SPECTRA AND DISSOCIATION CONSTANTS OF NICOTINAMIDE
173
NICOTINAMIDE
ULTRAVIOLET ABSORPTION SPECTRA AND DISSOCIATION CONSTANTS H. H. G . JELLINEK AND MARGARET G. WAYNE The Laboratories, J . Lyons and Company, Ltd., Hammersmith Rood, London W.14,England Received April df, 18.48
The ultraviolet absorption spectra of nicotinic acid and pyridine and their dissociation constants have been determined and communicated in two previous papers ( 2 ) . The hydrolysis of nicotinamide in hydrochloric acid solutions was also investigated (3), the analysis of the reaction mixtures being carried out polarographically. A reaction mechanism was proposed which accounts satisfactorily for the experimental results; it consists essentially of two reactions. An equilibrium is set up rapidly H+RCONH*
+ H+ e H+RCONH:
followed by a slow hydrolysis: H+RCONH:
+ H30++ H+RCOOH + NH: + H+
The dissociation constant for the formation of H+RCO?jH: has been measured spectroscopically in this work and found to be in good agreement with the dissociation constant which had to be assumed on the basis of the experimental results of the hydrolysis. In addition to the determination of the dissociation constant due to the amido group, the thermodynamic dissociation constant of the nitrogen in the ring has been evaluated spectroscopically. EXPERIMEKTAL
Xaterials: The nicotinamide used was a British Drug Houses product, m.p. 128°C. (corr.) A portion of this sample was recrystallized from benzene and dried. The absorption spectra of the two samples mere compared and found to be identical within experimental error. The original sample was therefore used in all experiments. The sodium acetate, acetic acid, sodium chloride, and hydrochloric acid were of Analar standard. Apparatus: The measurements were made on a Hilger Medium quartz spectrograph, using Kodak B.30 plates. pH values were determined by means of a glass electrode. Solutions: The concentration of nicotinamide in all final solutions was of the order of 0.0003 M . Aqueous stock solutions were made up freshly each week. The buffer solutions were mixtures, in different proportions, of a 0.01 M solution of acetic acid and a 0.01 M solution of sodium acetate except in the case of solutions with pH values smaller than 3, when hydrochloric acid was used. Sodium chloride was added to bring the final buffered solutions to an ionic strength of 0.01.
174
E, H. 0. JELLINEK AND MARQARET
(3.
WAYNE
ABSORPTION SPECTRA
The absorption spectrum of nicotinamide has been determined in water by Warburg, Christian, and Griese (5), who found a maximum absorption at 2600 .&.,with a molecular extinction coefficient of 4.5 X loa.In this work, the absorption spectra have been determined over a range of pH values. 6.5
I 0’
6.0
5.5 5.0 4.5
Y 0
5
4.0
V
r 3.5 B
t F 3.0 X w
2.5
5 2
2 2.0 z I. 5
I.o 0.5
>
2500
2600 WAVELENGTH
2700
2 800
f.
FIG.1. Ultraviolet absorption spectra of nicotinamide at various pH values
Figure 1 shows the spectra obtained for the following pH values: 5.72, 4.83, 3.83, 3.72, 3.48, 2.75, 2.05, 1.70, 1.30, 0.7, 0.3, 0.0, -0.4, -0.7, -0.8, -1.06. All pH values have been measured with the glass electrode except the negative values, which have been calculated from the concentrations of hydrochloric acid. The pH values from 5.72 to 3.48 have been obtained by means of an acetic acid-sodium acetate buffer. These buffered solutions were all adjusted to an ionic strength of 0.01 by the addition of requisite amounts of sodium chloride. The lower pH values were obtained by adding increasing amounts of hydrochloric acid. The concentration of nicotinamide in all solutions was of the order of 0.0003 M . The spectra of nicotinamide show absorption maxima at 2616 A,, the same
SPECl'RA AND DISSOCIATION CONSTANTS OF NICOTINAMIDE
175
wave length at which nicotinic acid also has its absorption maximum. The molecular extinctions of the maxima decrease as the pH values of the solutions increase, the maxima remaining a t 2615 A. The spectra fall into two groups, characterized by their isobestic points. However, only the two isobestic points for the larger wave lengths are fairly well developed, one lying approximately at 6 .I
-3 2
,x td
5.7.
-4
5.3.
w IA
s
4.9.,E,
EXPERIMENTAL POINTS
4.5.
CALCULATED CURVE
2
z
o
-
4.1.
q
3.7.
2
a 0
5 3.3.
r
0
2.9.,E, 2.54
. .
- 2 - 1
,
,
0
I
. .
2
3
,
,
4
5
.
6
PH
FIQ.2. Dissociation curve for
nicotinamide a t X = 2615
A.
2710 A., the other a t about 2720 A. The first one belongs, as will be shown below, to the dissociation equilibrium due to the nitrogen in the ring of the amide:
CONH2
ACONH2
+HzO=
I,!
+
OH-
(1)
Hf
The other belongs to the dissociation equilibrium due to the basic nature of the amido group:
CALCULATION O F DISSOCIATION CONSTANTS
Figuore 2 shows the molecular extinction coefficients for the peak values (X = 2615 A.) plotted against the pH values. Table 1 contains the corresponding pH and peak valuee. The dissociation constants are calculated as follows: For each of the two equilibria considered separately, the following relationships are
176
n.
H.
Q. JELLINEK AND MARQARET G. WAYNE
TABLE 1 Absorption mazima at different p H values 1
PH 5.72 4.85 3.83 3.72
~
1 ~
'
I
e (MOL--
a0l.l
EjmNmiott)
2.897 X 10' 2.897 3.355 3.508
PH 1.02 0.70 0.30 0.00
e "mi EXbCnON)
(U0-n
5.077 X 10' 5.291 5.495 5.698
and
where
el is the molecular extinction coefficient
of
H+ €2
is the molecular extinction coefficient of
'Iv/
Hf
e3
is the molecular extinction coefficient of
ncom2 "/
e.
is the molecular extinction coefficient of solutions having extinctions between 03 and €1, or €1 and €2, respectively, pK, is the negative logarithm of the dissociation constant of water, and pK1 and pK11 are the negative logarithms of the dissociation constants of equilibria I and 11, respectively.
SPECTRA AND DISSOCIATION CONSTANTS OF NICOTINAMIDE
177
pK values are obtained by putting experimental e= and corresponding pH values into equations l a and l b , respectively. The values for el, t2, e3 can be read from figure 2, and are as follows: =
2.89
x
103; f1 = 4.89
x
103; e 2 = 5.90
x
103
Those pK values are chosen which give the best agreement with the experimental points. The values found are (pK, has been taken as 14, the average temperature was 2OOC. f 2'): ~ K =I 10.60;
KI = 2.5 X lo-''
PKII = 13.50;
KII = 3.16 X lo-"
and The curve drawn through the experimental points in figure 2 has been calculated by using these pK values and the values for et, €2, €3 given above. The dotted lines show the dissociation curves which are obtained by considering each equilibrium separately. In point of fact, the equilibria interfere with one another appreciably in the region of pH values from cu. 1.5 to 2.5. This region of the curve has been calculated as follows: The degree of dissociation, (YI, with respect to
Pro", 'X/
H+ is given by
and that with respect to
n
CONH'
'S/
H+
by
Further c,
=
ale,
+
a?f2
+ (1 - a1 - ade3
178
H. H. 0. JELLINEK AND MARQARET G. WAYNE
K I , KII, [Hf], and K , are known; hence a1 and a2 can be calculated and put into equation 3, from which e= can then be obtained. This part of the calculated curve is also shown in figure 2. The thermodynamic dissociation constant for equilibrium I can be evaluated according to Debye and Huckel, as the ionic strength has been kept constant a t 0.01. The correction due to the activity coefficient of the ion
0
CONHe
H+
amounts to log = - 0.051 (compare paper on nicotinic acid (2)). Hence pK: = 10.65 and K; = 2.24 X lo-". DISCUSSION
As mentioned above, the ultraviolet spectra present further evidence in favor of the mechanism of the hydrolysis of nicotinamide in hydrochloric acid solutions discussed in a previous paper (3). The reaction scheme which accounts satisfactorily for the experimental results is as follows:
followed by
Hf
Hf
The experimental rate constants could be expressed as a function of the hydrogen-ion concentration by the equation:
Ks is the dissociation constant of the equilibrium expressed in equation 4. Its value at 76°C. was found to be 0.67. Equation 6a for 76°C. reads as follows;
k
=
7
x
10-~ 0'67[H7 [H80f]set:' 1 0.67[Hf]
+
The dissociation constant found spectroscopically a t 20°C. is 3.16 X lo-". This constant belongs to the same equilibrium as that of equation 4 only written somewhat differently:
SPECTRA AND DISSOCIATION CONSTANTS OF NICOTINAMIDE
179
+ OH-
(7)
Ka and KII are related by the following equation,
-.
~
__
THERSIODYSAYIC BASE CONST.\ST FOP NITROGEN IN PINO
~
1
THERMODYNAMIC ACW CONSTANT
I
.
P y r i d i n e . . . , . . , . .. . , . . Xicotinic a c i d . ... . . . . , Xicotinamide.. . , . . . . . .
1 . 3 2 X 10-@(2) 3 . 5 5 X 10-11 (2) 2 . 2 4 X 10-l'
1.17 X
(2)
~
BASE CONSTAhT OF AYIDO GROUP
i
i
3.16 X 3.1 X
180
LLOYD E. SWEARINGEN AND ALFRED F. SCHRAM
2. Two equilibria have been found whose dissociation constants have been evaluated. 3. The dissociation constant of the amido group has been found to be in good agreement with the dissociation constant derived previously from the kinetics of the hydrolysis of nicotinamide in hydrochloric acid solutions. The authors wish to thank Dr. L. H. Lampitt and Dr. E. B. Hughes for their interest in this work. They are also indebted to Mrs. B. A. Ambrose for help with part of the experimental work and to Messrs. J. Lyons Ltd. for permission to publish. REFERENCES (1) Handbook of Physics and Chemistry. Chemical Rubber Publishing Company, Cleveland, Ohio (1945). (2) HUGHES,JELLINEK, A N D AMBROSE: J. Phys. & Colloid Chem. 63,410,414 (1949). (3) JELLINEK AND GORDON: J. Phys. & Colloid Chem. 63, 996 (1949). (4) PAULING: The Nature of the Chemical Bond, 2nd edition, p. 208. Cornell University Press, Ithaca, New York (1945). ( 5 ) WARBURG, CHRISTIAN, AND GRIESE:Biochem. Z . 281, 157 (1935).
SUBSTITUTED AMIXES AS INHIBITORS IN THE ACID CORROSION OF STEEL1 LLOYD E . SWEARIRGEN
Department of Chemistry, University of Oklahoma, Norman, Oklahoma AND
ALFRED F. SCHRAM
Department of Chemistry, Southwestern Institute of Technology, Weatherford, Oklahoma Received January 3, 1950
The corrosion of iron and steel and methods of reducing the rate at which corrosion of these materials takes place have been the subject of much investigation. The corrosion process and also the corrosion rate depend upon a variety of factors, which may be divided into two principal groups. In one group, we may place those factors that are mainly associated with the metal itself, such as the homogeneity of the surface, its inherent ability to form a protective film, the chemical and physical character of the metal, and the hydrogen overvoltage on the metal surface. In another group, we may place those factors that are mainly associated with the environment of the metal, such as the kind and con1 This paper is an extract from a thesis submitted by Alfred F Schram to the Graduate College, University of Oklahoma, in partial fulfillment of the requirements for the Ph.D. degree.
