a , , , , *
HIGH-SCHOOL CHEMISTRY
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High-School Instruction the Electromotive Series WILLIAM T. HALL Rochester, Massachusetts Most textbooks of elementary chemistry err i n trying to make the subjecffarcinating and easy by neglecting simple mathematical relations. Host teachers of physics know that a studenf must have a goodfoundation i n algebra before studying physics, but feachers of chemistry lay less empharis upon mathematical training prior to the study of chemistry. This is unfortunafe, because mathematical reasoning isjust as important i n chemistry as it is i n physics and if is wrong to assume a lower mentality i n books on chemistry than i n books on physics. T h i s fendency to make chemistry booksgh,6ed primers har been regretted by men like W i h e l m Ostwald and Alesander Smith. The fendency to avoid mathematics and invent mythical reactions is illustrated by the fragmentary way i n which the elcctromofive series is taught to beginners. The elecfromotiue, or potential, series covers every care of oxidation and reduction, and high-school students should presumably have sufiienf mathematical training to benefit by the application of mathematical remoning to its consideration.
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HEN students enter college or technical school wlth Ideas that have to be changed the tendency is to be critical of the preliminary training. Some of us, however, who have taught both beginners and postgraduate students a t about the same time find that the older students have either forgotten or failed to comprehend previous instruction. Nevertheless it is important in high-school teaching not to introduce ideas that have to be changed later, and to remember that i t is much easier for youth to absorb new ideas a t the age of 16 than a t 21 or later, although with an enlarged background i t is easier to correlate knowledge. It seems almost incredible, but i t is the writer's experience that the minds of high-school boys and girls are often more alert than those of some mature college professors. One of the important fundamental theories of chemistry concerns oxidation and reduction. Most elementary texts of today acknowledge the fact that the most comprehensive definition of oxidation is loss of one or more electrons by the atom that is oxidized. In the same way, we say that an atom is reduced when i t accepts one or more electrons. The two reactions always take place simultaneously, for in our chemical reactions the only way that an element can lose electrons is by giving them to some other atom. Three decades ago most textbooks taught reactions such as the oxidation of an acidic solution of ferrous sulfate by potassium permanganate on the assumption that FeS04 = FeO.SOa and KMnO4 = K20.Mu20,. This is the old dualistic conception of Berzelius and I think it must please the spirit of this great Swedish chemist to know that his ideas still prevail in the minds of many chemists today. According to his dualistic theory the
following reactions can be assumed to be involved in the oxidation of FeSOl by KMnO, in the presence of HzSo4.
+ MnzO? M n ~ a= 2Mn0 + 5 0 10FeS0, = 10FeO + 10S08
2KMn0, = K 2 0
1OFeO
(1)
(2) (3)
+ 5 0 = 5FenO~
8Hd0' = 8HsO
(4)
+ 8SOs
(5)
+ SO* = &SO+ + 2SOo = ZMnSO, 5Fe20, + 15SOa = 5Fe&30Jn K20
(6) (7)
2Mn0
(8)
If we add up these eight equations and cancel out all molecules which are on both sides of the equality signs we get the correct equation
After a few years of practice, most of us became proficient in inventing these preliminary reactions and struggled with our classes to make them see how easy i t is to balance equations of oxidation-reduction. Today, however, we think of the Mu in KMn04 as having a valence of 7 and regard it as receiving 5 electrons from 5 atoms of Fe++, and we imagine these two reactions MnOt-
+ 5e + 8H++ 5Fe++
=
=
Mn++
5Fe++'
+
+ 4H20 5e
(reduction) (oxidation)
to take place, which add up to Mn04-
+ 5Fe++ + 8H+ = Mn++ + 5Fe+++ + 4H10
which expresses the same relation of Mn:Fe as in the above case, where eight fictitious reactions were im-
1. The table is merely qualitative, because no nuagined. The writer had fwo of his former teachers change to this latter method of writing oxidation-re- merical value is assigned to each metal to express its duction equations a t about 1915, just before the retire- actual electrolytic solution pressure or oxidation pokntial. ment of one of them. This professor found that he Some of these elements lie very close together, as K, could get the newer method into the minds of his fresh- Na, and Ba, but between Ca and Mg there is quite a man students in much less time than the laborious difference. 2. The reaction involved is not stated. In the case dualistic procedure. Essentially the same method, although the term negatiere valence is used instead of of some elements, such as Fe, the metal can appear in electron, is to be found in Prescott and Johnson's three different places according to whether the oxida"Qualitative Analysis," in an edition printed about 60 tion is Fe to Fe++, Fe to Fe+++, or Fe++ to Fe+++. All three oxidations should find a place in the table. years ago. It is important to state the electrode reaction. WHAT IS THE ELECTROMOTIVE SERIES? 3. The table gives no information with respect to Early in the 19th century Jons Jakob Berzelius, the concentration of the ions of the metal. This is with his friend Hisinger, camed out a series of experi- very important. Of course, it is easy to understand ments upon the electrolysis of salts which made a that, in accordance with the mass action law, it is great impression upon h i . He regarded the atoms harder to make an element form ions in solution when of the alkali metals as very positive in nature and the there are already many ions present, and it is easier to atoms of oxygen as "absolutely negative" with hydro- plate out a metal by electrolysis from a solution congen a t or near the neutral point. This was the begin- taining a high concentration of its ions than from a soluning of the potential or electromotive series derived tion containing a very low concentration. 4. The table should emphasize the fact that every from purely chemical considerations. According to reaction involved is a reversible one. When i t takes Berzelius, place in one direction i t represents an oxidation and "Every chemical compound depends entirely and alone upon when i t takes place in the opposite direction it is a retwo opposite forces of positive and negative electricity: every compound can be divided into two parts of which one is positive duction. More or less a t random, letus take 12 electrode reand the other negatively electric. For example, sulfate of soda i s not composed of sulfur, sodium, and oxygen, but of sulfuric actions and show how the table should be written. acid (SOJ and soda (NaaO)."
Bases, according to Berzelius, are positive oxides combined with water, and acids are negative oxides which are also hydrated. Berzelius, therefore, was the iirst to propose a rational electrochemical theory and suggest an electrochemical or potential series. The conception of Berzelius, however, was really qualitative in nature and not based upon quantitative measurements. His ideas were based on the theory that there were two kinds of electricity, positive and negative. Today we believe that the electric current represents a flow of electrons, or negatively charged particles, and that an element can be positive with respect to certain elements but negative toward others. We now think of the electromotive series as expressing the tendency of elements to become oxidized or to be reduced. In many textbooks of chemistry we find under the heading "Electromotive Series of the Metals" a vertical column in which the metals are placed in the following order: Li, Rb, K, Na, Sr, Ba, Ca, Mg, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, H, Sb, Bi, As, Cu, Hg, Ag, Pt, An. Such a table shows the relative tendencies of these metals to lose one or more electrons and become oxidized. The strongest reducing agents are the elements a t the top of the table and, in general, any element can reduce the ions of another element that follows it in the table. Such a table cannot be regarded as adequate for a t least four important reasons, and if it is corrected in these respects it becomes more useful and more intelligible to beginners.
++
L i e LiC e . . . . . . . . . . . . . . . . . . . . . -3.02 C a d CaC+ 2 e . .. . . . . . . . . . . . . . . . - 2 . 6
F ' $ ~ / ~ F ~. . +. . ~. . . . . . . . . . . . .
v.
+1.96
The term normal signifies in this table that in each reaction the concentration of every ion involved is molar. Thus in the case of Zn/Zn++ the value of -0.76 refers to the potential when metallic zinc is in contact with a molar solution of Zn++ ions. With the reaction for Fe++/Fe+++ the value holds only when the concentrations of both Fe++ and Fe+++are molar. In the Mn++/MnOa- reaction the concentrations of Mn++, MnOl-, and H+ are all molar; HzO is neglected in this last case because the quantity of HzO that enters into reaction is inappreciable with respect to the large quantity of water present in an aqueous solution. The significance of the temperature is this: The electrolytic solution tension, or tendency of the metal to form ions, acts in opposition to the osmotic pressure of the ions in solution, which is proportional to the absolute temperature. The value 30'6. (303' absolute) is chosen to simplify the application of the Nernst rule, to which we shall refer later. The significance of the sign before the voltage is a little confusing. Both physicists and chemists use the table but the attention of the physicist is focused upon
the direction of the current in the external wire, and he regards Cu positive with respect to Zn in the Daniell cell because the electrons flow in the external wire from the zinc electrode to the copper. The chemist, on the other hand, is thinking of movement within the solution and regards the zinc as positive with respect to copper. There is a complete circuit, and both the physicist and the chemist understand that the electron flow is from the zinc to the copper in the external wire and the anions go from the copper to the zinc in the solution. It is conventional to regard the electromotive series as one of oxidation potentials when the values are given as negative a t the top of the table and one of reduction potentials when positive signs are given opposite those metals which the chemist regards as most positive. In the case of oxidation potentials, the negative sign means that the electrode is negative with respect to the solution. It is absolutely arbitrary as to how the positive or negative potential is defined. It is also arbitrary to call the position of H/H+ as exactly zero. If there is an actual zero it is somewhat lower in the series but there is no absolute zero that is easy to measure exactly, and it suffices to use an arbitrary value. THE DANIELL CELL
cern us here. When the temperature is chosen as 30°C., and natural logarithms are converted into logarithms to the base 10, the Nernst formula can be writ0.060 ten Eao0= EO log c, where E" is the normal
+
electrode potential given in the table, cis the molar concentration of the ions, and n is the valence change or number of electrons shown in the electrode reaction. This expression holds, in this simplified form, when only one ion is involved in the electrode reaction. Since log 1 = 0, log 0.1 = - 1, log 0.01 = -2, we can see a t a glance that the Nernst formula serves to construct the following table of electrode potentials a t different dilutions: Molar Concenlration of Cui+ 10 M . . . . . . . . . . . . . . . 1 M... . .
. .. . . . .. . .
0.1 M . . . . . . . . . .. . . 0.01 M . . . . . . . . . . . . 0.001 M . . . . . . . . . . . 0.0001 M . . . . . . . . . .
E Value
++ 00.375 ". 3 4 5 v . (Eovalue) +O ,315". +0.285 v. 0.255 v. 0.225 v.
++
In the same way the Nernst formula shows that if Zn/Zn++ is -0.76 in a molar solution of Zn++ i t would be -0.73 in a 10 M solution. Since, however, only one mole of Zn++ can be formed while one mole of Cu++ is being reduced to Cu i t is clear that, while the two single potentials approach one another as the Daniell cell is being used, nearly all of the change will be a t the copper electrode.
The use of the table can be illustrated by considering a cell in which a strip of zinc is immersed in a molar solution of zinc ions and a strip of copper rests in a molar solution of cupric ions. When the two pieces of metal ELECTROLYTIC SEPARATIONS are connected by a metal wire and the two solutions are When a metal is deposited upon the cathode during connected by electrolytes, two reactions a t once take place. At the zinc electrode there is an oxidation, and electrolysis, i t is a case of reduction, and the lower the the reaction is Zn = Zn++ 2e; a t the copper elec- element in the list the easier i t is to accomplish the retrode there is a reduction, Cn++ 2e = Cu. Both re- duction. Thus gold plates out before platinum, copper actions take place simultaneously and the final result before bismuth, and hydrogen is evolved before any is Zn Cn++ = Zn++ +Cn, just as when a strip of lead is deposited in the presence of mineral acids which zinc is placed in a copper sulfate solution. If both are unbuffered. If we have a t the start 100 ml. of Cu++ and Zn++ solutions are molar a t the start the solution containing, say, 0.636 g. of Cu++, the solution initial voltage of the cell is the algebraic differencebe- is approximately 0.1 M in Cu++ and the electrode potween -0.76 and +0.345, as given in the table. This tential is about +0.315; when only 0.1 mg. of Cu++ is 1.01 v. As the cell is used the two electrode poten- remains in solution (as little as we can weigh on an ortials get closer together and the measured potential dinary chemical balance) the solution is about 0.000016 difference between the two electrodes becomes a little M and the electrode potential for Cu/Cu++ is about smaller, but the two potentials are so far apart that the +0.172. In general, therefore, we can deposit all the Cu+f reaction will continue until all the Cu++ ions have been from a solution without changing the electrode potential deposited or all the zinc metal has dissolved. as much as 0.2 v. Separations can be accomplished EFFECT OF ION CONCENTRATIONS ON ELECTRODE POTENelectrolytically if there is 0.2-v. difference in their norTIAIS mal electrode potentials. In an acid solution it is This is very important though seldom mentioned in possible to deposit all weighable copper without having elementary treatises. Any high-school student who any evolution of hydrogen gas. Cd/Cd++ has the knows what a logarithm is can understand it. In the normal electrode potential of -0.40 v. and cannot, first edition of Nernst's "Theoretische Chemie," pub- therefore, be deposited from solutions which are dislished in about 1893, Nernst traced the relationship be- tinctly acidic with mineral acid: Why is it then that tween electrolytic solution tension (or oxidation poten- cadmium can he deposited before copper from an alkatial) and the osmotic pressure of the ions formed. The line solution containingcyanide? In such a solution the Nernst equation can be arrived a t in a t least two dif- cadmium is present chiefly as Cd(CN)a-' ions and the ferent ways, but i t is done with the aid of integral cal- ratio of the concentration of the complex ion to that of culus. The derivation of the equation need not con- Cd++ is about 1017:l. The copper, on the other hand,
+
+
+
is present chiefly as C U % ( C N )ions ~ - ~ and the ratio of the complex to simple Cu+ ion in N KCN solution is about 1OZ6:1. If sufficient KCN is added, the cadmium is precipitable by H a or by electrolysis while the copper remains in solution. ELECTRODE POTENTIAL OF HYDROGEN
The table of normal electrode potentials shows us that hydrogen ions can be expected to oxidize (or dissolve) the metals above it. The reaction Zn
+ 2H+
ZnC+
+ HS
represents an oxidation of zinc by hydrogen ions when it takes place in the direction of left to right. If a piece of zinc is placed in a solution which is molar in Fe+++ and molar in H+, what will happen? It has sometimes been assumed that the zinc dissolves in the acid and the nascent hydrogen reduces the Fe+++. This is nonsense because the table of normal oxidation potentials shows that zinc is a much better reducing agent than H + because the Zn/Zn++ potential is much higher in the table. The table shows us that we would not expect dilute sulfuric or hydrochloric acid to dissolve copper, but since Fe++/Fe+++ is still lower in the scale we- can why an acidic solution of a ferric . . understand . .. Salt dissolves copper very easily. I n the electrolysis of sodium chloride between platinum electrodes, is i t easier to discharge Na+ or H+? are atElementary texts usually state that Na+ tracted toward the cathode, and this is absolutely true. On the other hand, the normal electrode potential of Na/Na+ is -2.72 v., and for a 10 M solution i t is -2.78 the Nernst rule. A lo of NaOH with OH- = 10 will give to H + the value 10-", for in aqueous solutions [H+J X [OH-] =
\
The electrode reaction a t the cathode is always the easiest possible reduction, as shown by the table of normal oxidation potentials, and the electrode reaction a t the anode is always the easiest possible oxidation, and it is silly to assume that the ion which carries most of the current must be discharged when it reaches the electrode. The reaction 2H20
+ 2e = 20H- + HZ
+
is an easier reduction than that of Na+ e = Na, and this last reaction cannot possibly take placein the presence of water. THE ELECTROMOTIVE SERIES APPLIED TO NONMETALS
The electromotive series applies to all reactions of oxidation-reduction and is not, as so many texts imply, applicable only to metal displacement. For the halogens we find the EOvalues to be
These values show that I- is a better reducing agent than Br-, cl-, or F- and that fluorine is a stronger oxidizing agent than chlorine, bromine, or iodine. In analiticd chemistry we have many iodimetric determinations. In sliehtlv basic solutions iodine will oai~-, dize in arsenite to an asenate, and we standardize an iodine solution by weighing out pure AsaOa. On the other hand, in strongly acidic solutions, alkali iodide will reduce quinquevalent arsenic to the trivalent condition. hi^ can be weby the application of the Nernst rule and considering the concentrations in reactionssuch as ~~~~
~
~
~
. . . . . . . . . . . .
lo-"
where [H+] denotes the concentration of H + in moles per liter and [OH-] is the concentration of OH- in moles per liter. Since the Nernst rule is Eta = Eo 0.06 log c this becomes for H+, E 3 = ~ 0.00 - 0.06 pH. Therefore in 10 M NaOH the electrode potential of H/H+ = -0.90 v. Even in this strongly basic solution i t is easier to discharge H + than Na+. The absurdity of assuming that Na+ ions are discharged and that sodium then reacts with water to form NaOH was pointed out by M. LeBlanc in a text translated into English by W. R. Whitney in 1895.
+
In nearly neutral solutions the upper reaction 'takes place completely, but in very strongly acidic solutions the lower reaction takes place. In reactions such as the latter and the very common MnO4-
+ 8H+ + 5e
+ MnC+
+ 4H10
the concentration of every ion involved must be taken into consideration. For the last reaction the Nernst equation becomes Es
=
E
[MnO,-]log [H+]e + 0.060 7log [Mn++
