SECTION 4 Electrolytic Cells

The precipitated copper from this reaction is an important constituent of the slime that collects at the bottom of the electrolytic cells. The accumulation of copper as well as of impurities such as nickel, arsenic, antimony, and bismuth is controlled by periodic bleed-off and treatment in the electrolyte purification section. 

Stress corrosion can arise in plain carbon and low-alloy steels if critical conditions of temperature, concentration and potential in hot alkali solutions are present (see Section 2.3.3). The critical potential range for stress corrosion is shown in Fig. 2-18. This potential range corresponds to the active/passive transition. Theoretically, anodic protection as well as cathodic protection would be possible (see Section 2.4) however, in the active condition, noticeable negligible dissolution of the steel occurs due to the formation of FeO ions. Therefore, the anodic protection method was chosen for protecting a water electrolysis plant operating with caustic potash solution against stress corrosion [30]. The protection current was provided by the electrolytic cells of the plant. 

The principles discussed in this chapter have a host of practical applications. Whenever you start your car, turn on your cell phone, or use a remote control for your television or other devices, you are making use of a voltaic cell. Many of our most important elements, including hydrogen and chlorine, are made in electrolytic cells. These applications, among others, are discussed in Section 18.6. ... 

AgsSBr, /3-AgsSI, and a-AgsSI are cationic conductors due to the structural disorder of the cation sublattices. AgsSI (see Fig. 5) has been discussed for use in solid-electrolyte cells (209,371, 374,414-416) because of its high silver ionic conductivity at rather low temperatures (see Section II,D,1). The practical use seems to be limited, however, by an electronic part of the conductivity that is not negligible (370), and by the instability of the material with respect to loss of iodine (415). 

The potential supplied to an electrolytic cell determines whether or not electrolysis can occur, but the current flow and the time of electrolysis determine the amount of material electrolyzed. Recall Equation from Section 19-1 ... 

A description of an electrolytic cell has already been given under cell features (Section 1.3.2, Fig. 1.1c). Another example is the cell with static inert electrodes (Pt) shown in Fig. 3.1 where an applied voltage (Eappl) allows a current to pass that causes the evolution of Cl2 gas at the anode and the precipitation of Zn metal on the cathode. As a consequence, a galvanic cell, (Pt)Zn 2 ZnCl2 Cl2 iPt+, occurs whose emf counteracts the voltage applied this counter- or back-emf can be calculated with the Nernst equation to be... 

In an earlier note (p. 9) we mentioned the occurrence of overvoltage in an electrolytic cell (and overpotentials at single electrodes), which means that often the breakthrough of current requires an Uappl = Eiecomp r] V higher than Ehack calculated by the Nernst equation as this phenomenon is connected with activation energy and/or sluggishness of diffusion we shall treat the subject under the kinetic treatment of the theory of electrolysis (Section 3.2). 

For Cl2 or 02 evolution the stability of ruthenium based electrodes is not sufficient on a technical scale. Therefore the possibility of stabilizing the ruthenium oxide without losing too much of its outstanding catalytic performance was investigated by many groups. For the Cl2 process, mixed oxides with valve metals like Ti or Ta were found to exhibit enhanced stability (see Section 3.1), while in the case of the 02 evolution process in solid polymer electrolyte cells for H2 production a mixed Ru/Ir oxide proved to be the best candidate [68, 80]. 

In this section, you learned that you can calculate cell potentials by using tables of half-cell potentials. The half-cell potential for a reduction half-reaction is called a reduction potential. The half-cell potential for an oxidation half-reaction is called an oxidation potential. Standard half-cell potentials are written as reduction potentials. The values of standard reduction potentials for half-reactions are relative to the reduction potential of the standard hydrogen electrode. You used standard reduction potentials to calculate standard cell potentials for galvanic cells. You learned two methods of calculating standard cell potentials. One method is to subtract the standard reduction potential of the anode from the standard reduction potential of the cathode. The other method is to add the standard reduction potential of the cathode and the standard oxidation potential of the anode. In the next section, you will learn about a different type of cell, called an electrolytic cell. 

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. 

Some metals are extracted in electrolytic cells. In section 11.3, you saw the extraction of sodium from molten sodium chloride in a Downs cell. Other reactive metals, including lithium, beryllium, magnesium, calcium, and radium, are also extracted industrially by the electrolysis of their molten chlorides. 

Other important parts of the cell are 1) the structure for distributing the reactant gases across the electrode surface and which serves as mechanical support, shown as ribs in Figure 1-4, 2) electrolyte reservoirs for liquid electrolyte cells to replenish electrolyte lost over life, and 3) current collectors (not shown) that provide a path for the current between the electrodes and the separator of flat plate cells. Other arrangements of gas flow and current flow are used in fuel cell stack designs, and are mentioned in Sections 3 through 8 for the various type cells. 

In this section the use of amperometric techniques for the in-situ study of catalysts using solid state electrochemical cells is discussed. This requires that the potential of the cell is disturbed from its equilibrium value and a current passed. However, there is evidence that for a number of solid electrolyte cell systems the change in electrode potential results in a change in the electrode-catalyst work function.5 This effect is known as the non-faradaic electrochemical modification of catalytic activity (NEMCA). In a similar way it appears that the electrode potential can be used as a monitor of the catalyst work function. Much of the work on the closed-circuit behaviour of solid electrolyte electrochemical cells has been concerned with modifying the behaviour of the catalyst (reference 5 is an excellent review of this area). However, it is not the intention of this review to cover catalyst modification, rather the intention is to address information derived from closed-circuit work relevant to an unmodified catalyst surface. 

Fei O, called wiistite, has been studied from the viewpoints of thermodynamics and physicochemical properties. As mentioned in Section 1.1, stoichiometric FeO cannot be prepared under the usual conditions. Many investigators have studied the thermodynamic properties of wustite by use of various kinds of techniques. Here we introduce a study carried out by Fender and Rileywho used a solid electrolyte cell (see Section 1.4.8) to determine the equilibrium oxygen pressure Por The following cell was utilized,... 

Fluorine. The distinguished chemist Henri Moissan first prepared fluorine by the electrolysis of a solution of potassium fluoride in liquid hydrogen fluoride. Because of the extreme chemical activity of this element, the electrolytic cell employed had to be made of platinum. At the present time, fluorine is produced in the laboratory and commercially by the electrolysis of fused potassium hydrogen fluoride (KHF2) in the manner described in the section on electrolysis. 

Let us consider an electric circuit consisting of a dynamo and metallic conductors connecting the poles of the dynamo to the two poles of an electrolytic cell. The same amount of electricity flows through every section of the circuit, but the mechanism by which the current passes in the different parts of the circuit is of three different kinds ... 

In this case the losses in voltage due to the resistance of the electrolyte cannot be calculated by the usual method i. e. from the specific resistance of the solution, the spacing of the electrodes and the area of their surfaces this difficulty is caused by the presence of gas bubbles suspended in the solution these bubbles decrease the actiial cross section between the electrodes, i. e. the area covered by the lines of force, and consequently the resistance of the bath rises. The decrease of conductance of the electrolyte caused by suspended bubbles is more marked when the current is higher, the height of electrodes greater (as bubbles accumulate mainly at the upper end of the electrode), the electrolytic cell narrower and the movement of the electrolyte slower. The amount of bubbles in one unit of volume also depends on the viscosity and the temperature of the electrolyte. 

So far, we ve focused our attention on voltaic cells, which rely on spontaneous chemical reactions to drive them. In this section, we will look more closely at a different type of cell—one that requires electrical energy from an external source to allow a nonspontaneous reaction to occur. This new type of reaction is known as electrolysis, and it takes place in an electrolytic cell. 

In the last section, we looked at an electrolytic cell containing molten NaCl. Very high temperatures are needed to melt sodium chloride, so you might ask, Why can t we just dissolve it in water and do the same thing Before answering that question, you may wish to consider a procedure that you may have seen demonstrated for you (or even performed) in an... 

In the section on voltaic cells, we saw that the anode lost mass over time (as the metals were oxidized and went into solution), while the cathode gained mass over time (as the cations were reduced and plated on the surface). The voltaic cell, however, requires spontaneous reactions in each half-cell, which limits the types of electrodes that can be used. In an electrolytic cell, because we are adding electric current to the cathode and the anode, we can force nonspontaneous reactions to occur. In some cases, this allows us to use electrolysis for purposes other than separating a molten compound or aqueous solution. One of the more common alternate uses is the purification of different metals. 

In an undivided electrolytic cell, the electro-Fenton process leads to the destruction of organics contained in wastewaters by simultaneous oxidation with OH formed at the anode surface from reaction (19.9) and in the medium from Fenton s reaction (19.12). Parallel slower reaction of pollutants with weaker oxidants such as H202, H02, S20g2-, and O3 formed from reactions (19.1), (19.4), (19.10), and (19.11), respectively, is also possible. In addition, final carboxylic acids can form complexes with iron ions that are difficult to be oxidized by OH. We will see the notable influence of the anode material (Pt or BDD) on the degradation of these compounds in further sections. 

The thermodynamics and compositions of dissociation are calculated here. CO is not out of court , in an equilibrium, pure water electrolyte, cell with no catalyst. The neutral water can be made slightly acid, hydro-nium ions, or slightly alkaline, hydroxyl ions. The hydroxyl alternative. Figure A.l, is selected, the opposite to Section A.2.9. 

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