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Determination of dissociation constants for coordination compounds of Cr(III) and Co(III) using potentiometric and spectrophotometric methods
J. Chem. Thermodynamics 40 (2008) 1290–1294
Contents lists available at ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Determination of dissociation constants for coordination compounds of Cr(III) and Co(III) using potentiometric and spectrophotometric methods A. Chylewska, D. Jacewicz, D. Zarzeczan´ska, L. Chmurzyn´ski * ´ sk, Sobieskiego 18/19, 80-952 Gdan ´ sk, Poland Department of Chemistry, University of Gdan
a r t i c l e
i n f o
Article history: Received 10 January 2008 Received in revised form 11 March 2008 Accepted 19 March 2008 Available online 29 March 2008 Keywords: Complexes of Cr(III) Complexes of Co(III) Acidity constants Potentiometry Spectrophotometry
a b s t r a c t The acid–base properties of analogous complex ions of chromium(III) and cobalt(III) in aqueous solution have been studied. The equilibrium constants for all metal complexes were determined by using potentiometric and spectrophotometric titration methods. First, dissociation constants for the studied complexes of Cr(III) and Co(III) were determined by means of the potentiometric titration method and using the STOICHIO computer programme. Then, pH-spectrophotometric titrations were performed and the OriginPro 7.5 computer programme was used to calculate the same constants. The measurements using both methods were carried out under the same conditions of temperature, T = 298.15 K, and over the same pH range 2.00–10.00, respectively. It turned out that the two methods used enabled us to obtain acidity constants in very good agreement. Ó 2008 Elsevier Ltd. All rights reserved.
1. Introduction Potentiometric and spectrophotometric titrations are the most useful methods to investigate equilibrium in solutions and to determine acid–base constants [1–3]. The potentiometric titration is used so often due to the simplicity of equipment and minimal time requirements [4]. However, this method does not cover all aspects of solution chemistry in which researchers have an interest. To have complete information about the species formed during the titration, spectrophotometric titrations are usually carried out simultaneously [5–7]. This technique shows how many equilibria exist in solution studied. Moreover, it can be applied to structural analysis of compounds and is neither limited to any pH range nor to aqueous solutions. Additionally, it is useful in discriminating between different models of equilibrium [8]. All these aspects prove that the spectroscopy is more powerful than the potentiometric method; however, it is advisable that they should be used together for a more detailed determination of acid–base interactions. In both types of titration, a controlled displacement of chemical equilibrium occurs during the addition of titrant. The determination of the constants of dissociation of the complex compounds of chromium(III) and cobalt(III) ions is extremely
* Corresponding author. Address: Department of General and Inorganic Chemistry, University of Gdan´sk, Poland. Tel.: +48 58 523 53 19; fax: +48 58 523 54 72. E-mail address: [email protected] (L. Chmurzyn´ski). 0021-9614/$ - see front matter Ó 2008 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2008.03.012
important in investigations executed by our research group. Based on results from the spectrophotometric and potentiometric measurements of analogous compounds, it can be concluded which complexes are more stable and active in biological and chemical materials. It is worth pointing that previously we synthesized the compound of chromium(III) with the aminodeoxysugar as ligand; cis-[Cr(C2O4)(AaraNH2)(H2O)2]+, where AaraNH2 stands for methyl 3-amino-2,3-dideoxy-a-D-arabino-hexapyranoside. This complex was so reactive that it caught the free radical of nitrogen dioxide generated from the pancreas of rats [9]. The exchange of the folded aminodeoxysugar’s ligand on simple ammonia molecules or the chloride anions can facilitate the reaction occurring for any gas (CO2, NO2 or SO2) uptake by coordination compounds studied. In our study, the complex compounds cis-[Cr(NH3)4(OH2)2](ClO4)3, cis-[Cr(NH3)4(OH2)Cl]Cl2, cis-[Co(NH3)4(OH2)2](ClO4)3 and cis½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O were used. We assumed that these analogous coordination compounds of Cr(III) and Co(III) would reveal similar acid–base properties. The only difference between the analogues species is the coordination centre. Based on determined acidity constants of the compounds studied, it is possible to define what impact the transition metal ion has on properties of complex compounds. In the present paper, the correlation between results obtained by two different experimental methods has been discussed. Regardless of the method, the determined values of the dissociation constants for the coordination compounds studied are similar, which suggests that these values are correct.
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2. Experimental 2.1. Materials 2.1.1. The synthesis of coordination compounds studied The complex compound of cis-[Cr(NH3)4(OH2)Cl]Cl2 was prepared according to the method of Glerup [10]. The synthesis of cis-[tetraamminediaquachromium(III)] perchlorate was described in paper [11]. The preparation of compound with cobalt(III) as the coordination centre binding four molecules of ammonia and bidentate carbonate ion was reported in [12]. The coordination compound, cis-[Co(NH3)4(OH2)2](ClO4)3, was produced from cis-[Co(NH3)4(O2CO)]NO3 0.5H2O according to the procedure described in reference [13]. All the coordination compounds studied were synthesized from commercial reagents (see table 1). 2.2. Instrumentation and measurement The measurements were carried out by means of potentiometric titrations at constant ionic strength using an automated system and applying the Microtitator programme. All probes, which were used in titrations, were prepared in nitrogen atmosphere to avoid CO2 contamination and the temperature was kept at (298.15 ± 0.1) K. A constant ionic strength of 0.1 M was maintained with NaClO4. The titration system consisted of a titration cell, a magnetic stirrer, and an automatic titrator with Hamilton’s syringe (0.5 cm3). The pH combined electrode was bought from the Metrohm firm. The electrode was calibrated using pH standard buffers [14]. The Perkin Elmer Lambda 800 UV–Vis double beam spectrophotometer, with automatic stirrer, was used for absorbance measurements with 1.00 cm quartz microcells T = 298.15 K and 0.1 M ionic strength (NaClO4). For each pH point, a known aliquot of solution was extracted and the absorption spectrum was recorded. The method of spectrophotometric determinations of acid–base constants was described in details elsewhere [15]. The profile of absorbances obtained at the wavelength of the absorption maximum versus pH in each run was used to obtain the equilibrium constants of the complex species formed in each system. 2.3. Determination of pK values by using potentiometry To calculate the acid–base constants from the potentiometric method, the STOICHIO computer programme was used. This programme is based on the algorithm of Kostrowicki and Liwo [16–18] and is able to handle any model of chemical equilibrium. The programme determines the equilibrium constants as parameters of the function to be minimized: U¼
ny n X X 1 b 1 b 1 2 ^i yi Þ2 ; ½ E i EðV i ; V; y; xÞ2 þ 2 ð V ðy i V iÞ þ 2 2 r r r yi E V i¼1 i¼1
ð1Þ where b E denotes the measured potential difference for the ith titration point, V is the measured volume of the solution at the ith titration point, vector x refers to unknown equilibrium constants and TABLE 1 Purity of the coordination compounds of transition metals checked by the elemental analysis Coordination compound
Calculated
Theoretical
cis-[Cr(NH3)4(OH2)2](ClO4)3 cis-[Co(NH3)4(OH2)2](ClO4)3 cis-[Cr(NH3)4(OH2)Cl]Cl2 cis-[Co(NH3)4(O2CO)]NO3 0.5H2O
12.35%N, 3.52%H 2.17%N, 3.53%H 22.86%N, 5.70%H 27.16%N, 5.10%H
12.33%N, 12.14%N, 22.90%N, 27.10%N,
For comparison theoretical values are included.
3.55%H 3.49%H 5.73%H 5.03%H
vector y represents calibration parameters, total concentrations of reagents in the stock solution, and titrants, etc. The minimization of the sum of the squares given in equation (1) is done using Marquardt’s method [19]. This method takes into account not only errors in the e.m.f., but also those in titration volume, stock-solution preparation, electrode calibration parameters, and reagent impurities, in addition to the equilibrium constants determined from other measurements [20]. The details of the calculations of equilibrium constants have been described previously [21,22]. 2.4. Determination of pK values by using spectrophotometry Based on the Lambert–Beer–Bouger law for spectrophotometry, viz., equation (2) Ak ¼ ekBH b0 þ ðekB ekBH Þ b ¼ ekB b0 þ ðekBH ekB Þb h;
ð2Þ
where b0 is the initial concentration of complex studied, [BH]; b is the concentration of deprotonated form of complex studied, [B]; h is the concentration of hydrogen ions; e is the molar absorptivity of the absorbing species; AkBH is the absorbance at the ‘‘acid end” of the titration (bh is the b0); AkB is absorbance at the ‘‘alkaline end” of the titration (b is the b0), one can obtain the Henderson–Hasselbach equation in the form of equation (3) pK 1 ¼ pH ln
b Ak AkBH ¼ pH ln : bh AkB Ak
ð3Þ
After transformation of equation (3) according to [23] to equation (4) Ak ¼
AkBH 10pH þ AkB 10pK 1 10pH þ 10pK 1
;
ð4Þ
this new form of expression was put to OriginPro 7.5 programme to calculate values of dissociation constants of complexes studied. To check how many dissociation constants we should calculate and how many equilibria existed in the solution studied, we had to construct relationships between absorbances at two, different wavelengths, which are called A-diagrams. These diagrams, created from the data from spectrophotometric titrations are criteria by which the presence of measurable overlap can be observed. These diagrams [24] constructed for all wavelength combinations must be linear in the presence of equilibrium conditions. For example, A-diagram is described always by a linear equation (5) for single step acid–base titration: A1 ¼
e1B e1BH A2 þ ðe1BH e2BH Þ b0 : e2B e2BH
ð5Þ
It should be pointed that the intersection point of two linear segments of the A-diagram corresponds to the presence of the deprotonated ampholytic form in the case when two equilibria are present in solution. All the titrations were performed under complete computer control. 3. Results and discussion 3.1. Spectrophotometric and pH-metric measurements The course of spectrophotometric titrations for all the coordination compounds studied cis-[Cr(NH3)4(OH2)2]3+; cis-[Cr(NH3)4(OH2)Cl]2+; cis-[Co(NH3)4(OH2)2]3+; cis-[Co(NH3)4(O2CO)]+ were registered. The identical conditions of measurements, which were applied in the potentiometric method, were kept. The results obtained are presented in graphical form as the relationship between absorbance and wavelength. Thus the relationships between absorbance (A) and wavelength (k), for two of coordination compound studied are presented in figures 1 and 2.
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FIGURE 1. Absorption spectra of cis-[Cr(NH3)4(OH2)2](ClO4)3 as a function of wavelength in the presence of increasing amounts of NaOH; [complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
FIGURE 2. Plot of absorption against pH at 538 nm cis-[Cr(NH3)4(OH2)2](ClO4)3 during spectrophotometric titration.
The spectrophotometric titrations are not limited to acid–base equilibrium systems. This method gives very important information about the amount of new species for each complex ion during titration. The data associated with a spectrophotometric titration can be also presented as a set of titration spectra acquired as a function of pH (figures 3 and 4). The absorption spectra of chromium(III) complexes {cis[Cr(NH3)4(OH2)2](ClO4)3 and cis-[Cr(NH3)4(OH2)Cl]Cl2} present two characteristic maxima (493 nm; 370 nm) and one minimum of absorption (416 nm). Increasing the pH (figure 1) moves the maxima towards long waves. During titration, the absorption rises at pH range (2.00–5.53). These changes are smaller and the absorption rises significantly again for pH values about 9.23 (figure 4). Small changes, the increase of absorption, can be seen after crossing this value of pH (9.23), too. The coordination compounds of cobalt(III) {cis-[Co(NH3)4(OH2)2](ClO4)3 and cis-[Co(NH3)4(O2CO)]NO3 0.5H2O} have two maxima in the acid medium (508 nm; 352 nm) and one minimum (434 nm). During titration (figure 2) when pH changes from acidic into basic, these maxima diminish and a new maximum appears (357 nm). Moreover, during this titration, the absorption rises over the pH range 2.00 to 6.22 (figure 4) and after crossing this value, the changes of absorption are smaller. The absorption rises significantly again for pH values about 9.05. After crossing this pH value,
FIGURE 3. Absorption spectra of cis-[Co(NH3)4(OH2)2](ClO4)3 as a function of wavelength in the presence of increasing amounts of NaOH; [complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
FIGURE 4. Plot of absorption against pH to illustrate three jumps of absorption at 522 nm for cis-[Co(NH3)4(OH2)2](ClO4)3 during spectrophotometric titration.
only the small increase of absorption can be seen to the end of titration. The calculation aimed at determining the values of pK are based on the Henderson–Hasselbach [25], which was presented in Section 2.2, equation (4). To get information about number of equilibrium constants, after a few transformations, A-diagrams can be obtained [25]. The corresponding A-diagrams (figure 5) show the presence of two strictly linear segments for various wavelength combinations in the region of 300–600 nm. Within the limits of error, all data points fall on one of the two linear segments. Thus, the overall titration system is non-overlapping. Consequently, based on A-diagrams for spectrophotometric titrations of cis-½CrðNH3 Þ4 ðOH2 Þ2 ðClO4 Þ3 and cis-½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O, it can be concluded that for both compounds studied two equilibria are present. It is worth noting that the A-diagram for the spectrophotometric titration of cis-[Cr(NH3)4(OH2)Cl]Cl2 complex ion is strictly linear through the entire range of pH. At every point on this spectrometric titration curve of a single-step acid–base titration, proton-transfer equilibrium is established. The A-diagram (figure 6) for all wavelength combinations over pH range of 2–10 consists of three linear segments, i.e., the system
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A. Chylewska et al. / J. Chem. Thermodynamics 40 (2008) 1290–1294 TABLE 2 Proposed equilibrium in the systems studied Lp. 1 2 3 a
4b 5b
Equilibrium in aqueous solution 3+
pK values 2+
[Cr(NH3)4(OH2)2] + OH ¡ [Cr(NH3)4(OH2)(OH)] + H2O [Cr(NH3)4(OH2)(OH)]2+ + OH ¡ [Cr(NH3)4(OH)2]+ + H2O [Cr(NH3)4(OH2)Cl]2+ + OH ¡ [Cr(NH3)4(OH)Cl]+ + H2O [Co(NH3)4(O2CO)]+ + 2 H+ + H2O ? [Co(NH3)4(OH2)2]3+ + CO2" [Co(NH3)4(OH2)2]3+ + OH ¡ [Co(NH3)4(OH2)(OH)]2+ + H2O [Co(NH3)4(OH2)(OH)]2+ + OH ¡ [Co(NH3)4(OH)2]+ + H2O
pK1 pK2 pK1 pK1 pK2
a
The reaction of carbon dioxide removing from coordination sphere of Co(III), after addition of acid. These equations are the same for two different compounds: cis-[Co(NH3)4(O2CO)]+ and cis-[Co(NH3)4(OH2)2]3+.
b
FIGURE 5. Plot of absorbances to show the relationship between absorbances at different wavelengths for cis-[Cr(NH3)4(OH2)2](ClO4)3 ion during spectrophotometric titration.
FIGURE 6. Plot of absorbance at 318 nm against absorbance at 522 nm for coordination compound cis-[Co(NH3)4(OH2)2](ClO4)3 during alkaline titration.
activity coefficients remain constant for all the species within the experiments. The titration curves obtained for cis-[Cr(NH3)4(OH2)2](ClO4)3 (figure 7) show three potential jumps and for cis[Co(NH3)4(OH2)2](ClO4)3 (figure 8) two potential jumps. Figures 7 and 8 present the comparison of experimental data (black squares) and data obtained from calculation by using STOICHIO programme (solid line). It can be seen that correlation of these two results is in
FIGURE 7. Plot of potential against volume of titrant NaOH during the potentiometric titration of cis-[Cr(NH3)4(OH2)2](ClO4)3; [Complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
involves three non-overlapping titration steps. Thus, it can be concluded that in the solution studied three equilibria exist. Two of them are proton-transfer type; one more is the equilibrium of geometric isomerisation of the cis-[Co(NH3)4(OH2)2](ClO4)3 to trans[Co(NH3)4(OH2)2](ClO4)3 coordination compound. 3.2. Potentiometric measurements and determination of models for equilibrium models The potentiometry may be considered as the most accurate technique for the evaluation of complex equilibrium constants. The spectrophotometric titrations are useful when the former method can not be employed, e.g. when insoluble species are formed at the concentrations typical of potentiometric titrations [26]. In the present study, the acid–base equilibrium, shown in table 2, has been in investigated for the systems in acidic solutions shown in table 2. The potentiometric technique requires the selection of a medium with the constant ionic strength in order to ensure that the
FIGURE 8. Plot of potential against volume of titrant NaOH during the potentiometric titration for cis-[Co(NH3)4(OH2)2](ClO4)3 complex compound; [complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
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TABLE 3 The values of dissociation constants for all complexes studied determined at T = 298.15 K from n titrations (I = 1.0 M NaClO4) Coordination compound 3+
cis-[Cr(NH3)4(OH2) 2]
2+
cis-[Cr(NH3)4(OH2)Cl] cis-[Co(NH3)4(OH2)2]3+ cis-[Co(NH3)4(O2CO)]+ a
Spectrophotometry
Potentiometrya
pK1 pK2 pK1 pK1 pK2 pK1 pK2
pK1 pK2 pK1 pK1 pK2 pK1 pK2
4.67 ± 0.05 7.30 ± 0.04 8.18 ± 0.03 5.70 ± 0.04 7.94 ± 0.01 5.54 ± 0.04 9.48 ± 0.02
4.73 ± 0.08 7.37 ± 0.10 8.28 ± 0.12 5.73 ± 0.06 7.97 ± 0.03 5.30 ± 0.07 9.69 ± 0.02
Values of pK only for proton-transfer equilibrium.
fact very good for each compound studied. Thus it can be concluded that spectrophotometric and automated potentiometric techniques are reliable, accurate and easily applicable and they are suitable methods for the determination of the dissociation constants of complexes studied. Based on these results, pK values for four coordination compounds have been determined. Spectrophotometric and potentiometric titrations for cis-[Cr(NH3)4(OH2)2](ClO4)3, cis-[Co(NH3)4 (OH2)2](ClO4)3 and cis-½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O allowed the calculation of two acidity constants for these compounds and only one constant for cis-[Cr(NH3)4(OH2)Cl]Cl2. Table 3 presents all constant values that were determined. The analysis of the values of the dissociation constants lead to the conclusion that the exchange of chromium(III) on cobalt(III) ion, in the same surroundings of simple ligands, causes the creation of the more basic complex compound and increases the values of pK1 and pK2. The similar values of dissociation constants for cis[Co(NH3)4(OH2)2]3+ and cis-[Co(NH3)4(O2CO)]+ show conversion of carbonate complex to a compound with two molecules of water in the coordination sphere of cobalt(III). The exchange of one molecule of water on the chloride anion (cis-[Cr(NH3)4(OH2)2]3+; cis-[Cr(NH3)4(OH2)Cl]2+) in the analogous complex compound leads also to obtaining the compound with more basic properties (table 3). 4. Conclusions In this paper, we present two different and independent methods, which are useful for determination of the dissociation constants in the case of complex compounds with simple ligands. Comparison of pK values for analogous coordination compounds, cis-[Cr(NH3)4(OH2)2](ClO4)3 and cis-[Co(NH3)4(OH2)2](ClO4)3, enables us to conclude that the first one is a stronger acid than with Co(III) as a coordinate centre. This conclusion is based on the value of pK1, which is higher for cis-[Co(NH3)4(OH2)2] (ClO4)3 than for cis[Cr(NH3)4(OH2)2](ClO4)3. The fact that the values of pK1 and pK2 for cis-[Co(NH3)4(OH2)2](ClO4)3 and cis ½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O, respectively, are very similar may be due to the fact that
in both solutions the same equilibrium exists. The carbonato complex of Co(III) after reaction with acid is transformed to cis-[Co(NH3)4(OH2)2](ClO4)3. The presence of sodium carbonate causes the increase of pH in the solution studied. Therefore, the higher values of both pK1 and pK2 have been determined for cis ½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O compound than for cis-[Co(NH3)4(OH2)2](ClO4)3. Acknowledgements This work was financially supported by Polish Ministry of Scientific Research and Education (years 2006–2009), Grants 2 P05A 055 30 and BW/8000-5-0375-7, as well as Foundation for Polish Science. References [1] D.J. Hawke, K.J. Powell, S. Sjorberg, Polyhedron 14 (1995) 377–381. [2] G.A. Ibanez, G.M. Escandar, Polyhedron 17 (1998) 4433–4437. [3] M. Kadar, A. Biro, K. Toth, B. Vermes, P. Huszthy, Spectrochim. Acta Part A 62 (2005) 1032–1038. [4] R.M. Dyson, S. Kaderli, G.A. Lawrance, M. Mader, A.D. Zuberbuhler, Anal. Chim. Acta 353 (1997) 381–393. [5] J. Saurina, S. Hernandez-Cassou, Analyst 120 (1995) 305–312. [6] R.M. Dyson, G.A. Lawrance, H. Macke, M. Maeder, Polyhedron 18 (1999) 3243– 3251. [7] C.L. Araujo, G.A. Ibanez, G.N. Ledesman, G.M. Escandar, A.C. Olivieri, Computer Chem. 22 (1998) 161–168. [8] H. Gampp, D. Haspra, M. Maeder, A.D. Zuberbuhler, Anal. Chem. 62 (1990) 2220–2224. [9] A. Da˛browska, D. Jacewicz, A. Łapin´ska, B. Banecki, A. Figarski, M. Szkatuła, J. Lehman, J. Krajewski, J. Kubasik-Juraniec, M. Woz´niak, L. Chmurzyn´ski, Biochem. Biophys. Res. Comm. 326 (2005) 313–320. [10] J. Glerup, Inorg. Chem. 15 (1976) 1408–1410. [11] J.G. Yu, J.C. Yu, M.K.P. Leung, W.K. Ho, B. Cheng, X.J. Zhao, J.C. Zhao, J. Catal. 217 (2003) 69–78. [12] G. Schlessinger, Inorg. Synth. 6 (1949) 65–69. [13] W.T. Jordan, L.R. Froebe, R.A. Haines, T.D. Leah, Inorg. Synth. 18 (2007) 96–102. [14] A.I. Vogel, Vogel’s Textbook of Quantitative Inorganic Analysis, fourth ed., Longman, London, 1978. [15] G.A. Ibanez, A.C. Olivieri, G.M. Escandar, J. Chem. Soc. Faraday Trans. 93 (1997) 545–551. [16] J. Kostrowicki, A. Liwo, Comput. Chem. 8 (1984) 101–107. [17] J. Kostrowicki, A. Liwo, Comput. Chem. 11 (1987) 195–199. [18] J. Kostrowicki, A. Liwo, Talanta 87 (1990) 645–650. [19] D.W. Marquard, J. Soc. Indust. Appl. Math. 11 (1963) 431–441. [20] E. Kaczmarczyk, K. Maj, L. Chmurzyn´ski, J. Chem. Thermodyn. 32 (2000) 901– 910. [21] Ł. Gurzyn´ski, A. Puszko, M. Makowski, J. Makowska, L. Chmurzyn´ski, J. Chem. Thermodyn. 38 (2006) 1584–1591. [22] Ł. Gurzyn´ski, A. Puszko, M. Makowski, L. Chmurzyn´ski, J. Chem. Thermodyn. 39 (2007) 1272–1278. [23] S. Ebel, W. Parzefall, Experimetelle Einführung in die Potentiometrie, Verlag Chemie, Weinheim, 1975. [24] G.L. Ellmann, Arch. Biochem. Biophys. 82 (1959) 70–77. [25] J. Polster, H. Lachmann, Spectroscopic Titrations: Analysis of Chemical Equlibria, VCH, 1992. [26] A.E. Martell, R.J. Motekaitis, The Determination and Use of Stability Constants, second ed., VCH, New York, 1992 (Chapter 3).
JCT 08-26
Contents lists available at ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Determination of dissociation constants for coordination compounds of Cr(III) and Co(III) using potentiometric and spectrophotometric methods A. Chylewska, D. Jacewicz, D. Zarzeczan´ska, L. Chmurzyn´ski * ´ sk, Sobieskiego 18/19, 80-952 Gdan ´ sk, Poland Department of Chemistry, University of Gdan
a r t i c l e
i n f o
Article history: Received 10 January 2008 Received in revised form 11 March 2008 Accepted 19 March 2008 Available online 29 March 2008 Keywords: Complexes of Cr(III) Complexes of Co(III) Acidity constants Potentiometry Spectrophotometry
a b s t r a c t The acid–base properties of analogous complex ions of chromium(III) and cobalt(III) in aqueous solution have been studied. The equilibrium constants for all metal complexes were determined by using potentiometric and spectrophotometric titration methods. First, dissociation constants for the studied complexes of Cr(III) and Co(III) were determined by means of the potentiometric titration method and using the STOICHIO computer programme. Then, pH-spectrophotometric titrations were performed and the OriginPro 7.5 computer programme was used to calculate the same constants. The measurements using both methods were carried out under the same conditions of temperature, T = 298.15 K, and over the same pH range 2.00–10.00, respectively. It turned out that the two methods used enabled us to obtain acidity constants in very good agreement. Ó 2008 Elsevier Ltd. All rights reserved.
1. Introduction Potentiometric and spectrophotometric titrations are the most useful methods to investigate equilibrium in solutions and to determine acid–base constants [1–3]. The potentiometric titration is used so often due to the simplicity of equipment and minimal time requirements [4]. However, this method does not cover all aspects of solution chemistry in which researchers have an interest. To have complete information about the species formed during the titration, spectrophotometric titrations are usually carried out simultaneously [5–7]. This technique shows how many equilibria exist in solution studied. Moreover, it can be applied to structural analysis of compounds and is neither limited to any pH range nor to aqueous solutions. Additionally, it is useful in discriminating between different models of equilibrium [8]. All these aspects prove that the spectroscopy is more powerful than the potentiometric method; however, it is advisable that they should be used together for a more detailed determination of acid–base interactions. In both types of titration, a controlled displacement of chemical equilibrium occurs during the addition of titrant. The determination of the constants of dissociation of the complex compounds of chromium(III) and cobalt(III) ions is extremely
* Corresponding author. Address: Department of General and Inorganic Chemistry, University of Gdan´sk, Poland. Tel.: +48 58 523 53 19; fax: +48 58 523 54 72. E-mail address: [email protected] (L. Chmurzyn´ski). 0021-9614/$ - see front matter Ó 2008 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2008.03.012
important in investigations executed by our research group. Based on results from the spectrophotometric and potentiometric measurements of analogous compounds, it can be concluded which complexes are more stable and active in biological and chemical materials. It is worth pointing that previously we synthesized the compound of chromium(III) with the aminodeoxysugar as ligand; cis-[Cr(C2O4)(AaraNH2)(H2O)2]+, where AaraNH2 stands for methyl 3-amino-2,3-dideoxy-a-D-arabino-hexapyranoside. This complex was so reactive that it caught the free radical of nitrogen dioxide generated from the pancreas of rats [9]. The exchange of the folded aminodeoxysugar’s ligand on simple ammonia molecules or the chloride anions can facilitate the reaction occurring for any gas (CO2, NO2 or SO2) uptake by coordination compounds studied. In our study, the complex compounds cis-[Cr(NH3)4(OH2)2](ClO4)3, cis-[Cr(NH3)4(OH2)Cl]Cl2, cis-[Co(NH3)4(OH2)2](ClO4)3 and cis½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O were used. We assumed that these analogous coordination compounds of Cr(III) and Co(III) would reveal similar acid–base properties. The only difference between the analogues species is the coordination centre. Based on determined acidity constants of the compounds studied, it is possible to define what impact the transition metal ion has on properties of complex compounds. In the present paper, the correlation between results obtained by two different experimental methods has been discussed. Regardless of the method, the determined values of the dissociation constants for the coordination compounds studied are similar, which suggests that these values are correct.
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2. Experimental 2.1. Materials 2.1.1. The synthesis of coordination compounds studied The complex compound of cis-[Cr(NH3)4(OH2)Cl]Cl2 was prepared according to the method of Glerup [10]. The synthesis of cis-[tetraamminediaquachromium(III)] perchlorate was described in paper [11]. The preparation of compound with cobalt(III) as the coordination centre binding four molecules of ammonia and bidentate carbonate ion was reported in [12]. The coordination compound, cis-[Co(NH3)4(OH2)2](ClO4)3, was produced from cis-[Co(NH3)4(O2CO)]NO3 0.5H2O according to the procedure described in reference [13]. All the coordination compounds studied were synthesized from commercial reagents (see table 1). 2.2. Instrumentation and measurement The measurements were carried out by means of potentiometric titrations at constant ionic strength using an automated system and applying the Microtitator programme. All probes, which were used in titrations, were prepared in nitrogen atmosphere to avoid CO2 contamination and the temperature was kept at (298.15 ± 0.1) K. A constant ionic strength of 0.1 M was maintained with NaClO4. The titration system consisted of a titration cell, a magnetic stirrer, and an automatic titrator with Hamilton’s syringe (0.5 cm3). The pH combined electrode was bought from the Metrohm firm. The electrode was calibrated using pH standard buffers [14]. The Perkin Elmer Lambda 800 UV–Vis double beam spectrophotometer, with automatic stirrer, was used for absorbance measurements with 1.00 cm quartz microcells T = 298.15 K and 0.1 M ionic strength (NaClO4). For each pH point, a known aliquot of solution was extracted and the absorption spectrum was recorded. The method of spectrophotometric determinations of acid–base constants was described in details elsewhere [15]. The profile of absorbances obtained at the wavelength of the absorption maximum versus pH in each run was used to obtain the equilibrium constants of the complex species formed in each system. 2.3. Determination of pK values by using potentiometry To calculate the acid–base constants from the potentiometric method, the STOICHIO computer programme was used. This programme is based on the algorithm of Kostrowicki and Liwo [16–18] and is able to handle any model of chemical equilibrium. The programme determines the equilibrium constants as parameters of the function to be minimized: U¼
ny n X X 1 b 1 b 1 2 ^i yi Þ2 ; ½ E i EðV i ; V; y; xÞ2 þ 2 ð V ðy i V iÞ þ 2 2 r r r yi E V i¼1 i¼1
ð1Þ where b E denotes the measured potential difference for the ith titration point, V is the measured volume of the solution at the ith titration point, vector x refers to unknown equilibrium constants and TABLE 1 Purity of the coordination compounds of transition metals checked by the elemental analysis Coordination compound
Calculated
Theoretical
cis-[Cr(NH3)4(OH2)2](ClO4)3 cis-[Co(NH3)4(OH2)2](ClO4)3 cis-[Cr(NH3)4(OH2)Cl]Cl2 cis-[Co(NH3)4(O2CO)]NO3 0.5H2O
12.35%N, 3.52%H 2.17%N, 3.53%H 22.86%N, 5.70%H 27.16%N, 5.10%H
12.33%N, 12.14%N, 22.90%N, 27.10%N,
For comparison theoretical values are included.
3.55%H 3.49%H 5.73%H 5.03%H
vector y represents calibration parameters, total concentrations of reagents in the stock solution, and titrants, etc. The minimization of the sum of the squares given in equation (1) is done using Marquardt’s method [19]. This method takes into account not only errors in the e.m.f., but also those in titration volume, stock-solution preparation, electrode calibration parameters, and reagent impurities, in addition to the equilibrium constants determined from other measurements [20]. The details of the calculations of equilibrium constants have been described previously [21,22]. 2.4. Determination of pK values by using spectrophotometry Based on the Lambert–Beer–Bouger law for spectrophotometry, viz., equation (2) Ak ¼ ekBH b0 þ ðekB ekBH Þ b ¼ ekB b0 þ ðekBH ekB Þb h;
ð2Þ
where b0 is the initial concentration of complex studied, [BH]; b is the concentration of deprotonated form of complex studied, [B]; h is the concentration of hydrogen ions; e is the molar absorptivity of the absorbing species; AkBH is the absorbance at the ‘‘acid end” of the titration (bh is the b0); AkB is absorbance at the ‘‘alkaline end” of the titration (b is the b0), one can obtain the Henderson–Hasselbach equation in the form of equation (3) pK 1 ¼ pH ln
b Ak AkBH ¼ pH ln : bh AkB Ak
ð3Þ
After transformation of equation (3) according to [23] to equation (4) Ak ¼
AkBH 10pH þ AkB 10pK 1 10pH þ 10pK 1
;
ð4Þ
this new form of expression was put to OriginPro 7.5 programme to calculate values of dissociation constants of complexes studied. To check how many dissociation constants we should calculate and how many equilibria existed in the solution studied, we had to construct relationships between absorbances at two, different wavelengths, which are called A-diagrams. These diagrams, created from the data from spectrophotometric titrations are criteria by which the presence of measurable overlap can be observed. These diagrams [24] constructed for all wavelength combinations must be linear in the presence of equilibrium conditions. For example, A-diagram is described always by a linear equation (5) for single step acid–base titration: A1 ¼
e1B e1BH A2 þ ðe1BH e2BH Þ b0 : e2B e2BH
ð5Þ
It should be pointed that the intersection point of two linear segments of the A-diagram corresponds to the presence of the deprotonated ampholytic form in the case when two equilibria are present in solution. All the titrations were performed under complete computer control. 3. Results and discussion 3.1. Spectrophotometric and pH-metric measurements The course of spectrophotometric titrations for all the coordination compounds studied cis-[Cr(NH3)4(OH2)2]3+; cis-[Cr(NH3)4(OH2)Cl]2+; cis-[Co(NH3)4(OH2)2]3+; cis-[Co(NH3)4(O2CO)]+ were registered. The identical conditions of measurements, which were applied in the potentiometric method, were kept. The results obtained are presented in graphical form as the relationship between absorbance and wavelength. Thus the relationships between absorbance (A) and wavelength (k), for two of coordination compound studied are presented in figures 1 and 2.
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FIGURE 1. Absorption spectra of cis-[Cr(NH3)4(OH2)2](ClO4)3 as a function of wavelength in the presence of increasing amounts of NaOH; [complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
FIGURE 2. Plot of absorption against pH at 538 nm cis-[Cr(NH3)4(OH2)2](ClO4)3 during spectrophotometric titration.
The spectrophotometric titrations are not limited to acid–base equilibrium systems. This method gives very important information about the amount of new species for each complex ion during titration. The data associated with a spectrophotometric titration can be also presented as a set of titration spectra acquired as a function of pH (figures 3 and 4). The absorption spectra of chromium(III) complexes {cis[Cr(NH3)4(OH2)2](ClO4)3 and cis-[Cr(NH3)4(OH2)Cl]Cl2} present two characteristic maxima (493 nm; 370 nm) and one minimum of absorption (416 nm). Increasing the pH (figure 1) moves the maxima towards long waves. During titration, the absorption rises at pH range (2.00–5.53). These changes are smaller and the absorption rises significantly again for pH values about 9.23 (figure 4). Small changes, the increase of absorption, can be seen after crossing this value of pH (9.23), too. The coordination compounds of cobalt(III) {cis-[Co(NH3)4(OH2)2](ClO4)3 and cis-[Co(NH3)4(O2CO)]NO3 0.5H2O} have two maxima in the acid medium (508 nm; 352 nm) and one minimum (434 nm). During titration (figure 2) when pH changes from acidic into basic, these maxima diminish and a new maximum appears (357 nm). Moreover, during this titration, the absorption rises over the pH range 2.00 to 6.22 (figure 4) and after crossing this value, the changes of absorption are smaller. The absorption rises significantly again for pH values about 9.05. After crossing this pH value,
FIGURE 3. Absorption spectra of cis-[Co(NH3)4(OH2)2](ClO4)3 as a function of wavelength in the presence of increasing amounts of NaOH; [complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
FIGURE 4. Plot of absorption against pH to illustrate three jumps of absorption at 522 nm for cis-[Co(NH3)4(OH2)2](ClO4)3 during spectrophotometric titration.
only the small increase of absorption can be seen to the end of titration. The calculation aimed at determining the values of pK are based on the Henderson–Hasselbach [25], which was presented in Section 2.2, equation (4). To get information about number of equilibrium constants, after a few transformations, A-diagrams can be obtained [25]. The corresponding A-diagrams (figure 5) show the presence of two strictly linear segments for various wavelength combinations in the region of 300–600 nm. Within the limits of error, all data points fall on one of the two linear segments. Thus, the overall titration system is non-overlapping. Consequently, based on A-diagrams for spectrophotometric titrations of cis-½CrðNH3 Þ4 ðOH2 Þ2 ðClO4 Þ3 and cis-½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O, it can be concluded that for both compounds studied two equilibria are present. It is worth noting that the A-diagram for the spectrophotometric titration of cis-[Cr(NH3)4(OH2)Cl]Cl2 complex ion is strictly linear through the entire range of pH. At every point on this spectrometric titration curve of a single-step acid–base titration, proton-transfer equilibrium is established. The A-diagram (figure 6) for all wavelength combinations over pH range of 2–10 consists of three linear segments, i.e., the system
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4b 5b
Equilibrium in aqueous solution 3+
pK values 2+
[Cr(NH3)4(OH2)2] + OH ¡ [Cr(NH3)4(OH2)(OH)] + H2O [Cr(NH3)4(OH2)(OH)]2+ + OH ¡ [Cr(NH3)4(OH)2]+ + H2O [Cr(NH3)4(OH2)Cl]2+ + OH ¡ [Cr(NH3)4(OH)Cl]+ + H2O [Co(NH3)4(O2CO)]+ + 2 H+ + H2O ? [Co(NH3)4(OH2)2]3+ + CO2" [Co(NH3)4(OH2)2]3+ + OH ¡ [Co(NH3)4(OH2)(OH)]2+ + H2O [Co(NH3)4(OH2)(OH)]2+ + OH ¡ [Co(NH3)4(OH)2]+ + H2O
pK1 pK2 pK1 pK1 pK2
a
The reaction of carbon dioxide removing from coordination sphere of Co(III), after addition of acid. These equations are the same for two different compounds: cis-[Co(NH3)4(O2CO)]+ and cis-[Co(NH3)4(OH2)2]3+.
b
FIGURE 5. Plot of absorbances to show the relationship between absorbances at different wavelengths for cis-[Cr(NH3)4(OH2)2](ClO4)3 ion during spectrophotometric titration.
FIGURE 6. Plot of absorbance at 318 nm against absorbance at 522 nm for coordination compound cis-[Co(NH3)4(OH2)2](ClO4)3 during alkaline titration.
activity coefficients remain constant for all the species within the experiments. The titration curves obtained for cis-[Cr(NH3)4(OH2)2](ClO4)3 (figure 7) show three potential jumps and for cis[Co(NH3)4(OH2)2](ClO4)3 (figure 8) two potential jumps. Figures 7 and 8 present the comparison of experimental data (black squares) and data obtained from calculation by using STOICHIO programme (solid line). It can be seen that correlation of these two results is in
FIGURE 7. Plot of potential against volume of titrant NaOH during the potentiometric titration of cis-[Cr(NH3)4(OH2)2](ClO4)3; [Complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
involves three non-overlapping titration steps. Thus, it can be concluded that in the solution studied three equilibria exist. Two of them are proton-transfer type; one more is the equilibrium of geometric isomerisation of the cis-[Co(NH3)4(OH2)2](ClO4)3 to trans[Co(NH3)4(OH2)2](ClO4)3 coordination compound. 3.2. Potentiometric measurements and determination of models for equilibrium models The potentiometry may be considered as the most accurate technique for the evaluation of complex equilibrium constants. The spectrophotometric titrations are useful when the former method can not be employed, e.g. when insoluble species are formed at the concentrations typical of potentiometric titrations [26]. In the present study, the acid–base equilibrium, shown in table 2, has been in investigated for the systems in acidic solutions shown in table 2. The potentiometric technique requires the selection of a medium with the constant ionic strength in order to ensure that the
FIGURE 8. Plot of potential against volume of titrant NaOH during the potentiometric titration for cis-[Co(NH3)4(OH2)2](ClO4)3 complex compound; [complex] = 0.0010 M, [NaOH] = 0.0494 M, [HCl] = 0.00393 M.
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TABLE 3 The values of dissociation constants for all complexes studied determined at T = 298.15 K from n titrations (I = 1.0 M NaClO4) Coordination compound 3+
cis-[Cr(NH3)4(OH2) 2]
2+
cis-[Cr(NH3)4(OH2)Cl] cis-[Co(NH3)4(OH2)2]3+ cis-[Co(NH3)4(O2CO)]+ a
Spectrophotometry
Potentiometrya
pK1 pK2 pK1 pK1 pK2 pK1 pK2
pK1 pK2 pK1 pK1 pK2 pK1 pK2
4.67 ± 0.05 7.30 ± 0.04 8.18 ± 0.03 5.70 ± 0.04 7.94 ± 0.01 5.54 ± 0.04 9.48 ± 0.02
4.73 ± 0.08 7.37 ± 0.10 8.28 ± 0.12 5.73 ± 0.06 7.97 ± 0.03 5.30 ± 0.07 9.69 ± 0.02
Values of pK only for proton-transfer equilibrium.
fact very good for each compound studied. Thus it can be concluded that spectrophotometric and automated potentiometric techniques are reliable, accurate and easily applicable and they are suitable methods for the determination of the dissociation constants of complexes studied. Based on these results, pK values for four coordination compounds have been determined. Spectrophotometric and potentiometric titrations for cis-[Cr(NH3)4(OH2)2](ClO4)3, cis-[Co(NH3)4 (OH2)2](ClO4)3 and cis-½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O allowed the calculation of two acidity constants for these compounds and only one constant for cis-[Cr(NH3)4(OH2)Cl]Cl2. Table 3 presents all constant values that were determined. The analysis of the values of the dissociation constants lead to the conclusion that the exchange of chromium(III) on cobalt(III) ion, in the same surroundings of simple ligands, causes the creation of the more basic complex compound and increases the values of pK1 and pK2. The similar values of dissociation constants for cis[Co(NH3)4(OH2)2]3+ and cis-[Co(NH3)4(O2CO)]+ show conversion of carbonate complex to a compound with two molecules of water in the coordination sphere of cobalt(III). The exchange of one molecule of water on the chloride anion (cis-[Cr(NH3)4(OH2)2]3+; cis-[Cr(NH3)4(OH2)Cl]2+) in the analogous complex compound leads also to obtaining the compound with more basic properties (table 3). 4. Conclusions In this paper, we present two different and independent methods, which are useful for determination of the dissociation constants in the case of complex compounds with simple ligands. Comparison of pK values for analogous coordination compounds, cis-[Cr(NH3)4(OH2)2](ClO4)3 and cis-[Co(NH3)4(OH2)2](ClO4)3, enables us to conclude that the first one is a stronger acid than with Co(III) as a coordinate centre. This conclusion is based on the value of pK1, which is higher for cis-[Co(NH3)4(OH2)2] (ClO4)3 than for cis[Cr(NH3)4(OH2)2](ClO4)3. The fact that the values of pK1 and pK2 for cis-[Co(NH3)4(OH2)2](ClO4)3 and cis ½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O, respectively, are very similar may be due to the fact that
in both solutions the same equilibrium exists. The carbonato complex of Co(III) after reaction with acid is transformed to cis-[Co(NH3)4(OH2)2](ClO4)3. The presence of sodium carbonate causes the increase of pH in the solution studied. Therefore, the higher values of both pK1 and pK2 have been determined for cis ½CoðNH3 Þ4 ðO2 COÞNO3 0:5H2 O compound than for cis-[Co(NH3)4(OH2)2](ClO4)3. Acknowledgements This work was financially supported by Polish Ministry of Scientific Research and Education (years 2006–2009), Grants 2 P05A 055 30 and BW/8000-5-0375-7, as well as Foundation for Polish Science. References [1] D.J. Hawke, K.J. Powell, S. Sjorberg, Polyhedron 14 (1995) 377–381. [2] G.A. Ibanez, G.M. Escandar, Polyhedron 17 (1998) 4433–4437. [3] M. Kadar, A. Biro, K. Toth, B. Vermes, P. Huszthy, Spectrochim. Acta Part A 62 (2005) 1032–1038. [4] R.M. Dyson, S. Kaderli, G.A. Lawrance, M. Mader, A.D. Zuberbuhler, Anal. Chim. Acta 353 (1997) 381–393. [5] J. Saurina, S. Hernandez-Cassou, Analyst 120 (1995) 305–312. [6] R.M. Dyson, G.A. Lawrance, H. Macke, M. Maeder, Polyhedron 18 (1999) 3243– 3251. [7] C.L. Araujo, G.A. Ibanez, G.N. Ledesman, G.M. Escandar, A.C. Olivieri, Computer Chem. 22 (1998) 161–168. [8] H. Gampp, D. Haspra, M. Maeder, A.D. Zuberbuhler, Anal. Chem. 62 (1990) 2220–2224. [9] A. Da˛browska, D. Jacewicz, A. Łapin´ska, B. Banecki, A. Figarski, M. Szkatuła, J. Lehman, J. Krajewski, J. Kubasik-Juraniec, M. Woz´niak, L. Chmurzyn´ski, Biochem. Biophys. Res. Comm. 326 (2005) 313–320. [10] J. Glerup, Inorg. Chem. 15 (1976) 1408–1410. [11] J.G. Yu, J.C. Yu, M.K.P. Leung, W.K. Ho, B. Cheng, X.J. Zhao, J.C. Zhao, J. Catal. 217 (2003) 69–78. [12] G. Schlessinger, Inorg. Synth. 6 (1949) 65–69. [13] W.T. Jordan, L.R. Froebe, R.A. Haines, T.D. Leah, Inorg. Synth. 18 (2007) 96–102. [14] A.I. Vogel, Vogel’s Textbook of Quantitative Inorganic Analysis, fourth ed., Longman, London, 1978. [15] G.A. Ibanez, A.C. Olivieri, G.M. Escandar, J. Chem. Soc. Faraday Trans. 93 (1997) 545–551. [16] J. Kostrowicki, A. Liwo, Comput. Chem. 8 (1984) 101–107. [17] J. Kostrowicki, A. Liwo, Comput. Chem. 11 (1987) 195–199. [18] J. Kostrowicki, A. Liwo, Talanta 87 (1990) 645–650. [19] D.W. Marquard, J. Soc. Indust. Appl. Math. 11 (1963) 431–441. [20] E. Kaczmarczyk, K. Maj, L. Chmurzyn´ski, J. Chem. Thermodyn. 32 (2000) 901– 910. [21] Ł. Gurzyn´ski, A. Puszko, M. Makowski, J. Makowska, L. Chmurzyn´ski, J. Chem. Thermodyn. 38 (2006) 1584–1591. [22] Ł. Gurzyn´ski, A. Puszko, M. Makowski, L. Chmurzyn´ski, J. Chem. Thermodyn. 39 (2007) 1272–1278. [23] S. Ebel, W. Parzefall, Experimetelle Einführung in die Potentiometrie, Verlag Chemie, Weinheim, 1975. [24] G.L. Ellmann, Arch. Biochem. Biophys. 82 (1959) 70–77. [25] J. Polster, H. Lachmann, Spectroscopic Titrations: Analysis of Chemical Equlibria, VCH, 1992. [26] A.E. Martell, R.J. Motekaitis, The Determination and Use of Stability Constants, second ed., VCH, New York, 1992 (Chapter 3).
JCT 08-26