The Chemistry Student THE STRUCTURE OF MATTER. V. PRIMARY CELLS OTTOREINMUTH, ASSOCIATEEDITOR
Last month, in connection with our discussion of ionization, we considered the phenomenon of electrolysis.' In other words, we noted how a solution of an electrolyte behaves when it is subjected to the action of an electric current of external origin. Let us now turn our attention to the manner in which electrolytic action can give rise to an electric current.
Solution I t will be necessary to begin with an apparent digression, for we must first examine briefly the phenomenon of solution. We know from experience and previous study that most solids have a certain tendency to dissolve in water. This tendency, for lack of a better name, has been called the solntion pressure. Opposed to the solution tendency of the solid is the tendency of dissolved matter to pass out of solution; this is called the osmotic pressure. In the case of a salt and water, the solution is said to be saturated when the two opposing tendencies equal and balance each other. A saturated solntion in the presence of excess solid salt is in a state of dynamic equilibrium-salt is passing into and out of solntion a t equal rates. This state of equilibrium may be reached from either side. Solid salt added to a supersaturated solution causes the crystallization of more salt nntil a state of equilibrium is reached. Solid salt added to an unsaturated solution dissolves nntil the solution becomes saturated. If we consider a strip of metal2 immersed in water, the case is somewhat similar, but there are very important differences. The metal, like the salt, has a certain tendency to go into solntion but, unlike the salt, i t is restricted in its manner of doing so. The salt may enter solution either in the form of its molecules or its ions. When it does either or both, the solution remains electrically neutral, for positive and negative ions go into solution in equal numbers. The metal, on the other hand, can enter the solution only in the form of positive metallic ions. Neither metallic molecules nor negative metallic ions go into solntion. The result is that as soon as metal begins to go into solution, the metal plate becomes negatively charged from the excess electrons left behind, and begins to attract 5, 1 6 3 9 4 6 (Dec., 1928). 'THISJOURNAL, It is understood that in the present article we shall confine our attention to metals which do not decompose water to any appreciable degree.
positive ions electrostatically. Likewise, the dissolved ions confer a positive charge on the solntion, which consequently opposes the entrance of further positive ions. A new balance (but not a true solution equilibrium) is set up. Now only a comparatively small difference in charge between plate and solution is necessary t o establish the new equilibrium. Consequently, a balance is reached when so few metal ions have gone into solution that no change in concentration is measurable a t all. We see therefore that while a salt will continue to go into solution until a state of satnration is reached, practically no metal a t all dissolves in water or in a solution of the metal salt. The equilibrium which now exists is not a true solution equilibrium. I n solntion equilibrium: osmotic pressure = solution pressure In the state of affairs which we have described: osmotic firessure electrostatic forces = solution pressure I t is as though a grocer were to cause a half-pound of sugar to balance a two-pound weight by affixing a magnet under one pan of his balance. The sugar and the weight would be balanced but they would not be in qavitational equilibrium. Let us scrutinize in a little more detail the equilibrium between a metal strip and isolution of its salt. We spoke a t the outset of the mutually opposing solution pressure and osmotic pressure of a salt. It is evident that a t any concentration between infinite dilution and saturation, the solntion pressure is thegreater of the two. We can therefore imagine an algebraic sum of the two opposing pressures which might be called the net solution pressure or tendency of the salt. Similarly, we might imagine a net solntion tendency for a metal strip immersed in water or in a solution of one of its own salts-provided the salt solution is unsaturated with respect t o the true solution equilibrium between the metal and its ions. At the risk of being tedious, it is necessary to emphasize the point that equilibrium is the true criterion of the saturation phenomenon. The impression that a solution is saturated with a substance when i t contains "all of the substance that it can h o l d is utterly misleading, and any idea of saturation based upon such an impression will be found hopelessly confusing in the explanation which is to follow. We must distinguish sharply between a salt solution which is saturated with respect to the solid salt and one in which there is satnration equilibrium between metal ions and the solid metal. A solution saturated with zinc sulfate (or any other zinc salt for that matter) is very far indeed from containing enough zinc ions t o be in saturation equilibrium with metallic zinc. A saturated solution of silver nitrate, on the other hand, is highly supersaturated from the standpoint of the metallic siluer-silver ion equilibrium. In the light of what we have said, it is evident that one of three condi-
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tions may prevail when a strip of metal is immersed in a solution of one of its own salts. (1) If the solution is unsaturated with respect to the true metal-metal ion solution equilibrium, the algebraic sum of the solution pressure and the osmotic pressure will be a net solution pressure. Metal ions will tend to pass into the solution from the metal strip. That tendency will quickly be counterbalanced by the negative charge built up on the zinc plate and the equivalent positive charge built up in the solution. It is apparent that the magnitudes of these charges are determined by the magnitude of the net solution pressure. A highly unsaturated solution gives rise to a large net solution pressure with a consequently large negative charge on the metal plate (2) If the number of metal ions present in the salt solution is precisely that necessary for true solution equilibrium with the metal, the algebraic sum of the solution and osmotic pressures will he zero. There will be no net tendency for metal ions either to dissolve or to leave the solution. In such a case no charge a t all is built up. (3) If the solution is supersaturated with metal ions, with respect to the true metal-metal ion equilibrium, there will be a net osmotic pressure and a tendency for metal ions to deposit on the metal plate. Since the metal ions are positively charged, their deposition results in the building up of a positive charge on the plate. At the same time, the solution itself becomes negatively charged because of the excess negative ions left behind. The osmotic pressure is therefore opposed and counterbalanced by the electrostatic attraction of the negatively charged solution for positive ions and the repulsion of the positively charged plate for additional positive charges, just as we previously saw that solution pressure was opposed and counterbalanced by the reverse electrostatic forces. In neither case are sufficient metal ions transferred to make a measurable difference in the concentration of the solution, and in neither case is a true solution equilibrium attained. The magnitude of the positive charge on the plate is obviously determined by the magnitude of the net osmotic pressure which it offsets and which is in turn dependent upon the concentration of the solution. The Concentration Cell We see, therefore, that we can vary the potential difference between a metal and a solution of one of its salts by varying the concentration of the salt solution. This fact gives us exactly the opportunity we need in order to be able to produce an electric current. Electrical energy, like heat energy, "runs down-hill" only.3 In order to accomplish a transfer of a This generalization applies, of course, t o all forms of energy. More elegantly expressed, it constitutes the First Law of Thermodynamics.
energy, we must first establish a difference in energy levels. This we can now do. I n order t o deal with a concrete example, let us suppose that Fignre 1~ represents a zinc plate immersed in a very dilute solution of zinc sulfate. Fignre 1~represents a zinc plateimmersed in a concentrated solution of zinc sulfate. Since both solutions are highly unsaturated with respect to the true Zn Zn+++2c solution equilibrium, it follows that in each system, the zinc plate is negatively charged and the solution is positively charged. But, since the solutions are of diierent concentrations, i t follows that CURE the charges in A and B will be of diierent magnitudes. If the previous discussion has not made this consequence obvious, a glance at the following equations will do so. For the system A we may write: (solution p r e ~ s u r e )= ~ (osmotic p r e ~ s u r e ) ~ (electrical charges)* For B we may write: (solution pressure), = (osmotic pressure), (electrical charges), Now since we are dealing with the same metal in both cases, (solution pressure), = (solution pressure), Hence: (osmotic p r e ~ s u r e ) ~ (electrical charges)* = (osmotic (electrical pre~sure)~
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=
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charge^)^ Since the osmotic pressure is directly proportional to the concentration of Zn++, (osmotic pressure), > (osmotic pressure), which necessitates that: (electrical charges)~> (elecFIGURE 2.-A CONCBNTRATION CELL trical charge^)^. If, now, we connect the two plates by means of a wire which will permit the transfer of electrons and the two solutions by means of a siphon or "salt-bridge" which will permit the transfer of sulfate ions, a current will flow. I n Figure 2 the movements of electrons and negative ions are indicated. The direction conventionally assigned to the current is opposite to the direction of electron transfer. (What would happen if the
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metal plates, only, were connected, the salt-bridge being omitted?) A cell of this type is known as a concentration cell. Obviously i t develops its greatest E.M.P. when the difference in concentration between the two solutions is greatest, and it "runs down" rapidly as it operates. When the two solutions have arrived a t the same concentration, no difference in charge exists between one side of the cell and the other and no further current flows. An interesting experiment illustrating the action of the concentration cell is described by Professor DILUTE M e l l ~ r . ~(See Figure 3.) SOLUTION
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A layer of a concentrated solution of stannous chloride in hydrochloric acid, about 10 cm. deep, is placed a t the CPNCENTRATTP bottom of a cylinder, and above this a layer of a dilute SOLUTION solution. A rod of metallic tin is fixed through a hole in the cork so that it is suspended axially in the liquid in the cylinder. The rod of tin thus represents both electrodes and FIDm 3,-A CONconnecting wire of a concentration cell. Tin is dissolved CgNTnArIoNceLL E ~ by the more dilute solution, and precipitated from the P ~ R I M E N T more concentrated solution. The diagram illustrates the appearance of the rod of tin after the vessel has stood a couple of days.
Zinc and zinc chloride do not work so well. Two-Metal Cells Since the concentration cell is incapable of delivering a uniform current and since it runs down so quickly, it is of no practical value. A re-examination of the equations (or -ELECTRONSrather, the equivalencies) on page 120 reveals the fact that in constructing our concentration cell we accomplished the necessary diierence in charge between the two half-cells by varying the osmotic pressures. The idea of accomplishing a like result by varying the solution pressures immediately sugF~cuns4.-THE Zmc-COPPERCELL gests itself. How can this be accomplished? So far we have said little about the fundamental nature of solution pressure-and for the very good reason that little is known on that subject. However, we do know that each metal has its own characteristic solution '"Modern Inorganic Chemistry," J. W. Mellor, Longmans, Green and Co., New York City, 1925, p. 438.
pressure and that these pressures vary greatly in magnitude from one metal to another. Incidentally, the solution pressures vary in the order of arrangement of the "electromotive se~ies."~ As one proceeds down the table, the solution pressures become smaller. Hence, to vary the solution pressures in our half-cells, i t is necessary only to choose different metals for our poles. Zinc and copper furnish good working examples. Now it is not true (as some textbooks stilte) that a strip of copper immersed in a solution of copper sulfate has a positive charge.' Nevertheless, the solution pressure of copper is very much smaller than that of zinc. Consequently, the negative charge on an immersed copper plate is much smaller than that on a zinc plate.' Hence, if we arrange a cell like that illustrated in Figure 4, current will flow as indicated. (It is not deemed necessary to discuss the operation of this cell in so much detail as was devoted to the concentration cell, since with the exception of the means of producing the diierence in charge, the principle of operation is identical.) The Daniel1 cell is one form of the copper-zinc cell. Here the necessity for an extra vessel and a salt-bridge is eliminated by placing the denser solution of concentrated copper sulfate in the bottom of a battery-jar and floating the less dense solution of dilute zinc sulfate on top of it. Since the operation of the cell involves the dissolving of metallic zinc and the deposition from solution of copper ions, the concentrations of the respective solutions are so chosen as to favor these processes. The zinc sulfate solution is either very dilute or is replaced by dilute sulfuric acid. The copper sulfate solution is saturated and is kept so by an excess of solid copper sulfate a t the bottom of the jar. The oxidation-reduction cell also falls in the class of primary cells, but we shall not discuss it a t this time. "ee the fourth article of this series, THISJOURNAL, 5, 1645 (Dec., 1928). 0 For a more detailed treatment of this point, see "Inorganic Physical Chemistry," C. H. Cartledge, Cinn and Company, New York City, 1925, pp. 311-6, but especially p. 314. 7 According to the conventional system of nomenclature the B plates in the cells illustrated in Figures 2 and 4 are positive and "current" flows from the B plates to the A plates. Actually, all the plates are negatively charged but the A plates have greater negative charges than the B plates. Electrons flow from the A plates to the B plates.
