Determination of Stability Constants of Copper(II)-GlycineComplex in Mixed Solvents by Copper(II)-Selective Electrode

Jing Fan,* Xuejing Shen, and Jianji Wang

School of Chemical and Environmental Sciences, Henan Normal University, Xinxiang, Henan, 453002, P. R. China

e-mail: jwang@mail.henannu.edu.cn

Received: August 2, 2000

Final version: December 6, 2000

Abstract

The first stability constants (in logb1 form) of the Cu(II)-glycine complex have been determined at 25 �C and ionic strength of 0.10 mol=Lin water and in mixed aqueous solutions of methanol (MeOH), ethanol (EtOH), dimethylsulfoxide (DMSO), N,N-dimethylformamide(DMF) and 1,4-dioxane (DOX) from pH and pCu measurements of cells containing a copper(II) ion-selective electrode. In general,the stability constants of the complex increase with increasing composition of the co-solvents in the order: DOX>DMF>DMSO>EtOH>MeOH>H2O. An approximate linear relationship between logb1 and the mole fraction of the co-solvent was observedfor the complex in all the mixed solvents except in DMFþH2O mixtures. The response of the ion selective electrode to copper(II) ion in themixed solvents was also investigated. The effects of solvent on the complex stability are discussed in terms of the polarity of the solvents.

Keywords: Ion-selective electrode, Mixed solvent, Stability constant, Glycine, Metal complex

1. Introduction

There has been an increased interest in the use of ion selectiveelectrodes (ISE) in nonaqueous and mixed solvents since the1980’s [1]. Among the recent work based on both thermo-

dynamic and analytical studies in this respect, some areconcerned with the determination of ionic solvation parameters[2–10], some are devoted to the determination of activity

coefficients of electrolyte [11–13], whereas others are interestedin the response of ISE to ions and its selectivity coefficients innonaqueous and=or mixed solvents [4, 7, 14–16]. In spite ofthese publications, the application of ISE to nonaqueous and

mixed solvents was rather limited.As a part of our continuing interest in the use of ISE in mixed

solvents [8, 9, 12, 17], we now present the result of the

first stability constants for the Cu(II) complex of glycine at 25 �Cand ionic strength I¼ 0.1 mol=L in water and in mixed aqueoussolutions of methanol (MeOH), ethanol (EtOH), dimethyl

sulfoxide (DMSO), N,N-dimethylformamide (DMF), or1,4-dioxane (DOX) determined from pH and pCu measurementsof cells containing Cu(II) ion solid membrane electrode. The

solvents were chosen to examine how such properties asdielectric constant, solvent polarity and tendency to solvationmight influence the stability of the complex. The response of theCu(II) ion selective electrode in these mixed solvents was also

investigated.

2. Experimental

2.1. Reagents and Solutions

MeOH, EtOH, DMSO, DMF, DOX and metallic copper(99.99 %) were obtained from Shanghai Chem. Reagent Co.,

China. Glycine, potassium nitrate (KNO3) and potassiumhydroxide (KOH) were purchased from Beijing Chem. ReagentCo., China. All chemicals were of analytical reagent grade unless

otherwise indicated. The organic solvents were used after dryingover 4A type molecular sieves. Glycine was recrystallized fromaqueous solution of ethanol and dried under vacuum. Other

chemicals were used as received. Copper(II) nitrate solution was

prepared from metallic copper. Stock solutions of copper(II)nitrate (0.1389 mol=L) and of glycine (0.0100 mol=L) were madein appropriate solvents. Generally, test solutions were prepared

by dilution of the respective stock solutions. The mixed solvents(waterþ organic solvent) were prepared by weight. The ionicstrength in all solutions was maintained constant at 0.1 mol=L byusing potassium nitrate as a background electrolyte. Deionized

and redistilled water was used throughout the experiment.

2.2. Apparatus

The potentiometric titrations were conducted at 25� 0.05 �Cin a water jacketed glass cell described previously [17, 18]. AJiangsu copper(II) ion solid membrane electrode (Model 306)

with CuSþAg2S mixture as active material was used togetherwith a Jiangsu saturated calomel reference electrode (Model 801)for measurements of Cu(II) ion activity in solution; the pH wasmeasured with a Shanghai pH glass electrode (Model 231)

against the same reference electrode. The cell potentials and thepH values were recorded, respectively, by means of a Chengduprecise pH-meter (Model pHs-2D). Before use, the surface of the

active membrane for Cu(II) ion selective electrode was polishedwith electrode polishing powder, then soaked in a 0.01 mol=Lcopper(II) nitrate solution; the glass electrode was immersed in a

mixed solvent for several days and then conditioned overnight in0.1 mol=L hydrochloric acid after washing with distilled water.Before measurements, the glass electrode was calibrated againstaqueous standard buffers. The Nernstian response of the Cu(II)

ion selective electrode was checked at constant ionic strength inaqueous copper(II) nitrate solutions.

2.3. Procedure

A series of standard solutions of copper(II) nitrate from 10ÿ2

to 10ÿ6 mol=L in a given solvent was prepared by successive

dilution of the respective stock solution with 0.1 mol=L potas-sium nitrate. Potentials (E) of copper(II) ion selective electrodewere measured in each of standard solutions. Readings were

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Electroanalysis 2001, 13, No. 13 # WILEY-VCH Verlag GmbH, D-69469 Weinheim, 2001 1040-0397/01/1309–1115 $17.50þ.50=0

taken when potential was constant within 0.5 mV for at least5 min. Then, a standard curve of pCu against the measured E was

plotted for each given solvent.A 25.00 mL glycine solution (0.0010 mol=L) and a 12.50 mL

copper(II) nitrate solution (0.0010 mol=L) were added to the cell.After thermal equilibrium was reached, the cell solution was

titrated with successive additions of 0.10 mol=L potassiumhydroxide up to pH > 6.3. The pH value and potential of the Cu(II)ion selective electrode were recorded after each titration. During

the pH and potential measurements, the test solution was stirredmagnetically. From the potential of the Cu(II) ISE, values of pCu insolution can be found from the standard pCu vs. E curves.

Furthermore, the measured pH values in the mixed solvents havebeen converted to hydrogen ion concentrations by followingclosely the procedure proposed by Van Vitiate and Haas [19].

It was observed that in order to obtain steady and reproducibleresults, thesolventfor0.1 mol=Lpotassiumnitrate intheexternalsaltbridge of the saturated calomel electrode must be the same as that forthe reaction solution. For example, if the reaction medium is an

aqueous solution of 44.1 wt% MeOH, a 0.1 mol=L potassium nitratein this mixed solvent would be the right filling solution. Otherwise, itis difficult to obtain stable values of pH and cell potential.

The response of copper(II) ion electrode in mixed solvents wasinvestigated over a range of concentrations of copper(II) ionextending typically from 10ÿ2 to 10ÿ6 mol=L. The procedure

used was similar to that described by Coetzee and Istone [5]. Theslope and linear range for copper(II) ion electrode in differentsolvents can be obtained from the standard pCu vs. E curvesmentioned above.

3. Results and Discussion

3.1. Determination of the First Stability Constants

The first stability constant of copper(II)-glycine complexreported are for the reaction

Cu2þ þ Lÿ Ð CuLþ

and are defined as

b1 ¼½CuLþ�

½Cu2þ�½Lÿ�

ð1Þ

where Lÿ stands for the conjugate base (H2NCH2COOÿ) ofglycine (H3NþCH2COOÿ), and where

½Lÿ� ¼C0

L ÿ ½CuLþ�

aLðHÞ

ð2Þ

and

aLðHÞ ¼ 1þ½Hþ�

K2

ð3Þ

here C0L is the initial concentration of glycine in the reaction

solution, aL(H) the side reaction coefficient of glycine in reaction

with hydrogen ion, and K2 the second dissociation constant ofglycine under given conditions.

Introducing Equations 2 and 3 into Equation 1 and takinglogarithms, it follows that

log b1 ¼ pCuþ logð1þ ½Hþ�=K2Þ

ÿ logfðC0L ÿ ½CuLþ�Þ=½CuLþ�g ð4Þ

Since C0L was known exactly in a given experiment and [Cu2þ]

and [Hþ] were determinable experimentally, the value of [CuLþ]

in Equation 4 can be calculated from the initial and freeconcentrations of copper(II).

Therefore, the final parameter required to calculate logb1 is thesecond dissociation constant K2 for glycine. Because the ionicstrength was kept at 0.1 mol=L in the present work, K2 valuesinvolved in Equation 4 should be those at I¼ 0.1 mol=L. Values

of K2 (I¼ 0.1 mol=L) for glycine in H2OþDMSO, H2OþDMFand H2OþDOX mixed solvents can be found directly fromliterature [20, 21]. However, values in H2O, H2OþMeOH and

H2OþEtOH solvents were reported [22–24] as the thermo-dynamic second dissociation constants ðKa2Þ. They havebeen corrected to the values at I¼ 0.1 mol=L using following

equation

K2 ¼Ka2

g2�

ð5Þ

with mean activity coefficients (g� ) calculated from DaviesEquation [25]

log g� ¼ÿAI1=2

1þ I1=2þ 0:3AI ð6Þ

where A is the Debye-Huckel constant given by

A ¼1:82646106

ðDT Þ3=2

ð7Þ

The required values of dielectric constant D for the mixed

solvents were interpolated from previous studies [26].The potentiometric titrations were carried out in pH 4–6.5 for

each experiment. Because of the limited stability of the 1:1(CuLþ) complex in acid solutions (pH< 5) and a possible

formation of a 1:2 (CuL2) complex when pH> 6 [17], only theexperimental data in the range of pH 5–6 were used to calculatethe stability constants in all cases. The final result is the average

of all the determinations in this pH range. As an example, Table 1lists the observed potentials of the copper(II) ion electrode, pCu,pH and the calculated logb1 values for the complexation in

aqueous solution of 44.1 wt% MeOH, along with the standarddeviation and coefficient of variation for logb1. Values of logb1

for the complex in water and in the mixed solvents are given inTable 2. Comparison with literature values, whenever available, is

also included in this table.The value of logb1 in water obtained in this work was 8.18,

which is in excellent agreement with the values 8.07, 8.15 and

8.36 reported previously [27–29]. Considering the difference(DpK2¼ 0.11 for the solvent containing 20.48 wt% DOX) in K2

Table 1. pH and pCu values for the determination of the first stabilityconstant of the Cu(II) complex of glycine in aqueous solution of44.1 wt% MeOH (25 �C, I¼ 0.1).

pHobs pHcor E(mV) pCu logb1

5.05 4.82 ÿ173.0 4.39 8.995.20 4.97 ÿ169.6 4.49 8.975.36 5.13 ÿ165.0 4.64 8.985.49 5.26 ÿ161.8 4.73 8.965.58 5.35 ÿ159.4 4.83 8.965.69 5.46 ÿ155.9 4.94 8.985.82 5.59 ÿ152.1 5.04 8.96

Mean 8.97Standard deviation 0.03Coefficient of variation 0.32 %

1116 J. Fan et al.

Electroanalysis 2001, 13, No. 13

used in the calculation, our logb1 values in aqueous solutions of

10.0 and 20.48 wt% DOX also agree very well (see Table 2) withthose reported by Zelano and co-workers [30]. No stabilityconstant data on the complex in H2OþMeOH, H2OþEtOH,

H2OþDMSO and H2OþDMF solvents has been reported in theliterature to the best of our knowledge.

3.2. Response of the Electrode

The potential response of the copper(II) ion selective elec-trode in water and in the mixed solvents was examined atconstant ionic strength of 0.1 mol=L and pH 4–6. The electrode

usually reached a steady state within 3 min. in response tochanges in the ion concentrations. The response of the electrodein water was faster than that in the mixed solvents. The slope

and the linear range for the electrode are summarized in Table 3.It can be seen that in the composition range investigated, theresponse was linear and Nernstian at copper(II) ion concentra-tion from 10ÿ6 (or 10ÿ5 for DOX co-solvent) to 10ÿ2. At higher

concentrations, the response became increasingly sub-Nernstianwith increasing concentration. At lower concentrations, theresponse became increasingly sub-Nernstian with decreasing

concentration. Very similar results have been reported byCoetree and Istone [5] from a copper(II) ion selective electrodein several mixed solvents.

It is noted that the slope of the response varies slightly with thecomposition of the co-solvent in the composition range indicatedin Table 3. However, when the co-solvent composition goesbeyond this range, the slope and linear range for the electrode

decrease rapidly with increasing composition of the co-solvent.Furthermore, when 30 wt% or more DOX was presented in themixed solvent, the potentials were never steady. A similar

phenomenon has been observed by Cheng et al. [16] for lead(II)ion selective electrode in this mixed solvent.

In addition, we tried to examine the response of the copper(II)

ion electrode in H2OþAN (acetonitrile) mixed solvent in thesame way as above. Unfortunately, the potentials drift seriouslyand no reproducible results can be obtained even in aqueous

solution of 10 wt% AN. So, no further studies were carried out inthese mixed solvents.

3.3. Solvent Effect on the Complex Stability

It is evident from Table 2 that for copper-glycine complex,values of the first stability constants (logb1) in the mixed solventsare greater than that in water, and they increase with increasing

composition of the co-solvent in the order: DOX>DMF>DMSO>EtOH>MeOH>H2O. In their stability studies oncertain metal complexes (nickel(II), zinc(II) and manganese(II))with glycine in aqueous mixed solvents, Mui and McBryde [21]

found a similar solvent effect at lower composition of the co-solvent.

The relationships between logb1 and the mole fraction (X2) of

the co-solvent in the mixed solvents are illustrated in Figure 1. Ascan be seen, there is actually a linear relationship between logb1

and X2 in all cases studied except in aqueous DMF solutions. A

similar relationship is also reported between data obtained fromthe study of complexation equilibria of a number of ligand-metalion systems in MeOHþH2O [31], ANþH2O [32] and

EtOHþH2O [33, 34] mixed solvents.It is known that the ligand must compete with solvent mole-

cules for the cation in the complexation process. Thus,

Table 2. The first stability constants (in logb1 form) of the Cu(II)-glycine complex in mixed solvents (25 �C, I¼ 0.1).

MeOHþH2O EtOHþH2O DMSOþH2O DMFþH2O DOXþH2O

MeOH wt% logb1 EtOH wt% logb1 DMSO wt% logb1 DMF wt% logb1 DOX wt% logb1

0.0 8.18 0.0 8.18 0.0 8.18 0.0 8.18 0.0 8.188.0 8.34 10.0 8.33 11.8 8.32 19.2 8.56 10.0 8.39

16.4 8.43 20.0 8.43 18.6 8.41 38.7 8.90 8.33 [30]25.2 8.63 30.0 8.53 32.5 8.60 48.6 9.09 20.5 8.7534.4 8.79 50.0 9.01 8.53 [30]44.1 8.97

Table 3. Response of the solid membrane electrode to copper(II) ionaqueous solutions of different co-solvent at 25 �C and I¼ 0.1 mol=L.

Co-solventSlope(mV=decade)

Linear range(mol=L)

Compositionrange ofco-solvent (wt%)

H2O 29.3� 0.4 10ÿ2–10ÿ6 –MeOH 29.8� 0.5 10ÿ2–10ÿ6 44.1EtOH 29.5� 0.6 10ÿ2–10ÿ6 50.0DMSO 30.0� 0.4 10ÿ2–10ÿ6 32.5DMF 30.2� 0.7 10ÿ2–10ÿ6 38.7DOX 29.0� 0.8 10ÿ2–10ÿ5 20.5

Fig. 1. Variation of the stability constants (logb1) with mole fraction(X2) of the co-solvents in mixed solvents: —�— MeOH; —m— EtOH;—.— DMSO; —r— DMF; —þ— DOX.

Stability Constants of Copper(II)-Glycine Complex 1117

Electroanalysis 2001, 13, No. 13

variation of the solvent is expected to change the apparentbinding properties of the ligand. Water is a solvent of high

polarity with ET(30) (an empirical solvents polarity parameter) of63.1 [35], which can strongly compete with glycine for Cu2þ.Therefore, it is reasonable to expect an increase in the stabilityconstants on addition of the co-solvent studied, whose ET(30)

values are lower than that of water, to the reaction media. Theorder of ET(30) polarity parameter for pure solvents is found tobe [35]:

H2O > MeOH > EtOH > DMSO > DMF > DOX

This is the same as the order of the solvent effect given above.

Moreover, the lower dielectric constant of the co-solvent incomparison with that of water would also cause the electrostaticcontributions to the bond formation to increase with increasingconcentration of the co-solvent in the mixed solvents. However,

the order of the dielectric constants for the co-solvents [26] doesnot exactly follow the order of relative solvent effects as notedabove. This indicates that a continuum model of solvent based on

coulomb interactions is only limited in interpreting the presentresults.

4. Conclusions

It can be concluded from this study that ISE can be usedsuccessfully for the determination of stability constants of metalcomplexes in aqueous mixed solvents, if attention is given to the

slope and the detection limit of the electrodes under givenconditions. This method is characterized by good accuracy,reproducibility and simplicity of the measurement techniques. Byusing ISE, two parameters, pH and pM (M refers to metal ion)

can be measured. Therefore, the determination of stabilityconstants is greatly simplified and the results obtained shouldbe more reliable compared with the conventional pH-titration

technique.

5. Acknowledgement

The authors are grateful to the Natural Science Foundation ofHenan Province for financial support.

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