Electrolysis, of aqueous solutions

Solutions of compounds can also be electrolysed, but here the situation is more complicated. This is because there are extra ions present from the water in which the compound is dissolved. 

Water contains hydrogen and hydroxide ions, because a few of the water molecules ionise. 

If you electrolyse a solution of salt, NaCl, for example, there will be two positive ions present, hydrogen ions (H+) from the water and sodium ions (Na+) from the salt. Both of these travel to the cathode. Only the hydrogen ions are discharged, however, because it takes too much energy to force a sodium ion to accept an electron. Electron gain is also called reduction, the oxidation number has gone down from +1 to 0. 

In the same way there will be two negative ions present, hydroxide from the water and chloride from the salt both travel to the anode. What happens when they get there depends on the concentration of the solution if it is very dilute, it will be almost only the hydroxide ions that are discharged  

In all of these reactions the compounds must be ionic in nature and so contain ions. When electricity is passed through the melted compound or the solution, the ions move because they are attracted to the electrode which has opposite charge to their own. Remember this migration can 

This nonspontaneous reaction is the reverse of that used in the fuel cell. There are two aspects of the reaction to note. First, pure water is a poor conductor of electricity, and dilute solutions, often of acids such as sulphuric acid, H2SO4, are used. Second, because of overpotentials, it is necessary for a voltage greater than 2 V to be applied to the electrodes for the reactions to take place. These are as follows. 

When other anions and cations are present they can be discharged in preference if the energy requisite is less than that of the reactions above. For example, in the presence of sulphate ions, SO4 (aq), these may also give up electrons at the anode and be oxidised. However, this needs a potential of 2.05 V, which is more than the competing water reaction and so does not occur to any extent. 

This indicates that hydrogen gas will be evolved at the cathode rather than sodium metal. The two anode reactions are  

It would be expected, therefore, that oxygen gas would be evolved at the anode. However, the overpotential for oxygen production is substantial and in 

The chloralkali process, which involves the electrolysis of brine, is widely used for the production of sodium hydroxide and chlorine gas. During electrolysis it is necessary to keep the sodium hydroxide separate from the chlorine, to prevent the formation of sodium hypochlorite, NaOCl, and this determines cell design. In older processes, the cathode used was flowing mercury. At this electrode, sodium is formed, and this dissolves in the mercury to form a sodium amalgam. The sodium amalgam is removed continually from the cell and reacted with water to produce hydrogen gas and 

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Chlorine and caustic soda are coproducts of electrolysis of aqueous solutions of sodium chloride [7647-14-5] NaCl, (commonly called brine) following the overall chemical reaction... 

Electrolysis of Aqueous Solutions. The electrolytic process for manganese metal, pioneered by the U.S. Bureau of Mines, is used in the Repubhc of South Africa, the United States, Japan, and beginning in 1989, Bra2il, in decreasing order of production capacity. Electrolytic manganese metal is also produced in China and Georgia. 

Electrolysis. Electrowinning of zirconium has long been considered as an alternative to the KroU process, and at one time zirconium was produced electrolyticaHy in a prototype production cell (70). Electrolysis of an aH-chloride molten-salt system is inefficient because of the stabiUty of lower chlorides in these melts. The presence of fluoride salts in the melt increases the stabiUty of in solution, decreasing the concentration of lower valence zirconium ions, and results in much higher current efficiencies. The chloride—electrolyte systems and electrolysis approaches are reviewed in References 71 and 72. The recovery of zirconium metal by electrolysis of aqueous solutions in not thermodynamically feasible, although efforts in this direction persist. 

Table 6. Production of Metals by Electrolysis of Aqueous Solutions ...

Explosion Hazards. The electrolysis of aqueous solutions often lead to the formation of gaseous products at both the anode and cathode. Examples are hydrogen and chlorine from electrolysis of NaCl solutions and hydrogen and oxygen from electrolysis of water. The electrode reactions. 

At present about 77% of the industrial hydrogen produced is from petrochemicals, 18% from coal, 4% by electrolysis of aqueous solutions and at most 1% from other sources. Thus, hydrogen is produced as a byproduct of the brine electrolysis process for the manufacture of chlorine and sodium hydroxide (p. 798). The ratio of H2 Cl2 NaOH is, of course, fixed by stoichiometry and this is an economic determinant since bulk transport of the byproduct hydrogen is expensive. To illustrate the scde of the problem the total world chlorine production capacity is about 38 million tonnes per year which corresponds to 105000 toimes of hydrogen (1.3 x I0 m ). Plants designed specifically for the electrolytic manufacture of hydrogen as the main product, use steel cells and aqueous potassium hydroxide as electrolyte. The cells may be operated at atmospheric pressure (Knowles cells) or at 30 atm (Lonza cells). 

Ions not solvated are unstable in solutions between them and the polar solvent molecules, electrostatic ion-dipole forces, sometimes chemical forces of interaction also arise which produce solvation. That it occurs can be felt from a number of effects the evolution of heat upon dilution of concentrated solutions of certain electrolytes (e.g., sulfuric acid), the precipitation of crystal hydrates upon evaporation of solutions of many salts, the transfer of water during the electrolysis of aqueous solutions), and others. Solvation gives rise to larger effective radii of the ions and thus influences their mobilities. 

Cataldo F (1992) Effects of ultrasound on the yield of hydrogen and chlorine during electrolysis of aqueous solutions of NaCl or HC1. J Electroanal Chem 332 325-331... 

The same group, in a previous work, reported on the realization of a hybrid anode electrode [197]. An appreciable improvement in methanol oxidation activity was observed at the anode in direct methanol fuel cells containing Pt-Ru and Ti02 particles. Such an improvement was ascribed to a synergic effect of the two components (photocatalyst and metal catalyst). A similar behavior was also reported for a Pt-Ti02-based electrode [198]. Another recent study involved the electrolysis of aqueous solutions of alcohols performed on a Ti02 nanotube-based anode under solar irradiation [199]. 

The electrolysis of aqueous solutions may not yield the desired products. Sir Humphry Davy (1778-1829) discovered the elements sodium and potassium by electrolyzing their molten salts. Before this discovery, Davy had electrolyzed aqueous solutions of sodium and potassium salts. He had not succeeded in reducing the metal ions to the pure metals at the cathode. Instead, his first experiments had produced hydrogen gas. Where did the hydrogen gas come from ... 

It is generally agreed that electrolysis of aqueous solutions offers the best prospect for the production of hydrogen from water, because of the easy separation of the H2 and O2 products and because of the relatively low energy consumption if catalytically active metal electrodes are used. Thus, the minimum energy requirements are those for which water, hydrogen and oxygen, each at 1 atmosphere pressure, are in equilibrium ... 

The main reason for avoiding water as a solvent is the fact that the electrolysis of aqueous solutions of alkali and alkaline-earth metal salts commences at 1.7-2.0 volts (depending on the electrode material) and results in the evolution of O2 and H2. If the cell itself has a higher voltage, internal electrolysis can, but not always does occur, accompanied by the evolution of H2 and O2 and by self-discharge (117). However, this fact does not preclude attempts to create moist primary batteries with Li, Na or Ca, if the activity of H2O is kept sufficiently low. 

Further examples of recent attempts to reduce the consumption of electrical energy are the electrolysis of aqueous solutions of methanol (but CO2 is still produced at the anode) [78, 79] and water electrolysis using ionic liquids as electrolytes [80]. In the latter case, the authors claimed the possibility of obtaining high hydrogen production efficiencies using an inexpensive material such as low-carbon steel. 

Table 7.3 Electrolysis of aqueous solutions of ammonia experimental parameters."...

Thermal phenomena at the Cd electrode during the electrolysis of aqueous solutions of Cd(II) were investigated [219]. The determination of heat flux, heat quantity, and temperature gradient at the... 

Since chemical reduction means gain of electrons, electrolysis is the most direct way of recovering a metal from its ores, as long as these can be handled in a fluid state.6 Consideration of E° values for reactive metal halfcells such as Na+(aq)/Na(s), Mg2+(aq)/Mg(s), and Al3+(aq)/Al(s) (-2.71, -2.36, and -1.67 V, respectively) shows that these metals can never be obtained by electrolysis of aqueous solutions of their salts, as H2 would be produced instead, but they can often be obtained by electrolysis of suitable molten salts such as NaCl and MgCl2 ... 

Other industrial processes involve the electrolysis of aqueous solutions. To help you to understand what is happening in these processes, we will first consider the electrolysis of dilute sulfuric acid. 

Electrolysis of aqueous solutions of the following using inert electrodes sodium chloride, copper(n) sulfate, sodium sulfate and sodium hydroxide. 

The dilute acid solution contained hydrogen ions supplied by the solute and also by the slight ionization of water. Particular attention should be directed to the fact that the solution contained two anions (i.e., a very low concentration of hydroxyl ions) and a relatively high concentration of sulfate ions. Despite the fact that the sulfate ions were present at much greater concentration than the hydroxyl ions, only the latter were discharged at the anode. From this fact alone, it may be concluded that not all anions are discharged on electrolysis of aqueous solutions and that in such cases the anion of the solvent is involved in the reaction at the anode. 

In view of these results, one may recognize that not all cations are discharged during the electrolysis of aqueous solutions and that the cation of the solvent may be discharged at the cathode in place of the cation of the solute. 

On the basis of the preceding generalizations, it is possible to predict the products of electrolysis of aqueous solutions of simple salts. It is not... 

Similar predictions may be made readily with regard to the electrolysis of aqueous solutions of other salts. 

Since all soluble acids furnish hydrogen ions, the electrolysis of aqueous solutions of acids yields hydrogen gas at the cathode. Consequently, the only question that remains is whether the anion of the particular acid is dischargeable. If not, oxygen gas is liberated. There are two possible types of behavior. One is illustrated by hydriodic acid ... 

The following subsections discuss a few chemical processes that involve the electrolysis of aqueous solutions or fused salts. It is not intended to provide here an exhaustive treatment of the subject but rather to select a few typical cases that serve to acquaint the student with the nature, the scope, and the importance of these industries. 

Chlorine is produced almost entirely by the electrolysis of aqueous solutions of alkali metal chlorides (Fig. 1), or from fused chlorides. Brine electrolysis produces chlorine at the anode and hydrogen along with the alkali hydroxide at the cathode. At present, three types dominate the industry the diaphragm cell, the membrane cell, and the mercury cell, and there are many variations of each type. 

Aluminum is present in most rocks and is the most abundant element in the earth s crust (eight percent by weight.) However, its isolation is very difficult and expensive to accomplish by purely chemical means, as evidenced by the high E° (-1.66 v) of the A13+/A1 couple. For the same reason, aluminum cannot be isolated by electrolysis of aqueous solutions of its compounds, since the water would be electrolyzed preferentially. And if you have ever tried to melt a rock, you will appreciate the difficulty of electrolyzing a molten aluminum ore Aluminum was in fact considered an exotic and costly metal until 1886, when Charles Hall (U.S.A) and Paul Herault (France) independently developed a practical electrolytic reduction process. 

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