ELECTROCHEMISTRY
SUBMITTED TO: Mr. DEEPAK AGGARWAL H.O.D CHEMISTRY
SUMIT SINGH CLASS: XII “SCIENCE”
ACKNOWLEDGEMENT First and foremost I thank my teacher Mr. DEEPAK AGGARWAL who has assigned me this project to bring out my creative capabilities. I express my gratitude to my parents for being a continuous source of encouragement for all their financial aid. My heartfelt gratitude to my class-mates and for helping me to complete my work in time.
Electrochemical Cells Many Many oxida oxidati tion on-r -red educt uctio ion n reac reacti tions ons occu occurr sp spont ontan aneo eousl usly, y, gi givi ving ng off off energy. An example involves the spontaneous reaction that occurs when zinc metal is placed in a solution of copper ions as described by the net ionic equation shown below. Cu+2 (aq) + Zn (s) -------> Cu(s) + Zn+2 (aq) The zinc metal slowly "dissolves" as its oxidation produces produc es zinc ions which enter into solution. At the same time, the copper ions gain electrons and are converted into copper atoms which coats the zinc metal or sediments to the bottom of the container. The energy produced in this reaction is quickly dissipated as heat, but it can be made to do useful work by a device called, an electrochemical cell. This is done in the following way. An electrochemical cell is composed to two compartments or half-cells, each composed of an electrode dipped in a solution of electrolyte. These hal half-ce f-cellls are desi design gned ed to cont contai ain n the the oxida xidattion half half-r -rea eact ctio ion n and and reduction half-reaction separately as shown below.
The The halfhalf-ce cell ll,, call called ed the anode, is the site which the oxidat oxidation ion of zinc zinc as shown below.
at occurs occurs
Zn (s) ----------> ----------> (aq) + 2e-
Zn+2
During During the oxidati oxidation on zinc, th e zinc electr electrode ode will will slowly slowly disso di ssolv lve e to produ produce ce ions (Zn+2), whic which h into the solution containing Zn+2 (aq) and SO4-2 (aq) ions.
of
zinc zinc ente enterr
The half-cell, called the cathode, is the site at which reduction of copper occurs as shown below. Cu+2 (aq) + 2e- -------> Cu (s) When the reduction of copper ions (Cu+2) occur ccurs s, coppe pper atoms oms accumulate on the surface of the solid copper electrode.
The reaction in each half-cell does not occur unless the two half cells are connected to each other. Recall that in order for oxidation to occur, there must be a corresponding reduction reaction that is linked or "coupled" with it. Moreover, in an isolated oxidation or reduction half-cell, an imbalance of electrical charge would occur, the anode would become more positive as zinc cations are prod produce uced, d, and and the the cath cathod ode e woul would d becom become e more more negat negativ ive e as coppe copperr cations are removed from solution. This problem can be solved by using a "salt bridge" bridge" connect connecting ing the two cells cells as shown in the diagram below. A "salt bridge" is a poro porous us barr barrie ierr whic which h preve prevent nts s the spontaneous mixing of the the aque aqueou ous s sol solutio utions ns in each each comp ompartment, nt, but but allows the mig igra rati tion on of ions ions in both both directions directions to maintain maintain electrical electrical neu neutrality. As the oxidationreduct ductiion reaction occu ccurs, cations ( Zn+2) from from the the anode anode mig igra rate te via via the sal salt bri bridge dge to -2 the cathode, while the anion, (SO4) , migrates in the opposite direction to maintain electrical neutrality. The two half-c half-cell ells s are also also connect connected ed extern externall ally. y. In this this arrange arrangemen ment, t, electrons provided by the oxidation reaction are forced to travel via an external circuit to the site of the reduction reaction. The fact that the reac reacti tion on occu occurs rs sp spon onta tane neou ousl sly y once once thes these e half half cell cells s are are conn connec ecte ted d indicates that there is a difference in potential energy. This difference in potential energy is called an electomotive force (emf) and is measured in terms of volts. The zinc/copper cell has an emf of about 1.1 volts under standard conditions.
Any electrical device can be "spliced" into the external circuit to utilize this potential energy produced by the cell for useful work. Although the energy energy availabl available e from from a single single cell cell is relati relativel vely y small small,, electr electroche ochemi mical cal cells can be linked in series to boost their energy output. A common and useful application of this characteristic is the"battery". An example is the lead-acid battery used in automobiles. In the lead-acid battery, each cell has a lead metal anode and lead (IV) oxide (lead dioxide) cathode both of which are immersed in a solution of sulfuric acid. This single electrochemical cell produces about 2 volts. Linking 6 of these cells in seri series es prod produc uces es the the 12 12-v -vol oltt batt batter ery y foun found d in most most cars cars toda today. y. One One
disadvantage of these "wet cells" such as the lead-acid battery is that it is very very heav heavy y and and bulk bulky y. Howe Howeve verr, like ike many any othe otherr "wet "wet cell cells" s",, the the oxidation-reduction reaction which occurs can be readily reversed via an external current such as that provided by an automobile's alternator. This prolongs the lifetime and usefulness of such devices as an energy source. Standard Hydrogen Electrode (SHE) The SHE is the univer universal sal referen reference ce for reporti reporting ng relative half-cell potentials. It is a type of gas electr electrode ode and was widely widely used used in early early studie studies s as a refe refere renc nce e elec electr trod ode, e, and and as an indi in dica cato torr electrode for the determination of pH of pH values values.. The SHE could could be used as either either an anode anode or cath cathod ode e depen dependi ding ng upon upon the the natu nature re of the the half half-c -cel elll it is us used ed with with.. The The SHE SHE cons consis ists ts of a platin platinum um electr electrode ode immerse immersed d in a soluti solution on with with a hydr hydroge ogen n ion ion conce concent ntrat ratio ion n of 1.00 1. 00M. M. The The pl plat atin inum um elec electr trod ode e is made made of a smal smalll square of platinum foil which is platinized (kno (known wn as plati platinu num m bl black ack). ). Hydr Hydrog ogen en gas, gas, at a press pressur ure e of 1 atmo atmosph spher ere, e, is bubbl bubbled ed arou around nd the platinum electrode. The platinum black serves as a large surface area for the reaction to take place, and the stream of hydrogen keeps the solut solutio ion n satur saturat ated ed at the the elec electr trod ode e site site with with respe respect ct to the the gas. gas. It is interesting to note that even though the SHE is the universal reference standard, it exists only as a theoretical electrode which scientists use as the definition of an arbitrary reference electrode with a half-cell potential of 0.00 volts. (Because half-cell potentials cannot be measured, this is the perfe perfect ct elect electro rode de to allo allow w scien scienti tist sts s to perfo perform rm theor theoret etic ical al resea researc rch h calculations.) The reason this electrode cannot be manufactured is due to the fact that no solution can be prepared that yields a hydrogen ion activity of 1.00M.
hydrogen electrode is made by adding platinum black to platinum wire or a platinum plate. It is immersed in the test solution and an electric charge is applied to the solution and platinum black with hydrogen gas. The hydrogen-electrode method is a standard among the various methods for measuring pH. The values derived using other methods become trust rustwo wort rthy hy onl nly y whe when they hey match atch those measured using hydrogen electrode method. Howeve ever, this
method is not appropriate for daily use because of the effort and expense involv involved, ed, with with the inconv inconvenie enience nce of handlin handling g hydrog hydrogen en gas and great great influence of highly oxidizing or reducing substances in the test solution.
Fuel cell A fuel cell is a device that generates electricity e lectricity by a chemical reaction. Every fuel cell has two electrodes, one positive and one negative, called, respectively, the anode and cathode. The reactions that produce electricity take place at the electrodes. Every fuel cell also has an electrolyte, which carries electrically charged particles from one electrode to the other, and a catalyst, which speeds the reactions at the electrodes. Hydrogen is the basic fuel, but fuel cells also require oxygen. One great appeal of fuel cells is that they generate electricity with very little pollution—much of the hydrogen and oxygen used in generating electricity ultimately combine to form a harmless byproduct, namely water. One detail of terminology: a single fuel cell generates a tiny amount of direct current (DC) electricity. In practice, many fuel cells are usually assembled into a stack. Cell or stack, the principles are the th e same.
How do fuel cells work? The purpose of a fuel cell is to produce an electrical current that can be directed outside the cell to do work, such as powering an electric motor or illuminating a light bulb or a city. Because of the way electricity behaves, this current returns to the fuel cell, completing an electrical circuit. (To learn more about electricity and electric power, visit “Throw The Switch” on the Smithsonian website Powering a Generation of Change.) The chemical reactions that produce this current are the key to how a fuel cell works. There are several kinds of fuel cells, and each operates a bit differently. But in general terms, hydrogen atoms enter a fuel cell at the anode where a chemical reaction strips them of their electrons. The hydrogen atoms are now “ionized,” and carry a positive electrical charge. The negatively charged electrons provide the current through wires to do work. If alternating current (AC) is needed, the DC output of the fuel cell must be routed through a conversion device called an inverter.
Oxygen enters the fuel cell at the cathode and, in some cell types (like the one illustrated above), it there combines with electrons returning from the electrical circuit and hydrogen ions that have traveled through the electrolyte from the anode. In other cell types the oxygen picks up electrons and then travels through the electrolyte to the anode, where it combines with hydrogen ions. The electrolyte plays a key role. It must permit only the appropriate ions to pass between the anode and cathode. If free electrons or other substances could travel through the electrolyte, they would disrupt the chemical reaction. Whether they combine at anode or cathode, together hydrogen and oxygen form water, which drains from the cell. As long as a fuel cell is supplied with hydrogen and oxygen, it will generate electricity. Even better, since fuel cells create electricity chemically, rather than by combustion, they are not subject to the thermodynamic laws that limit a conventional power plant (see “Carnot Limit” in the glossary). Therefore, fuel cells are more efficient in extracting energy from a fuel. Waste heat from some cells can also be harnessed, boosting b oosting system efficiency still further.
So why can’t I go out and buy a fuel cell? The basic workings of a fuel cell may not be difficult to illustrate. But building inexpensive, efficient, reliable fuel cells is a far more complicated business. Scientists and inventors have designed many different types and sizes of fuel cells in the search for greater efficiency, and the technical details of each kind vary. Many of the choices facing fuel cell developers are constrained by the choice of electrolyte. The design of electrodes, for example, and the materials used to make them depend on the electrolyte. Today, the main electrolyte types are alkali, molten carbonate, phosphoric acid, proton exchange membrane (PEM) and solid oxide. The first three are liquid electrolytes; the last two are solids. The type of fuel also depends on the electrolyte. Some cells need pure hydrogen, and therefore demand extra equipment such as a “reformer” to purify the fuel. Other cells can tolerate some impurities, but might need higher temperatures to run efficiently. Liquid electrolytes circulate in some cells, which requires pumps. The type of electrolyte also dictates a cell’s operating temperature–“molten” carbonate cells run hot, just as the name implies. Each type of fuel cell has advantages and drawbacks compared to the others, and none is yet cheap che ap and efficient enough to widely replace
traditional ways of generating power, such coal-fired, hydroelectric, or even nuclear power plants. The following list describes the five main types of fuel cells. More detailed information can be found in those specific areas of this site.
Different types of fuel cells. Alkali fuel cells operate on compressed hydrogen and oxygen. They generally use a solution of potassium hydroxide (chemically, KOH) in water as their electrolyte. Efficiency is about 70 percent, and operating temperature is 150 to 200 degrees C, (about 300 to 400 degrees F). Cell output ranges from 300 watts (W) to 5 kilowatts (kW). Alkali cells were used in Apollo spacecraft to provide both electricity and drinking water. They require pure hydrogen fuel, however, and their platinum electrode catalysts are expensive. And like any container filled with liquid, they can leak. Molten Carbonate fuel cells (MCFC) use high-temperature compounds of salt (like sodium or magnesium) carbonates (chemically, CO3) as the electrolyte. Efficiency ranges from 60 to 80 percent, and operating temperature is about 650 degrees C (1,200
degrees F). Units with output up to 2 megawatts (MW) have been constructed, and designs exist for units up to 100 MW. The high temperature limits damage from carbon monoxide "poisoning" of the cell and waste heat can be recycled to make additional electricity. Their nickel electrodecatalysts are inexpensive compared to the platinum used in other cells. But the high temperature also limits the materials and safe uses of MCFCs—they Drawing of a molten carbonate would probably be too hot for home use. Also, carbonate ions from the electrolyte cell are used up in the reactions, making it necessary to inject carbon dioxide to compensate. Phosphoric Acid fuel cells (PAFC) use phosphoric acid as the electrolyte. Efficiency ranges from 40 to 80 percent, and operating temperature is between 150 to 200 degrees C (about 300 to 400 degrees F). Existing phosphoric acid cells have outputs up to 200 kW, and 11 MW units have been tested. PAFCs tolerate a carbon monoxide concentration of about 1.5 percent, which broadens the choice of fuels they can use. If gasoline is used, the sulfur must be removed. Platinum electrode-catalysts are needed, and internal parts must be able to withstand the corrosive acid. .
Proton Exchange Membrane (PEM) fuel work with a polymer electrolyte in the form thin, permeable sheet. Efficiency is about 40 to percent, and operating temperature is about degrees C (about 175 degrees F). Cell outputs generally range from 250 kW. The solid, electrolyte will not leak crack, and these cells operate at a low temperature to make them suitable for homes and cars. But their
cells of a 50 80
50 to flexible or enough fuels
must be purified, and a platinum catalyst is used on both sides of the membrane, raising costs. Solid Oxide fuel cells (SOFC) use a hard, ceramic compound of metal (like calcium or zirconium) oxides (chemically, O 2) as electrolyte. Efficiency is about 60 percent, and operating temperatures are about 1,000 degrees C (about 1,800 degrees F). Cells output is up to 100 kW. At such high h igh temperatures a reformer is not required to extract hydrogen from the fuel, and waste heat can be recycled to make additional electricity. However, the high temperature limits applications of SOFC units and they tend to be rather large. While solid electrolytes cannot leak, they can crack. More detailed information about each fuel cell type, including histories and current applications, can be found on their specific parts of this site. We have also provided a glossary of technical terms–a link is provided at the top of each technology page.
Mercuric-Oxide Zinc Cell The mercuric-oxide zinc cell (mercury cell) is a primary cell that was developed during World War II. Two important assets of o f the mercury cell are its ability to produce current for a long period of time and a long shelf life when compared to the dry cell shown in figure 2-4.The mercury cell also has a very stable output voltage and is a power source that can be be made in a small physical size. With the birth of the space program and the development of small transceivers and miniaturized equipment, a power source of small size was needed. Such equipment requires a small cell which is capable of delivering maximum electrical energy at a constant discharge voltage. The mercury cell, which is one of the smallest cells, meets these requirements. Present mercury cells are manufactured in three basic types as shown in figure 2-5. The wound-anode type, shown in figure 2-5 view A, has an anode composed of a corrugated zinc strip with a paper absorbent. The zinc is mixed with mercury, and the paper is soaked in the electrolyte which causes it to swell and press pre ss against the zinc and make positive contact. This process ensures that the electrolyte makes contact with the anode.
Figure 2-5. - Mercury cells. In the pressed-powder cells, shown in figure 2-5 views B and C, the zinc powder for the anode is mixed prior to being pressed into shape. The absorbent shown in the figure is paper pa per soaked in the electrolyte. The space between the inner and outer containers provides passage for any gas generated by an improper chemical c hemical balance or impurities present within the cell. If the anode and cathode of a cell are connected together without a load, a SHORT CIRCUIT condition exists. Short circuits (shorts) can be very dangerous. They cause excessive heat, pressure, and current flow which may cause serious damage to the cell or be a safety hazard to personnel. WARNING Do not short the mercury cell. Shorted mercury cells have exploded exp loded with considerable force.
PRIMARY DRY CELL The dry cell is the most popular type of primary cell. It is ideal for simple applications where an inexpensive and noncritical n oncritical source of electricity is all that is needed. The dry cell is not actually dry. The electrolyte is not in a liquid state, but is a moist paste. If it should become totally dry, it would no longer be able to transform chemical energy to electrical energy.
The construction of a common type of dry cell is shown in figure 2-4. These dry cells are also referred to as Leclanche' cells. The internal parts of the cell are located in a cylindrical zinc container. This zinc container serves as the negative electrode (cathode) of the cell. The container is lined with a nonconducting material, such as blotting paper, to separate the zinc from the paste. A carbon electrode is located loca ted in the center, and it serves as the positive terminal (anode) of the cell. The paste is a mixture of several substances such as ammonium chloride, powdered coke, ground carbon, manganese dioxide, zinc chloride, graphite, and water.
Figure 2-4. - Cutaway view of the general-purpose dry cell. This paste, which is packed in the space between the anode and the blotting paper, also serves to hold the anode rigid in the center of the cell. When the paste is packed in the cell, a small space is left at the top for expansion of the electrolytic paste caused by the depolarization action. The cell is then sealed with a cardboard or plastic seal. Since the zinc container is the cathode, ca thode, it must be protected with some insulating material to be electrically isolated. Therefore, it is common practice for the manufacturer to enclose the cells in cardboard and metal containers. The dry cell (fig. 2-4) is basically the same as the simple voltaic cell (wet cell) described earlier, as far as its internal chemical action is concerned. The action of the water and the ammonium chloride in the paste, together with the zinc and carbon electrodes, produces the voltage of the cell. Manganese dioxide is added to reduce polarization when current flows and zinc chloride reduces local action ac tion when the cell is not being used. A cell that is not being used (sitting on the shelf) will gradually deteriorate because of slow internal chemical changes (local action). This deterioration is usually very slow if cells are properly stored. If unused cells are stored in a cool place, their shelf sh elf life will be greatly preserved. Therefore, to minimize deterioration, they should be stored in refrigerated spaces.
The blotting paper (paste-coated pulpboard separator) serves two purposes - (1) it keeps the paste from making actual contact with the zinc container and (2) it permits the electrolyte from the paste to filter through to the zinc slowly. The cell is sealed at the top to keep air from entering and drying the electrolyte. Care should be taken to prevent breaking this seal. NickelNic kel-Cad Cadmium mium Cell Cel lThe nickel-cadmium cell is a secondary cell, and the electrolyte is pot potassium hydroxide. Thenegative electrode is made of nickel hydroxide, and the positive electrode is made of cadmiumhydroxide. Th e nominal voltage of a nickelcadmium cadmium cell is 1.25 1.25 volts. volts. The nickel-cadmi nickel-cadmiumbatt umbattery ery has the advantage of being a dry cell that is a true storage batter battery y with with a revers reversibl ibleche echemic mical al reacti reaction on (i.e. (i.e.,, it can be recha recharge rged). d). The nickel-cadmi nickel-cadmium um battery battery is a rugged, dependablebatte dependablebattery. ry. It gives dependable service under extreme conditions of temperature, shock, andvibr andvibrati ation. on. Due to its dependa dependabil bility ity,, it is ideall ideally y suited suited for use in portable communicationsequipment
Salt Bridge The two compartments of a cell must be separated so they do not mix, but but cann cannot ot be com compl plet etel ely y sepa separa rate ted d with with no way way for ions ions to be transferred. Because of this, a salt bridge is an important part of a concentration cell. It solves the major problem of electrons beginning to pile up to much in the right beaker. beaker. This build up is due to electrons electrons moving from the left side, or left beaker, to the right side, or right beaker. beaker. The salt bridge itself can be in a few different different forms, such as a salt solution in a U-tube or a porous barrier (direct contact). It evens the charge by moving ions to the left side, or left beaker. In the written expression which shows what is occurring in specific reactions, the salt bridge is represented by the double d ouble lines. An example of this would be: The double lines between the Zn2+ (1M) and the Cu2+ (1M) is the salt bridge. The single lines do not represent bridges, they represent the different phase changes. If there is a comma where you would expect to see a single line, this is not incorrect. Instead it means that no phase changes occurred.
Standard Electrode Potentials In an electrochemical cell, cell, an electric potential is created between two dissimilar metals. This potential is a measure of the energy per unit charge which is available from the oxidation/reduction reactions to drive the reaction. It is customary to visualize the cell reaction in terms of two half-reactions, an oxidation half-reaction and a reduction half-reaction. Reduced species -> oxidized Oxidation at anode species + neOxidized species + ne- -> reduced Reduction at cathode species The cell potential (often called the electromotive force or emf ) has a contribution from the anode which is a measure of its ability to lose electrons - it will be called its "oxidation potential". The cathode has a contribution based on its ability to gain electeons, its "reduction potential". The cell potential can then be written Ecell = oxidation potential + reduction potential If we could tabulate the oxidation and reduction potentials of all available electrodes, then we could predict the cell potentials of voltaic of voltaic cells created from any pair of electrodes. Actually, tabulating one or the other is sufficient, since the oxidation potential of a half-reaction is the negative of the reduction potential for the reverse of that reaction. Two main hurdles must be overcome to establish such a tabulation 1. The electrode electrode potential potential cannot be determine determined d in isolation, isolation, but in a reaction with some other electrode. 2. The electrode electrode potential potential depends upon the concentra concentrations tions of the substances, the temperature, and the pressure in the case of a gas electrode. In practice, the first of these hurdles is overcome by measuring the potentials with respect to a standard hydrogen electrode. It is the nature of electric potential that the zero of potential is arbitrary; it is the difference in potential which has practical consequence. consequen ce. Tabulating all electrode potentials with respect to the same standard electrode provides a practical working framework for a wide range of calculations and predictions. The standard hydrogen electrode is assigned a potential of zero volts. The second hurdle is overcome by choosing standard thermodynamic conditions for the measurement of the potentials. The standard electrode potentials are customarily determined at solute concentrations of 1 Molar, gas pressures of 1 atmosphere, and a standard temperature which is
usually 25°C. The standard cell potential is denoted by a degree sign as a superscript.
° Cell
E
1. 2. 3. 4.
Measured Measured against against standard standard hydroden electrod electrode. e. Concent Concentrat ration ion 1 Molar Molar Pressu Pressure re 1 atmosph atmosphere ere Temper Temperatur ature e 25°C 25°C
Electrochemical corrosion of iron Corrosion often begins at a location ( 1) where the metal is under stress (at a bend or weld) or is isolated from the air (where two pieces of metal are joined or under a looselyadhering paint film.) The metal ions dissolve in the moisture film and the electrons migrate to another location (2) where they are taken up by a depolarizer . Oxygen is the most common depolarizer; the resulting hydroxide ions react with the Fe2+ to form the mixture of hydrous iron oxides known as rust .
CONTROL OF CORROSION Since both the cathodic and anodic steps must take place for corrosion to occur, prevention of either one will stop corrosion. The most obvious strategy is to stop both processes by b y coating the object with a paint or other protective coating. Even if this is done, there are likely to be places where the coating is broken or does not penetrate, particularly if there are holes or screw threads. A more sophisticated approach is to apply a slight negative charge to the metal, thus making it more difficult for the reaction M M2+ + 2 e– to take place. →
Sacrificial coatings One way of supplying this negative charge is to apply a coating of a more active metal. Thus a very common way of protecting steel from corrosion is to coat it with a thin layer of zinc; this process is known as galvanizing.The zinc coating, being less noble than iron, tends to corrode selectively. Dissolution of this sacrificial coating leaves behind electrons which concentrate in the iron, making it cathodic and thus inhibiting its dissolution.
The effect of plating iron with a less active metal provides an interesting contrast. The common tin-plated can (on the right) is a good example. As long as the tin coating remains intact, all is well, but exposure of even a tiny part of the underlying iron to the moist atmosphere initiates corrosion. The electrons released from the iron flow into the tin, making the iron more anodic so now the tin is actively ac tively promoting corrosion of the iron! You have probably observed how tin cans disintegrate very rapidly when left outdoors.
Cathodic protection A more sophisticated strategy is to maintain a continual negative electrical charge on a metal, so that its dissolution as positive ions is inhibited. Since the entire surface is forced into the cathodic condition, this method is known as cathodic protection. protection. The source of electrons can be an external direct current power supply (commonly used to protect oil pipelines and other buried structures), or it can be the corrosion of another, more active metal such as a piece of zinc or aluminum buried in the ground nearby, as is shown in the illustration of the buried propane storage tank below. b elow.
HYDROGEN ECONOMY It seems like every day there is a new announcement in the news about automobiles powered by fuel cells. cells. The promises are tantalizing, since fuel cells have the potential to very quickly double theefficiency the efficiency of cars while significantly reducing air pollution. At the same time, there have been news stories for decades about the problems associated with petroleum petroleum.. Everything from oil spills to ozone alerts to gl global obal warminggets warminggets blamed on our dependence on fossil fuels. These two forces are leading the world toward what is broadly known as the hydrogen economy. If the predictions are true, over the next several decades we will all begin to see an amazing shift away from the fossil fuel economy we have today toward a much cleaner hydrogen future. Can society actually make this shift, or will the technological, economic and political barriers keep us bound to petroleum p etroleum and other fossil fuels for the next century and beyond? In I n this article, you will learn about the benefits of a hydrogen economy, along with its potential problems. We will also examine some of the technology that would make the transition possible.
