Solvent extraction of pyridine complexes of copper(II)

J. inorg,nucl.Chem., 1969, Vol. 31. pp. 513 to 525. PergamonPress. Printedin GreatBritain

SOLVENT EXTRACTION OF PYRIDINE COMPLEXES OF COPPER(II) J. A G G E T T and M. W. B I L L I N G H U R S T Chemistry Department, University of Auckland, N e w Zealand

(Received 20 May 1968) Abstract--Copper(l l) is extracted into chloroform by pyridine in the presence of bromide, chloride, perchlorate, and salicylate ions. At pyridine concentrations between 10 -2 and 10 -t M the extracted species have the formula Cupy2Xz. At higher pyridine concentrations both CupyaX2 and CupyvY2 appear to be extracted. The partition data confirms the existence of the dimer Cu2py4(OH)~ 2+ in the aqueous phase at pH > 7. INTRODUCTION

Pm/vIous papers from this laboratory[l, 2] have been concerned with the solvent extraction of salicylate complexes. This paper is concerned with the extraction of copper(lI) into chloroform by pyridine in the presence of salicylic acid, a system which was used by Gordieyeff[3] as the basis ofa colorimetric method for analysis of copper. Although details of the conditions necessary for quantitative extraction were published, the nature of the extracted species was not determined. The research reported in this paper was undertaken to determine the role of both pyridine and salicylic acid in the extraction process. Preliminary experiments showed that in the presence of bromide, chloride, and perchlorate ions copper(l I) was extracted by pyridine even in the absence of salicylic acid. Since these inorganic systems were considered to be less complicated than the salicylate system in that the distribution of the anion between chloroform and aqueous solution is independent of pH and also because these anions are unlikely to form complexes with copper(ll) in the aqueous phase as might salicylate ions under the experimental conditions the copper(ll)-pyridine-bromide system was also studied in detail As in the previous papers [ 1,2] the species C7HrOa, C7H503-, and C7H4032are denoted by H~Sal, HSaI-, and Sal2- respectively. EXPERIMENTAL Chloroform was used as the organic solvent for all experiments except those reported in Section ( f ) in which cyclohexane was used. Distribution measurements on both chloroform-aqueous and cyclohexane-aqueous systems were made by equilibrating equal volumes (usually 5 ml) of aqueous and organic phases for I hr at 25°C; the phases separated on standing. Copper(l !) was determined colorimetrically by one of two methods. The first method was based on the formation of bis-8-hydroxyquinolinatocopper(ll)[4]. Copper in the organic phases was converted to bis-8-hydroxyquinolinatocopper(ll) by addition of an equal volume of 2 x 10-z M 8-hydroxyquinoline in chloroform. The absorbance of this solution was measured using 10 -2 M 8-hydroxyquinoline as blank. There was no interference from salicylic acid concentrations ~< 10-j M. The I. 2. 3. 4.

J. Aggett and P. Crossley, J. inorg, nucl. Chem. 29, 1113 (1967). J. Aggett, D. Evans and R. Hancock,J. inorg, nucl. Chem. 30, 2529 (1968). U. A. Gordieyeff, Analyt. Chem. 22, 1166 (1950). G. H. Morrison and H. Freiser. Solvent Extraction in Analytical Chemistry. Wiley, N e w York (1957). 513

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J. A G G E T T and M. W. B I L L I N G H U R S T

aqueous copper(ll) concentration was obtained by subtraction from the total copper(ll)concentration in the system. The second method was based on the absorbance of copper(ll) at 385 n ~ in conc. hydrochloric acid. Aqueous phases were analysed by diluting 1-10 ml with conc. hydrochloric acid. Organic phases were back-extracted twice with 4 ml aliquots of conc. hydrochloric acid; the extracts were diluted to 10 ml with conc. hydrochloric acid for absorbance measurements. The distribution ratio was determined from the formula

Distribution ratio (q)

Total conc. of copper(I I) in organic phase Total conc. of copper(ll) in aqueous phase "

In the pyridine distribution experiments pyridine was determined spectrophotometrically using its absorbance at 255 mtz. The distribution of salicylic acid between water and chloroform was examined similarly using its absorbance at 295 m/~. Solubility measurements were made by evaporating 5 ml samples to dryness then redissolving the complex in an NHa-NH4CI buffer and titrating the copper(II) with EDTA. Organic materials were all purified by the methods of Perrin, Amarego, and Perrin[5]. It was particularly necessary to remove traces of benzene from the cyclohexane used in the pyridine distribution experiments as benzene absorbs light at 255 m/.L.This was done by passing cyclohexane through a column of activated silica [5]. The complexes Cu(C5 H~N )2Br2 and Cu(C5 HsN )~(H Sal)~ were precipitated from aqueous copper( 11) solutions on addition of the other reagents. Analysis figures for the bromide complex were 16.65% Cu (16.65% calc.), 42.1% Br (42.0% calc.). The salicylate complex contained 12"75% Cu (12-80% calc.). RESULTS

(a) Distribution of copper(lI) as a function of pH Typical pH profiles for the chloride, bromide, perchlorate, and salicylate systems are shown in Fig. 1. The profiles for the systems containing inorganic anions are similar; the value of logq rises with a slope of between 1 and 2 and reaches a maximum at about pH 5.5; it remains constant to about pH 7 and then drops sharply. The salicylate system differs in that there is no horizontal region between pH 5.5 and 7.0 and the fall in the value of q at higher pH levels occurs in two stages.

(b) Variation of anion concentration At constant total copper(lI) concentration (10 -s M), and pyridine concentration (1 M), the total bromide concentration was varied between 2.5 x 10-z M and 2 × 10-1 M. The results (Figs. 2 and 3) show a second order dependence of q on the total bromide concentration at constant pH. Figure 4 shows pH profiles of systems with different total salicylic acid concentration. Complete interpretation of this data requires a knowledge of the distribution of salicylic acid (Section (e)) and also the degree of complex formation between copper(II) and salicylic acid in the aqueous phase. Since calculations showed that the latter effect was negligible at pH 3.0 in solutions containing ~< 10-~ M total salicylic acid the dependence of q on the total aqeuous salicylate concentration was determined at this pH in a separate series of experiments in which the total salicylic acid concentration was varied between 10-2 M and 10-1 M while the total pyridine concentration was maintained constant at 10-1 M. The plot of log q vs. log H2Saltot~l.aqueous(Fig. 5) has a slope of 1.95. 5. D. D. Perrin, W. L. F. Amarego and D. R. Perrin, Purb'ication of Laboratory Chemicals. Pergamon Press, Oxford (1966).

Solvent extraction of pyridine complexes of copper(! !)

515

O

-I

-2 4

5

6

7

8

Fig. 1. pH profiles. Curve l: Cur = I0-a M,pyr = 1 M, Clr = 10-1 M. Curve 2: Cur = 10-3 M,pyr = 1 M, BrT= 10-1M. Curve 3: Cur= 10-3 M,pyr= 1 M, CIO4r= 10-~ M. Curve 4: Cur = 10-a M,pyr = 10-~ M, H2Salr = 10-z M.

(c) Variation of pyridine concentration At constant total copper(II) concentration (10 -a M) and bromide concentration (2 × 10-~ M), the total pyridine concentration was varied between 0.25 and l M. The data are shown in Fig. 6. The results of pyridine variation studies in the salicylate system are shown in Fig. 7.

(d ) Variation of copper( ll ) concentration The total copper(lI) concentration was varied between l 0 -a M and 10 -2 M in solutions containing constant total pyridine (1 M) and bromide (2 × 10-1 M) concentrations. T h e s e results (Fig. 8) show that at constant pH the distribution ratio is independent of the total metal ion concentration below about pH 5.5, but that above this p H log q decreases with increasing total copper(l I) concentration at constant pH.

(e) Distribution of pyridine and salicylic acid between chloroform and water The results for the pyridine system are shown in Fig. 9, and those for the salicylic acid system in Fig. 10. According to H o k [ 6 ] the dependence of the distribution of salicylic acid on the concentration of the acid is due to dimerisation in the organic phase, the equilibria involved being 6. B. Hok, Svensk. Kern. Tidskr. 65, 182 (1953),

516

J. A G G E T T and M. W. B I L L I N G H U R S T D

(1)

HsSaI~ ~ HsSalo k2

(2)

2HsSalo ~ (H4Sals)o. These combined with Equation (3) [nsSaltotal]o = [HsSal]o-q-

[H4Sals]o

(3)

give the relationship [H2Saltotal]0 [H2Sal]a = D + 2k2D2[H2Sal]a.

(4)

From the plot of [H2Saltotag]0

[HsSal]a

vs. [HsSal]a

values of D, and ks were found to be 3.0 and 40 respectively. These values are in good agreement with those of Hok[6] viz. D = 3.0, ks = 42. ( f ) Distribution of pyridine between cyclohexane and water in the presence o f copper(ll) In order to obtain further information of the species present in the aqueous phase at pH 6-9 concentration of pyridine extracted into cyclohexane at pH 8.0 1.0

0 o~ 0 0 ._1

1.0

2.(3

4

5

6

7

Fig. 2. Variation of bromide concentration. Curve 1 : Bromide.r ----2x 10-1 M. Curve 2: Bromider = 10-1 M. Curve 3: B r o m i d ~ = 5 x 10-2 M. Curve 4: BromideT = 2-5 × 10-2 M.

8

Solvent extraction of pyridine complexes of copper( I I)

517

I

O cP (.9 O ..J

-I

-2

LOG [N B-:I I

I

-1"5

I

-I'O

-0"5

Fig. 3. Logq vs. log [NaBr]. Curve I : pH 5. Curve 2: pH 6, 7. Curve 3: pH 7-5.

was determined for a series of solutions containing constant total pyridine concentration (2 × 10-1 M) and different total copper(l !) concentrations. Copper(l 1) was added as copper(lI) sulphate. As no anion capable of effecting extraction of copper(ll) was present in the system uncomplexed pyridine was the only source of pyridine in the organic phase. The results of this experiment were used to calculate the average number of pyridine molecules coordinated to each copper(11) ion (t~) in the aqueous phase, h was also determined over the pH range 7.5-9-0 for solutions containing a constant copper(ll) concentration (5 × 10-2 M). The value of a was found to be 1-9 __+0-1 in all these experiments.

(g) Absorption spectra and solubility measurements The visible absorption maxima of the extracts are listed in Table 1. For the systems containing inorganic anions these spectra showed no significant variation with pH. In the salicylate system the spectra were also independent o f p H except at low pyridine concentrations (10 -2 M) when a second species was extracted at about pH 6. This species had an absorption maximum at 410 rn/z in addition to the normal copper(l I) spectrum for octahedral configuration. However in both the bromide and salicylate systems the spectra varied with the total pyridine concentration. In both cases this change in spectra occurred at pyridine concentrations t> 10-1 M. The same variation in spectra was observed in chloroform solutions of Cupy2Br2 and Cupy2HSaI2 containing added pyridine (Table 2).

J. AGGE'I-F and M. W. B I L L I N G H U R S T

518

O

-I

-2 3

4

5

6

7

8

Fig. 4. Variation of salicylate concentration. Curve 1: HzSalr = 10-1 M, Curve 2: H~Salr = 3 × 10-2 M. Curve 3: H2Salr = 10-~ M, O

0

-I

-2

I

-2"4

I

I

-2.0

I

-1"6

Fig. 5. Logq vs. log [H2Sal]total aqueousat pH 3.

T a b l e 3 c o n t a i n s d a t a f o r t h e s o l u b i l i t y o f Cupy2Br2 an d Cupy2HSaI2 in c h l o r o f o r m as a f u n c t i o n o f p y r i d i n e c o n c e n t r a t i o n .

Solvent extraction of pyridine complexes of copper{ I I )

I [ C u T = 1(33M

ii 4

N~B~ = 2X IO-IM

6

5

7

8

Fig. 6. Variation of pyridine concentration in bromide system. Curve I: Pyridiner= I M. Curve 2: Pyridiner = 5 × 10-1 M. Curve 3: Pyridiner = 2-5 × 10 -~ M. Table I. Spectra of chloroform extracts [pyridine]total (mole/I.)

Bromide system hm~, (m/x)

Salicylic acid system ~-max(m/x)

2.5 × 10-2 5 × 10-2 1 × 10-~ 2-5 × 10-j 5 x 10 -~ 1M

755 750 750 745 735 720

640 640 645 655 660 665

Table 2. Spectra of Cupy2Br2 and Cupy2HSal2 in pyridine/ chloroform [pyridine] (mole/L)

hmax(Cupy2 Br2) (mg)

hmax(Cupy2HSal2) (mp.)

2'5 × 10-2 5 × 10-2 1 × 10 1 2"5 × 10-z 5 x 10-~ 1M

760 750 750 740 735 720

650 645 650 655 660 670

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J. A G G E T T and M. W. B I L L I N G H U R S T

520

Table 3. Solubility of Cupy2Br2 and Cupy~HSal~ in chloroform in the presence of pyridine

[pyridine] (mole/L)

Solubility of Cupy2Brz (mmole/l.)

Solubility of CupyzHSal2 (mmole/l.)

1-0 X 10.2 2-5 x 10-2 5.0 X 10-2 1.0 × 10-1

2.30 2-25 2.30 3"51

-54.0 54.2 54-0

2"5 x lO -I 5-0× 10 -I 7.5 × 10 -I IM

9.10 24"I 49.0 72-0

55"0 61-7 70.2 78"2

0

-I

-2 3

4

5

0

7

8

Fig. 7. Variation of pyridine concentration in salicylic acid system. Curve 1: P y f i d i n e r - 10 -1 M. Curve 2: PyridineT = 3 x 10-2 M. Curve 3: PyridineT = 10-2 M.

Solvent extraction of pyridine complexes of copper(11)

521

DISCUSSION

Since in general only uncharged species are extracted into organic solvents the second order dependence of q on bromide ion concentration is not unexpected. It does show however that there is no association of bromide ions with the copper(11)pyridine species in the aqueous phase. The observed second order dependence of q on aqueous salicylate concentration at pH 3 is of more interest because it shows that the salicylate is functioning as a monovalent ligand in the extracted species. The solubility data suggest that in chloroform solutions containing/> 10-1 M pyridine the copper(lI) species contains more than 2 molecules ofpyridine. In the bromide system the total concentration of copper(If) in solution can be expressed as

Cur = [Cupy2Br2] + [Cupy3Br2] +

[Cupy4BF2]

where Cur is the total concentration of copper(II) in solution. This can be rearranged to Cur

[CupyzBrz] = 1 + k3 [ py ] JI- k3k4 [ py ] 2. where k3 and k4 are defined as follows [Cupy3 Br=] k3 = [Cupy2Br2][py ]

[ Cupy4 Br2] k4 = [Cupy3Brz][py] "

1.0 N,',Br = 2X IO'IM

P~T =IM

O t~ t9

o

.-I

-I'(

-2'0 4

pH ,

i

,

5

6

7

8

Fig. 8: Variation of copper(l 1) concentration in bromide system. O Cu r = 10-a M. • Cur = 3 x 10-a M. Cur = 10-2 M.

522

J. A G G E T T and M. W. B 1 L L I N G H U R S T

I00

6ot°

,ol.S pH

3

I

I

I

4

5

6

Fig. 9. Distribution of pyridine. © Pyridiner = 10-2 M. tD Pyridiner = 2.5× 10-1 M. Pyridiner = 5 x 10-1 M. (ll Pyridiner = 1 M.

I00

/o/

/~

i o

I

7/

'< 40 .z

/~( 2 F / oY/ ,D

-~ ~ J

~o . ~ / / !

2

3

pH i

4

!

5

Fig. 10. Distribution of salicylic acid. Curve 1 : H~Salr = 10-' M. Curve 2: H2SalT = 10-1 M.

Solvent extraction of pyridine complexes of copper(l I )

523

2

=o O

LOG I

I

-0.8 ,

PYRIDINE CONC. -0"6

r

[Cur]

Fig. l I. log l[Cu--P-y2Br2]

I

-0"4

I

-0-2

1] VS. log [pyridine]. /

Since the function Cur/[Cupy2Br2] does not vary linearly with either [py] or the plot of log[(Cur[[Cupy~Br2])-1] vs. log [py] has a slope of 1.75 (Fig. 11) it appears that both Cupy3Br~ and Cupy4Br2 are present in chloroform solutions with pyridine concentrations between 10-1 and 1 M. The changes in absorption spectra also appear to be associated with this change in coordination of pyridine. And although this argument has been applied to isolated chloroform solutions it also holds for chloroform solutions in equilibrium with aqueous phases; this is shown by the similarity of the spectra of chloroform phases from extraction experiments with those of the isolated solutions containing the same pyfi.dine concentration. A similar argument can be applied to the solubility and spectral data for the salicylate system although the accuracy is lower because of the smaller differences in solubility. Again, at pyridine concentrations ~< 10-1 M the extracted species appears to be Cupy2(HSAL)2 and at higher pyridine concentrations both Cupy3(HSal)z and (cupy4(HSal)2 appear to be extracted. It is interesting to note that whereas the spectrum of the bromide system undergoes a hypsochromic shift with increasing pyridine concentration that of the salicylate system undergoes a bathochromic shift. This difference is probably caused by steric hindrance in the latter system where the presence of the bulky salicylate ions does not permit the third and fourth pyridine molecules to approach the copper(II) ion as closely as in the corresponding bromide complexes. According to the treatment of solvent extraction systems by Irving, Rossotti,

[py]Z and

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J. A G G E T I " and M. W. B I L L I N G H U R S T

and Williams[7] the dependence of the distribution ratio on the copper(ll) concentration (Fig. 8) is an indication of polymer formation in one of the phases; the difference in the degree of aggregation in the two phases is given by the expression dlogq d log M r

(r~o- r~a) (1 + q) r~o - - m a q

where mo, ma are the polymerisation numbers in the organic and aqueous phases. Evaluation of this indicates that copper(II) exists as a dimer in the aqueous phase if it is assumed that the extracted species is monomeric. The results of the pyridine distribution experiments show that there are two pyridine molecules complexed to each copper(II) ion in this dimer, while the change in q with pH in the vicinity of pH 8 shows that the species contains one hydroxyl ion per copper(I 1) ion i.e. the species is [Cuzpy4(OH)2] 2+. This confirms the work of Leussing and Hansen[8] who reported a formation constant (/342) of 10~4"7for the reaction 2Cu 2+ + 2OH- + 4py ~.~ Cthpy4(OH)~2+. However, the data in this paper suggests that the value of this constant is higher than that reported by Leussing and Hansen. Using Sun and Brewer's [9]

'° L~o

3

uJ

0'8

045 V-

0.4

0.2

3

4.

5

6

7

8

Fig. 12. Calculated fractions of species in 10-1 M pyridine-10 -2 M salicylic acid system. Curve 1: Cupy '+. Curve 2:Cupy2 2+. Curve 3: CuSal. Curve 4:CuSal22-. 7. H. Irving, F. J. Rossotti and R. J. P. Williams, J. chem. Soc. 1906 (1955). 8. D. L. Leussing and R. C. Hansen, J. A m . chem. Soc. 79, 4270 (1957). 9. M. S. Sun and D. G. Brewer, Can. J. C h e m . 4S, 2729 (1967).

Solvent extraction of pyridine complexes of copper(I I)

5 25

formation constant data for copper(II)-pyridine complexes we have calculated the concentration of each species present in the systems under study; the fall in q at higher pH corresponds with the calculated formation of Cu2py4(OH)~ 2+ if fl42 has the value 6 × 102~. N o w that the species present in both phases are established it is possible to explain the shapes of the pH profiles. The copper(II)-pyridine-bromide was studied at pyridine concentrations > 10 -1 M. Under these conditions the species extracted are Cupy2Br2, Cupy3Br2, and Cupy4Br2. The rise in q between pH 3-5 and the horizontal region correspond with the formation of Cupy2~+, Cupy32÷, and Cupy4~+ in the aqueous phase; the drop in q at pH > 7 is caused by the formation of Cuzpy4OH~2÷ in the aqueous phase. The copper(ll)-pyridine-salicylic acid was studied at pyridine concentrations ~< 10-1 M; the major species extracted is CupyzHSal2. And the formation of extractable species in the aqueous phase occurs in competition with the formation of the chelated salicylate complexes CuSal, and CuSal~ 2-. Using known formation constants for copper(I I)-pyridine [9] and copper(ll)-salicylate [ 10] complexes, and Cuzpy4OH2z+ (/342 --- 6 × 1025), the concentration of the species in the aqueous phase was calculated as a function of pH. This calculation showed that the extraction profile corresponds with the formation of Cupy2z÷ in the aqueous phase; the initial fall in q at about pH 5 corresponds with formation of CuSal, and the further fall at pH > 7 with the formation of CuSal22- (Fig. 12). The calculation also showed that less than 1 per cent of the copper(l 1) in the aqueous phase existed as Cu2py4(OH)22+ at pH 8. 10. D. D. Perrin, Nature, Lond. 182, 741 (1958).