The Colors of the Copper Salts. - The Journal of Physical Chemistry

THE COLORS OF THE COPPER SALTS WILDER D. BANCROFT

AND

HARRIS WHITE ROGERS

Department of Chemistrv, Cornell University, Ithaca, New York Received August 21, 1933 INTRODUCTION

In 1892 Ostwald (1) pointed out that according to the electrolytic dissociation theory the absorption spectrum of a dilute salt solution must depend exclusively on the absorption spectrum of the solvent, the cation, and the anion. From this it follows that we shall have identical absorption spectra for dilute solutions of different salts with the same colored ion. In accordance with this, dilute solutions of copper sulfate, nitrate, chloride, and bromide should have the same blue color, and it is an experimental fact that they do. The undissociated salt may have a different color and Ostwald called attention to the fact that anhydrous cupric chloride is yellowish-brown, cupric bromide blackish-violet and cupric sulfate gray. Since nobody was especially interested in those days in hydrated ions, the behavior of the copper salts was taken to mean that the copper ion is blue. We see now that that is not necessarily true. All that the experiments really show is that the copper ion common to, and occurring in, those dilute solutions is blue. Nobody paid any attention a t that t>imeto the earlier observation by Vogel (2) that “copper sulphate is one of the few substances which shows the same absorption spectrum in the solid state and in solution. It absorbs the red very strongly up to wave length 620; from there on the absorption decreases rapidly and ends in the yellow-green. Green, bright blue, and dark blue are practically not absorbed, but there is some absorption in the violet.” Since 1908 Hantzsch (3) has taken the ground that there is no necessary relation between color and ionization. He states that colored substances are those which contain an element or group capable of forming a complex of the Werner type. If the complex is saturated completely, all the coordination places being filled, then the color is constant, regardless of the conditions surrounding the complex. Dissociation, salt formation, solution, change of temperature or of solvent, will have no effect on the color. Conversely, if a given color persists through t’hechanging of such conditions, we may regard the saturated complex as persisting through the changes. For those complex states which may not be entirely saturated but become so through a change in conditions, we shall expect an accompanying change 1061

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WILDER D. BANCROFT AND HARRIS WHITB ROGERS

in color. If, by any dehydration or substitution, the complex is altered, this will also result in a change of color. Anhydrous salts, which are unsaturated but which become saturated by the addition of water or ammonia molecules or other groups, will change color until complete saturation is reached. After this point there will be no further change in color. If a saturated complex cannot be isolated as a solid or is stable only in presence of a large excess of the solvent, there may be intermediate complexes which will explain the color of certain solutions. From a study of the chloroplatinate salts of the alkalies, Hantzsch concluded that neither change of dissociation nor change of temperature had any effect on the color. Change of solvent had only a slight effect. An investigation of chromate solutions (4)showed that the acid and its potassium salt had identical absorption, which was independent of dilution and of temperature change. Addition of sulfuric acid up to 10 normal had only a slight effect, accounted for by the formation of some H&r3010. Solutions of chromic acid in water, alkalies, and methyl alcohol were optically identical in all dilutions examined, except in the water solutions where the slight variation was accounted for by some dichromate ion. The dichromate and chromate solutions were of course unlike. Hantzsch states that the color group of the acid solution is the complex Crz07,and of the ,chromate solution Cr04, irrespective of whether these groups are joined to hydrogen or to an alkali metal. The degree of electrolytic dissociation has no effect on the color because the color of the ion is the same as that of the undissociated complex. Hantzsch points out that, years before, Sabatier (5) had found that “the absorption exercised by potassium dichromate dissolved or solid is practically the same as that produced by the chromic acid in the salt.” In a study of alkali permanganates Hantzsch and Clark (6) confirmed these assumptions. Change of temperature or of alkali metal had no effect and the effect of adding sulfuric acid was but small. There was no effect due to change of concentration and the slight effect produced by change of solvent was due to a partial reduction of the permanganates. Hantzsch found that the absorption of solid copper sulfate pentahydrate was very similar to that of its aqueous solution, thus confirming to that extent the earlier observation by Vogel. Hantzsch concluded that the color of both ions is due to the complex, Cu.4H20. He believed that this group is also present in dilute solutions of cupric chloride, while the green concentrated solutions contain the unsaturated complex, Cu. 2H20, which is also present in the solid hydrated salt. Hantzsch ascribes the intense blue color of the copper ammines to the saturated complex, Cu 4NH3. Investigation of the copper ammines by Hantzsch and Robertson (7) showed the color to be independent of the anions. The same was true for pyridine solutions, but the complex is

.

COLORS OF COPPER SALTS

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apparently less stable in the pyridine solutions, as more is required to produce the blue color, Hantzsch suggests the presence of such groups as Cu.3Py.Hz0 or Cu.2Py.2Hz0, as being likely. In the case of both solid and solution the saturated complex is the chromophoric group, such as Cu.4Hz0 or Cu.4NH3. This chromophoric group determines the color regardless of the degree of electrolytic dissociation of the compound in which it appears. Some of Hantzsch’s conclusions had been reached independently and somewhat earlier by Bjerrum (8). “I will next emphasize that it is permissible to assume that the ions of an electrolyte have the same color whether in the free or the bound state, so long as no new complex is formed. This assumption was put forward by me as probable in the spring of 1907 (9). I came to this conclusion by combining the optical data now being presented with the newer views on the constitution of inorganic salts. In the fall of 1907 Hantzsch formulated the same generalization and confirmed it by new experimental data. Just recently some studies on the color of the chromic salts has given me a very good confirmation of this law. I propose to discuss these investigations more fully elsewhere and I will only mention the following here. All normal hexaquochromic salts have exactly the same color, the color of the hexaquochromic ion, even at concentrations a t which the salts are only very slightly dissociated into ions.” A possible explanation of the practical identity of color of crystallized copper sulfate pentahydrate and of copper sulfate solutions is given by’ the theory of complete dissociation. If copper sulfate is 100 per cent ionized under all conditions, there is no reason for any change in color with changing concentration. The difficulty with this is that the monohydrate, CuSOd.HzO, is green and the anhydrous copper sulfate is colorless. From this it follows that either anhydrous copper sulfate is not dissociated at all or that the copper ion is not blue. Roscoe and Schorlemmer (10) state that when copper sulfate pentahydrate or trihydrate is heated for some time to 100°C., the monohydrate, CuS04.Hz0, remains as a bluish-white powder. This is not correct. It was known at least a century ago that monohydrated copper sulfate is green. Graham (11) says that “the sulphate of copper with one atom [molecule] of water was also obtained in a crystallized state by Dr. Thomson and called by him green sulphate of copper.” Muller (12) found that the monohydrate is greenish, and we have confirmed his results. Muller looks upon the green color of the monohydrate and the lack of color of the anhydrous sulfate as quite inexplicable. SPECIAL ASSUMPTIONS

We adopt the general theory of Hantzsch; but the facts now at our disposal make necessary some minor changes in the wording. For instance,

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WILDER D. BANCROFT AND HARRIS WHITE ROGERS

the experiments of Dewar (13) show that apparently all colored substances become paler with falling temperature. “The optical properties of bodies cooled to the temperature of boiling liquid air will require long and patient investigation. An interesting fact, easily observed, is the marked change in colour of various bodies. Thus, for instance, oxide, sulphide, iodide of mercury, bichromate of potash, all become yellow or orange; while nitrate of uranium and the double chloride of platinum and ammonium become white. Chromic acid, dilute solutions of iodine in alcohol, strong solutions of ferric chloride and other coloured solutions become greatly changed. Such facts are sufficient to prove that the specific absorption of many substances undergoes great changes at the temperature of - 190°C.” When the change with falling temperature involves an apparent change from yellow to white, as is the case with heated zinc oxide, this could easily be considered as contradicting the wording of Hantzsch’s theory even though it does not violate the real principle, because the change is only or chiefly one of intensity. In the discussion of the colors of the cupric salts, so far as affected by water and ammonia, we make the following assumptions:’ 1. The color of the chromophoric group does not vary appreciably with varying degree of ionization. 2. The color of the chromophoric group becomes paler with falling temperature. 3. The anhydrous cupric ion is not blue. It is probably colorless; but the possibility of its being red is not yet excluded. 4. Cupric copper with one or two molecules of water attached is green. 5. Cupric copper with three or more molecules of water attached is blue. 6. The NH, group has an effect similar to that of a water molecule in the copper complex, though the actual blue is different. 7. Copper oxide is blue and not black. 8. Some double salts have an effect which cannot be predicted a t present, anhydrous potassium copper sulfate being blue. I

GREEN AND BLUE SALTS

Anhydrous cupric sulfate, fluoride, selenate, and perrhenate are colorless. Anhydrous cupric nitrate has been reported by Ditte (14) as nearly colorless, though with a slight greenish tint. It is probable that this discoloration is due to incomplete drying or to slight decomposition, and that the pure anhydrous nitrate is colorless like the other salts cited. If these salts are ionized appreciably, then the anhydrous cupric ion is colorless. If the salts are not ionized appreciably, cupric copper in this form is colorless. We thought, at first, that Hantzsch’s theory required that anhydrous cupric ion should be colorless, but there are some anhydrous copper double


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salts which are red. It would follow equally wel1,from this that anhydrous cupric ion is red. Until we know whether the red or the colorless salts contain the simple cupric group, the question of the color of the anhydrous cupric ion must remain undecided. We shall discuss the question more in detail when considering the brown and red salts. The important thing from our point of view is that the blue ion in hydrated copper sulfate and nitrate cannot be the anhydrous cupric ion. TABLE 1 Green salts

TABLE 2 Blue salts Cu(N03)~.4NHa CUL.5"s CU(NOa)z.6NH, CuIz.6NH3 CuIz.4NHa Hz0 CU(N0z)z 4NH3 Cu(I03)z.4NH3 CuW04.4NHa CuM004.2NHs. Hz0 c U ( I o 3 ) z . 4NHa *HzO CuSiFe.6Hz0 CU(CNS), .4NH3 CU(CNS) z ' 5NH3 CuClz. 4NH3 CuClz.6NHa Cu(CNS)z.GNHa CU(C103)z.4HzO cUso1.MzS04. 6Hz0* CU(CIOa)z.6Hz0 Cu(Cz04) .5NH3 CU(CIOa)2.4NH3 CuCzOr. 2"s. HzO CU(ClOS)2.6NH3 CU(HCOZ)Z.~NH~ CU(HC0z)z SZCeHaN. HzO CuBrz.5NH3 CuBrZ.6NH3 CU(CHaCO:) 2.4"s C U ( B ~ O ~ ) ~ . ~ N H ~CU(CHBCOZ)L. 4CsHsN 3

* M is K, Rb, Cs, NH4. In table 1 is given a list of some cupric salts which are green when crystallizing with one or two molecules of water, ammonia, hydroxylamine, or pyridine. The list is not exhaustive. The free acids and the sodium salts of the acids in table 3 are colorless, whether hydrated or not, so that there is no reason to believe that the water or ammonia influences the color by being attached to the anion. Where the literature was conflicting as to the color, the salts have been made. Unless the color was perfectly obvious, the crystals were immersed in a

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liquid of approximately the same index of refraction t o eliminate diffuse reflection. The color wa’s then checked by several observers. In table 2 is given a list of some of the blue cupric salts with three or more molecules of water, ammonia, hydroxylamine, or pyridine. Chuard (15) claims to have obtained a blue cupric chloride with three molecules of water by working below zero. This would be a welcome confirmation of the theory, but no one else has yet been able t o isolate this hydrate. A conspicuous exception to the assumption that cupric salts with three or more molecules of water are blue is CuClz.CdC12.4H20,a green salt. We know that cadmium salts in general tend to be hydrated, so that it is probable that part of the water in the crystal is attached to the cadmium, and that the formula should be written CuClz.2Hz0.CdClz. 2Hz0. When the red crystal, CuClz.LiC1.2Hz0,is discussed later, it will be shown that some of the water is undoubtedly attached to the lithium. BROWN AND RED SALTS

Hantzsch’s theory enables us to draw some interesting conclusions in regard to some of the brown and red cupric salts. Anhydrous cupric chloride is a yellow-brown; but Bancroft and Weiser (16) reported that its vapor is violet-red. This has been confirmed by volatilizing the anhydrous chloride rapidly either by itself or in an atmosphere of chlorine to prevent, dissociation. In both cases the vapors were distinctly violet-red and condensed to the brown solid. From Hantzsch’s theory it follows that the vapor must have a different constitution from the solid. Ward (17) has already shown that cupric chloride and cupric bromide are abnormal, because addition of an excess of concentrated sulfuric acid precipitates anhydrous yellow-brown cupric chloride and anhydrous black cupric bromide, respectively, instead of converting these salts into sulfates. Hantzsch and Carlson (18) state quite definitely that cupric chloride is a pseudo salt. Qualitatively, solid anhydrous cupric chloride is similar in color to one constituent of a solution of cupric chloride in aqueous hydrochloric acid. Donnan and Bassett (19) showed that the yellow-brown color in these solutions is due to an anion containing copper, CuCb or something of that type. Kohlschutter (20) extended the work of Donnan and Bassett, and found that a brown color moved to the anode in concentrated solutions of cupric chloride. In a dilute solution all the copper went to the cat,hode, but increase of concentration caused some and then more copper to go t o the anode. On these facts we conclude that anhydrous cupric chloride is really the copper salt of a chlorocupric acid, and that the formula should be written Cu(CuCl4) or Cu(CuCl&. On the other hand, some of the anhydrous double salts are red, such as CuCl2.NH4C1,CuClz.CsC1, and CuC12.KC1. In these salts we undoubtedly have the cupric chloride having the same chromophoric group as the

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vapor. It is probable that an x-ray analysis would show a distinct difference between the copper atoms in anhydrous cupric sulfate and those in anhydrous potassium cupric chloride. This is the more plausible because Hendricks and Dickinson (21) have examined some of the hydrated double chlorides and have found evidence indicating a CuC12.2H20group in the ammonium, potassium, and rubidium salts of this type. This confirms our application of Hantzsch’s theory. It is a pity that Hendricks and Dickinson did not also study the salt which we write CuC12.2H20.CdC12.2H20. The hydrated double salt of lithium and copper chlorides has the composition, CuC12.LiC1.2Hz0, and is red. Consequently the water must be attached to the lithium chloride and not to the cupric chloride. This is additional evidence for the formula which we have given for hydrated copper cadmium chloride and should be confirmed by an x-ray study. Engel (22) and Sabatier (23) have made a red compound of cupric chloride, hydrogen chloride, and water. Engel wrote the formula CuClz.HC1. 3H20and considered that the compound was derived from the hypothetical trihydrate of cupric chloride. Sabatier wrote the formula CuClz 2HC15H20 and .did not care how it came to pass. Since neither man had a method of analysis that was worth anything, nobody knows what the composition really was; but the color was red-Sabatier called it a hyacinth red-and changed reversibly from red t o green and back again as hydrogen chloride (and water) evaporated off and was put back again. With the idea of testing Hantzsch’s conclusions, anhydrous cupric chloride was added to molten potassium chloride and to molten sodium chloride. In both cases the melt was a t once colored red. In time the melt decolorized, presumably owing to volatilization of the cupric chloride. This agrees with what Hantzsch would predict if the cupric chloride vapor is normal, and means apparently that anhydrous cupric ion is red. This contradicts the conclusion drawn from the colorless anhydrous salt. Either copper sulfate, fluoride, selenate, nitrate, and perrhenate form colorless complexes in which copper is not the basic radical, or copper is not the basic radical in cupric chloride vapor. Until some independent evidence is forthcoming, it is impossible to determine which conclusion is right. On general principles it seems probable that cupric sulfate and cupric nitrate are normal and not pseudo salts. There is no reason why the vapor of cupric chloride should not be abnormal to some extent. We know that it condenses to a pseudo salt. Sabatier (24) found that cupric bromide is yellowish-red in absolute alcohol and is purple in concentrated solutions of hydrobromic acid, potassium bromide, sodium bromide, lithium bromide, and calcium bromide. There are anhydrous double salts which are red, but apparently no single salts. According to Bodtker (25) anhydrous cupric chloride in absolute alcohol is brown, although the salt which crystallizes is green and has the composition CuClz.2CH3CHz0H.

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The organic cupric compounds are blue, green, or brown, but apparently never red. Copper acetylacetonate (26) is cobalt-blue when anhydrous and sky-blue when crystallizing with two molecules of water, Copper acetoacetic ethyl ester (27) is green, while the basic methylate is blue. Wislicenus wrote the formulas: CHs CHa c ~ o ~ c u ~ o ~ c CH

CH

COzCzHs

COzCgH,

Copper acetoacetic ester (green)

CHa C . 0 .CU*O* CHa

CH COzCzHs Basic methylate (blue)

The blue form is stable in methyl alcohol and the green form in benzene. The copper derivative of benzoylcamphor (28) is greenish-yellow in alcohol and greenish-brown in chloroform. Since the red salts occur only under conditions where one would expect pseudo salts, we feel that it is probable that the violet-red of the cupric chloride vapor is abnormal and that the anhydrous cupric ion is colorless. In fused ammonium acetate a t 90°C. cupric chloride, sulfate, nitrate, chlorate, acetate, and oxide are blue, undoubtedly because of the formation of the C U ( N H ~group. )~ In fused ammonium sulfate a t 160°C. they are all green, undoubtedly because of the formation of the C U ( N H ~or )~ Cu(NH8) group. In fused ammonium nitrate at 160°C. all were green except the acetate and oxide, which were blue. All the salts gave a black color in a fused mixture of potassium and sodium nitrates a t 300°C. This is apparently due to the formation of undissolved, black, cupric oxide. SOLUTIONS

Addition of considerable quantities of sulfuric acid to a copper sulfate solution causes the color to shift from blue to green. This is what should happen if the chromophoric group was being dehydrated. Gladstone (29) says that “even where different salts of a base have the same colour, the same amount of the base does not give the same intensity of colour. Thus if equal portions of oxide of copper be dissolved respectively in acetic, hydrochloric, nitric, and sulphuric acids, and equally diluted, the acetate will be found to be far deeper in colour than the sulphate, and this again far deeper than the chloride. On being convertedinto ammoniacal salts, these four approximate more nearly, but are still far from identical in colour.” The acetate did not seem to be entirely trustworthy because of the possible presence of basic salt. Dr. F. H. Getman was kind enough to run spectrophotometric determinations on copper nitrate and copper sulfate

COLORS OF. COPPER SALTS

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solutions. He confirmed Gladstone’s observations, finding that solutions of copper nitrate are a deeper blue than the equivalent solutions of copper sulfate. This does not apply of course to very dilute solutions. Since copper nitrate crystallizes with six molecules of water instead of with five, the natural hypothesis is that the copper in undissociated copper nitrate is more highly hydrated than the copper in the corresponding copper sulfate solutions. As yet there is no independent proof of this. We hoped to find in the literature that the solid solutions of copper and zinc sulfates cont,aining seven molecules of water per molecule of salt were much bluer per unit of copper in a given thickness than the corresponding solid solutions with five molecules of water. The literature seems to destroy this hypothesis. Friend (30) says that copper sulfate forms pale blue monoclinic crystals of the series, CuS04.7Hz0.ZnSOa.7Hz0, but dark blue triclinic crystals of the series of solid solutions, CuS04 5Hz0.ZnS04 5Hz0. This is so inherently improbable that we suspect that proper allowance has not been made for the low copper concentration. This point will be checked as soon as possible. In the field of non-aqueous solutions the evidence is rather contradictory. Guthrie (31) heated glycerol with cupric sulfate, obtaining an emeraldgreen solution. This is what one would expect if there were dehydration. Griin and Bockisch (32) say that copper sulfate pentahydrate with dry glycerol gives a dark blue solution. On precipitation with alcohol they obtained a blue oil which could be dried to a blue glass. They write the formula for this :

-

. . .HO.CHz

bU(

. . .HO .CH.CHzOH

)

SOd*H%O]

8

On repeating Guthrie’s experiments we got a green solution as he did. We were not using absolutely anhydrous glycerol any more than Guthrie was. It is possible, though not proved, that anhydrous glycerol gives a blue solution and glycerol with small amounts of water gives a green one. As has been stated, cupric chloride gives a brown solution in absolute alcohol (33) and a green diethylate crystallizes from it. Ley (34) found cupric chloride to be green in alcohol and blue in pyridine. Mason and Mathews (35) reported both green and blue solutions in pyridine just as in water. They suggest that the green is due to cuprous chloride, as this salt is known to be green in pyridine; but this does not yet seem to be necessary. Sammis (36) found that both cupric acetate and cupric formate are blue in pyridine. Mathews and Benger (37) found only two stable solvates of cupric acetate and pyridine, a green one with one molecule of pyridine and a blue one with four molecules. Hantzsch and Robertson (38) state that all the cupric salts give the same blue in dilute

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solution in aqueous pyridine. They suggest that the blue complex may contain water as well as pyridine. MISCELLANEOUS CASES

We are still in the descriptive stage with regard to the copper salts. It seems fairly certain that the addition of one or two molecules of water to the copper atom will give a green color, while the addition of three or more will give a blue color; but we do not know why the one is green and the other blue rather than two other colors. Consequently we cannot tell as yet what other grouping will make a copper compound green, blue, or red. The alkaline cupric tartrate solution contains a blue anion, but we could not have predicted that. Covellite, cupric sulfide, is blue and so apparently is finely divided cupric oxide. In borate and silicate glasses cupric oxide is blue. Here there can be no question of hydrated ions, and the turning green of these glasses when heated too hot is due to the formation of cuprous oxide (39) and has nothing whatsoever t o do with the green of such salts as hydrated cupric chloride. Dioptase, HzCuSiOl, is green, and crysocolla, HzCuSi04.HzO, is blue. Egyptian blue (40) is crystallized CaO CuO .4sio2, and is blue. At 850900°C. it changes reversibly to a green glass. This is undoubtedly due to the partial decomposition of the cupric salt, but this has not yet been shown analytically. It is possible, though not proved, that dioptase may contain some cuprous salt. Copper pentammino metachloroantimonate, c ~ ( S b C 1 ~5NH3, ) ~ . is blue (41) as it should be, but it becomes green on losing ammonia and that could not have been predicted. The analyses of Scheele’s green, Paris green, etc., are so conflicting (42) that one cannot tell anything about them; but copper metarsenite, Cu(AsO&, is said to be colorless and the dihydrated salt, C U ( A S O ~ ) ~ . ~ is H said ~ O , to be green. That is satisfactory if true. It is impossible to make any definite statement at present about the basic salts. Basic copper sulfate, CuSOI.3Cu02H2, is green (43) and CuS04.Cu0 is said to be orange. Malachite, CuC03.Cu02H2, is green, and azurite, 2CuCOa.Cu02H2, is blue. The double sulfate of potassium and copper is a very trying case. Cowan and Ferguson (44) say that the dihydrate, KzS04-CuS04.2H20,is bluer than the hexahydrate. The anhydrous salt is blue. One does not like to postulate that potassium sulfate is equivalent to three or more molecules of water in its chromophoric properties and yet Graham (45) put forward what is practically our working hypothesis nearly one hundred years ago. “This double salt [K2S04.CuS04]retains its blue color after being fused at a red heat and cooled, and does not become white like the sulphate of cop-


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per. Indeed it appears that to be colored, the salts of the oxide of copper require the addition of some other constituent, such as saline water [what is now called water of crystallization], sulphate of potash, or ammonia. Hence if the absolute sulphate of copper could be obtained in a crystallized state [as has since been done], it would be a colorless salt.” A century of progress has brought us back to Graham. Our conclusions and those of Hantzsch would also be acceptable to Gladstone (46). “Every observation made on this salt [cupric chloride] is perfectly explicable on the supposition that the proper color of chloride of copper is brown, and that it forms hydrates which are green or blue, just as the white sulphate of copper becomes blue when hydrated. Some of the facts, too, are more easily explained on this view. “Thus it is hard to imagine that if green be the color of CuCl [CuC12], and blue that of CuO, HC1 [CuO.2HCl-HzO or CuO.2HC11, the addition of more HC1 should render it green; while it is readily conceivable that the hydrochloric acid should replace a portion of the water in the blue hydrated chloride of copper and form a green double chloride, CuCl, 2HC1 [CuCI2. xHCl*yHzO]. “If the change of color is to be taken as evidence that crystallized chloride of copper becomes, when treated with a considerable amount of water, CuO, HC1, a parity of reason should lead us to conclude that the bluishgreen crystals contain none of the yellowish-brown CuCl [CuCI2]; yet, if we suppose that these crystals actually contain the oxide, we can give no consistent account of the subsequent change of color on solution.” CONCLUSIONS

1. The cupric salts confirm Hantzsch’s theory of the identity of color of solid, solution, and vapor when the chromophoric groups are the same in the three states. Conversely, a difference in the color connotes a difference in the chromophoric groups. 2. The color of a chromophoric group is practically independent of the degree of ionization but becomes less intense with falling temperature. 3. Cupric copper with one or two molecules of water is green; with three or more molecules of water it is blue. 4. The ammonia molecule has practically the sqme effect in a copper complex as has the water molecule, though the shade of the blue is different. 5. Anhydrous cupric sulfate, fluoride, selenate, nitrate, and perrhenate are colorless, which means that the anhydrous cupric ion is colorless if these salts are normal. 6. We have confirmed the statement by Bancroft and Weiser that cupric chloride vapor is red. If the vapor is normal, the theory of Hantzsch requires that the color of the anhydrous cupric ion is red. It is probable that the vapor contains a pseudo salt.

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7. Anhydrous cupric chloride is yellow-brown, which is the color of the anion in aqueous hydrochloric acid solutions of cupric chloride and is not th& color predicted by the theory of Hantzsch. Consequently anhydrous cupric chloride must contain copper as part of the acid radical. Hantzsch has himself pointed out that anhydrous cupric chloride must be a pseudo salt. 8. The anhydrous double salts, CuCL NH4Cl, CuClz.CsC1, and CuClz.KCI, are red, and the chromophoric groups are therefore similar to that in the vapor of cupric chloride. 9. The hydrated double chloride, CuClz.LiC1~2H20,is red and must therefore have the two molecules of water attached to the lithium chloride and not to the copper. 10. The hydrated salt, C u C 1 ~ . C d C l ~ ~ 4 H is ~green 0 , and not blue. Its formula should therefore be written CuClz.2H20.CdClz.2H20. 11. By means of x-ray analysis Hendricks and Dickinson have shown the probable existence of the group, CuClz.2Hz0, in some of the green double salts. I t is very much to be desired that a systematic study should be made. 12. The anhydrous cupric ion is probably colorless, but may possibly be red. 13. Copper produces a blue color in the borate and silicate glasses and in crystallized Egyptian blue. 14. We do not know why copper is green with one or two molecules of water and blue with three or more molecules of water and consequently we can not predict what other groupings will give blues, greens, or reds. 15. The basic salts are rather hopeless for the present. 16. The anhydrous double sulfate, K2S04.CuS04,is blue. It was pointed out by Graham in 1835 that the effect of potassium sulfate is similar to that of water or ammonia. REFERENCES (1) OSTWALD: Z. physik Chem. 9, 579 (1892). (2) VOGEL:Ber. 11, 913 (1878).

(3) HANTZSCH: Ber. 41, 1216, 4328 (1908); Z. physik. Chem. 63, 367 (1908); 74, 362 (1910); 84,321 (1913); 86, 624 (1914); Z. anorg. allgem. Chem. 160,5 (1927). (4) HANTZSCH: 2. physik. Chem. 72, 362 (1910). (5) SABATIER:Compt. rend. 103, 49 (1886). (6) HANTZSCH AND CLARK:Z. physik. Chem. 63,367 (1908). (7) HANTZSCH AND ROBERTSON: Ber. 41,4328 (1908). (8) BJERRUM:Z. anorg. Chem. 63, 146 (1909). (9) BJERRUM: Kgl. Danske Videnskab. Selskabs Skrifter [71 4, 27 (1907). (IO) ROSCOEA N D SCHORLEMMER: A Treatise on Chemistry, Vol. 2, p. 451. Macmillan and Co., London (1927). (11) GRAHAM: Phil. Mag. [3] 6, 418, (1835). (12) MULLER:Ann. Physik [4] 12, 777 (1903).


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