PowerStream Battery Chemistry FAQ

Batteries come in a lot of different varieties, and many years of work at universities, government labs, industrial workshops and inventor's basements have been put into investigating thousands of different chemistries. The most common are carbon-zinc, alkaline, lead acid, nickel metal hydride, nickel cadmium, lithium iron phosphate and lithium ion. But there are many other commercially available battery chemistries, each with their own advantages and disadvantages. Below 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
  • How to Charge Batteries
  • Standard Cylindrical Battery Size Chart
  • Standard Sealed Lead Acid Battery Size Chart
  • Battery Chemistry Performance Chart
  • How to Store Batteries
  • Battery Pack Engineering
  • Bulk Batteries for Sale

Definitions

1. Anode: The electrode where oxidation (loss of electrons) takes place. The anode of a battery is the negative electrode when discharging, and when charging it is connected to the negative electrode of the charger. So when the anode is charging it technically receives electrons, making it the cathode, but no-one ever makes this distinction due to the confusion it can cause.
2. 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 would have an amp of marbles.
3. 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 a constant rate of current. But note that usually if you discharge faster than the rate at which the the amp hours were specified you will get fewer amp-hours out, see the definition of Peukart Effect below.
   1. 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 needingrecharge." 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, but technically it is a combination of two or more.
• 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. A discharge rate of C/10 would be 0.5 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. The same confusion applies when charging as the anode, but the cathode remains the positive side of the battery to avoid confusion.
• 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 as ampere-hour, voltage, or 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. This is often a water solution, but in lithium batteries it consists of organic solvents with dissolved salts.
• 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 versus 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 (watts/liter). Specific power generally refers to the ratio of power to mass (watts/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, rectangular or any other such flattened shape.
• Peukart Effect: When a battery is discharged extremely quickly it will have less capacity than expected. This is the Peukart effect, which is very strong for lead acid batteries and many primary cells, but much less so for nickel cadmium, NiMH and lithium batteries. Peukert's equation is Iⁿ · t = C, where I is the discharge rate, t is the discharge time and C is the capacity. The exponent "n" depends on the battery chemistry and the temperature. A log-log plot of discharge time versus discharge load will have a slope of "n"
• Reserve cell: A primary cell that may be kept inactive and which is activated by adding an electrolyte or electrode, or melting an electrolyte which is normally 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.

Parameters for Battery Performance

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:

1. Voltage: Normal voltage during discharge, maximum and minimum permissible voltages, discharge curve profile. Note that secondary cells are characterized by their average voltage during a full discharge, whereas primary cells are characterized by their peak voltage. Thus an alkaline cells is rated at 1.5V, though its average voltage during discharge is 1.2V. A NiMH cell is rated at 1.2 volts, but its discharge range is 1.4 to 1 volt.
2. Duty cycle: Conditions the battery experiences during use. Type of discharge and current drain, e.g.., continuous, intermittent, continuous with pulses, etc. A 50% duty cycle means that the cell is off for half the cycle, and on for the other half of the cycle.
3. Temperature: In storage and in use. Temperatures that are too high or too low can greatly reduce battery capacity.
4. Shelf life: How rapidly the cell loses potential while unused.
5. 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%.
6. Physical restrictions: These include dimensions, weight, terminals, etc.
7. Maintenance and resupply: Ease of battery acquisition, replacement, charging facilities, disposal.
8. Safety and reliability: Failure rates, freedom from outgassing or leakage; use of toxic components; operation under hazardous conditions; environmentally safe
9. Cost: Initial cost, operating cost, use of expensive materials
10. Internal resistance: Batteries capable of a high-rate discharge must have a low internal resistance.
11. 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.
12. 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 .

1. 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 line—zinc chloride; broken line—Leclanché (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.

Alkaline Cells

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.

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)

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 electrode—a 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 220–300 Wh/kg (compared to 99–123 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.

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:

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

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 Lithium–Sulfur 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 240–280 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 3.3–3.5 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

These cells are used in coin cells for watches and memory-back up, nuclear missile batteries, the space shuttle safety system, and other governmental and space applications. 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. This is not a stoichiometric reaction, so the proportion of fluorine atoms can vary between 0.8 and 1.2, which is why the half-reactions are also not stoichiometric. This is interesting because it allows a lithium-fluorine reaction, which is probably the most energetic possible by safely storing the fluorine atoms in a graphite matrix. This is similar to how lithium ions are stored in lithium ion batteries. The carbon intercalation makes it safe, but it also reduces the voltage and lowers the electrical current capability.

A typical electrolyte is lithium tetrafluoroborate (LiBF₄) salt in a solution of propylene carbonate (PC) and dimethoxyethane (DME).

Anode	Li —> Li+ + e
Cathode	CFx + xe —> xC + xF
Overall	CFx + xLi —> xLiF+ xC

Note that one of the reaction products is carbon, which lowers the resistance of the cell as the battery is discharged. 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

LiI₂—Lithium 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· (n–1)I2 + 2LiI
Overall	2Li + P2VP· nI2 —> P2VP· (n–1)I2 +2LiI

LiI is formed in situ by direct reaction of the electrodes.

Secondary batteries

Lead–acid Cells

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 re-combiners 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.

Nickel/Cadmium Cells

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.

v

Nickel/Hydrogen cells

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.

Nickel/Metal Hydride (NiMH) Cells

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.

v

Sodium/Sulfur Cells

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 contact—with 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

Lithium Ion Cells

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, NiNi_0.3Co_0.7O₂, LiNiO₂, LiV₂O₅, LiV₆O₁₃, LiMn₄O₉, LiMn₂O₄, LiNiO_0.2CoO₂.

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 relation—not 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 battery—the graphite is conductive, dilutes the lithium for safety, is reasonably cheap, and does not allow dendrites or other unwanted crystal structures to form.

v

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.

v

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 acronym—RAM (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.

v

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.

v

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.

v

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

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).

Supercapacitors

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,

E_cap = 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 electrode—or 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 V—remember 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.

v

References and Other Helpful Sources

1. Berndt, D. Maintenance-Free Batteries. New York:: John Wiley & Sons, 1997.
2. Crompton, T. R. Battery Reference Book. London: Butterworth–Heinemann, 1990.
3. Linden, D. (Ed), Handbook of Batteries. Maidenhead: McGraw–Hill, 1995.
4. Linford, R. G. (Ed), Electrochemical Science and Technology of Polymers. New York: Elsevier, 1990.
5. Ovshinsky, S. R., Fetcenko, M. A., and Ross, J. A. "A Nickel Metal Hydride Battery for Electric Vehicles", Science 260: 1993, 176–81.
6. Rechargeable Batteries Applications Handbook. Stoneham: Butterworth–Heinemann, 1992.
7. Wells, A. F. Structural Inorganic Chemistry. Oxford: Clarendon Press, 1975.

v