The major differences between the spectra of these two polymers are due to the greatly increased concentration of piperazine rings in the poly(ethy1eneamines). The piperazine rings are supported by absorptions at 935 cm-l, 1010 cm-’, 1160 cm-l, and 1340 cm-1. An estimation of their concentration is made possible by ratioing the absorption of the 1010 cm-l ring mode band to the -NH2 absorption at 1605 cm-l. As previously stated, the absorptivities of these two bands and a knowledge of the primary amine content of the polymer are
requisites for quantitative analysis. The infrared analysis indicated a piperazine ring concentration of 27% in the poly(ethyleneamine) product shown in Figure 2. ACKNOWLEDGMENT The author is grateful to G. E. Ham for helpful discussions during the course of this work.
RECEIVED for review November 14, 1968. Accepted April 10,1969.
Spectrophotometric Studies of Iron(l1) Complexes of PyridyI Benzodiazepin-2-ones J. D. Sabatino, 0. W. Weber, G . R. Padmanabhan, and B. Z. Senkowski Analytical Research Laboratories, Hoffmann-La Roche Inc., Nutley, N . J . 07110
Ferrous ions form purple complexes with 7-bromo-1,3dihydro-5-(2-pyridyl)-2H 1,4- benzodiazepin-2-one (A) and 7-bromo-1,3-dihydro-l-(3-dimethylaminopropyl)-5(2-pyridyl)-2H-1,4-benzodiazepin-tone di hydrochloride (B). Compounds A and B have the basic dipyridyl type bond structure which has excellent metal complexing properties. Iron(l1) complexes with A and B are pH dependent, and the optimum pH range for the formation of these complexes is between 5 and 7. The mole ratio of ligand to metal for these complexes is 3 to 1. The formation constants of the A and B complexes are 8.9 X 10” and 3.2 x lo”, respectively. The effect of various anions and cations on the formation of these complexes i s discussed. These complexes are adapted to the determination of iron. Some advantages of compounds A and B over the dipyridyl type are presented.
-
COMPOUNDS A (I) and B (2) have the following structures: H
Table I. Preparation of Stock Buffer Solutions Final volume, PH Components ml 1 20 ml of 0.5M oxalic acid 3.6 ml of 100 1.OM potassium hydroxide 100 2 20 ml of 0.5M oxalic acid $8.5 ml of 1.OM potassium hydroxide 100 3 50 ml of 0.1M potassium acid phthalate 22.3 ml of 0.1M hydrochloric acid 100 4 50 ml of 0.1M potassium acid phthalate 0.1 ml of 0.1M hydrochloric acid 100 5 50 ml of 0.1M potassium acid phthalate 22.6 ml of 0.1M sodium hydroxide 100 6 136 grams of sodium acetate 6 ml of glacial acetic acid 100 7 10 ml of 0.5M THAMa 46.6 ml of 0.1M hydrochloric acid 8 10 ml of 0.5M THAM 29.2 ml of 100 0.1M hydrochloric acid 100 9 10 ml of 0.5MTHAM 5.7 ml of 0.1M hydrochloric acid 4 Tris (hydroxymethyl) aminomethane.
+
+ + +
+
+ +
+
dN A
B
These compounds are similar to a,a’-dipyridyl in having the following basic structure: -N
N-
Because of this structural similarity, it was of interest to investigate the nature of the iron(I1) complexes of A and B, compare these complexes with those obtained with the (1) R. Ian Fryer, R. A. Schmidt, and L. H. Sternbach, J. Pharm. Sci., 53, 264 (1964). (2) L. H.Stembach, G. A. Archer, J. V. Earley, R. Ian Fryer, E. Reeder, N. Wasyliw, R. 0. Randall, and R. Banziger, J. Med. Chem., 8 , 815 (1965).
dipyridyl type compounds (3-IO), and evaluate their relative merits for the determination of iron. EXPERIMENTAL Apparatus. Optical absorption measurements were made with Beckman Model DU and Model DK spectrophotometers. A Corning Model 12 pH meter was employed for pH determinations. Solutions. Stock buffer solutions were made as described in Table I. (3) W. W. Brandt, F. P. Dwyer, and E. C . Gyarfas, Chem. Rev., 54, 959 (1954). (4) H. Diehl and G. F. Smith, “The Iron Reagents: Bathophenanthroline, 2,4,6-Tripyridyl-s-Triazine,Phenyl-2-Pyridyl Ketoxime” G. Frederick Smith Chemical Co., Columbus, Ohio, 1960. ( 5 ) D. D. Bly and M. G. Mellon, ANAL.CHEM.,36, 1276 (1964). (6) T.S. Lee, I. M. Kolthoff, and D. L. Leussing, J. Amer. Chem. Soc., 70,3596 (1948). (7) G.Anderegg, Helv. Chim. Acta, 46, 2397 (1963). (8) J. H.Baxendale and G. George, Narure, 162,777(1948). (9) A. A. Schilt and G. F. Smith, J. Phys. Chem., 60,1546(1956). (10) M. A. McKenzie, Ausr. J. Chem., 8 , 569 (1955). VOL. 41, NO. 7, JUNE 1969
905
0.5
-
0.4
-
W
2 0.3
m K 0 v)
:0.2 -
0.1
0.0
-
I 1 360
I 400
I 450
I
I
I
I
500
550
600
650
WAVELENGTH
mu
Figure 1. The effect of pH on the absorption spectra of iron(1I) A complex Cre = 2 X IO-'Mand CA = 1 X 10-'Mh u and b CF. = 2 x 10-6Mand CA = 1 X 1 0 - a M hc through g pH of u = 1.5, b = 2.0, c = 3.0, d = 4.0,e = 5.0,f = 6.9, and g = 8.1
0.5
0.4
W
y
0.3
v)
9
0.2
0.I
0.0 360
4 00
450
500 WAVELENGTH
550
600
650
mu
Figure 2. The effect of pH on the absorption spectra of iron(1I) B complex CF. = 2 X IO-'M and CB = 1 X 10-'M h 0 and b C F ~= 2 X 10-6Mand CB = 1 X 10-aMin c g pH of a = 1.5, b = 2 . 2 , ~= 3.0, d = 4.1, e = 5.0,f = 6.9, and g = 8.1
A lOmM stock solution of compound A (1) was prepared in a 100-mlvolumetric flask by dissolving 316.0 mg of the solid sample in 50 ml of methanol and diluting to volume with water. A lOmM solution of compound B (2) was prepared by dissolving 474.3 mg of the solid sample in 100 ml of water. A lOmM stock solution of ferrous sulfate was freshly prepared by weighing accurately 278 mg of ferrous sulfate (Maltinckrodt analytical reagent) and dissolving in 100 ml of water. Further dilutions were made to obtain the required concentrations and the resulting solutions were immediately used. Solutions were prepared by taking 1 ml of ferrous sulfate, 1 ml of ligand solution, and 8 ml of the appropriate stock buffer solution. Reference blank solutions were also prepared 906
ANALYTICAL CHEMISTRY
substituting 1 ml of water for 1 ml ferrous sulfate. The solutions were mixed well and their absorbances measured after 10 minutes against the respective blank solutions. Effect of pH on the Complexes. In a series of solutions, the pH was changed from 1 to 8 maintaining an iron concentration of 1 x 10-6M and ligand concentrations of 1 X 10-aM, respectively, Because the absorption of solutions at pH 1 and 2 were low, ferrous ion concentration was increased to 2 X lO-4M. The spectrum was recorded from 650 mp to 360 mp. The results are shown in Figures 1 and 2. Determination of Ligand to Metal Ratio. JOB'S METHOD (11). In a series of solutions, the total concentration of metal and ligand was kept constant at 1 X lO-'Mand the individual (11) P. Job, Ann. Chim. (Paris), 9,113 (1928).
04001
0,4ml
0.360
0 360
0.320 ’ 0.260
g
I
0.240 ’
sIn h 4
42
0.200
0,160
L
s 0.120 0.040
0.000 0.0
1.0
MCUR MOLAR CONCENTRATION OF F E R R O U S ION
x io5
Figure 3. Job’s curve. Effect of metal-ligand mole ratio on iron(I1) A complex
+
The ratio CA:CF~(II) was varied from 19:l to l:O; (CFa(I1) CA) was maintained constant at 1.0 X lO-*M. The solutions were buffered at pH 5. The curve showed the maximum absorbance to be at ligand to metal ratio of 3.1
concentrations were varied. The absorbances measured were plotted against the molar concentration of ferrous ion. The mole ratio at the maximum absorbance gave the mole ratio of the complex formed (Figures 3 and 4). MODIFIEDYOE AND JONES’METHOD(12). The metal ion concentration was maintained constant in a series of solutions and the Concentration of ligand was varied. Readings were taken at three different concentrations of metal ion (3.6 X IO-SM, 7.2 X 10-5M, and 10.8 X 10-5M). A pH of 5 was maintained for all solutions. The absorbances us. the ligand concentrations were plotted (Figure 5 ) . The absorption varied linearly at low ligand concentrations and remained constant at high ligand concentrations. The ratio of ligand to metal at the point where the two linear sections intersected gave the ratio of ligand to metal in the complex. Determination of Stability Constant. A series of solutions was prepared varying the metal concentration from 2 X 10-6M to 7 X 10-6Mand the ligand concentration from 6 X 10-6M to 21 X lO-5M. These solutions were all buffered at pH 5 and the absorbances were measured ten minutes after preparation of the solutions. The concentrations of iron and ligand were chosen so that the amount of complex formed was in the order of 1 X 10m5Mto 5 X lO-5M. The absorbances of these solutions ranged from 0.2 to 0.8. The concentrations of the complexes were calculated assuming that the absorbance at 580 mk was due to only the 1 :3 complex. The free metal and free ligand concentrations were calculated by difference. The conditional formation constant K was calculated using the formula: K = [C,I/[CF,I[CLJ~. The calculated values are given in Table XI. Analysis of Iron Using Compound A and B. A standard calibration curve was prepared from the measurements of a series of solutions varying in ferrous ion concentration from 2 x 10-6M to 1 X 10-4M. A ligand concentration of 10-aM and a pH of 5 were maintained for all solutions. The ab(12) J. H. Yoe and A. L. Jones, IND.ENG.CHEM., ANAL.ED., 16, 111 (1944).
21)
3.0
4.0
5.0
&O
C O N C E N T R A ~ NOF FERROUS ION x
ZQ
8.0
SO
lo5
Figure 4. Job’s curve. Effect of metal-ligand mole ratio on iron(1I) B complex The ratio CB:CF~(II) was varied from 19:l to 1:19; (CF ~ ( I I ) + CB) was maintained constant at 2.0 X 10-‘M. The solutions were baered at pH 6. The curve showed the maximum absorbance to be at a ligand to metal ratio of 2.8
IdO
t
/
CONCENTRATION OF LIGAND IN mM
Figure 5. Modified Yoe and Jones Method. Effect of ligand A concentration on the complex. All solutions were buffered at pH 5 A. C F = ~ 10.8 X 10”M; CA:C F = ~ 4.6 B. C F = ~ 7.2 x 10*M; CA: CFe = 5.4 C. Cpe = 3.6 X 10”M: CA:CF,= 8.1
sorbances of the solutions were measured at 580 mp against appropriate blanks. Sample solutions containing unknown quantities of iron(I1) were treated similarly with buffer and VOL. 41,NO. 7, JUNE 1969
907
Table 11. Calculations for K Fe (II) A Complex
Sample
10‘
x
CFe‘
2.06 3.09 4.12 5.15 6.18
1 2 3 4 5
x CLB) 6.00 9.00 12.00 15.00 18.00
106
IOS x [ C J d 1.15 2.09 2.99 3.91 4.83
Ac 0.211 0.384 0.555 0.720 0.889
10‘
x
[CFe]’
0.913 1.003 1.131 1.237 1.348
los X [CL$ 2.56 2.74 3.03 3.26 3.50
X Kg 7.49 10.1 9.50 9.13 8.36 8.87
10-13
~
Fe @) B Complex 10’ x CFe 4.28 5.35 6.47 7.49
Sample 1 2 3 4
106 x CLg 12.00 15.00 18.00 21 .oo
A 0.148 0.252 0.365 0.520
10’
x [CCI 0.85 1.45 2.10 2.99
10‘
x
[CFe]
3.43 3.90 4.32 4.50
10‘ X [CLJ 9.45 10.65 11.7 12.03
CFe = total molar concentration of iron. CLg = total molar concentration of ligand. c A = absorbance at 580 ml.
lo-” X K 2.94 3.08 3.04 3.82 3.20
a
d 6
[C,] = calculated molar concentration of the complex formed. = calculated molar concentration of free iron.
[CFe]
f
[CLg]= calculated molar concentration of free ligand.
0
K
= formation constant.
ligand reagents. From the absorbance measurements at 580 mp, the concentrations of ferrous ion present were determined from the standard calibration curve.
Table 111. Molar Absorptivities and Conditional Formation Constants of Iron(II) Complexes
RESULTS AND DISCUSSION
Effect of pH on the Iron(I1) Complexes of A and B. The effects of pH on the spectra of the iron complexes are shown in Figures 1 and 2. The absorbance at 580 mp showed minimum change in the pH range of 4 to 7. Hence pH 5 was chosen for further studies, The molar absorptivities at pH’s below 4 and at pH’s above 7 were found to be lower than in the pH range 4 to 7. The effect of low pH on the complexes was similar to the effects observed by Nielands (13) in the case of the ferric hydroxamate complex and by Krumholz (14) in the case of the iron(I1) dipyridyl complex. Nielands pointed out that at low pH’s, the hydroxamic acids tended to form a 1 :1 complex with iron due to the protonation of ligands, and the complexes dissociated to free metal and ligand at high acid concentration. Lee, Kolthoff, and Leussing (6) presented the following mechanism for the pH effects on the iron(I1) o-phenanthroline complexes :
+ (phen) e Fe (phen)2+ Fe (pher02+ + (phen) S Fe (phen)?+ ki Fe (phen)2+ + (phen) S Fe (phen)2+ Fez+
(phen)
+ H+
S H phen+
At low pH’s free o-phenanthroline is protonated thus decreasing the rate of formation and concentration of the complex. At high pH’s complexation is inhibited due to depletion of ferrous ion by the formation of the ferrous hydroxide precipitate. Krumholz (14) and Burgess and Prince (15) proposed a mechanism for the dissociation of iron(I1) dipyridyl com(13) J. B. Nielands, Science, 156,1443 (1967). (14) P. Krurnholz, Nature, 163, 724 (1949). (15) J. Burgess and R. H. Prince, J. Chern. SOC.,1%5, 6061. 908
0
ANALYTICAL CHEMISTRY
Ligand cu,a’-Dipyridyl(7) 2,2’-Dipyrimidines (5) o-Phenanthroline (9) A B
Molar absorptivity
log K
8,820 4,540 11,100 18,400 17,400
17.1 7.5 21.3 13.9 11.5
plexes in acid medium which involved the decomposition of the protonated species F e Dipy,H3+]. The effect of adding acid to the iron(I1) complexes of compound A, a,a’-dipyridyl and o-phenanthroline is presented in Figure 6. This figure showed that once formed, the compound A complexes were exceptionally stable in acid medium. Determination of Molar Absorptivity of the Complexes. Table I11 gives the values obtained for the molar absorptivities for the iron(I1) complexes of A and B. Because the molar absorptivities of the A and B complexes with iron are almost twice that of the dipyridyl type complexes, the sensitivity of photometric ferrous ion determinations with ligands A and B is nearly doubled. Determination of Molar Ratio of Ligand to Metal. JOB’S METHOD (11). Figures 3 and 4 show that the ligand to metal mole ratio for the iron(I1) complex of both A and B is 3. It should be pointed out that Job’s method of calculating the mole ratio is reliable only for a 1 :1 type complex. YOEAND JONES’METHOD(12). The results obtained by this method are presented in Figure 5. The ratio of ligand to metal for the iron (11) A complex varied from 8 to 4.8 when the iron concentration was changed from 3.6 X 10-‘M to 10.8 X lO-5M. When the apparent ligand to metal ratio was plotted against the reciprocal of metal concentration and the plot extrapolated to I / c F e l + = 0, the ratio of ligand to metal was found to be 2.8 for the A complex.
TIME IN MINUTES
Figure 6. The effect of adding acid to iron(I1) complexes of compound A, o-phenanthroline, and qat-dipyridyl Two and one-half milliliters of lO-*M was added to 2.5 ml of lO-'M ligands. The solutions were allowed to stand for 10 minutes. Two and one-half milliliters of 1N HCI was added to each and the solutions diluted to 25-ml volume with water. The absorbance of each solution was read at the appropriate maximum, at various time intervals. Curve A represents absorbance readings of compound A complex at 580 mp, Curve B, o-phenanthroline complex at 515 mp,and Curve C, a-a'-dipyridyl complex at 525 mr
The following structure was ascribed to the 1 : 3 iron(I1) complexes of pyridyl benzodiazepinones.
where R is aliphatic Formation Constants of Iron(I1) Complexes of A and B. Table I1 shows the obtained data and the calculated value of K for the iron(I1) complexes with ligands A and B. Even though the molar absorptivities of the A and B complexes are much higher than those of a,a'-dipyridyl and o-phenanthroline, the formation constants are lower. The lower K values for A and B are probably due to the size of the ligand molecules. However, the accuracy of the ferrous ion determination is not affected by this lower K value; because the ligand concentration is always in excess, complete complexation is obtained. Determination of Iron Using Compounds A and B. As mentioned earlier, compounds A and B offered higher sensitivity for the determination of iron. These complexes followed Beer's law from 2.5 to 30 pg of iron per 100 ml of final solution. An ammoniacal solution of these complexes
was extracted with dichloromethane. The dichloromethane extractions obeyed Beer's law from 5 to 30 pg of iron per 25 ml of final solution. The water solubility of the dihydrochloride compound B eliminated the introduction of organic solvents to effect solution. Thus B offered an additional advantage over compound A, a,a'-dipyridyl, and o-phenanthroline. The bathochromic shift found with the iron complexes of compounds A and B, reduces potential interferences due to extraneous color common at lower wavelengths. The effect of the anions (phosphate, tartrate, citrate, and borate) on the complexes of iron with A and B were investigated. No interferences were observed with these anions. As might be expected, in the presence of EDTA, no complexing of iron with compounds A or B was observed. The cations Zn2+ (996 pg/ml), Pb2+ (1,282 pg/ml), Ala+ (266 pg/ml), Mn2+(566 pg/ml), and Ni2+(476 pg/ml) did not form colored species absorbing at 580 mp. Cobalt (2+) and copper ( + l and +2) formed colored complexes with compound A. However, if the concentrations of Co2+,Cu2+, and CUI+were less than 50 pg/ml, 77 pglml, and 135 pg/ml, respectively, no interference was observed. ACKNOWLEDGMENT
The authors acknowledge the helpful discussions held with Mr. H. L. Newmark and Dr. R. I. Fryer. RECEIVED for review January 15, 1969. Accepted March 18, 1969.
VOL. 41, NO. 7, JUNE 1969
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