T. L. Rose Texas ARM University College Station. 77843 and R. J. Seyse Howard College at Big Spring Big Spring, Texas 79720
Determination of the Stability Constants of Nickel (Il)-Cysteine An undergraduate experiment
The study of the stability constants of transition metal complexes with amino acids involves many different areas of chemistry such as the physical chemistry principles of complex equilibria, the inorganic area of metal chelates, and the biochemical importance of metal ion-amino acid interactions. Cysteine and other sulfur containing amino acids have recently been thought to play a role as antioxidants in the cell and may be related to the aging process. Despite this general relevance to chemistry, we know of only two undergraduate experiments devoted to stability constants of amino acid chelates: Ni(glycinate), using the acid-base titration technique ( I ) and Cu(glycinate), using polarography (2).We wish to describe an experiment performed in our upper level integrated laboratory course (3) in which the stepwise stahility constants of Ni(I1)-cysteine chelate are determined. This system is particularly interesting because of the different possible structural complexes formed by Ni(I1) and because cysteine, being triprotic, has in principle several ways of bonding with the nickel. This experiment also fits in verywell with the determination of the microscopic ionization constants of cysteine ( 4 ) . The equilihrium scheme for the Ni(I1)-cysteine system can he represented as follows [Ni(Cys)l Ni2+ + Cy& e Ni(Cys) KI = (1) [Ni2+][CysZ-] INi(C~s)2~-1 (2) Ni(Cys) + Cys2- = Ni(Cys)s2- Ks = [Ni(Cys)l[C~s~-I where K, and K2 are the stepwise stability constants, Cys2stands for -SCH,CH,COOI
and the Ni(Cys)& complex has the proposed square planar geometry (5,6).
termine K1, Kz, and K, if the ionization constants are known. These data are conveniently obtained by recording the p H from titrations with a strong base of a solution of cysteine with and without Ni present. The derivation of the general equations used to relate the p H values to the stability constants is given by Albert (a), and only the results will he summarized here. The concentration of free cysteine [Cys] present in the solution a t any point in the titration is given by log [Cys] = lag 13[Cyso]- [KOH] - [Hf]I - logo (5) where [Cyso] is the initial concentration of. the cysteine present; [KOH] is the moles of base added divided by the solution volume; [H+] is'the hydrogen ion c6ncentration calculated from the p H reading^,^ and = [H+l+ 2[H+IZ 3[H+13 K.8 KazKn3 KolKo~Kaa (6) ~h~ 3 hefore the [cyso] in eqn. (5) arises from the fact that there are three hydrogen ions liberated for each molecule of ligand which is bound to the metal atom, ~h~ value of [cys] is related to ri, the average number of molecules of ligand bound by one atom of metal, by the equation +
[Crso1- &ysI [Ni"] where [Nio] is the initial concentration of nickel and A=
(7)
a = [H' - l+ p
[H+12 + [H+13 + (8) Ka3 KnzKo3 Ko~KozKo3 The values for K1 and K2 can then either he determined from a plot of n versus log [Cys] a t the point where f i = % and %, respectively (I), or from the relationships
KI =
(A
- 1)
(9)
(2 - ?i)[CysI
a t several values of ri and averaged. A check on the results can be made by determining the value of K, when ri = 1 from the relationship
For the high ligand to metal ratios (51)used in this study the importance of polynuclear species (7), such as N i d C ~ s ) 3 ~ - , can be neglected (6). The overall stability constant K, is obtained from K1 and K2
.
.. . .
Since the metal ion is competing with hydrogen ions for honding with the cysteine ligand, hydrogen ions are released when the chelate is formed thus perturbing the acid-base equilihrium
where K,,, K,z, and Ko3are the macroscopic proton ionization constants.' Measurement of the hydrogen ion concentration in solutions of cysteine and Ni can, therefore, be used to de728 1 Journal of Chemical Education
log K, = -2 log [Cys] (10) and comparing the resulting value with that calculated by eqn. (3). In our laboratory we provide a computer program to handle the calculations.
I Cys- represents a mixture of two structures, one with the thiol group ionized and the other with the ammonium group neutralized, whose ratio at any pH can be calculated from the microscopic ionization constants (4). Since the pH meter measures the activity of the hydrogen ion a"+, accurate results require the conversion of o ~ to+ [Ht]. This is done by determining the mean activity coefficient y. at the ionic strength p of the experiment. These quantities can he calculated from the fallowing equations (9)
-logy+ = 0.50 12122)
{w ,'LIZ
o.2fi]
where 21, and ~2 are the charges on the ions of the supporting electrolyte and ml and ms their molar concentration.
Table 1. Stability Constants Determined From Student Data BY Comouter P r o o r a d
Experimental
Three titrations are run. The first is a standardization of the sodium hydroxide solution with potassium acid phthalate. The second and third are titrations with the base of cysteine solutions without and with the nickel metal ion. At Texas A&M a Sargent automatic titrator and a Leeds and Northrup 7413 pH meter connected to a Mosely strip chart recorder were used to trace out the titration curves. Equally satisfactory results, however, were ohtained a t Howard College using a conventional buret and manually recording the p H readings. Because of cysteine's propensity for oxidation to cystine all titrations involving the amino acid were done with nitrogen flowing through the solution, and solutions were prepared using air-free water. L-Cysteine hydrochloride monohydrate (Eastman Kodak Co.) and Ni(NOa)z6Hz0(Fischer Certified Reagent) were used directly from the bottles. The nickel nitrate solution (0.01 N) was standardized usingdimethylgyloxime. The pH meter was calibrated atpH of 4,7, and 9, using commercially available buffers. Duplicate determinations were made to determine the normality of the sodium hydroxide. Then a blankof 200ml of 0.1N KN0s was titrated with NaOH with nitrogen flowingthrough the solution. About 0.9 gm (5 mmole) of cysteine was accurately weighed out, dissolved in exactly 200 ml of 0.1 N KNOa and titrated with the base to a p H of about 11.Finally another 5 mmole portion of cysteine was weighed out, exactly 100ml of the standardized Ni2+solutionsdded,and the solution volume brought up to 200 ml with 0.2 N KNOs. The titration was continued until about 30ml of base had been added.3Durinethe
1.
4.11-4.90 5.20-5.60
0.17-0.5 1.28-1.88
7 5
2.
4.94-5.27 5.54-6.0
0.14-0.82 1.15-1.4
7
3.
4.70-5.10 5.36-5.90
6 6
4.
4.91-5.28
0.20-0.80 1.28-1.82 0.15-0.81 1.15-1.37 0.224.87
8 9
5.56-6.00 4.49-5.11 5.25-5.84
5.
~
Results
Table 1presents a summary of the student results. Points in the data set without ~hvsicalsignificance were eliminated, and the remaining pointsaverag& to give the tabulated p K values along with the standard deviations. The acid dissociation constants used in the calculations were the students' It should he noted that since we are using a computer program to analyze the data, corrections of the concentrations due to dilution by the volume of the added base are easily handled. If the calculations are going to be done by hand, it is more convenient to use a stranger
3
1.24-1.90
=
9.90 f
=: 20.54
0.20
0.22
10.09 t 0.28 = 8.50 t 0.16 = 18.59 = 10.25 f 0.18
20.3
=
19.0
= 9.20 f 0.09
= 19.45 p4= 10.77 0.25 p K = 9.15 t 0.21
19.6
i
p ~ pK, pK2 pKs
=f 19.92 = 10.44 i 0.19 = 9.50 i 0.13 = 19.94
20.5
19.7
U ~ a t ref a 1 war taken at Howard College. T h e other retr were by students at ~ e x a A&M r University. Temperature war
24 f 2 ' ~ .
Table 2.
Literature Valuesfor the Nickel (1I)Cystelns Stability Constants Temp.
PC1 25
0.1 M
(101
9.57
20.07 19.77 (20.10)
25 (15)
(61
9.40
19.04
KN03 0.2 M KNOI
25
0.1 M KNO,
(5)
PK,
9.82
10.25
10.20 (10.36)
9.64 0.28
ref
pKS
PKI
...
Calculations are carried out by a Fortran IV program. T h e program is entirely general and would work for any ligand. ~. -parameters are the hydrogen activity, metal s ~ s t e mInput the initial metal ion, ligand, and base concentrations, and the initial volume. The protonation constants of the acid are then entered. These may he literature or student determined values. The titration data are entered as ml of base versus pH. Appropriate dilution corrections are made for each data point by the program. For the nickel-cysteine system, most useful data are collected hetween a p H of 4 and 7 recorded a t intervals of 0.1 p H . The computer output consists of volume of base added, p H , log [Cys] from eqn. (5), [H+],ri from eqn. (7), and log K I and K Zfrom eqn. (9) for each data point. Only points where n < 1are used to calculate K1, and where ri > 1are used to calcul a t e K ~The . nrogram has been designed to perform the tedious calklatiin oTdata reduction, while leaving plenty of room for student interpretation. The student must select the meaningful data, interpret trends, calculate the final values for K , and KI bv averapin~selected data or using- graphical techniques, aid-estimate the error limits of the constants ohtained.
7
= 10.64 t
acquired
8.98 t
Calculations
5
pK, pK p ~ pK, pK9 pK8 pK, pK$ pK
(9.74)
.. .
11.18
19.3 20.16 i
0.01
20 20
~r
0 : ; ~ NaCIO.
(81 (71
values determined from the titration of cysteine with base in the absence of the Ni2+ ion. The hydrogen activity coefficient in these systems was 0.8L5 With the exception of pKz from data set 2, the values all fall in the ranee of those renorted in the literature. which are aiven in Table Data set 5 appears to be best. I n general themost consistent results are obtained for ri = 0.2-0.8 and 1.2-1.8 for pK1 andpK2, respectively, a s is expected since in these ranges the suhstractions ineqn. (9) lead to the smallest relativeerrors. Determination of K, using eqn. (10) (last column of Table 1) shows good agreement with that calculated from the individual stabilit; constants (eqn. (3)). A further check on the data is the value of n a t high pH. This value should level off a t 2.0, since the complex consists of 2 Ligand molecules per nickel ion. Data sets 1,.3, and 5 level off between 1.9 and 2.2, whereas data set 2 and 4 level off a t about 1.45. The reason for this low value for 2 and 4 is not readily apparent; perhaps it is due to faulty standardization of the p H meter or misreading of data, etc.
5
Additional Experiments
The experiment described has a number of variations and extensions. Students can determine the stability complex for metal ions other than Ni(II), such as Mn(II), Zn(II), Ph(II), or Hg(II).'j If thermostated cells are available, the thermodynamic functions of the complex formation may be determined (7).The bonding and geometry of the complex can be analyzed by the uv-visible absorption spectrum of the chelate. Literature Cited ( I ) An~lici.B.J.."SynthesisandTechniqu~~inlnorganieChamisfry~Saunderi,Phitadetphia,1969.p. 105. (2) Betulheim. F. A . "Expecimental Physical Chemistry," Saunden. Philadelphia, 1971.
base (1.5 N), add only 10 ml, and neglect the concentration changes (1).
4 A listing of
the program may be obtained from T. L. R. "The calculated stability constants are relatively insensitive to chanees in this osrameter.
,.",",. 171 Perrin.D. D..andSavce.
1. G..J.
Chsm. Soc. (A). 53. (19681
Volume 53, Number 11. November 1976 / 729
