Galvanic cells Batteries Cell potential

Galvanic cells in which stored chemicals can be reacted on demand to produce an electric current are termed primaiy cells. The discharging reac tion is irreversible and the contents, once exhausted, must be replaced or the cell discarded. Examples are the dry cells that activate small appliances. In some galvanic cells (called secondaiy cells), however, the reaction is reversible that is, application of an elec trical potential across the electrodes in the opposite direc tion will restore the reactants to their high-enthalpy state. Examples are rechargeable batteries for household appliances, automobiles, and many industrial applications. Electrolytic cells are the reactors upon which the electrochemical process, elec troplating, and electrowinning industries are based. 

Lithium metal had few uses until after World War II, when thermonuclear weapons were developed (see Section 17.11). This application has had an effect on the molar mass of lithium. Because only lithium-6 could be used in these weapons, the proportion of lithium-7 and, as a result, the molar mass of commercially available lithium has increased. A growing application of lithium is in the rechargeable lithium-ion battery. Because lithium has the most negative standard potential of all the elements, it can produce a high potential when used in a galvanic cell. Furthermore, because lithium has such a low density, lithium-ion batteries are light. 

FIGURE 2.3 Potential distribution in galvanic cells functioning as a battery (a) and as an electrolyzer (b) the dashed lines are for the zero-current situation. 

In this section, you learned how to identify the different components of a galvanic cell. Also, you found out how galvanic cells convert chemical energy into electrical energy. You were introduced to several common primary batteries that contain galvanic cells. In the next section, you will learn more about the cell potentials of galvanic cells. 

Many types of rechargeable batteries are much more portable than a car battery. For example, there is now a rechargeable version of the alkaline battery. Another example, shown in Figure 11.20, is the rechargeable nickel-cadmium (nicad) battery. Figure 11.21 shows a nickel-cadmium cell, which has a potential of about 1.4 V. A typical nicad battery contains three cells in series to produce a suitable voltage for electronic devices. When the cells in a nicad battery operate as galvanic cells, the half-reactions and the overall cell reaction are as follows. 

In this section, you learned about electrolytic cells, which convert electrical energy into chemical energy. You compared the spontaneous reactions in galvanic cells, which have positive cell potentials, with the non-spontaneous reactions in electrolytic cells, which have negative cell potentials. You then considered cells that act as both galvanic cells and electrolytic cells in some common rechargeable batteries. These batteries are an important application of electrochemistry. In the next two sections, you will learn about many more electrochemical applications. 

Tower, Stephen. All About Electrochemistry. Available online. URL http //www.cheml.com/acad/webtext/elchem/. Accessed May 28, 2009. Part of a virtual chemistry textbook, this excellent resource explains the basics of electrochemistry, which is important in understanding how fuel cells work. Discussions include galvanic cells and electrodes, cell potentials and thermodynamics, the Nernst equation and its applications, batteries and fuel cells, electrochemical corrosion, and electrolytic cells and electrolysis. 

The commercial galvanic cells we know as batteries are currently the subject of intense research by scientists who see their potential to solve environmental, health, communication, and transportation problems. 

A battery (or galvanic or voltaic cell) is a device that uses oxidation and reduction reactions to produce an electric current. In an electrolytic cell, an external source of electric current is used to drive a chemical reaction. This process is called electrolysis. When the electric potential applied to an electrochemical cell is just sufficient to balance the potential produced by reactions in the cell, we have an electrochemical cell at equilibrium. This state also occurs if there is no connections between the terminals of the cell (open-circuit condition). Our discussion in this chapter will be limited to electrochemical cells at equilibrium. 

Galvanic cell reactions, snch as those involving the everyday use of batteries, follow the same eqnations as electrolysis cells do. When we mea-snre cell potentials, however, we do not allow the reactions to proceed, becanse if they did, their potentials wonld change as the concentrations changed. The potentials are measured without a complete circuit. 

A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series, where the potentials of the individual cells add to give the total battery potential. Batteries are a source of direct current and have become an essential source of portable power in our society. In this section we examine the most common types of batteries. Some new batteries currently being developed are described at the end of the chapter. 

Referring to a list of standard electrode potentials, such as in Table 8.3, one speaks of an electrochemical series, and the metals lower down in the se-ries(with positive electrode potentials) are called noble metals. Any combination of half-reactions in an electrochemical cell, which gives a nonzero E value, can be used as a galvanic cell (i.e., a battery). If the reaction is driven by an applied external potential, we speak of an electrolytic cell. Reduction takes place at the cathode and oxidation at the anode. The reduction reactions in Table 8.3 are ordered with increasing potential or pe values. The oxidant in reactions with latter pe (or E°) can oxidize a reductant at a lower pe (or ) and vice versa for example, combining half-reactions we obtain an overall redox reaction ... 

Even in a galvanic cell with a salt bridge, there is some leakage of ions across the liquid junction, which causes the battery to lose its chemical potential over time. Commercial cells use an insoluble salt to prevent this from happening. 

Nernst equation An equation that correlates chemical energy and the electric potential of a galvanic cell or battery. Links the actual reversible potential of an electrode (measured in volts), E, at nonstandard conditions of concentration or pressure, to the standard reversible potential of the electrode couple, EO, which is a thermodynamic value. The Nernst equation is named after the German physical chemist Walther Nernst. 

Oxidation reduction reactions occur at two electrodes. The electrode at which oxidation occurs is called the anode the one at which reduction takes place is called the cathode. Electricity passes through a circuit under the influence of a potential or voltage, the driving force of the movement of charge. There are two different types of interaction of electricity and matter. Electrolysis is when an electric current causes a chemical reaction. Galvanic cell action is when a chemical reaction causes an electric current, as in the use of a battery. 

Electrons participating in the intercalation/deintercalation reaction (Equation (5.1)) can be represented by a current-producing system. Second, it is characteristic that the current-producing system reversibly operated by a self-driven (galvanic) cell (discharging the battery) performs the electrical useful work AG = —zFE (where E is the EMF of the cell), because electrical potential difference is spontaneously developed between two electrodes. By contrast, when the cell is short-circuited - that is, when the two electrodes are not separated from each other but are directly in electrical contact - electrons do not appear explicitly but rather participate in corrosion (or permeation in the case of solid electrolyte cells). They perform no electrical useful work because the two electrodes have the same electrical potential. 

The dry cell battery is a typical example of galvanic corrosion, or two metal corrosion as it is otherwise called. When two dissimilar metals are immersed in a conductive or corrosive medium, there is always the potential for a change in them. Once these metals are connected this difference induces electron flow between them. The less corrosion resistant metal is attacked more than the more resistant metal. This is an electrochemical process. In the case of a dry cell battery, the carbon electrode acts as the cathode (the more resistant materials) and zinc as the corroding anode. The natural phenomenon of corrosion is used in this case for producing electricity. 

One of the well-known applications of the electrochemistry is the use of galvanic cells in batteries. A battery is in principle just a group of galvanic cells in series, in which the potential of each cell is summed up to give a higher voltage across the batteiy. Batteries are used for a variety of purposes in our daily life. There are several different principles of how a battery may be build. In the following examples we will look at three types of batteries. 

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