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The complexing ability of 2-pyridinecarboxaldehyde toward cupric ions
L inorg, nucl. Chem., 1976, Vol. 38, pp. 1533-1539. Pergamon Press.
Printed in Great Britain
THE COMPLEXING ABILITY OF 2-PYRIDINECARBOXALDEHYDE TOWARD CUPRIC IONS A. E. EL-HILALY and M. S. EL-EZABY* Department of Chemistry, Kuwait University, Kuwait (Received 25 September 1975) Abstract--The reactions of 2-pyridinecarboxaldehyde with copper(II) ions were examined by potentiometric and spectrophotometric methods. Stability constants, measured at 25°C and 0.5 M KCI as well as 0.1 M KNO3 ionic strength, have been determined under various experimental conditions. A 1 : 2 neutral solid complex has been isolated from aqueous solutions. The IR spectra showed that the deprotonated hydrated species of the ligand was involved in complex formation.
INTRODUCTION IN SPITE of the enormous amount of stability constant data available in the literature, very little has been done to calculate those of pyridinecarboxaldehydes. Metal complexes with 2-pyridinealdoximes rather than 2pyridinecarboxaldehydes were known and their solution equilibria were studied[l]. In this work the complexes of copper(II) with 2-pyridinecarboxaldehyde were studied in solution and in the solid form.
EXPERIMENTAL Materials. 2-Pyridinecarboxaldehyde was obtained from Fluka, and was redistilled under a reduced pressure of nitrogen and stored under nitrogen at 0°C [2]. Stock aqueous solutions of 1.0 M 2-pyridinecarboxaldehyde in 1.0M HC1 or 1.0M HNO3 were stored at 0°C and were stable by UV spectral criteria (Amax235 nm (em~ = 6.1 × 103), for equilibrated aldehyde+hydrate, pH 6-10, /z = 0.5 M KCI) over a period of a month. Stock solutions of CuCI2 or Cu(NO3)2 (0.1 M) were checked complexometrically by EDTA titration [3]. Measurement o[ pH, pCu and spectra. Potentiometric measurements were done at ~ = 0.1 M KNO3 and the spectrophotometric study at tz = 0.5 M KCI. All measurements were carried at 25°C. Radiometer pH meter model 63 (accurate to -0.01 pH unit) equipped with combined glass electrode type (GK 2301 C) was used to record the hydrogen-ion concentration. The titrations were carried out in 50 ml thermostated cell. Sodium hydroxide titrant (0.118M) was added in small increment through Jurgens microburrette accurate to -+0.005 ml. Purified nitrogen was bubbled in the solutions before and during titrations. The pH meter was calibrated as previously reported [4]. Radiometer pH meter model 63 provided with both extinsion copper selectrode (type F 1112 Cu) and calomel electrode were used to determine the cupric-ion concentration. Calomel electrode was immersed in saturated solution of KC1 connected to the solution in the titrating cell through KNO3 salt bridge. A calibration curve of cupric ion concentration vs mV was established using standardized solutions of Cu(N03)2 of different concentrations at p, = 0.1 M KNO3. The potential of a given cupric ion concentration in absence of the ligand was constant in the pH range 1.5-5.5. Spectrophotometric measurements in visible and UV regions were recorded using Unicam SP 8000 and 700, spectrophotometers. The instruments were calibrated as previously reported [4]. The solid complex of copper(II) with 2-pyridinecarboxaldehyde was isolated by simple addition of Cu(NO3)2 to a solution of the ligand. The pH was adjusted to -8.0. The precipitated pink complex was washed by successive portions of water followed by ethanol and dried in vacuo. A blue solid complex was also isolated when the solution was kept for a week at pH 7.0. Authentic samples of Cu picolinate was also prepared and compared with the above two complexes. The solid complexes were analysed for C, H, N and Cu. The
analysis was 46.16, 3.95, 8.99 and 20.15% for C, H, N and Cu respectively, for the pink complex. The IR spectra were taken for the solids in KBr matrix and for the liquid 2-pyridinecarboxaldehyde between two transparent NaCI discs. IR spectra were measured in the 300-4000 cm ' region with a Beckman IR 12 double beam spectrophotometer, previously calibrated. RESULTS Ligand dissociation constants The acid-base equilibria of 2-pyridinecarboxaldehyde and their corresponding acid dissociation constants may be represented as follows: HP.H:O +. P.H20
(1)
• POH- = H+; kl = [H][POH] [P.H20]
(2)
where HP.H20*, P.HzO and POH- are the hydrated species of the protonated and unprotonated species and the deprotonated form of the hydrated species, respectively. The dissociation constants are calculated by spectrophotometric[5] and potentiometric [6] methods. Figure 1 shows the variation of absorbance as function of pH at various wavelengths. The equation used in calculating the dissociation constants is k , - [H],-M[H],+j M- 1
(3)
where M - As~+j - Asi-, [H]~_~- [H]~ As~ - As,-t [H]~_~- [H],_j ' Figure 2 shows the variation in a, the number of moles of base added per mole of ligand, as function of pH. The first dissociation step may be considered separately since the 2nd dissociation step is expected to take place at higher pH. The equation used is k2 = [H](aTL + [H]) (1 - a)TL - [H]
(4)
where TL is the analytic concentration of the ligand. The value of pK~ obtained from spectrophotometric method was little higher than that obtained from potentiometric method (Table 1). This was attributed to the effect of ionic strength factor used in both techniques.
1533 JINC Vol. 38, No. 8--H
"
" P . H 2 0 + H + ; k ~ - [H][P.H20] [HP.H20]
A. E. EL-HILALYand M. S. EL-EZABY
1534
Stability constants of metal complexes (a) pH metric study. The titration curve (Fig. 2) for a
0.8 %
0.6 -
°~X260 nm
0.4
X260nm k275nm
+c
0.2
1 : 1 metal to ligand ratio shows slight inflection at a = 2 and steep inflection at a -- 3. The reaction equilibria and their equilibrium constants which fit the observed experimental data may be represented as follows:
-\
•
. . . . :. . . . [CuP'H20] ' LuI"'I'-IzU , ~ , = [ ~ ]
P ' H 2 0 + Cu :+ .
f 5 / x 275nm
O0 2
(5)
I
I
I
|
I
I
4
6
8
I0
12
14
CuP.H202+ .
"CuPOH + + He; K2
[CuPOH][H] [CuP.H20] (6)
pH P O H - + C u 2+
Fig. 1. Absorbanceof 10-"M solutionof2PAas functionofpH.
"CuPOH÷; K3
"
. . . . . . . . .
CuPOH ÷ + O H - .
t.uunrun;
[CuPOH]
(7)
[Cu][POH]
[CuOHPOH] ~4 = [ ~ ] .
(8) .,I. /
,o
/
The stability constants were evaluated algebraically. The concentration of different species were determined with the aid of the following equations
/'/
TM = [Cu '+] + [CuP'H202÷] + [CuPOH ÷] + [CuOHPOH]
(9) ,.¥,o
TL = [HP .H20 +] + [P .H~O] + [POH] + [CuP'H202+] + [CuPOH +] + [CuOHPOH].
+,,.o 6
/
,
In these equations CuPOH, CuOHPOH and POH were ignored at pH's lower than 3.4, where Cu2+ and CuPH20 and HP.H20 are mainly present in considerable concentration. From the electroneutrality principle
4;'1 ,.+<,+
[Na ÷] + [H ÷] + [HP'H20 ÷] + [Cu2+] + [CuP'H202+]
2
l I
l 2
= [ N O r ] + [OH ]
l :3
Q
Fig. 2. Titration curves of Cu-2PA system. 2PA, 2 x 10-+ M, lxl0-2M, 5x10-3M, 3×10-3M. Cu, II, l×10-2M; Q, l x 10-2 M; +,5 x 10-~M;A,3 x 10-~M.
Spectrophotometrl¢
methods
Reaction ,LL . O. 5 M KCI
p. H20 PB20
+
~= .... ~
P.H20
~=====~
POH-
Cu 2÷
H+
Cu P0H +
+
Cu POH +
Cu (08) (P0~)
~ ..... ~
Cu (POH) +
*
Cu 2+
~. . . . . •
+
2EH20
~'====~
Cu(PB20)~ +
Cu
+
3~H20
~. . . . . t
Cu (P~20~;
Cu
+
4PH20
~=====~
C~(PH20)~+
R. . . . . •
(POH) (PH20)
~. . . . .
-12.61
+
0.18
3.19
+
0.02
OH-
Cu
- 3.66
+ 0.06
2.72
_+ O.O6
(POH) 2
The error limits reported here are
2.94
4.66 10.67
_+ 0.04
- 7.13
_+ 0.03
-'20.5 4.68
+
0.02
4.75 6.32 8.52
Cu (P~B) (Pa20) =
0.08
Cu (POH) 2
÷
Cu
+
P o t e n t i o m e t r i c methods Glass electrode Copper selectrode
///= 0.i M KNO 3
-
Cu
Gu (P~20)2
- 3.96
H÷
~.==.=a
2POH-
*
+
Cu 2+
POB-
÷
H+
Cu PH202÷
~='===~
~=====~
Cu PH20
+
+ +
S+ B+
± ~ o" .
(11)
where Na ÷ is the concentration of NaOH used after each addition and NO3 is the concentration of nitrate ions other than that used for adjusting the ionic strength. From eqns (9)-(11) and introducing the first dissociation
Table 1. Log of the acid dissociation constants of 2-pyfidinecarboxaldehydeand its copper(II)complexes*
Hp. H2 O÷
(10)
"
- 5.55 - 7.72
+
0.06
The complexingability of 2-pyridinecarboxaldehydetoward cupric ions constant of the ligand we have [P'H20] =
150
(1--a)TL + [OH]- [H] [H÷l/k,
÷
I00
(13)
TM
In a similar way, the species CuPH20 and CuOHPOH were ignored in the pH range 3.4-5.5. A new set of equations were derived taking in consideration the new electroneutrality relation combined with eqns (9) and (10). [POH] =
(2 - a)TL - [HI + [OH] 2[H]: [H] klk2
k~_
[H]
[H]
~
(12)
and a, the average ligand number, is
ti =
1535
\
g
,
g_ (14)
5O
and TL - [POH] (k~V~+--~-2+1 ) vi T,
(15)
From pH 6.0 the value of a increases owing to the formation of probably hydroxo complex species. In such a case the species CuPOH ÷ and probably only CuOHPOH coexist. The stability constants were evaluated from the following relations: TL = TM= [CuPOH ÷] + [CuOHPOH]
(16)
\ 2
3
4
5
pH Fig. 3. Potential dependence of Cu-2PA system on pH. +, lx10-3-2x10-2M; L l x 1 0 3 - 5 x 1 0 - 2 M ; ©, 1x10-3-8x 10-z M; x, 1 x 10-3-1 x 10-' M; A, 1 x 10-2-1.5 × 10 ' M.
To. - [OH ] = [CuOHPOH] where ToK = concentration of hydroxyl ions at any point along the titration curve (above pH 6) and OH was evaluated from the corresponding pH values and K~ (= 10 '37~[7]). The average number of OH groups bound per mole of the first complex may be expressed by aoH -
To~- [OH] _ K4[OH] TM 1 + K4[OH] "
(17)
Numerical methods (8) were used to calculate the stability constants of the equilibria 5, 7 and 8. The value of K2 was calculated from the knowledge of the equilibrium constants K1 and K3. The titration curve for 1 : 2 metal to ligand ratio shows steep inflection at a = 2 (Fig. 2) indicating that the ligand species is POH . Two stability constants were determined for the following complex equilibria: CuPOH + , Cu(POHh,
' Cu 2+POH ' CuPOH ++POH .
,oo
Io'
,I
.% 10"2
\
(18)
(19)
The stability constants are shown in Table 1. (b) Potentiometric study using copper(H) selectrode. Figure 3 illustrates the dependence of the potential of the copper(II) selectrode as function of pH. The ligand concentration was always in excess with respect to the cupric ion concentration. A unique formation curve of o~o([Cu]/TM)as function of P.H20 was obtained for various initial concentrations of 2-pyridinecarboxaldehyde and cupric ions (Fig. 4). The free ligand concentration was
IO-3
oe the 10-3
i
l
I i l lll] I0-2
[PHzO] Fig. 4. The plot of log ~ o as functions of log [PH20].
1536
A.E. EL-HILALYand M. S. EL-EZABY
obtained from the following relation, [PH20] = = ITL
(20)
where
2PA
2 . 5 . 1 0 -I M
i
[H]J-' I I k, I=0
I
j =o
j
0.6
[HI'-J I-I k, 1=0
Approximate values of the stability constants of the following equilibria Cu+ P'H20. Cu+ 2P.H20. Cu+3P'H20. Cu+4P.H20.
" CuP'H20 ' Cu(P.H20)z ' Cu(P.H20)3 " Cu(P'H20)4
(21) (22) (23) (24)
,~ 0.4
were obtained using the following equation 0.2
1
= o- N 0
[Cu] TM
(25)
/3. [PH20]"
where/3 is the overall stability constant. The values so obtained were refined further by successive approximation and depicted in Table 1. Back calculation of the formation curve was made utilizing the refined values of the four/3's. (c) Spectrophotometric study. Figure 5 shows the spectra of Cu-2-pyridinecarboxaldehyde system at different pH values. A regular blue shift was observed as pH increases. Graphical methods [9] were used to determine the number of absorbing species present in solution. A conclusion was drawn that in the pH range 1.4-2.7 and 3.0--4.5 two absorbing species were present in both cases and a single absorbing species may exist in the pH range 5.0--8.9. Figure 6 shows the pH dependence of absorbance at the wavelength 670 nm. The absorbance is dependent on both the initial concentration of the ligand and metal ion. At constant metal ion concentration and various initial concentrations of the ligand the absorbance is dependent upon pH and the concentration of the ligand in the pH range 1.0-4.4. At pH's higher than 4.4 the absorbance is solely pH dependent. A unique formation curve was obtained when the average molar absorptivity [~ = (As/TM)] was plotted as function of the free ligand species (P.H20) in the pH range 1.0-4.2 (Fig. 7). The free ligand species was calculated using eqn (20). The absorbance data at 670 nm were analysed by the least square procedures. In this procedure, the sum of the squared deviations, s, given by eqn (27) s = ~ (~obs- ~c~o)z
0.0
500
700
900
11(30
X nm
Fig. 5. Absorptionspectraof Cu-2PAsystem.
1.6
\
!
/
1.2
/
\
\ \
io8F ,
(27)
-i-
should be minimum, where gobsand ~ o are the observed and calculated values of the average molar absorptivities and N~ is the number of the experimental points considered. In other way,
0.0 ~ / " l
N
(28) o
o
I
I
I
L
3
5 pH
7
9
Fig. 6. Variation of absorbance of Cu-2PA systemas functionof pH. x, 2xl0-L2.5xl0-'M; O, 3x10-L2.5x10-~M; [3, 3x10-3-4x10-1M; A, 3xl0-L5xl0-tM; @, 4x10-L2.5x 10-1 M; +, 8 x 10-L2.5x l0-1 M; ;t, 670rim; l, 4 cm.
The complexingability of 2-pyridinecarboxaldehydetoward cupric ions
1537
60
•
~" 4"0
2o
..5
-
..y7 0
t
t
i
I
I
I I II
I
I
I
I
I
I I I[
I0-2
10-3
I
I
[
I
[
I0 -I
[PH20] Fig. 7. The dependence of f on log [PHzO]. O, 2 x 10-3-2.5× 10-~ M; A, 3 × 10 3-2.5× 10 ' M; @, 3 × 10-~4 × 10-' M; V1,3 × 10-3-5× 10-' M; +, 4 × 10-3-2.5x 10 ~M. where E, the molar absorptivity of the nth complex and a stands for the free ligand, P.H20, calculated from eqn (20). Initially guessed values of the stability constants were used to solve for the N + I unknown ~ by Gauss-Jordan elimination method. Since the initially guessed values of the stability constants are not the optimum ones, it is necessary to adjust their values to reduce s. The method of the steepest descents[4, 11] was used to optimize the values of the stability constants. The molar absorptivities and the sum of the squared residuals then recalculated. The whole process may be repeated till refined/3's could be obtained from the initial guesses. The plot of s vs, e.g./3, or log/3, give the optimum value of/3, which can be used to get/3m from the approximate linear relation
/3,. =a,./3. +b,.
described in previous report [4], where am and b~ could be obtained from the plot of/3m vs/3,. The reactions (21) and (22) were assumed to take place in the pH range 1.0-4.2. Table 2 shows the initially guessed/3's and their rapid convergence values. Table 1 shows/3, and/35 obtained by this method. At pH's higher than 5, the following reactions were assumed to take place Cu(P.H~O)22+ .
' Cu(P.H20)(POH) + + H +; K ~ - [Cu(P.H20)(POH)][H] [Cu(P.H~.O)2]
Cu(P.H20)(POH)
+ .
• Cu(POHff
+ H+; [Cu(POHh][H] K6 = [Cu(P .H20)(POH)] '
n~m
Table 2. Initially guessed and rapid convergence values of the stability constants in 2-pyridinecarboxaldehyde copper(II) system* Initially guessed values
Rapid convergence values
~i
/32
Sigma
/31
/92
Sigma
0.50 x 102
0.50 x 102
1.2026
0.50 x 102
0.50 x 102
i x 102
i x 102
1.339
0.64 x 102
0.72 x 102
1.1715
5 x 102
5 x 102
2.8782
1.48 x i02
1.93 x 102
0.9425
1 x 103
1 x 103
3.529
3.27 x 102
6.22 x 103
0.80205
5 X 103
5 X 103
4.5515
6.02 x 102
1.40 x 104
0.74467
1,2026
I x 104
1 x 104
4.7741
1.67 x i05
4.56 x 104
0.74797
5 x 104
5 x 104
4.9897
7.94 x 103
2.46 x 105
0.81976
1 x 105
1 x 105
5.0198
1.90 x 104
6.10 x 105
0.84392
5 X 105
5 x 105
5.0445
7.25 x 104
2.38 X 106
0.85993
i x 106
I x 106
5.0476
1 . 6 7 x 105
5 . 4 7 x 106
0.86328
I x I07
i x 107
5.0505
1.47 x 106
4.80 x 107
0.86570
TM
8 x 10-3 M
TL
2,5 x I0-I M ffi 670 nm -
~X
4 cm
ffi 0.5 M KCI
Computer p r o g r ~ MONONUC - IBM 370
A. E. EL-HILALYand M. S. EL-EZABY
1538
It was assumed that at pH -4.5, Cu(P.H:O)2 is completely formed. If this was the case, eqn (3) could be used to calculate K5 and K6 (Table 1).
IR spectra of the copper complex The analysis of the metal complex prepared in this study was found to be compatible with the empirical formula, Ct2Ht204N:Cul. Assignments for the IR spectrum of the copper complex can be made by a comparison with the assignments for the parent ligand (Table 3). The vanishing of the C=O stretching band frequency at 1723 cm-l from the spectrum of the copper complex with the appearance of OH stretching band frequency at I 3500 cm-', indicates the presence of -C-OH grouping in I I place of -C=O. The appearance of C-O stretching band frequencies in the region of 982-1100 cm-l was a further I confirmation of the absence of --C=O group. The general shift of the pyridine band frequencies to higher values and the appearance of Cu--O and Cu-N stretching band frequencies at 360 and 510 cm-l, respectively, proves the involvement of nitrogen and oxygen atoms in chelation.
obtained from copper selectrode and spectrophotometric measurements, where excess ligand was used, are different since less hydrolysis and olation of the hexaaquocopper(II) ion may be expected specially in acidic medium. The stability constants obtained in this study for the formation of some complex species were compared with complexes having similar potential ligand (Table 4). It is clear that 2-pyridinecarboxaldehyde may act as a unidentate ligand and bind the copper(II) ion by the hetrocyclic nitrogen. On the other hand, it could act as a bidentate ligand and chelate the metal ion by the heterocyclic nitrogen and the oxygen of the aldehydic group. In the latter case the stability constants should be little higher than in the former. The pH-metric and spectrophotometric studies indicates that the deprotonated hydrated species may also form with the cupric ions a quite stable complex species in solution. We have not found to the best of our knowledge that such ligand species may be involved in complex formation with metal ions. The
I
O--Cu--O
DISCUSSION The 2-pyridinecarboxaldehyde copper(II) system exhibits different species under different experimental conditions. Results from potentiometric studies indicate that hydrolysis and olation of hexaaquocopper(II) ion would be expected to be considerable. However, results
I
I I
I
.ify
Fig.8.
Table 3. IR spectraof 2-pyridinecarboxaldehydeand its coppercomplex Frequency modes
llgand,
Pyridlne ring
cm-I
Copper complex,
cm-I
413(sh,w), 620(sh,w)670(sh,m)
432(sh,m),638(sh,m),670(sh,w),
770(sh,s),838(sh,s),lOOO(sh,s),
780(sh,s),820(sh,m)1445(sh,s),
1050(sh,m)lO95(sh,w),1445
1480(sh,s),1575(sh,w),1607
(sh,m),1478(sh,m),1593(sh,w)p
(sh,s),2720(br,m),2850(br,s),
2850(br,w),2850(br,m).
3040(shd),3080(sh,s).
~iphatic C-H
1375(sh,m).
1260(sh,m),1280(sh,m),1315(sh,m).
C-O
1723(sh,s)
?-c
l155(sh,w),1220(sh,s)
l160(sh,m),1220(sh,m) 980(sh,m),lO20(sh,m),lO50(sh,s),
C-0
1070(sh,s).
Cu - N
360(br,w). 510(sh,m).
O-H
3500(br,w).
Cu - 0
ah, w, s, br, m, shd stands for sharp, weak, strong, broad, medium and shoulder
Table 4. Stabilityconstantsof some copper complexes Co~o.nd
B1
/32
~3
/~
5.80
6.70
6.32
8.52
Pyridine*
2.6
4.54
P£colinlc acid** (chelate)
7.95
14.95
2-Pyridlnecarboxaldehyde
2.94 Cu(PH20) 10.67 Cu(POH)
4.75 Cu(PH20) 2 20.5 Cu(POH) 2
The complexingabilityof 2-pyridinecarboxaldehydetoward cupric ions analysis has established the absence of anions indicating that the complex is neutral with 1 : 2 metal to ligand ratio. Moreover, the loss of C=O stretching band frequency and the appearance of OH in addition to the C-O band frequencies in the IR spectrum of the complex is a strong evidence of the involvement of the deprotonated hydrated ligand species in the metal complex (Fig. 8). This complex may be deceived with copper picolinate complexes. Samples of Cu picolinate complex was prepared and their IR spectrum was compared with that of 2pyridinecarboxaldehyde copper(II) complex. The spectrum of the latter was quite different in many respects from the former. The C=O stretching band frequency in the former exist at slightly lower frequency ( 1725 cm ~~ 1650 cm ~) as compared to that of picolinic acid. Moreover, the stretching frequency band of the OH of the carboxylic group is lost upon complex formation.
REFERENCES
t. L. G. Sill6n and A. E. Martell, Stability Constants of Metal-Ion Complexes, Spec. Publ. No. 17 (1964); Supplement
1539
No. 1, Spec. Publ. No. 25 (1971). The Chemical Society, London. 2. H. Diebler and R. N. F. Thomeley, J. Am. Chem. Soc. 95(3), 896 (1973). 3. G. Schwarzenbach and H. Flaschka, Complexometric Titrations (Translated by H. M. N. H. Irving), 2nd English Edn. Methuen, London (1%9). 4. M. S. El-Ezaby and N. Gayed, J. Inorg. Nucl. Chem. 37, 1065 (1975). 5. M. Osman, T. M. Salem and M. S. E1-Ezaby,J. Chem. Soc. (A), 1401 (1971). 6. R. L. Gustafson and A. E. Martell, Arch. Biochem. Biophys. 68, 485 (1957). 7. R. G. Bates, Determination of pH, p. 74. Wiley, New York (1964). 8. F. J. C. Rossotti and H. Rossotti, The Determination of Stability Constants. McGraw-Hill,New York (1%1). 9. J. S. Coleman, L. P. Varga and S. H. Mastin, Inorg. Chem. 9, 1015 (1970). 10. S. Feldherg, P. Koltz and L. Newman, Inorg. Chem. It, 12, 2860 (1972). 11. W. E. Grove, Brief Numerical Methods. Prentice-Hall, London (1966). 12. M. S. Sun and D. G. Brewer, Can. J. Chem. 45, 2729(1967). 13. G. Anderegg, Heir. Chim. Acta 43, 414 (1960).
Printed in Great Britain
THE COMPLEXING ABILITY OF 2-PYRIDINECARBOXALDEHYDE TOWARD CUPRIC IONS A. E. EL-HILALY and M. S. EL-EZABY* Department of Chemistry, Kuwait University, Kuwait (Received 25 September 1975) Abstract--The reactions of 2-pyridinecarboxaldehyde with copper(II) ions were examined by potentiometric and spectrophotometric methods. Stability constants, measured at 25°C and 0.5 M KCI as well as 0.1 M KNO3 ionic strength, have been determined under various experimental conditions. A 1 : 2 neutral solid complex has been isolated from aqueous solutions. The IR spectra showed that the deprotonated hydrated species of the ligand was involved in complex formation.
INTRODUCTION IN SPITE of the enormous amount of stability constant data available in the literature, very little has been done to calculate those of pyridinecarboxaldehydes. Metal complexes with 2-pyridinealdoximes rather than 2pyridinecarboxaldehydes were known and their solution equilibria were studied[l]. In this work the complexes of copper(II) with 2-pyridinecarboxaldehyde were studied in solution and in the solid form.
EXPERIMENTAL Materials. 2-Pyridinecarboxaldehyde was obtained from Fluka, and was redistilled under a reduced pressure of nitrogen and stored under nitrogen at 0°C [2]. Stock aqueous solutions of 1.0 M 2-pyridinecarboxaldehyde in 1.0M HC1 or 1.0M HNO3 were stored at 0°C and were stable by UV spectral criteria (Amax235 nm (em~ = 6.1 × 103), for equilibrated aldehyde+hydrate, pH 6-10, /z = 0.5 M KCI) over a period of a month. Stock solutions of CuCI2 or Cu(NO3)2 (0.1 M) were checked complexometrically by EDTA titration [3]. Measurement o[ pH, pCu and spectra. Potentiometric measurements were done at ~ = 0.1 M KNO3 and the spectrophotometric study at tz = 0.5 M KCI. All measurements were carried at 25°C. Radiometer pH meter model 63 (accurate to -0.01 pH unit) equipped with combined glass electrode type (GK 2301 C) was used to record the hydrogen-ion concentration. The titrations were carried out in 50 ml thermostated cell. Sodium hydroxide titrant (0.118M) was added in small increment through Jurgens microburrette accurate to -+0.005 ml. Purified nitrogen was bubbled in the solutions before and during titrations. The pH meter was calibrated as previously reported [4]. Radiometer pH meter model 63 provided with both extinsion copper selectrode (type F 1112 Cu) and calomel electrode were used to determine the cupric-ion concentration. Calomel electrode was immersed in saturated solution of KC1 connected to the solution in the titrating cell through KNO3 salt bridge. A calibration curve of cupric ion concentration vs mV was established using standardized solutions of Cu(N03)2 of different concentrations at p, = 0.1 M KNO3. The potential of a given cupric ion concentration in absence of the ligand was constant in the pH range 1.5-5.5. Spectrophotometric measurements in visible and UV regions were recorded using Unicam SP 8000 and 700, spectrophotometers. The instruments were calibrated as previously reported [4]. The solid complex of copper(II) with 2-pyridinecarboxaldehyde was isolated by simple addition of Cu(NO3)2 to a solution of the ligand. The pH was adjusted to -8.0. The precipitated pink complex was washed by successive portions of water followed by ethanol and dried in vacuo. A blue solid complex was also isolated when the solution was kept for a week at pH 7.0. Authentic samples of Cu picolinate was also prepared and compared with the above two complexes. The solid complexes were analysed for C, H, N and Cu. The
analysis was 46.16, 3.95, 8.99 and 20.15% for C, H, N and Cu respectively, for the pink complex. The IR spectra were taken for the solids in KBr matrix and for the liquid 2-pyridinecarboxaldehyde between two transparent NaCI discs. IR spectra were measured in the 300-4000 cm ' region with a Beckman IR 12 double beam spectrophotometer, previously calibrated. RESULTS Ligand dissociation constants The acid-base equilibria of 2-pyridinecarboxaldehyde and their corresponding acid dissociation constants may be represented as follows: HP.H:O +. P.H20
(1)
• POH- = H+; kl = [H][POH] [P.H20]
(2)
where HP.H20*, P.HzO and POH- are the hydrated species of the protonated and unprotonated species and the deprotonated form of the hydrated species, respectively. The dissociation constants are calculated by spectrophotometric[5] and potentiometric [6] methods. Figure 1 shows the variation of absorbance as function of pH at various wavelengths. The equation used in calculating the dissociation constants is k , - [H],-M[H],+j M- 1
(3)
where M - As~+j - Asi-, [H]~_~- [H]~ As~ - As,-t [H]~_~- [H],_j ' Figure 2 shows the variation in a, the number of moles of base added per mole of ligand, as function of pH. The first dissociation step may be considered separately since the 2nd dissociation step is expected to take place at higher pH. The equation used is k2 = [H](aTL + [H]) (1 - a)TL - [H]
(4)
where TL is the analytic concentration of the ligand. The value of pK~ obtained from spectrophotometric method was little higher than that obtained from potentiometric method (Table 1). This was attributed to the effect of ionic strength factor used in both techniques.
1533 JINC Vol. 38, No. 8--H
"
" P . H 2 0 + H + ; k ~ - [H][P.H20] [HP.H20]
A. E. EL-HILALYand M. S. EL-EZABY
1534
Stability constants of metal complexes (a) pH metric study. The titration curve (Fig. 2) for a
0.8 %
0.6 -
°~X260 nm
0.4
X260nm k275nm
+c
0.2
1 : 1 metal to ligand ratio shows slight inflection at a = 2 and steep inflection at a -- 3. The reaction equilibria and their equilibrium constants which fit the observed experimental data may be represented as follows:
-\
•
. . . . :. . . . [CuP'H20] ' LuI"'I'-IzU , ~ , = [ ~ ]
P ' H 2 0 + Cu :+ .
f 5 / x 275nm
O0 2
(5)
I
I
I
|
I
I
4
6
8
I0
12
14
CuP.H202+ .
"CuPOH + + He; K2
[CuPOH][H] [CuP.H20] (6)
pH P O H - + C u 2+
Fig. 1. Absorbanceof 10-"M solutionof2PAas functionofpH.
"CuPOH÷; K3
"
. . . . . . . . .
CuPOH ÷ + O H - .
t.uunrun;
[CuPOH]
(7)
[Cu][POH]
[CuOHPOH] ~4 = [ ~ ] .
(8) .,I. /
,o
/
The stability constants were evaluated algebraically. The concentration of different species were determined with the aid of the following equations
/'/
TM = [Cu '+] + [CuP'H202÷] + [CuPOH ÷] + [CuOHPOH]
(9) ,.¥,o
TL = [HP .H20 +] + [P .H~O] + [POH] + [CuP'H202+] + [CuPOH +] + [CuOHPOH].
+,,.o 6
/
,
In these equations CuPOH, CuOHPOH and POH were ignored at pH's lower than 3.4, where Cu2+ and CuPH20 and HP.H20 are mainly present in considerable concentration. From the electroneutrality principle
4;'1 ,.+<,+
[Na ÷] + [H ÷] + [HP'H20 ÷] + [Cu2+] + [CuP'H202+]
2
l I
l 2
= [ N O r ] + [OH ]
l :3
Q
Fig. 2. Titration curves of Cu-2PA system. 2PA, 2 x 10-+ M, lxl0-2M, 5x10-3M, 3×10-3M. Cu, II, l×10-2M; Q, l x 10-2 M; +,5 x 10-~M;A,3 x 10-~M.
Spectrophotometrl¢
methods
Reaction ,LL . O. 5 M KCI
p. H20 PB20
+
~= .... ~
P.H20
~=====~
POH-
Cu 2÷
H+
Cu P0H +
+
Cu POH +
Cu (08) (P0~)
~ ..... ~
Cu (POH) +
*
Cu 2+
~. . . . . •
+
2EH20
~'====~
Cu(PB20)~ +
Cu
+
3~H20
~. . . . . t
Cu (P~20~;
Cu
+
4PH20
~=====~
C~(PH20)~+
R. . . . . •
(POH) (PH20)
~. . . . .
-12.61
+
0.18
3.19
+
0.02
OH-
Cu
- 3.66
+ 0.06
2.72
_+ O.O6
(POH) 2
The error limits reported here are
2.94
4.66 10.67
_+ 0.04
- 7.13
_+ 0.03
-'20.5 4.68
+
0.02
4.75 6.32 8.52
Cu (P~B) (Pa20) =
0.08
Cu (POH) 2
÷
Cu
+
P o t e n t i o m e t r i c methods Glass electrode Copper selectrode
///= 0.i M KNO 3
-
Cu
Gu (P~20)2
- 3.96
H÷
~.==.=a
2POH-
*
+
Cu 2+
POB-
÷
H+
Cu PH202÷
~='===~
~=====~
Cu PH20
+
+ +
S+ B+
± ~ o" .
(11)
where Na ÷ is the concentration of NaOH used after each addition and NO3 is the concentration of nitrate ions other than that used for adjusting the ionic strength. From eqns (9)-(11) and introducing the first dissociation
Table 1. Log of the acid dissociation constants of 2-pyfidinecarboxaldehydeand its copper(II)complexes*
Hp. H2 O÷
(10)
"
- 5.55 - 7.72
+
0.06
The complexingability of 2-pyridinecarboxaldehydetoward cupric ions constant of the ligand we have [P'H20] =
150
(1--a)TL + [OH]- [H] [H÷l/k,
÷
I00
(13)
TM
In a similar way, the species CuPH20 and CuOHPOH were ignored in the pH range 3.4-5.5. A new set of equations were derived taking in consideration the new electroneutrality relation combined with eqns (9) and (10). [POH] =
(2 - a)TL - [HI + [OH] 2[H]: [H] klk2
k~_
[H]
[H]
~
(12)
and a, the average ligand number, is
ti =
1535
\
g
,
g_ (14)
5O
and TL - [POH] (k~V~+--~-2+1 ) vi T,
(15)
From pH 6.0 the value of a increases owing to the formation of probably hydroxo complex species. In such a case the species CuPOH ÷ and probably only CuOHPOH coexist. The stability constants were evaluated from the following relations: TL = TM= [CuPOH ÷] + [CuOHPOH]
(16)
\ 2
3
4
5
pH Fig. 3. Potential dependence of Cu-2PA system on pH. +, lx10-3-2x10-2M; L l x 1 0 3 - 5 x 1 0 - 2 M ; ©, 1x10-3-8x 10-z M; x, 1 x 10-3-1 x 10-' M; A, 1 x 10-2-1.5 × 10 ' M.
To. - [OH ] = [CuOHPOH] where ToK = concentration of hydroxyl ions at any point along the titration curve (above pH 6) and OH was evaluated from the corresponding pH values and K~ (= 10 '37~[7]). The average number of OH groups bound per mole of the first complex may be expressed by aoH -
To~- [OH] _ K4[OH] TM 1 + K4[OH] "
(17)
Numerical methods (8) were used to calculate the stability constants of the equilibria 5, 7 and 8. The value of K2 was calculated from the knowledge of the equilibrium constants K1 and K3. The titration curve for 1 : 2 metal to ligand ratio shows steep inflection at a = 2 (Fig. 2) indicating that the ligand species is POH . Two stability constants were determined for the following complex equilibria: CuPOH + , Cu(POHh,
' Cu 2+POH ' CuPOH ++POH .
,oo
Io'
,I
.% 10"2
\
(18)
(19)
The stability constants are shown in Table 1. (b) Potentiometric study using copper(H) selectrode. Figure 3 illustrates the dependence of the potential of the copper(II) selectrode as function of pH. The ligand concentration was always in excess with respect to the cupric ion concentration. A unique formation curve of o~o([Cu]/TM)as function of P.H20 was obtained for various initial concentrations of 2-pyridinecarboxaldehyde and cupric ions (Fig. 4). The free ligand concentration was
IO-3
oe the 10-3
i
l
I i l lll] I0-2
[PHzO] Fig. 4. The plot of log ~ o as functions of log [PH20].
1536
A.E. EL-HILALYand M. S. EL-EZABY
obtained from the following relation, [PH20] = = ITL
(20)
where
2PA
2 . 5 . 1 0 -I M
i
[H]J-' I I k, I=0
I
j =o
j
0.6
[HI'-J I-I k, 1=0
Approximate values of the stability constants of the following equilibria Cu+ P'H20. Cu+ 2P.H20. Cu+3P'H20. Cu+4P.H20.
" CuP'H20 ' Cu(P.H20)z ' Cu(P.H20)3 " Cu(P'H20)4
(21) (22) (23) (24)
,~ 0.4
were obtained using the following equation 0.2
1
= o- N 0
[Cu] TM
(25)
/3. [PH20]"
where/3 is the overall stability constant. The values so obtained were refined further by successive approximation and depicted in Table 1. Back calculation of the formation curve was made utilizing the refined values of the four/3's. (c) Spectrophotometric study. Figure 5 shows the spectra of Cu-2-pyridinecarboxaldehyde system at different pH values. A regular blue shift was observed as pH increases. Graphical methods [9] were used to determine the number of absorbing species present in solution. A conclusion was drawn that in the pH range 1.4-2.7 and 3.0--4.5 two absorbing species were present in both cases and a single absorbing species may exist in the pH range 5.0--8.9. Figure 6 shows the pH dependence of absorbance at the wavelength 670 nm. The absorbance is dependent on both the initial concentration of the ligand and metal ion. At constant metal ion concentration and various initial concentrations of the ligand the absorbance is dependent upon pH and the concentration of the ligand in the pH range 1.0-4.4. At pH's higher than 4.4 the absorbance is solely pH dependent. A unique formation curve was obtained when the average molar absorptivity [~ = (As/TM)] was plotted as function of the free ligand species (P.H20) in the pH range 1.0-4.2 (Fig. 7). The free ligand species was calculated using eqn (20). The absorbance data at 670 nm were analysed by the least square procedures. In this procedure, the sum of the squared deviations, s, given by eqn (27) s = ~ (~obs- ~c~o)z
0.0
500
700
900
11(30
X nm
Fig. 5. Absorptionspectraof Cu-2PAsystem.
1.6
\
!
/
1.2
/
\
\ \
io8F ,
(27)
-i-
should be minimum, where gobsand ~ o are the observed and calculated values of the average molar absorptivities and N~ is the number of the experimental points considered. In other way,
0.0 ~ / " l
N
(28) o
o
I
I
I
L
3
5 pH
7
9
Fig. 6. Variation of absorbance of Cu-2PA systemas functionof pH. x, 2xl0-L2.5xl0-'M; O, 3x10-L2.5x10-~M; [3, 3x10-3-4x10-1M; A, 3xl0-L5xl0-tM; @, 4x10-L2.5x 10-1 M; +, 8 x 10-L2.5x l0-1 M; ;t, 670rim; l, 4 cm.
The complexingability of 2-pyridinecarboxaldehydetoward cupric ions
1537
60
•
~" 4"0
2o
..5
-
..y7 0
t
t
i
I
I
I I II
I
I
I
I
I
I I I[
I0-2
10-3
I
I
[
I
[
I0 -I
[PH20] Fig. 7. The dependence of f on log [PHzO]. O, 2 x 10-3-2.5× 10-~ M; A, 3 × 10 3-2.5× 10 ' M; @, 3 × 10-~4 × 10-' M; V1,3 × 10-3-5× 10-' M; +, 4 × 10-3-2.5x 10 ~M. where E, the molar absorptivity of the nth complex and a stands for the free ligand, P.H20, calculated from eqn (20). Initially guessed values of the stability constants were used to solve for the N + I unknown ~ by Gauss-Jordan elimination method. Since the initially guessed values of the stability constants are not the optimum ones, it is necessary to adjust their values to reduce s. The method of the steepest descents[4, 11] was used to optimize the values of the stability constants. The molar absorptivities and the sum of the squared residuals then recalculated. The whole process may be repeated till refined/3's could be obtained from the initial guesses. The plot of s vs, e.g./3, or log/3, give the optimum value of/3, which can be used to get/3m from the approximate linear relation
/3,. =a,./3. +b,.
described in previous report [4], where am and b~ could be obtained from the plot of/3m vs/3,. The reactions (21) and (22) were assumed to take place in the pH range 1.0-4.2. Table 2 shows the initially guessed/3's and their rapid convergence values. Table 1 shows/3, and/35 obtained by this method. At pH's higher than 5, the following reactions were assumed to take place Cu(P.H~O)22+ .
' Cu(P.H20)(POH) + + H +; K ~ - [Cu(P.H20)(POH)][H] [Cu(P.H~.O)2]
Cu(P.H20)(POH)
+ .
• Cu(POHff
+ H+; [Cu(POHh][H] K6 = [Cu(P .H20)(POH)] '
n~m
Table 2. Initially guessed and rapid convergence values of the stability constants in 2-pyridinecarboxaldehyde copper(II) system* Initially guessed values
Rapid convergence values
~i
/32
Sigma
/31
/92
Sigma
0.50 x 102
0.50 x 102
1.2026
0.50 x 102
0.50 x 102
i x 102
i x 102
1.339
0.64 x 102
0.72 x 102
1.1715
5 x 102
5 x 102
2.8782
1.48 x i02
1.93 x 102
0.9425
1 x 103
1 x 103
3.529
3.27 x 102
6.22 x 103
0.80205
5 X 103
5 X 103
4.5515
6.02 x 102
1.40 x 104
0.74467
1,2026
I x 104
1 x 104
4.7741
1.67 x i05
4.56 x 104
0.74797
5 x 104
5 x 104
4.9897
7.94 x 103
2.46 x 105
0.81976
1 x 105
1 x 105
5.0198
1.90 x 104
6.10 x 105
0.84392
5 X 105
5 x 105
5.0445
7.25 x 104
2.38 X 106
0.85993
i x 106
I x 106
5.0476
1 . 6 7 x 105
5 . 4 7 x 106
0.86328
I x I07
i x 107
5.0505
1.47 x 106
4.80 x 107
0.86570
TM
8 x 10-3 M
TL
2,5 x I0-I M ffi 670 nm -
~X
4 cm
ffi 0.5 M KCI
Computer p r o g r ~ MONONUC - IBM 370
A. E. EL-HILALYand M. S. EL-EZABY
1538
It was assumed that at pH -4.5, Cu(P.H:O)2 is completely formed. If this was the case, eqn (3) could be used to calculate K5 and K6 (Table 1).
IR spectra of the copper complex The analysis of the metal complex prepared in this study was found to be compatible with the empirical formula, Ct2Ht204N:Cul. Assignments for the IR spectrum of the copper complex can be made by a comparison with the assignments for the parent ligand (Table 3). The vanishing of the C=O stretching band frequency at 1723 cm-l from the spectrum of the copper complex with the appearance of OH stretching band frequency at I 3500 cm-', indicates the presence of -C-OH grouping in I I place of -C=O. The appearance of C-O stretching band frequencies in the region of 982-1100 cm-l was a further I confirmation of the absence of --C=O group. The general shift of the pyridine band frequencies to higher values and the appearance of Cu--O and Cu-N stretching band frequencies at 360 and 510 cm-l, respectively, proves the involvement of nitrogen and oxygen atoms in chelation.
obtained from copper selectrode and spectrophotometric measurements, where excess ligand was used, are different since less hydrolysis and olation of the hexaaquocopper(II) ion may be expected specially in acidic medium. The stability constants obtained in this study for the formation of some complex species were compared with complexes having similar potential ligand (Table 4). It is clear that 2-pyridinecarboxaldehyde may act as a unidentate ligand and bind the copper(II) ion by the hetrocyclic nitrogen. On the other hand, it could act as a bidentate ligand and chelate the metal ion by the heterocyclic nitrogen and the oxygen of the aldehydic group. In the latter case the stability constants should be little higher than in the former. The pH-metric and spectrophotometric studies indicates that the deprotonated hydrated species may also form with the cupric ions a quite stable complex species in solution. We have not found to the best of our knowledge that such ligand species may be involved in complex formation with metal ions. The
I
O--Cu--O
DISCUSSION The 2-pyridinecarboxaldehyde copper(II) system exhibits different species under different experimental conditions. Results from potentiometric studies indicate that hydrolysis and olation of hexaaquocopper(II) ion would be expected to be considerable. However, results
I
I I
I
.ify
Fig.8.
Table 3. IR spectraof 2-pyridinecarboxaldehydeand its coppercomplex Frequency modes
llgand,
Pyridlne ring
cm-I
Copper complex,
cm-I
413(sh,w), 620(sh,w)670(sh,m)
432(sh,m),638(sh,m),670(sh,w),
770(sh,s),838(sh,s),lOOO(sh,s),
780(sh,s),820(sh,m)1445(sh,s),
1050(sh,m)lO95(sh,w),1445
1480(sh,s),1575(sh,w),1607
(sh,m),1478(sh,m),1593(sh,w)p
(sh,s),2720(br,m),2850(br,s),
2850(br,w),2850(br,m).
3040(shd),3080(sh,s).
~iphatic C-H
1375(sh,m).
1260(sh,m),1280(sh,m),1315(sh,m).
C-O
1723(sh,s)
?-c
l155(sh,w),1220(sh,s)
l160(sh,m),1220(sh,m) 980(sh,m),lO20(sh,m),lO50(sh,s),
C-0
1070(sh,s).
Cu - N
360(br,w). 510(sh,m).
O-H
3500(br,w).
Cu - 0
ah, w, s, br, m, shd stands for sharp, weak, strong, broad, medium and shoulder
Table 4. Stabilityconstantsof some copper complexes Co~o.nd
B1
/32
~3
/~
5.80
6.70
6.32
8.52
Pyridine*
2.6
4.54
P£colinlc acid** (chelate)
7.95
14.95
2-Pyridlnecarboxaldehyde
2.94 Cu(PH20) 10.67 Cu(POH)
4.75 Cu(PH20) 2 20.5 Cu(POH) 2
The complexingabilityof 2-pyridinecarboxaldehydetoward cupric ions analysis has established the absence of anions indicating that the complex is neutral with 1 : 2 metal to ligand ratio. Moreover, the loss of C=O stretching band frequency and the appearance of OH in addition to the C-O band frequencies in the IR spectrum of the complex is a strong evidence of the involvement of the deprotonated hydrated ligand species in the metal complex (Fig. 8). This complex may be deceived with copper picolinate complexes. Samples of Cu picolinate complex was prepared and their IR spectrum was compared with that of 2pyridinecarboxaldehyde copper(II) complex. The spectrum of the latter was quite different in many respects from the former. The C=O stretching band frequency in the former exist at slightly lower frequency ( 1725 cm ~~ 1650 cm ~) as compared to that of picolinic acid. Moreover, the stretching frequency band of the OH of the carboxylic group is lost upon complex formation.
REFERENCES
t. L. G. Sill6n and A. E. Martell, Stability Constants of Metal-Ion Complexes, Spec. Publ. No. 17 (1964); Supplement
1539
No. 1, Spec. Publ. No. 25 (1971). The Chemical Society, London. 2. H. Diebler and R. N. F. Thomeley, J. Am. Chem. Soc. 95(3), 896 (1973). 3. G. Schwarzenbach and H. Flaschka, Complexometric Titrations (Translated by H. M. N. H. Irving), 2nd English Edn. Methuen, London (1%9). 4. M. S. El-Ezaby and N. Gayed, J. Inorg. Nucl. Chem. 37, 1065 (1975). 5. M. Osman, T. M. Salem and M. S. E1-Ezaby,J. Chem. Soc. (A), 1401 (1971). 6. R. L. Gustafson and A. E. Martell, Arch. Biochem. Biophys. 68, 485 (1957). 7. R. G. Bates, Determination of pH, p. 74. Wiley, New York (1964). 8. F. J. C. Rossotti and H. Rossotti, The Determination of Stability Constants. McGraw-Hill,New York (1%1). 9. J. S. Coleman, L. P. Varga and S. H. Mastin, Inorg. Chem. 9, 1015 (1970). 10. S. Feldherg, P. Koltz and L. Newman, Inorg. Chem. It, 12, 2860 (1972). 11. W. E. Grove, Brief Numerical Methods. Prentice-Hall, London (1966). 12. M. S. Sun and D. G. Brewer, Can. J. Chem. 45, 2729(1967). 13. G. Anderegg, Heir. Chim. Acta 43, 414 (1960).