LETTERS To the Editor:
which was printed in the Journal of Chemical EducaIn the January issue of the JOURNAL OF CHEMICALtion rather than in the Journal of Physical Chemistry EDUCATION there appears a statement by Professor in order to reach more teachers of qualitatiue analysis. William T. Hall criticizing the conclusions of an Somefewyears ago, the top ranking student ofhisclass article on the absence of coprecipitation of harium atWestPointwassuhsequentlysentto theMassachusetts ion with the sulfides of Group I1 [J.Phys. Chem., 47, Institute of Technology to he trained as a chemical 264 (1943)l. Professor Hall points out several ways engineer. At the Institute, as a t West Point, he in which harium may he lost during the course of a proved the brightest and most painstaking man of qualitative analysis. However, the conclusions of the the class. He was given a solution to analyze which article in question are: "No coprecipitation (adsorp- contained 3 mg. of Ba++ and somewhat more As04'. tion, occlusion, or postprecipitation) of barium ion When he made his report, the instructor told him that takes place wheu the sulfides of the copper and tin he had missed Ba++. I t was easy to prove by direct groups are precipitated in the presence of ammonium test that the Ba++ was present but this did not satisfy ion," and "No coprecipitation of barium ion takes the lieutenant and he worked assiduously for over 10 place wheu the chlorides of the silver group are pre- hours trying to find where he lost the Ba++ and the cipitated." While Professor Hall's statements re- instructor finally came to me for advice. I talked garding the loss of barium by the formation of known with the student and found that he had followed very insoluble barium compounds may he correct, they carefully the excellent text of the late Arthur A. Noyes have no significance with respect to the conclusions who recommended making the solution acidic with quoted above. HNOa before introducing H2Sand subsequently heating The conclusions of the article under discussion do not to see if there was any indication of arsenate's being state that barium is never precipitated with the present. Subsequent study showed that every time copper and tin groups. Professor Hall has read this he did this, the Ba++ was precipitated together with meaning into them. Therefore, his statement about sulfides of the copper-tin group. Now the question is the conclusions is incorrect. whether the precipitation of Ba++ in Group I1 of the LEOLEHRMAN qualitative scheme of systematic analysis can he attributed to what is called coprecipitation. Tne COLLBCE OF THE CITYOF NEWYORK NEW YORKCITY The word in question is used by Dr. I. M. Kolthoff and his students and I find in my copy of KolthoffSandell on page 103, "The contamination of a preTo the Editor: In the letter of Professor Leo Lehrman he protests cipitate by substances that are normally soluble under that I cannot understand the English language as he the conditions of the precipitation is called coprecipiwrites it. The trouble all lies with respect to the word tation." Since Ba+f does not normally precipitate coprecipitation which is a recently coined word and is with the copper-tin group and will not be precipitated not in my dictionaries. by H2S unless the solution contains arsenic (and the The paper of Lehrman and Mandel was published in solution is evaporated with HCI until fairly strong an excellent journal, which, unfortunately, few analyti- aqua regia is formed), this would appear to be a case of cal chemists read regularly. It does not, as a rule, coprecipitation. To my mind, i t is just as much copreconcern itself with the ordinary problems of the cipitation as when Ba++ is precipitated by NH40H analytical chemist nor with those of instructors in when Cr+++ is present. In this case barium chromite elementary qualitative analysis. It seemed, however, is probably formed which, to be sure, is not a "suhto the editors of Chemical AGstracts that the paper was stance normally soluble under the conditions of the really concerned with qualitative analysis so it was precipitation." A careful reading of the chapter in sent to me to abstract. The impression that the Kolthoff-Sandell makes me think that my esteemed paper made upon me, which may be entirely wrong, friend Kolthoff would probably agree with Professor as I haven't seen it for about eight months, was that it Lehrman in the case under discussion. Much as I was written as the first part of a careful research origi- admire Kolthoff I do not always agree with him. Thus nated with the purpose of finding how and where be- he distinguishes between iodometry and ipdimetry ginners often miss Ba++ in qualitative analysis. The although he seems to be one o'f a very few chemists conclusion, as I read it, seemed to show that there was making such a distinction. no danger of losing Ba++ by precipitation of the silver Professor Lehrman's statement that "no precipitation group with CI- (to which I agree) or of the copper-tin of barium ion takes place when the sulfides of the group with HzS. This last is contrary to my own ex- copper and tin groups are precipitated in the presence perience. Now i t is a rigid rule of Chemical Abstracts of ammonium ion" appears to be correct because in that abstractors must give the content of papers with- the heating and evaporating of a nitric acid solution out any criticism or correction on the part of the ah- the NH,+ undoubtedly is oxidized to N1 before the stractor and that is my excuse for writing my note Aspss or As& is dissolved.
I am glad that Professor Lehmlan voiced his objection to my ignorance with respect to the English language for it gives me opportunity to emphasize once more that there is occasional danger of precipitation of Ba++ with the sulfides of the Cu-Sn group. ROCHESTER. M ~ s s ~ c ~ u s ~ r r s WILL- T. HALL Nan-stoichiometric Equations
To the Editor: Under this title, Otto F. Steinbach' points out that an equation may balance algebraically and yet not represent the truth. This fact is one that probably has bothered every student of chemistry a t one time or another, and every teacher frequently has students ask why a certain equation is marked wrong in spite of the fact that i t balances. Chemical equations are commonly written for two purposes: (1) to show the probable reaction products, and (2) to show the quantities of materials that enter into reaction. If an equation does not fulfill the latter requirement it can be termed "nonstoichiometric." The reactions that are used in quantitative analysis must be stoichiometric or they have no value. With most reactions of organic chemistry the equations are idealistic and nonstoichiometric, but nevertheless have value. Recognizing this fact, it is the common practice of all teachers of organic chemistry to pay little attention, for example, to balancing equations of oxidation and reduction. Thus if it is a matter of oxidizing an alcohol to an aldehyde or an acid with permanganate, the instructor merely writes the graphic formula of the alcohol and shows that the element oxygen, or the radical OH, can replace hydrogen in the formula to give the desired product. This practice is, of course, related to the old dualistic conception of oxidation by permanganate which was based upon the fact that the acid anhydride, Mn20,, was unstable and easily lost oxygen. The formula of permanganate was commonly written, in the good old days, K20.Mnz0, rather than KMnO.. Today it is common practice to consider the oxidation by permanganate as related to the reduction of manganese from a valence of seven to a valence of two (or four in some cases). From the valence-change standpoint, few students of organic chemistry can balance equations for the oxidation of organic compounds, and there isn't much use in doing it unless the reaction is stoichiometric. I n such cases the valence-change method is easy to apply if one is willing to admit-as a few organic chemists are-that the valence bonds of the carbon atom may be either positive or negative, as illustrated by the compounds CCla and CHI. The non-stoichiometric equations cited by Dr. Steinbach are all equations of oxidation and reduction involving hydrogen peroxide, sodium hypobromite, chloric acid, potassium hypochlorite, and potas. -
sium chlorate. These compounds, for the most part, are not very stable and, as he points out, may give rise to more than one reduction product. In the case of hydrogen peroxide, however, the statement that the reaction between it and potassium permanganate is a good example of variable coefficients is absolutely wrong and cannot pass unchallenged. The reaction between KMuO, and HzOz in the presence of H2S04 has been known to be stoichiometric for many, many years and was studied by C. F. Schonbein (1799-1868) who found that the reaction as expressed in our present nomenclature, is In accordance with Schonbein's theory, which is not bad even today, half of the oxygen evolved comes from the KMn04 and half from the HzOZ. All the other equations which Dr. Steinbach suggests are obviously incorrect. However, since he is bothered by the fact that he can write a t least nine different equations to express such a reaction and many college freshmen and high-school students are troubled in the same way, just a word or two may be said in explanation. Hydrogen peroxide in dilute acidic solutions is quite stable. I t slowly decomposes, however, into HzO and 0,. During the course of a titration of a solution of hydrogen peroxide with permanganate there is no appreciable decomposition of this kind and the volume of oxygen evolved as compared with the volume of permanganate reduced is perfectly definite. The hydrogen peroxide content of an aqueous solution can be determined accurately either by measuring the volume of permanganate solution required to oxidize the peroxide or the volume of the oxygen evolved. In terms of the old ideas of oxidation, the correct equation may be written 5Hz0z 50 = 5H10 50s
+
+
and the equivalent weight of HzOz(or the weight equal to an oxygen atom of oxidizer) is the molecular weight divided by two. I can write equations such as these:
but they are wrong and have no value because they do not correspond to the equivalent weight of hydrogen peroxide as has been determined experimentally. In other words, the eight incorrect equations written by Dr. Steinbach do not represent the empirical value of H202as a reducing agent; it is equivalent in reducing power to two atoms of hydrogen. This has been proved experimentally, but the astounding fact is that by adding oxygen to the water molecule we get H202which will, (1) by adding a little sodium hydroxide and heating, decompose into water and oxygen, (2) act as an efficient oxidizing agent, or (3) act as an efficient stoichion.etric reducing agent with permanganate and with oxides such as Rhoz.
