Feb. .j,1%8
SPECTROPHOTOMETRIC INVESTIGATIONS OF COPPER PYROPHOSPHATE COMPLEXES
the results on racemization of trioxalatochromate (111) reported by Bushra and Johnson.1o Table I1 gives the values calculated for these reactions. TABLE I1 a.s-t3 a&#
kcal.
cal./deg.
AI:%m* kcal.
trans-cis Isomerization dioxalatodiaquochromate( 111) -15.3 Exchange of H2O for C1 in dichlorotetraaquochromium( 111) -16.3 Racemization of trioxalatochromate-24.5 (111)
1 7 . 5 22.1
14.4 1 9 . 3 14.9 22.3
Long" has conclusively demonstrated that the racemization of trioxalatochromate(II1) does not involve the separation of oxalate ions from the complex. I t is likely that this separation of oxalate ions does not occur in the isomerization of the dioxalatochromate(II1). It may be postulated that the reaction could occur upon collision of the trans complex ion with a properly oriented water molecule. A reaction intermediate would be formed in which the incoming water molecule would occupy a position in the chromium octahedron. The end of the oxalate thus released could be held loosely by the three water molecules, possibly through hydrogen bonding. If the central water were released the trans isomer would be recovered, but if either of the other water molecules were released by rebonding the oxalate, the cis isomer would be the product. This reaction may be represented by trans
+ H20
ki
k3
intermediate
k?
cis
+ H?O
kr
(10) E , Bushra and C. H. Johnson, J. Chem. Soc., 1937 (1939). (11) F. A. Long, THISJOURNAL, 61, 370 (1939).
[COSTRIBUTION FROM
THE
61 1
In order to explain the results obtained, both k2 and k3 would have to be very large with respect to kl, and to explain the final equilibrium being shifted very far in the direction of the cis isomer kl must be very large with respect t o kk. If the slow steps are those forming the intermediate there will be no appreciable amount of the intermediate present a t any time as i t will be reacting very rapidly to form the end-products. The negative entropy of activation obtained is in agreement with this postulate, since in addition to requiring a properly oriented water molecule, which will require the water molecule t o lose its rotational freedom, there is the formation of the additional bond in the intermediate. The exchange of a water molecule for a chloride in the dichlorotetraaquochromium(II1) has essentially the same entropy of activation and therefore can be assumed to involve the same type process. The greaterepart of the negative entropy of activation in both cases can be explained by the loss of freedom of rotation of a water molecule. In the case of the racemization of the trioxalatochromate (111) the large negative entropy of activation may involve either the requirement of two properly oriented water molecules or a transmission coeficient which is considerably less than unity. The lower heat of activation in the case of the racemization of the trioxalatochromate may be explained by the fact that this is a triply negative charged ion instead of a singly negative charged ion as in the case of the dioxalatochromate. This larger negative charge would make i t easier t o separate the bonds to give the reaction intermediate despite the fact that there are two bonds t o be separated. SALTLAKECITY1, UTAH
CHEMISTRY DEPARTMENT, THEOHIOSTATE UNIVERSITY ]
Spectrophotometric Investigation of the Complexes Formed between Copper and Pyrophosphate Ions in Aqueous Solution BY JAMES I. WATTERS AND ARNOLD AARON RECEIVED SEPTEMBER 7, 1951 A series of complex species, [Cu(P207)2l6-, [ C U P Z O ~ ][~C- ,U Z P ~ and O ~ ][~C U ~ P ~ Oexist ~ ] ~in+solutions containing the corresponding ratios of copper to pyrophosphate ions. The last two, which have not been reported before, were detected only in very dilute solutions due to the low solubility of precipitates such as Na2Cua( PpO7h.Aq. Dipyrophosphatocuprate( 11)ion is the predominant complex of copper(I1) in solutions containing a moderate excess of pyrophosphate ion, in the PH range of about 10 to 7. If the pyrophosphate ion concentration is decreased or if the solution is made weakly acidic, the equilibrium shifts to form an increasing proportion of monopyrophosphatocuprate( 11) complex. With sufficient acid added to yield a PH in the range of 5.5 to 2, a precipitate such as NanC~3(P207)~Aq may form. In more acidic solutions no precipitate forms and acidic complexes containing hydrogen ion are obtained. At an ionic strength of unity, the first and second instability and 2 X IO-O. constants for the stepwise dissociation of dipyrophosphatocuprate(I1) are 2 X
It has long been evident that copper and pyrophosphate ions form complexes in aqueous solution since the blue color of the mixture is considerably more intense than that due t o the hydrated copper ion alone. Bassett, Bedwell and Hutchinson' observed that the intensely blue color of the copper complex in solutions persisted in only one crystal, namely, the moderately soluble crystalline NaCCu(P207)2*16H20. On the basis of this observation
they concluded that a dipyrophosphatocuprate(I1) complex was present in this particular crystalline phase but that the copper ions were hydrated with variable amounts of water in other crystals containing copper and pyrophosphate ions. Rogers and Reynolds,2 in an investigation of the complexes of copper and other metals with alkali pyrophosphates by potentiometric, conductometric and amperometric titrations, found that the equivalence
(1) H. Bassett, W. L. Bedwell and J. B, Hutchinson, J . Chem. Soc.,
(2) L. B. Rogers and C. A. Reynolds, THISJOURNAL, 71, 2081 (1949).
1412 (1936).
(i12
JAMES
point varied between one and two pyrophosphate ions for each copper ion depending on the dilution. Similar experiments and conclusions were reported without data by Haldar." The complexes have been investigated polarographically by Eriksson4 and by Laitinen and Onstott.6 Both papers reported reversible waves for the complex only if the pH was below 5 . Laitinen and Onstott concluded that the complex, Cu(HP207)24-, was present in solutions having a pH smaller than 5 . Characteristic extinction curves were obtained in the present investigation for the various copper pyrophosphate complexes. Consequently spectrophotometric measurements in conjunction with polarographic and 9H measurements have permitted an extensive interpretation of the equilibria. Throughout this paper, Beer's law is in the form loglo lo/1 = E cl, where loglo 10/1is called the extinction, c is the molar concentration, 1 is the length of cell in cm. and E is the molar extinction coefficient.
Experimental Apparatus.-The spectrophotometric measurements were made in the Beckman D. U. spectrophotometer with l-cm. glass cells unless 10-cm. cells are specified. The pH measurements were made with the Beckman model G PH meter standardized with the Clark and Lubs borate buffer having a PH of 10. The solutions were plaoced in a thermostatically controlled bath kept a t 25 f 0.1 The polarographic measurements were made with a manual polarograph. The design of the instrument and experimental technique are described by Watters.6 The capillary had a drop time of 4.76 seconds and delivered 1.38 mg. of mercury per second a t the potential of the saturated calomel electrode. Reagents.-0.2 M sodium pyrophosphate was prepared by weighing Na4PzOrlOH20 and diluting to exact volume. The molarity was verified by an electrometric PH titration with standardized hydrochloric acid; 0.1 M copper sulfate was prepared in a similar manner using CuS01.5H20 and was standardized gravimetrically by electroplating copper. Approximately 1 perchloric acid was prepared by diluting the 60% solution. It was standardized against sodium carbonate. Solid IGa2CuFe(CN)6was prepared as follows: 25 g. of CuS01.5HzO and 85 g. of Na4P207.10H20 were dissolved in 700 ml. of water and heated to about 50'. Then 25 g. of KdFe(CS)a, dissolved in 200 ml. of water, was added to this solution with vigorous stirring. The precipitate, which consisted of very small reddish-brown colored crystal\,
.
< 600
I
mo
800
sa,
Yol. it?
I. FVATTERS AND ARNOLD.AARON
am
I*'o
WAWLCNGTH. my.
Fig. 1.-Effect of varying the ratio of copper-to-pyrophosphate ions. All solutions contained 0.020 Jf CuSO, with the following concentrations of SaaPlOj: 1, 0.04 X; 2 , 0.036 J I ; 3, 0.032 .If; 4, 0.028 111;5, 0.021 Ai; 6, 0.02 ,If; 7 , none. (3) B. C. Haldar, Science and Cullurc, 14, 341 (1949). (4) Erik Eriksson, Kirngl. Lanfbrzrks Hogskolans onn., 16, 72 (1949). ( 5 ) H. A. Laitinen and E. I. Onstott. THIS JOURNAL, 74, 4720 (1950). ( 6 ) J. I. Watters, Rev. Sci. IttXt71L"2.Y, 22, 851 (1951).
was washed nine times with 200-ml. portions of 0.3 M KaSOs by vigorously mixing the slurry, centrifuging, and finally decanting off the wash solution. The precipitate was found to contain copper and ferrocyanide ions in the mole ratio of exactly 1: 1 by spectrophotometric measurements of its air free solution in concentrated sodium pyrophosphate a t 370 and 740 mp. At 370 mp the molar extinction coefficient of ferrocyanide ion was found to be 53.4 h1-l cm.-l while those for the copper complexes investigated were practically zero. The dipyrophosphatocuprate(I1) was determined a t 740 mp where the extinction of the copper complex reached a maximum and that due to the ferrocyanide ion was negligible. Sodium was determined spectrographically, Copper-to-Pyrophosphate Ratio Experiments.-The general procedure consisted of transferring various volumes of 0.2 Jf sodium pyrophosphate by pipet to 100-ml. volumetric flasks. Then. after adding most of the water, 20 ml. of 0.1 it1 cUso4 was added slowly with stirring. The solution was diluted to the mark, mixed well, and a portion was transferred to a I-cm. cell for the extinction measurements shown in Fig. 1. The extinction curves for solutions containing more pyrophosphate than the 1:2 ratio are not shown since they practically coincide with that for the 1:2 ratio. N o extinction curves are shown for less than the 1: 1 ratio since a precipitate formed immediately. Consequently, a second series of extinction curves was obtained using more dilute solutions which were one millimolar in copper ion with greatly varied concentrations of pyrophosphate ion. These solutions were prepared by mixing, from burets, various proportions of solutions, all of which contained one millimolar copper ion but different pyrophosphate concentrations. Precipitates also formed in some of these solutions on standing several hours. Several of the extinction curves obtained using 10-cm. cells are shown in Fig. 2 .
WAVELENGTn mp
Fig. 2 -Effect of increasing the pyrophosphate concentration in very dilute solutions. All solutions were 0.001 111 in CuS04 with the following concentrations of h'a4P20j, (10 cm cells were used): 1,O.l M; 2,0.002 M; 3, 0.0016 M; 4, 0 001 ill; 5, 0 0008 M; 6, 0 0004 M; 7,O 0002 At; 8. none The p H data are plotted in Fig. 4. Curve 2 is a tenfold enlargement of the first part of curve 1. Data from these experiments were used to obtain the continuous variation results7 shown in curves 2 to 6, Fig. 5 . Curve 1, Fig. 5 , was obtained by continuously varying a solution which was 25 millimolar in copper and pyrophosphate ions with a second solution which was 25 millimolar in pyrophosphate ion. Acidity Variation Experiments .-Copper sulfate, sodium pyrophosphate in a known small excess of that needed for complex formation, and varying amounts of 0.87 N perchloric acid were mixed as in the ratio experiments. The PH data reported in the first two columns of Table I and the extinction curves shown in Fig. 3 were obtained immediately. A very pale blue precipitate formed in the PH range of about 5.5 to 2 and a soluble complex, curve 8, Fig. 3, remained in solution below a PH of 1.70. A preliminary analysis summarized below indicates that, in the absence of potassium ion, the formula of this precipitate is NatCus(P207)n.Aq. Spectrographic analysis of the pale blue precipitate obtained in the presence of potassium ion revealed considerable potassium as well as sodium. This compound has been reported by Fleitmann and Hennberf (7) W. C. Vosburgh and G (1941).
R. Cooper, THISJ O U R N A L , 63, 437
(S) T. Fleitmann and I\' Hennberg, Anit , 65, 387 (1848).
Feb. 5 , 1953
SPECTROPHOTOMETRIC INVESTIGATIONS OF COPPER PYROPHOSPHATE COMPLEXES ti13
Fig. 3.-Effect of increasing acidity. All solutions were 0.025 M in CuSO4, 0.075 M in NaaPzO7, and the ionic strength was adjusted t o unity with KNO3. The total concentrations of hydrogen ion added were: 1, none; 2,0.04 M : 3, 0.05 M ; 4, 0.07 M i 5, 0.08 M; 6, 0.09 M; 7, 0.10 M; 8, 0.18 M; 9, 0.16 M . Solutions were filtered for curves 4, 5, 6, 7 and 9 since precipitates formed in these solutions. to have crystals that are somewhat efflorescent a t room temperature but that 3.5 H20 are retained a t 100'. The sample was dried and water of crystallization was removed by heating t o 260' for three hours. Copper was determined electrolytically and sodium was determined gravimetrically as sodium zinc uranyl acetatesg Pyrophosphate was determined gravimetrically as magnesium pyrophosphate after hydrolyzing to phosphate by boiling the acidified solution. Anal. Found: Na, 8.37; Cu, 32.8; P z O ~57.9. , Calcd. for N ~ z C U ~ ( P Z ONa, ~ ) Z7.88; : Cu, 32.6; Pz07, 59.6. After standing several weeks these crystals had been replaced by larger sky-blue crystals which were considerably less soluble. The properties of the latter correspond to X~ZCUPZO~.~/~H~O
TABLE I EFFECTOF ADDINGACID TO SOLUTIONSCONTAINING A KNOWNEXCESSOF PHOSPHATE 2.5 millimoles CuSO4. 7.5 millimoles NarP207, 100 millimoles KNOl in 100 ml. of solution H + added, millimoles
#H
X"
X'b
KiiC
NarCuFe( CN)BEquilibration Experiments.-Portions of the precipitate, in excess, were shaken for 2 hours with oxygen-free solutions containing various concentrations of
TABLE I1 EQUILIBRATION OF SODIUMCOPPERFERROCYANIDE WITH VARIOUS CONCENTRATIONS OF COMPLEXINC AGENT IPecn ( ~ N i s l c - complex
molar molar concn. concn. (at equil.) (at equil.)
O . O Z M N ~ P I O T $ O . ~ M N ~ N0.00214 O~ .05 M N a P , O v + 0 . 8 M NaNOa ,00691 .10 M N ~ P ~ O T .0212 . l o M NtUPZO? ,0233 . 2 5 M NarCzOa .00110
0.00215 .00695 .0220 ,0223 .001082
Av.
PKt
(calcd.)
12.48 12.59 12.78 12.75
...
1 2 . 6 5 & 0.17
(9) H. H. Barber and I. M. Kolthoff, THISJ O U R N A L , SO, 1623 (1928).
- arrm
I I L . o'r CaJuw IOL. Io%.
I
OUJI
*.
X 10'
2.99 6.39 0.10 0.10 1.5 3.98 5.85 .26 .26 1.5 5.07 5.53 .42 .42 1.3 6.06 5.35 .59 .58 1.5 6.52 5.23 .63 1.2 precipitate formed 8.06 in this range 17.11 acidic complex formed 18.10 1.42 in this pH range 21.09 a Mole fraction [CuPzO7]-z based on equation (1). Mole fraction [CUPZO~] -z based on equation (6). e See equations (2) and (3).
Complex electtolyte molar concn. (initial)
~ I o I y I P H 4 l Zco"Am
Fig. 4.-pH titration of 0.001 M CuSOd with NarP207 without dilution. Curve (2) is a tenfold enlargement of the first portion of curve 1. pH values for curve (2) are indicated on right margin and numerical values of the pyrophosphate concentration are tenths millimolar per millimolar concentration of copper ion.
X.
Fig. 5.-Continuous variations: curve 1, s o h A, 0.025 M CuSO4, 0.025 M Pu'adPzO7; soln. B, 0.025 M Na4PzO7, 740 m p . In curve 1 multiply ordinate values by 3.33. X is fraction of solution B in all curves. Curve 2, soln. A, 0.001 Jf CuSOa, 0.002 M Na4P~07; soh. B, 0.001 MCuSO,, 780 mp. Curve 3, same as curve 2 a t 820 mp. Curve 4, soln. A, 0.001 M CuSO4, 0.001 M NarPzOT; s o h B, 0.001 M CuSO,, 820 mp. Curve 5, soln. A, 0.001 MCuSO4; soln. B, 0.001 M CuSO4, 0.0005 M Na4P~07, 1000 mp. Curve 6, same as curve 5 a t 1100 mp. sodium pyrophosphate and sufficient sodium nitrate to make the ionic strength approximately unity. In preliminary experiments equilibrium was approached from both directions by diluting some of the solutions to their final concentrations and repeating the procedure after first following this procedure a t higher pyrophosphate concentrations. The equilibrium concentration of ferrocyanide and dipyrophosphatocuprate( 11) in solution, Table 11, were determined spectrophotometrically as described before.
Discussion The existence of an isosbestic point at 760 mp in Figs. 1 and 2 for dilute and more concentrated solutions, respectively, indicates the presence of two copper complexes other than the aquo ion in solutions containing copper-to-pyrophosphate ion con-
centration ratios greater than 1 : 1 in pyrophos- pyrophosphatodicopper(II), having the formula phate. The tendency for precipitates to form in [ C U ~ P ~ O The ~ ] ~ maximum . in curve 3 a t X = solutions having ratios less than 1 : 1 indicates that 0.7.5 and in curve 4, Fig. 5 , a t X = 0.50 prove the one of these is monopyrophosphatocuprate(I1) and existence of [ C U ~ P ~ OThe ~ ] ~bulge . in curve 3 at the similarity of the extinction coefficients for the X = 0.30 is caused by the formation of [ C U P Z O ~ ] ~ dilute solutions, curve 4, Fig. 2 and more concen- which was confirmed by curve 2. ,4 model having trated solutions, curve 6, Fig. 1, indicates that this two copper ions in six-membered chelated rings ion is very stable. The similarity of the extinction with a single pyrophosphate ion can be readily concurves for the 1 : 2 ratio in the more concentrated structed. I t is probable that in this complex each solutions, curve I , Fig. 1, with those containing copper ion is coordinated with two water molemore pyrophosphate, which practically coincide cules. The pH of 3.59, in a one millimolar solution with curve 1, indicates that the second complex is of the complex, Fig. 4, indicates that the complex dipyrophosphatocuprate(I1j. The continuous vari- hydrolyzes to the same extent as does the aquo ation results shown in curves 1 and 2, Fig. 3, con- complex, which has a pH of 5.58. Curve 2 , Fig. 4, has a minimum pH of 5.38 in firm the existence of these two complex ions since the maxima a t X = 0.5 correspond to the formulas the region where the copper-to-pyrophosphate concentration ratio is 4:1. Since the partial con[Cu(P207)2]6- and [CUP~O~}", respectively. having a pH of 5.58 to Since the extinction curve for the dilute 1 :2 solu- version of [CU(H~O,)]~+ O ~ ] a pH of 3.39 cannot account for tion, curve 2, Fig. 2 , is not completely shifted to the [ C U ~ P ~having position for the solution containing an excess of the increase in acidity, the existence of a complex pyrophosphate, curve 1, Fig. 2, it follows that the cation containing four copper ions for each pyrodipyrophosphatocuprate(I1) is partially dissociated phosphate ion is indicated. The small maxima a t to the monopyrophosphatocuprate(I1). The de- X = 0.3 in the continuous variation curves 5 and gree of dissociation a t various dilutions, Table 111, ti, Fig. 5 , confirm the existence of the complex will be used to evaluate the first instability constant cation, [Cu4P207l4+. The effects are too small to verify the complex containing three copper ions oi the dipyrophosphato complex. which probably also exists. T A B L E 111 I t is possible that this complex cation, [CuqP2EFFECTOF DILUTIONo s THE DECREEOF DISSOCIATION OF O7I4+, contains two of the copper ions in sixDIPYROPHOSPHATOCUPRATE( 11) membered rings and retains the additional copper (KSOa added to produce ionic strength of unity) ions in four-membered rings containing one phosComplex Kit Cumplex Kii phorus atom and two oxygen atoms, one of which a X 10' X 104 concn., m concn., tic a may be shared with a copper ion in a six-membered 0. 13 1. f i 1.7 ( J . 008 0 , 0005 0.-13 ring. The tendency of the coordinated water 1 .cj 02 .09 1.8 0008 .:31i molecules to lose hydrogen ions which is manifested 025 .08 1.7 ,001 .&I I 7 by the low pH of 5.38 may be due to the relatively 0Oq5 .17 1.7 large positive charge resulting from the proximity I t is reasonable that, in both of these complexes, of two copper ions sharing an oxygen atom. The anomalous drop in pH observed by Rogers the copper is bonded with the pyrophosphate by six-membered chelate rings of the type proposed by and Reynolds? a t the beginning of titrations of Kolthoff and Watters'O~" for the trivalent manga- metal ions such as copper, zinc and cadmium by nese complex. Any remaining coordination bonds pyrophosphate ion was undoubtedly caused by the of copper in any of the complexes with pyrophos- formation of complex cations similar to [Cu4P2phate are probably satisfied by water. In the di- O7I4 Furthermore, the abnormally large capacity pyrophosphato complex, copper's usual coordina- of pyrophosphate for complexing iron and other tion number of four is satisfied by the two pyrophos- metal ions in hard water, which is well known, can be explained by the formation of complex cations of phate ions. T h a t the pyrophosphate ions in both the mono- this type. In the acidity variation experiments, the added and dipyrophosphato complexes hydrolyze to a negligible extent is evident from the pH values of hydrogen ion, in addition to combining with free 7.6 and 8.6, respectively, measured in solutions 25 pyrophosphate ions, might compete with copper rnillitnolar in these ions. This is ,reasonable since a ion for pyrophosphate ions or it might associate pyrophosphate ion complexed with a positive ion, with pyrophosphate ions in the complex. The siniicopper in this case, must have less tendency to larity of curves 1 to 3, Fig. 3, to curves 1 to 6, Fig. associate with hydrogen ions than does a free pyro- 1, and the existence of the isosbestic point a t '760 mp phosphate ion. indicate that the dipyrophosphato complex is conAdditional complex species were discovered in verted to the monopyrophosphato complex in the solutions containing an excess of copper ion. These PH range of about '7 to 5 . This conclusion will be complex ions were detected only in very dilute solu- confirmed by equilibrium calculations. A Iretions due to the insolubility of precipitates such as cipitate, such SI NazCua(P2O.i),.Xq, may form N a2Cua(P2C);)e.Aq. The continuous variation data sorilewhere iri the pH rauge of 5.5 to 2, depending of curves 3 and 4, Fig. 5 , prove the existence, ill on the cornposition of the solution. Below a pH of very dilute solutions, of a neutral complex, mono- 2 or :;, tlir solution coiit:tiiiS soluble complexes which (IC)) I . 11. Kolthofi' and J . 1 . Wallrrs, liiil. I
