Electrical Energy from Oxidation-Reduction Reactions

4 Electrical Energy from Oxidation-Reduction Reactions 

However, if we place a Cu metal strip in a Zn solution, nothing will happen. The reaction does not run spontaneously in the reverse direction because Cu does not lose electrons as easily as Zn. 

We can determine the direction of a spontaneous reaction from the activity series, which ranks the metals and H2 in terms of how easily they lose electrons. 

In the activity series, the metals that lose electrons most easily are placed at the top, and the metals that do not lose electrons easily are at the bottom. Thus the metals that are more easily oxidized are above the metals whose ions are more easily reduced (see Table 15.3). Active metals include K, Na, Ca, Mg, Al, Zn, Fe, and Sn. In single replacement reactions, the metal ion replaces the H in the acid. Metals listed below H2( ) will not react with H from acids. 

According to the activity series, a metal will oxidize spontaneously when it is combined with the reverse of the half-reaction for any metal below it on the list. We use the activity series to predict the direction of the spontaneous reaction. Suppose we have two beakers. In one, we place a Mg strip in a solution containing Ni ions. In the other, we place a Ni strip in a solution containing Mg ions. Looking at the activity series we see that the half-reaction for the oxidation of Mg is listed above that for Ni, which means that Mg is the more active metal and loses electrons more easily than Ni. Using the activity series table, we write these two half-reactions as follows  

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Oxidation-Reduction Reactions Oxidation Numbers Balancing Oxidation-Reduction Equations Using Half-Reactions Electrical Energy from Oxidation-Reduction Reactions Oxidation-Reduction Reactions That Require Electrical Energy Oxidation of Alcohols Extended Topic... 

ELECTRICAL ENERGY FROM OXIDATION-REDUCTION REACTIONS 535... 

Section 15.4, which is titled Electrical Energy from Oxidation-Reduction Reactions, now begins with the Activity Series followed by Voltaic Cells. 

Chapter 17 emphasizes the principles associated with obtaining electrical energy from electron-transfer reactions in solution. This chapter emphasizes what happens when electrical energy is applied to solutions in the operation of electrolytic cells. The oxidation and reduction processes that take place in an electrolytic cell are called electrolysis. We focus on determining what products are obtained and how much energy is required. 

The transfer of a single electron between two chemical entities is the simplest of oxidation-reduction processes, but it is of central importance in vast areas of chemistry. Electron transfer processes constitute the fundamental steps in biological utilization of oxygen, in electrical conductivity, in oxidation reduction reactions of organic and inorganic substrates, in many catalytic processes, in the transduction of the sun s energy by plants and by synthetic solar cells, and so on. The breadth and complexity of the subject is evident from the five volume handbook Electron Transfer in Chemistry (V. Balzani, Ed.), published in 2001. The most fimdamental principles that govern the efficiencies, the yields or the rates of electron-transfer processes are independent of the nature of the substrates. The properties of the substrates do dictate the conditions for apphcability of those fimdamental... 

The chemical process that produces an electrical current from chemical energy is called an oxidation-reduction reaction. The oxidation-reduction reaction in a battery involves the loss of electrons by one compound (oxidation) and the gain of electrons (reduction) by another compound. Electrons are released from one part of the batteiy and the external circuit allows the electrons to flow from that part to another part of the batteiy. In any battery, current flows from the anode to the cathode. The anode is the electrode where positive current enters the device, which means it releases electrons to the external circuit. The cathode, or positive terminal of the battery, is where positive current leaves the device, which means this is where external electrons are taken from the external circuit. 

Electron-transfer reactions occur all around us. Objects made of iron become coated with mst when they are exposed to moist air. Animals obtain energy from the reaction of carbohydrates with oxygen to form carbon dioxide and water. Turning on a flashlight generates a current of electricity from a chemical reaction in the batteries. In an aluminum refinery, huge quantities of electricity drive the conversion of aluminum oxide into aluminum metal. These different chemical processes share one common feature Each is an oxidation-reduction reaction, commonly called a redox reaction, in which electrons are transferred from one chemical species to another. 

Electrochemical cells produce electrical energy from a spontaneous chemical reaction. In electrolysis, the process is reversed so that electrical energy is used to carry out a nonspontaneous chemical change. A cell arranged to do this is called an electrolytic cell. An electrolytic cell is similar to an electrochemical cell except that an electrolytic cell s circuit includes a power source, for example, a battery. The same electrochemical cell terminology applies to electrolytic cells. Reduction occurs at the cathode and oxidation at the anode. 

Virtually all energy transductions in cells can be traced to this flow of electrons from one molecule to another, in a downhill flow from higher to lower electrochemical potential as such, this is formally analogous to the flow of electrons in a battery-driven electric circuit. All these reactions involving electron flow are oxidation-reduction reactions one reactant is oxidized (loses electrons) as another is reduced (gains electrons). 

So we see that with the proper setup it is possible to harness electrical energy from an oxidation-reduction reaction. The apparatus shown in Figure 11.8 is one example. Such devices are called voltaic cells. Instead of two containers, a voltaic cell can be an all-in-one, self-contained unit, in which case it is called a battery. Batteries are either disposable or rechargeable, and here we explore some examples of each. Although the two types differ in design and composition, they function by the same principle two materials that oxidize and reduce each other are connected by a medium through which ions travel to balance an external flow of electrons. 

Disposable batteries have relatively short lives because electron-producing chemicals are consumed. The main feature of rechargeable batteries is the reversibility of the oxidation and reduction reactions. In your car s rechargeable lead storage battery, for example, electrical energy is produced as lead dioxide, lead, and sulfuric acid are consumed to form lead sulfate and water. The elemental lead is oxidized to Pb2+, and the lead in the lead dioxide is reduced from the Pb4+ state to the Pb2+ state. Combining the two half-reactions gives the complete oxidation-reduction reaction ... 

We learn much about chemical reactions from the study of electrochemistry. The amount of electrical energy consumed or produced can be measured quite accurately. All electrochemical reactions involve the transfer of electrons and are therefore oxidation-reduction reactions. The sites of oxidation and reduction are separated physically so that oxidation occurs at one location, and reduction occurs at the other. Electrochemical processes require some method of introducing a stream of electrons into a reacting chemical system and some means of withdrawing electrons. In most applications the reacting system is contained in a cell, and an electric current enters or exits by electrodes. 

A fundamental understanding of oxidation-reduction reactions is vital to the inorganic chemist in contexts ranging from energy transduction - chemical to electrical and the converse, in technical matters in corrosion processes and metallurgy, redox processes in environmental chemistry and metalloenzymes and metallo-proteins involved in electron transfer. Electron-transfer reactions of transition metal complexes are accompanied by a change in the oxidation state of the metal... 

A battery uses the energy from an oxidation-reduction reaction to produce an electric current. This is an important illustration of electrochemistry, the study of the interchange of chemical and electrical energy. 

In Chapter VI we defined oxidation as the loss of electrons and reduction as the gain of electrons, and we showed that oxidation-reduction reactions that involve ions can generally be made to produce an electric current. Chemical energy is thereby transformed into electrical energy, and the electromotive force of the cell is a measure of the free energy of the reaction. Conversely, we can make an electric current produce a chemical reaction. If a current is passed through an electrolyte—a conducting solution or molten salt—oxidation takes place at the anode, where electrons are withdrawn from the solution and reduction takes place simultaneously at the cathode, where the electrons enter. 

To obtain useful electrical energy from an oxidation-reduction process, we must set up the reaction in such a way that the oxidation half-reaction and the... 

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