Drossman/Veirs EV112

Block 5, 2002

ENERGY

Batteries

Reading:
Silberberg, pp. 923-928; (EText, pp. 689-694)
ACS, Chapter 8, pp. 315-322; (EText, pp. 813-820)

All of the electrical generating technologies discussed up to now have been based on the electromechanical generator principles discovered by Michael Faraday. However, more than 30 years before Faraday produced electricity by moving a wire in a magnetic field, Alessandro Volta in 1800 discovered a different way of generating electricity. Volta found that two dissimilar metals separated by an electrolyte produced a small voltage that could drive a current through an external circuit connecting the two metal electrodes. Thus was invented the electrochemical cell, predecessor to the modern battery. Different combinations of electrode materials produce different voltages, typically in the range of 1­2 v. Higher voltages are achieved by stacking several electrochemical cells in series to form a battery. The voltage of an electrochemical system has a constant polarity, so, current always flows in a single direction. This type of current flow is called a direct current (DC), in contrast to the alternating current produced by electromechanical generators.

The electrochemical cells in common use today (more commonly called batteries) deliver relatively low levels of energy and power suitable for small electronic appliances. Larger batteries, such as those used in automobiles, provide greater amounts of energy (with currents in excess of 100 amps) for short periods. Eventually, however, all batteries become depleted and must be replaced, including rechargeable batteries.

A battery, strictly speaking, is a self-contained group of voltaic cells arranged in series (plus-to-minus-to-plus, and so on), so that their individual voltages are added together. In everyday speech, however, the term may also be applied to a single voltaic cell.

Most batteries convert chemical energy into electrical energy with an efficiency of about 90%. This may be compared with the much lower efficiencies that typically characterize the conversion of heat to work (30­40%). However, it is important to remember that considerable energy is required to manufacture electrochemical cells. Metals and minerals must be mined and processed, and the various components manufactured and assembled. Moreover, a battery has a finite life. Sooner or later, the chemical reaction will reach completion, the voltage will drop to zero, and electrons will no longer flow. The battery will be "dead" and ready for disposal, and disposal is a formidable problem. In February 1993, National Geographic reported that some two and a half billion household batteries are purchased each year in the United States at a total price of $3.3 billion. Of these, over 90% are single-use batteries that find their way into landfills or incinerators.

Batteries are ingeniously engineered devices that house rather unusual half-reactions and half-cells, but they operate through the same electrochemical principles we've been discussing. There are several classes of batteries. A primary battery cannot be recharged, so it is thrown away when the battery is "dead," that is, when the components have reached their equilibrium concentrations. In contrast, when a secondary, or rechargeable, battery runs down, it is recharged by supplying electrical energy to reverse the cell reaction and re-form reactant. In other words, in this type of battery, the voltaic cells are periodically converted to electrolytic cells to restore nonequilibrium concentrations. A fuel cell, or flow battery, is one that is not self-contained.

Types of Batteries (1, 3)

 

Dry Cell (1)
Invented in the 1860s, the common dry cell, or Leclanché cell, has become a familiar household item. A zinc anode in the form of a can
houses a mixture of MnO2 and an electrolyte paste, consisting of NH4Cl, ZnCl2, H2O, and starch. Powdered graphite improves conductivity.
The cathode is an inactive graphite rod. Its uses include portable radios, toys and flashlights. The advantagesare that it is inexpensive, safe, available in many sizes. The disadvantages are that at high current drain, NH3(g) builds up, causing voltage drop. Short shelf-life because zinc
anode reacts with the acidic NH4 ions.

Zn(s) --> Zn2+(aq) + 2e-

2MnO2(s) + 2NH4+ (aq) + 2e --> Mn2O3(s) + 2NH3(aq) + H2O(l)

 

Alkaline Battery (1)
The alkaline battery is an improved dry cell. The half-reactions are essentially the same, but the electrolyte is a KOH paste. The basic electrolyte
eliminates the build-up of gases and maintains the Zn electrode. Its uses are the same as the dry cell. However, there is no voltage drop
and longer shelf life than dry cell. Safe, many sizes. Its major disadvantage is that it is more expensive than common dry cell. These are teh common batteries that you use for AAA, AA , C amd D sizes. All have a potential of ~ 1.5 V.

Zn(s) + 2OH(aq) --> ZnO(s) + H2O(l) + 2e

MnO2(s) + 2H2O(l) + 2e --> Mn(OH)2(s) + 2OH-(aq)

Mercury and Silver (Button) Batteries (1)
The mercury battery and the silver battery are quite similar. Both use a zinc anode (reducing agent) in a basic medium. One employs HgO as
the oxidizing agent, the other Ag2O, and both use a steel cathode. The solid reactants are compacted separately with KOH, with moist paper for
a salt bridge. These baterries are most commonly used in watches and calculators, the silver cell in cameras, heart pacemakers, and hearing aids They both are very small, with relatively large voltage. The silver cell has a very steady output and is nontoxic. Discarded mercury cells release the toxic metal. Silver cells are expensive.

Zn(s) + 2OH(aq) --> ZnO(s) + H2O(l) + 2e

HgO(s) + H2O(l) + 2e --> Hg(l) + 2OH(aq)

Ag2O(s) + H2O(l) + 2e --> 2Ag(s) + 2OH(aq)

 

 

Lead-Acid Battery (1)
A typical 12-V lead-acid car battery has six cells connected in series, each of which delivers about 2 V. Each cell contains two lead grids packed with the electrode materials: the anode is spongy, powdered Pb, and the cathode is powdered PbO2. The grids are immersed in an electrolyte solution of 4.5 M H2SO4. Fiberglass sheets between the grids prevent shorting by chance physical contact. When the cell discharges, it generates electrical energy as a voltaic cell.

Uses: In automobiles and trucks.

Advantages: Provides a large burst of current to the engine starter motor; reliable, long life; effective at low temperature.

Disadvantages: 1. Loss of capacity: PbSO4, which is required in the recharging stage, coats the battery grids after the battery discharges. Mechanical strain and normal bumping can dislodge PbSO4 and reduce battery capacity. If enough PbSO4 is lost, the cell cannot be recharged. 2. Safety hazard: Older batteries have a cap on each cell to monitor electrolyte density and replace the water lost during discharging. During recharging, some water can electrolyze to H2 and O2 and, if sparked, the gases can explode and splatter H2SO4. Modern batteries use a lead alloy that inhibits electrolysis and reduces water loss, so the cells are sealed.

Pb(s) + SO42-(aq) --> PbSO4(s) + 2e-

PbO2(s) + 4H(aq) + SO42-(aq) 2e- --> PbSO4(s) + 2H2O(l)

 

 

 

Nickel-Cadmium (Nicad) Battery (1)
The nickel-cadmium battery has an anode half-reaction that oxidizes cadmium in a basic (NaOH or KOH) electrolyte, while nickel(III) as NiO(OH) is reduced at the cathode.

Uses: Cordless razors, photo ·ash units, and power tools.
Advantages: Lightweight.
Disadvantages: Disposal of toxic cadmium.

Cd(s) + 2OH-(aq) --> Cd(OH)2(s) + 2e-

2NiO(OH)(s) + 2H2O(l) + 2e- --> 2Ni(OH)2(s) + 2OH-(aq)

 

Lithium Solid-State Battery (1)
The lithium battery is a recent design that incorporates an Li anode and a transition metal oxide (or sulfide) cathode (e.g., MnO2, V6O13, or TiS2). The electrolyte is a polymer that allows Li ions to flow but not electrons.

Li(s) --> Li(in solid electrolyte) + e

MnO2(s) + Li + e --> LiMnO2(s)

 

Aluminum-Air Battery (1)
Air batteries, of which this is one type, are similar to ·ow batteries in that the components must be supplied periodically, if not continuously. The aluminum anode is oxidized and oxygen from flowing moist air is reduced at an inactive porous graphite cathode. An aqueous NaOH electrolyte circulates through the cell. With such basic conditions, the half-reactions are

4[Al(s) + 4OH-(aq) --> Al(OH)4-(aq) + 3e-]

3[O2(g) + 2H2O(l) + 4e --> 4OH-(aq)]

 

Sodium-Sulfur Battery (1)
In contrast to the traditional battery, which houses solid electrodes and liquid (or slurried) electrolyte, this design incorporates liquid electrodes separated by a solid electrolyte. Molten Na (mp 98°C) is the anode, and molten S8 (mp 113°C) is the cathode (mixed with powdered graphite to improve conductivity). The Na loses electrons, which pass through the external circuit (wire and stainless-steel electrode housing) and reduce the S8 to polysulfide ions. The electrolyte is b-alumina, a mixture of metal (Na, Mg, Al) oxides that allows Na ions to pass through and reach the nS2 ions.

Advantages: Na-S batteries can provide four to five times the energy per unit mass and undergo about three times as many dischargerecharge
cycles as lead-acid batteries. Battery life may be as long as 125,000 miles.

Disadvantages: Moderate speed and short discharge time: a BMW powered by a 265-kg Na-S battery had a top speed of 60 mi/h and a
cruising range of 100 mi; earlier versions had long recharge times (16 h). A temperature of 350°C must be maintained to keep reactants and products molten and a high Na conductivity in the b-alumina.

2Na(l) --> 2Na(l) + 2e-

S8(l) + 2e- --> nS2(l)

 

 

 

Battery Powered Cars-A Pollution Cure or Problem? (3)

Electrical power generation is notoriously inefficient. Less than half of the energy released from burning fuel is converted into electricity. Furthermore, power plants that draw their energy from fossil fuels release sulfur dioxide, nitrogen oxides, and carbon dioxide. In fact, calculations indicate that the SO2 and NOx emitted from power plants generating the electricity to keep a fleet of battery-powered cars operational will exceed the amount of these two gases that would have been released by the gasoline-powered cars that have been replaced. The overall CO2 emission does decline when electric cars are substituted for internal combustion automobiles, but by less than 50%.

Another serious criticism of "pollution free" electric cars was published in Science in May 1995 by Lester Lave, Chris Hendrickson, and Francis McMichael. These three Carnegie Mellon University professors argue that a switch to cars powered exclusively by lead storage batteries would dramatically increase the amount of lead released into the environment. Their calculations include estimates of lead dispersed in mining, processing, and battery manufacture; and they conclude that with current technology, 1.34 grams of the toxic metal would be emitted per kilometer traveled by an electric car. This corresponds to 2.16 grams of lead per mile or 47.5 pounds in 10,000 miles, a typical annual mileage driven in the United States. Ironically, this amount of lead is 60 times that released over the same distance by a car burning leaded gasoline. To be sure, critics have questioned some of the assumptions made by Lave, Hendrickson, and McMichael and argue that their conclusions greatly exaggerate the problem. Moreover, power plants, lead mines, lead refineries, and battery factories are point, not mobile, sources of pollution.

A compromise, perhaps an ideal one, that would get past the limitations of depending solely on batteries would be a car with the convenience and range of a gasoline-powered car combined with the environmental advantages of an electric vehicle. Such a vehicle, called a hybrid car, has been developed by Toyota, and is in development by other auto manufacturers as well. The Toyota Prius has been available in Japan since 1998. With a 1.5-liter gasoline engine sitting side-by-side with nickel-metal hydride batteries, an electric motor, and an electric generator, the Prius does not need to be recharged. It consumes only about half the gasoline, emits 50% less carbon dioxide and far less nitrogen oxides than a conventional car, while delivering 66 miles per gallon of gasoline (650 miles per tankful). The electric motor draws power from the batteries to get the car moving or when it is traveling at low speeds. Using a process called regenerative braking, the kinetic energy of the car is transferred to the generator, which charges the batteries during deceleration and braking. The gasoline engine assists the electric motor during normal driving, with the batteries boosting power when extra acceleration is needed.

References

1) Figures and text excerpted from Silberberg, Chapter 17, Electrochemistry

2) Figures and text excerpted from Rubin, Chapter

3) Figures and text excerpted from ACS, Chemistry in Context, Chapter 8


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