the working materials are silver oxide and metallic iron. The high cost of these batteries have long been
a problem, but an ounce of silver in a cell phone battery would probably cost less than an ounce of the rare earths now used in
some NiMH batteries.
v
Redox (Liquid Electrode) Cells
These consist of a semipermeable membrane having different liquids on either side. The membrane permits ion
flow but prevents mixing of the liquids. Electrical contact is made through inert conductors in the liquids. As the ions flow
across the membrane an electric current is induced in the conductors. These cells and batteries have two ways of recharging.
The first is the traditional way of running current backwards. The other is replacing the liquids, which can be recharged in
another cell. A small cell can also be used to charge a great quantity of liquid, which is stored outside the cells. This is an
interesting way to store energy for alternative energy sources that are unreliable, such as solar, wind, and tide. These
batteries have low volumetric efficiency, but are reliable and very long lived.
Electrochemical systems that can be used are FeCl3 (cathode) and TiCl3 or
CrCl2 (anode). Vanadium redox cells: A particularly interesting cell uses vanadium oxides of different oxidation
states as the anode and cathode. These solutions will not be spoiled if the membrane leaks, since the mixture can be charged as
either reducing or oxidizing components.
Unlike batteries, which store energy chemically, capacitors store energy as an electrostatic field. Typically,
a battery is known for storing a lot of energy and little power; a capacitor can provide large amounts of power, but low
amounts of energy. A capacitor is made of two conducting plates and an insulator called the dielectric, which conducts
ionically, but not electrically. In a capacitor,
Ecap = qV = ½CV2
where the capacitance, C, is directly proportional to the surface area of the plates and inversely
proportional to the distance between them.
So in other words, as the plate surface area increases and the distance between the plates decreases, the
energy you can store in a capacitor increases. Normal every-day capacitors have capacity on the orders of millifarads per cubic
foot. Aluminum electrolytics are about a farad per cubic foot. But for useful energy storage we need farads per cubic inch.
That is where supercapacitors come in.
First let's see how clever we can get to obtain a big surface area in a small volume. Imagine a polymer foam
cleaning sponge. It has a tremendous amount of surface area in a small area because of all the crenulations (OK, nooks and
crannies). Now, put it in a furnace, excluding the oxygen and bake it until only the carbon is left. You now have a conductive
carbon surface with an incredible surface area in a small volume.
But to get a high capacitance there has to be two plates. You can't just go in there and create complimentary
surface as the other electrodeor can you? Yes, just fill it with a conductive liquid (e.g., an aqueous acid or salt
solution). The last thing you need is an ultra-thin insulator on the carbon. Ultra thin to get high capacitance, and insulator
so the carbon and the liquid don't short out. This is also easy, you can electrochemically deposit an insulator on the carbon
surface (or electrochemically deposit something that could be turned into an insulator upon baking).
Now attach one electrode to the carbon, one to the liquid, and you can have a capacitor that can have Farads
of capacitance per cubic inch. Very nice.
Most practical supercapacitors have low voltage (2 to 5 Vremember that insulator is ultra-thin and so
can break down at low voltages), which is a problem for energy storage, since the stored energy is proportional to the square
of the voltage. Also, conduction through an ionic liquid is slow, so these capacitors cannot be discharged quickly compared
with standard capacitors, but can be discharged very quickly compared to batteries!
Typical numbers for capacitors and batteries are given below:
device |
volumetric energy density Wh/L |
power density W/L |
number of charge/discharge cycles |
discharge time s |
batteries |
50-250 |
150 |
1 - 103 |
> 1000 |
capacitors |
0.05 - 5 |
105 - 108 |
105 - 106 |
<1 |
Supercapacitors have several advantages over batteries: they can experience virtually indefinite number of
cycles (charging and discharging), they are maintenance free, they work well in high-rate discharge, they recharge quickly, and
they have no negative environmental impact.
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- Berndt, D. Maintenance-Free Batteries. New York:: John Wiley & Sons, 1997.
- Crompton, T. R. Battery Reference Book. London: ButterworthHeinemann, 1990.
- Linden, D. (Ed), Handbook of Batteries. Maidenhead: McGrawHill, 1995.
- Linford, R. G. (Ed), Electrochemical Science and Technology of Polymers. New York: Elsevier,
1990.
- Ovshinsky, S. R., Fetcenko, M. A., and Ross, J. A. "A Nickel Metal Hydride Battery for Electric Vehicles",
Science 260: 1993, 17681.
- Rechargeable Batteries Applications Handbook. Stoneham: ButterworthHeinemann, 1992.
- Wells, A. F. Structural Inorganic Chemistry. Oxford: Clarendon Press, 1975.
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Last Update: 17 August 2003 |
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