THE EQUILIBRIUM RELATIONS IN A WATER SOLUTION OF CUPRIC BROMIDE
The use of cupric bromide solutions to demonstrate chemical equilibrium and the common-ion effect is criticized by showing that i t will not consistently do this. The actual results obtained are discussed and a n explanution based upon the Werner theory i s advanced. This explanation i s for the instructor only. It i s suggested that our texts should choose a better system to demonstrate the processes mentioned above. Such e system i s ferric thiocyanate dissolved, or formed, in water.
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The use of an aqueous solution of cupric bromide to demonstrate chemical equilibrium to the student is fraught with difficulties unless the instructor is careful. It is, of course, common knowledge that when water is added to brown cupric bromide crystals a green color results and changes to blue as more water is added. This color change is attributed to the ionization of CuBra into Cu++ and 2Br-, in two of the better books on inorganic chemistry, ( I ) , (2). To quote Smith (1)directly: CuBn (brown) % Cu++(blue)
+ 2Brr.
The student is now told that if the concentration of bromide ions is increased by the addition of potassium bromide, i t will cause the equilibrium to be shifted to the left with the formation of more undissociated molecules of cupric bromide and that this increae in the number of undissociated molecules of cupric bromide produces the brown color. That can be readily shown the student by adding potassium bromide crystals to the solution of cupric bromide. So far all is well but the equilibrium should also be shifted to the left by the addition of copper ions (from some soluble copper salt) to the solution of cupric bromide. However, the instructor will be embarrassed if he tries this because he can saturate the cupric bromide solution with cupric chloride and only obtain a darker shade of green: The same is true if he uses other soluble copper salts such as cupric acetate. He may mumble something about cupric chloride not being as soluble as potassium bromide and so not furnishing enough copper ions to shift the equilibrium all the way to the left but if some wide-awake student adds equivalent amounts of copper ion and bromide to two separate, but identical, portions of cupric bromide solution, he will observe a great difference in effects. It is the purpose of this paper to offer an explanation of this anomaly in terms of experimental work done by the author as well as other investigators. This explanation presumes a reading knowledge of Werner's theory (3) and is not advisable for elementary students of chemistry but is offered for two reasons: first, to give the instructor a plausible explanation based upon experimental facts; secondly, to suggest that since the system 1457
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is not so simple as i t first appears it is a poor illustration to use in demonstrating to elementary students such general processes as chemical equilibrium, ionic equilibrium, and the common-ion effect. In the first place, it is erroneous to write as a symbol for cupric ion in water solutions simply, Cu++, because most chemists agree that anhydrous cupric ion is not hlue. It acquires its characteristic hlue color after i t is dissolved in water and students of Werner's hypothesis know that the more probable composition is represented by:
To rewrite the equation for the solution of cupric bromide in water according to the Werner system we would have something like this:*
DlAQUODIBROMOCUPRIIM
TRIAQUO~~ONOBROMOCUPROBROMIDE
TBTRAAQUOCWROBROMIDE
It is readily seen that the addition of water to the original cupric salt results in the replacement of bromo groups by water (aquo) groups, and that the addition of hromide ion to the tetraaquocuprohromide results in the replacement of aquo groups by bromide ion to form bromo groups. It should be evident from this that any source of bromide ion will produce a brown color if enough can be obtaitied to change the aquo complexes to E the diaquodibromocuprum salt. So much for the bromide ion; now for the copper ion. If the demonstrator chooses to add cupric acetate to his solution of cupric bromide he gets a darker green color but not so dark as when he adds cupric chloride to the solution of cupric hromide, even though he add equivalent amounts of each salt. This can be shown even more strikingly by the following experiment. Take ten cubic centimeters of a solution of cupric hro'mide in water that has snfficientcupric bromide dissolved to give a dark green color. Divide this into two equal portions and to one add one gram of sodium acetate; then compare the two colors. The one to which the sodium acetate was added is now lighter in color. If this experiment is repeated with cupric acetate instead of sodium acetate the same result is obtained but is less striking. But this is seemingly contrary to all theory because we would expect the addition of any cupric salt to drive the equilibrium back in the direction of the brown color (to the left). It has been shown by Gustavson (4) and Stiasny (5) that in solutions of basic chromium salts the substitution of one anion for another directly attached to the central chromium ion follows a regular order. The author * For the nomenclatureof Werner compounds see (3).
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also found this to be true for solutions of basic aluminum salts (6). The order is: NO1- < Br- < SO4- < CHaCOO- < C20,The position of H1O in this series is difficult to place because all work was done in water solutions but it is certain that they are replaceable and can also replace other groups because so many inorganic isomers can be prepared by varying the conditions; e. g., CrC1&H20 exists in three isomeric forms but all have the same empirical formula. They differ in the number of aquo and chloro groups directly attached to the central chromium ion. If the reader will refer to the anion series just given above, he will see that acetate ion is more powerful than bromide ion in forming undissociated complex ions because it will break up the bromocupro complex and form an acetocupro complex instead. Now it should be apparent why acetate ion will prevent the formation of diaquodibromocuprum and so the brown color.
I+
Br-
TRIAQUOMONOBROMOCUPROBROMlDE
+ Br
TRIAQ~~MONOACETOCUPROBROMIDE
Now suppose a large amount of cupric chloride, or cupric acetate, is added to a solution of cupric bromide. These crystals of the cupric salt will act like cupric bromide when it came in contact with water, that is, aquo groups will tend to replace chloro groups in the salt. This removal of water from the bulk of the solution into the new cupric complex ion just as effectively removes water from bromide io~las though the water were boiled off as steam. Therefore, the concentration of bromide ion is increased even though its actual amount is unchanged. This increase in bromide-ion concentration and decrease in aquo groups tends to favor the replacement of aquo groups by bromo groups in the triaquomonobromocuprobromide and thus to add more brown to the solution. This addition of a brownish tinge to the green solution naturally makes the shade of green appear darker to the observer. But please note that although the removal of water by the cupric chloride favors the formation of more diaquodibromocuprum, a t the same time this tendency is offset to some extent by the return of aquo,groups to the bulk of the solution as they are pushed out by the bromide ions from their place around the central cupric ion, and this keeps the equilibrium from ever going very far to the left. Therefore, no matter how much cupric chloride one adds, he never gets a definitely brown . coloration in the solution. Now to cover one more point. Mellor (7) and Hopkins (8)note that the color of cupric bromide solutions may be due to the formation of some CuBr&-ions. This is not unlikely for many salts of copper are reported by Werner (9) in which the copper is part of a complex anion; e. g., Na2Cu-
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(CN)4, NaCuCbHzO, N%Cu(CzOa)z, and others. This means that sufficient anion would be present to replace all of the aqua groups accordingto the scheme:
P O T A S STETRAI~ BROMOCUPROATE
If the brown color really is due to potassium tetrabromocuproate the reasoning given previously in this paper is unaffected for, as you will see, it only extends the system in equilibrium one step farther to the left:
It should now be apparent that the simple system pictured as CuBrz is really a much more complicated affair than a t first appears. It is not a good example for the purposes usually imposed upon it, namely, to demonstrate the common-ion effectand chemical equilibrium. And why not? Because the addition of iodide ion and cupric ion in equivalent quantities does not produce the results that the theory predicts. Therefore we must admit to our advanced students that we tricked them when they were freshmen and this tends to suggest to them that they are still being tricked. Certainly this will not increase their respect for chemistry as a science and for the teacher of chemistry. The substitution of some system like th& of ferric thiocyanate in water lends itself more successfully to demonstrating chemical equilibrium, ionic equilibrium, and common-ion effect. It has the additional advantage of causing fewer pangs of conscience to the teacher who likes to play fair with his students. Literature Cited
i Cu++ + 2Br-
(I) SMITH,"Inorganic Chemistry," revised by KENDALL, The Century Co., New York City. 1926. 1030 pn. .. (20 IIOPKIXS, 'GenPrill Ch~tnistr).,''I). C. Heath 8; Co New York City. 1930, X 7 pp. (. J.) WIXNER."NCW1de;u on Inorganic Chen.irrry." Longn~ani,(;rccn and Cu Sew York City, 1911,268 pp: (4) GUSTA~SON, J. Am. Leathn Chem. Assoc., ZZ,68 (Mar., 1927).
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(5) STIASNY,Collegium, No. 694, 73 (1928). (6) WHITEHEAD, T. H., "Ion Interchanges in Aluminum Orychloride Hydrasols," Dissertation, Columbia University, New York, 1930.
(7) MELLOR,"A Comprehensive Treatise on Inorganic and Theoretical Chemistry," Vol. 111, p. 196, Longmans, Green and Co., Few York City, 1923,927 pp. (8) HOPSUNS."General Chernism," see ref. (Z),footnote p. 237.
(9) WERNER,"Neuere Anschauungen auf dem Gebiete der anorganischen Chemie," Friedr. Vieweg & Son, Akt.-Ges., Braunschweig. 1923, revised by P. PEE~PFER, 444 PP.
