DETERMINATION OF EXCITED STATEpK, VALUESUSINGPHOTOPOTENTIOMETRY
3857
Determination of Excited State pKa Values Using Photopotentiometry
by Donald D. Rosebrook’ and Warren W. Brandt Chemistry Department, Kanaaa State University, Manhattan, Kansas
(Received June 6 , 1966)
Photopotentiometry is herein defined as the measurement of the potential developed between one illuminated electrode and one dark electrode in a solution. This potential, which is primarily a function of the various species in solution, was used to produce data leading to excited state pK, values. The values found were -2.89 and 12.3 for the 2naphthylamine cation and anion, respectively; 13.5 for the 1-naphthylamine anion; and 4.37 and 9.5 for the 3-pyridinol anion and cation, respectively. The presence of ethanol in the aqueous systems acted to shift the position of the excited-state pK, values and to enhance the potentials. The first effect is attributed to differences in solvating power while the cause of the second effect is not clear.
I. Introduction Forster was the pioneer in the study of pK, values of the excited state. His studies consisted of the examination of hydroxypyrene derivatives2” and naphthalene derivatives.2b Wellera-5 has given extensive consideration to the determination of the physical constants of protolytic reactions of the excited states of the mononaphthols. Derkacheva6 and Hercules and Rogers’ have reported pK, values for excited naphthalene diols; however, there is little agreement in their results. The pK,* values determined by these workers vary by as much as 1.8 units, and neither worker has found values which are consistently high or low. All of the above studies produced the same result: bases become weaker in the excited state by 5-10 pK units and acids become stronger in the excited state by 5-10 pK units. Kokubuns reported the excited state pK, for acridone and found indications that the heterocyclic nitrogen became more basic in the excited state, as well as indications that in this type of molecule the hydroxy group became a stronger acid in the excited state. Wellerg noted that the heterocyclic nitrogen in acridine also becomes a stronger base in the excited state. Bartok, et al.,1° measured the dissociation constants of some excited phenols. They also measured pK,* for phenol in glycerol and found that the value of pKa* - pK, was significantly larger in the nonaqueous solvent than in water. Up to this time excited singletrstate pK, values have
been measured by fluorescence. This paper presents a new technique which we have called photopotentiometry. The technique has been applied to the determination of excited singlet-state pK, values. Levin and White”-l3 first showed that a positive correlation existed between the photopotential and the wavelength of the incident light such that a plot of photopotential vs. wavelength bore a very strong resemblance to the absorption spectrum. They were unsuccessful in an attempt to correlate fluorescence and what they called “photovoltaic” behavior Surash14 later showed that in deoxygenated solutions the primary reaction was a photoreduction and that (1) To whom correspondence should be addressed at Midwest Research Institute, Kansas City, Ma. (2) (a) T. Forster, 2. Elektrochem., 54, 42 (1950); (b) ibid., 54, 531 (1950). (3) A. Weller, ibid., 56, 662 (1952). (4) A. Weller, 2. Physik. Chem. (Frankfurt), 3 , 238 (1955). (5) A. Weller, ibid., 17, 224 (1958). (6) L. D. Derkachevil, Opt. Spectry., 9, 209 (1960). (7) D . %I. Hercules and L. B. Rogers, Spectrochim. Acta, 393 (1959). (8) H. Kokubun, 2. Elektrochem., 62, 599 (1958). (9) A. Weller, ibid., 61, 956 (1957). (10) W. Bartok, P. J. Lueehesi, and N. S. Snider, J. A m . Chem. SOC., 84, 1842 (1962). (11) I. Levin and C. E. White, J . Chem. Phys., 18, 417 (1950). (12) I. Levin and C. E. White, ibid., 19, 1079 (1951). (13) I. Levin and C. E. White, ibid., 21, 1654 (1953). (14) J. J. Surash, Ph.D. Dissertation, Lehigh University, Bethlehem, Pa., 1960.
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DONALD D. ROSEBROOK AND WARREN W. BRANDT
this produced negative potentials. Pitts, et a1.,16 have since utilized negative photopotentials to show the occurrence of photoreduction. Tsepalov and Shlyapintokh16 have employed photopotential measurements to study qualitatively the kinetics of the lowtemperature photoreduction of eosin, erythrosin, and rose bengal.
11. Theory The absorption of energy sufficient to cause an electronic transition perturbs the ground-state electronic configuration of the molecule. The electronic distribution then changes to conform to the new total energy. This change usually results in an alteration of the acid-base properties of a molecule. Many methods have been utilized to study the new molecular properties of the excited molecule. Foremost among the methods employed has been fluorescence because of the relative ease of obtaining meaningful data. For this reason nearly all of the pertinent data on excited singlet-state acid-base characteristics has resulted from the interpretation of fluorescent measurements. However, we propose to show that a second technique is now available which provides data pertaining to the acid-base characteristics of excited molecules. The validity of the interpretation of these data will be shown by an analogy to the situation in fluorescence. Consider a simple series of reactions involving a base and its conjugate acid BH+ and their respective excited states.
+ hv1+ BH+ + h ~ 2 B
B* +B
I+
B*
(1)
BH+*
(2)
+ hva
PI
(4)
BH+* +BH+ /-+
-'
(3)
+ huq
P2 B*
(5)
(6)
+ H+
(7)
Because the total concentration of species in the excjted state is very small, it is correct to refer to the point at which the concentration of B* equals the concentration of BH+* sls pK,*. The normal procedure js to measure the intensity of the fluorescence of either species at different pH values and to designate pK,* to be the pH value at which half-intensity occurs. The analogy is this: one of the alternative means of deactivation of the excited state is by a chemical reaction as shown in eq 4 and 6. In the case of reaction 4, the product PI is produced as an alternative to or coinThe Journal of Physical Chemistry
cidentally with fluorescence h v , which is indicative of the species B*. I n the simplest situation the reaction leading to PI is
B -% PI Thus we may postulate the Nernstian expression
The potentials EA and E," found in the Nernst equation depend on the concentration and the identity of the various materials in solution. We shall show later that neither E," or the identity of PI is important in this application of photopotentiometry. So long as the concentrations of B and PI remain constant the value of potential A will remain constant regardless of the pH. When the pH is such that B is consumed to form BH+, three alternatives present themselves: (1) B is excited to B* from whence PI is formed. B, disappearing by reaction, is supplemented from the equilibrium involving BH+. In this case, the ratio of PI/B is the same, but the reaction takes longer due to the competing equilibrium with BH+. (2) P1 is formed in the following sequence of reactions: (2) followed by (7) followed by (4). The result here is the same as above because there will always be a finite concentration of B produced through reaction 3. (3) P, is formed by reaction 6; this necessitates the introduction of a second Nernstian expression
-RT nF
-In-
P 2 1
[BH+]
=
Ec
- Ecn
C
The general considerations for the magnitude of potential C are similar to the arguments for potential A . The over-all effect is to change the absolute value of the potential until the point where consideration of A is no longer significant. Graphically this looks as shown in Figure 1. Under these circumstances the nature of the reactions and of the reaction products is not important as long as B* always forms product PI and BH+* always forms product P2. In general, any time the extent of a photoprocess is being measured at two conditions the nature of the reaction is unimportant, as is the nature of the product, so long as the identity of the product is the same under both sets of conditions. The analogy in this case is to a concentration cell. To apply this technique to the determination of pK,*, (15) J. N. Pitts, Jr., H. W. Johnson, Jr., and T. Kuwana, J. Phy8. Chem., 66, 2456 (1962). (16) V. F. Tsepalov and J. Ya. Shlyapintokh, Izu. Akad. Nauk SSSR,Otd. Khim. Nauk, 637 (1959).
DETERMINATION OF EXCITED STATE pK, VALUESUSINGPHOTOPOTENTIOMETRY
-n e-
3859
I I1
- I n
?,
I
Ha
ARC
ELECTRODE
\ -&-
ELECTRODE
Figure 2. Core of the photopotentiometric apparatus.
b+I
+
Figure 1.
one has only to record the photopotential as a function of pH. So long as one product, PI, is being produced and the ratio P1/B is constant, the photopotential will remain constant. When the excited-state equilibria shift and a different product, P, (or no product a t all), is produced, the photopotential will change. The pH a t the midpoint of the change between the two constant values of the photopotential may be assumed to be numerically equal to the pK,*.
111. Experimental Section Photopotential is measured between one illuminated and one dark electrode in a solution and is defined as the actual electrometer reading of the illuminated electrode at any time us. ground. For experimental purposes the term photodeflection is defined as the photopotential us. ground at photoequilibrium minus the potential vs. ground at dark equilibrium. The apparatus used to obtain photopotentials is shown in Figure 2. The sample was placed in the quartz cell which was wrapped in black electrical tape except for a window of 1 cm2 area. The solutions were deaerated 10 min with nitrogen from which oxygen had been removed by bubbling the gas through two successive solutions of 0.08 M chromous sulfate. The gas was also passed through sulfuric acid and over Drierite to remove water, then over Ascarite to remove carbon dioxide. The pH of the solutions was measured after deaeration. The pH was also checked after irradiation to ensure that the potential change was not due to a pH change. The electrodes were made from platinum wire, and during the deaeration period they were cleaned in another cell by anodizing and cathodizing for 5 min in 1 M HC101. A Nicro Tek HCP-500 power supply provided 3 ma at 3 v for this purpose. I n order to
clean both electrodes simultaneously a third electrode was provided as well as provisions for connecting the indicating electrode and the reference electrode together. The “cleaning” electrode was on a swivel and was removed from the solution and grounded when not in use. After several days of use the electrodes were soaked in 72% HC104 for 1 hr. The electrodes were then rinsed with deionized water and dried with absorbant tissue. This treatment was found to be sufficient to give reproducible dark potentials. The cell with the deaerated sample was placed in the cell holder such that the indicating electrode was against the side of the cell at the window. The reference electrode was in a dark portion of the cell. The electrodes were allowed to equilibrate in the dark until the potential change was less than 0.3 mv/min for at least 10 min. (The dark potential was not always zero because of the difficulty in maintaining two electrodes exactly alike. This necessitated the determination of a dark potential and the introduction of the relative term photodeflection.) After determination of the dark equilibrium potential the shutter was opened and the unfiltered radiation of the mercury arc was allowed to strike the solution around the indicating electrode. The change of the potential of the indicating electrode was measured with a Keithley 610A electrometer and recorded as a function of time. The reference electrode was connected to an earth ground and to the ground terminal of the electrometer. When the deflection dropped below 0.3 mv/min, the shutter was closed and pH was again recorded. pH measurements were made with a Beckman Expanded Scale pH meter. When a normal glass electrode was used, the potassium error was calculated according to the equation of Jordan.“ I n some of the work a Beckman 0-14 glass electrode was utilized. (17) D. 0.Jordan, Trans. Faraday Soc., 34, 1305 (1938).
Volume 70,Number 1.2 December 1966
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DONALD D. ROSEBROOK AND WARREN W. BRANDT
0
photoequilibrium. The position of photoequilibrium was not affected.
IV. Results and Discussion - 10
2-Naphthylamine. Forster2b has measured pKa* for the excited-state reaction BNH3+*
- 20 8
'"
*
8 -30 %
B*
d - 40
- 50
-1
-3
-2
-4
-5
Ha.
Figure 3. 2-Naphthylamine in acid solution.
I n concentrated H2S04the hydrogen ion activity was interpolated from the concentration data of Paul and Long.lS When necessary a correction for ethanol was applied according to the recommendations of Gutsbezahl and G r u n ~ a l d . ' ~ The solutions were not buffered because the use of buffers in many cases introduced spurious potentials that resulted in unintelligible data. The magnitude of the photopotential was found to be directly related to the purity of the solute in question. Therefore all solids were purified by four successive vacuum sublimations. Ethanol was treated according to the instructions of Lappin and Clark20 to remove carbonyl impurities and distilled; the fraction boiling at 77" (uncorrected) was collected and chromatographed through a segmented column of absorption alumina and charcoal. "Blank" experiments were performed and the measured photodeflections varied from $0.5 mv for absolute ethanol to -8 mv for some aqueous acids. These potentials in all cases were negligible, relative to the potentials measured in the corresponding samples. The effect of the intensity of the Hg arc was examined qualitatively by defocusing the radiation. The only effect noted was a decrease in the rate of attaining The Journal of Physical Chemistw
BNHZ*
+ H+
(8)
and obtained a value of approximately -2. This determination was completed in aqueous H2S04 and Hammett acidity values, Ho, were used to denote hydrogen ion activities. Figure 3 shows the same equilibrium determined photopotentiometrically. I n this case, a pKa* of approximately -2.9 was found. When the cationic species are irradiated, only a small fraction of the ions are excited, owing to the low absorptivity of the solution at the wavelengths of the mercury arc. This low absorptivity is reflected in the small photodeflections depicted in Figure 3. Since the photodeflection is small in the region from pH 2 to Ho = -2, where the species BNH2* was produced according to reaction 8, and then it increases to a maximum at Ho = -4, where the equilibrium favors BNH3+*, the indication is that the reduction of the species BNH3+* is more efficient than the reduction BNH2* because the solutions absorb the same amount of energy in both cases. I n alkaline solution another break in the photodeflection vs. pH curve occurs as is shown in Figure 4. This break was assigned to the pK,* for formation of the amide ion in the excited state according to the following reaction BNH2*
BNH-*
+ H+
(9)
The value of this pK,* was found to be 12.3. Forster2b observed the same phenomenon by following the change in the wavelength of fluorescence and assigned the value 12.2 to this pKa*. Forster also states that the parallel ground-state reaction is only observed in a strong alkaline solvent such as liquid ammonia. In alkaline solution the production of the species BNH-* via reaction 9 occurs at the expense of BNH2* thus effectively reducing the observed photodeflection. When the solvent system was changed from water to 42.2 wt yo ethanol in water, the break in the photodeflection us. pH curve appeared at 8.95. The shift of pKa* for 2-naphthylamine from 12.3 in water to 8.95 in aqueous ethanol is in general agreement with the work of Bartok, et aZ.,"J who found that pKa* for phenol (18) M.A. Paul and F. A. Long, Chem. Rev., 5, 71 (1957). (19) B. Gutsbezahl and E. Grunwald, J. Am. Chem. SOC., 75, 505 (1953). (20) G.R. Lappin and L. C. Clark, Anal. Chem., 23, 541 (1951).
DETERMINATION OF EXCITED STATEpK, VALUESUSINGPHOTOPOTENTIOMETRY
- 50
- 60
-70
-80
8 E
.c.
g
-90
B
z
4
z0
a, -100
- 110
--*O -130
t I
11
I
I
I
I
13
12
I 14
PH.
Figure 4. 2-Naphthylamine in alkaline solution.
goes from 5.7 in water to 3.2 in glycerol. This shift in pK,* is attributed to the difference in the solvating power of the various solvents. The effect of solvating power is such that a polar solvent such as water exerts a stronger pull on the nonbonding electrons of tht: amine (or hydroxy group) than does the ethanol or glycerol. This solvation pull reduces the interaction of the nonbonding electrons with the P cloud of the aromatic ring according to the polarity of the solvent. The most polar solvent induces the greatr est part of the electron density due to the n electrons to remain on the atom where those electrons originate. A higher concentration of electron density on an acidic or basic group Ln an aromatic molecule causes that group to be less acidic or more basic. This argument is true in the ground state and generally holds for the excited state except that the dielectric relaxation time of the solvent cage may cause some anomalous shifts in pK,* values. I-Naphthylamine. No pK,* could be determined for 1-naphthylamine in acid solution. The magnitude of the photodeflection remained at - 5 mv from pH 2 to Ho = -7.4. pK,* for this compound for reaction 8 has not been determined by fluorescence either because of fluorescence quenching which occurs at the ground-state pK,.
3861
ForsterZb also found this type of situation with 2naphthol. He postulated that the rearrangement time of 2-naphthol was severely limited by the short lifetime of the excited state of this molecule. (By rearrangement t,ime is meant the time required for protonation or deprotonation in the excited state.) The limited rearrangement time resulted in the appearance of the neutral molecule fluorescence at a pH corresponding to a point just less than the groundstate dissociation constant and far above the excitedstate pK, where fluorescence of the anion should have been observed. The analogy in this case is that the protonated naphthylamine should have exhibited the fluorescence of the neutral molecule and should have undergone the photoreduction of the neutral molecule at pH values below the pK, and above the pK,*. Since it did not do this, the protonated species must not have had time to rearrange during the excited-state lifetime. In alkaline solution a pK,* was again observed for reaction 9. The potential change over the break in the curve of photodeflection vs. pH was 40 mv and quenching of the fluorescence was visually observed. Weller4 concluded that pK,* for 1-naphthylamine in alkaline solution fell somewhere between 13 and 14. This study showed the pK,* = 13.5 i 0.15. The high uncertainty in this study is due to the difficulty in calculating the potassium error at this pH. The Eflect of Ethanol. In aqueous solution the photodeflection from 2-naphthylamine was relatively low (- 120 mv) while in 8.7 wt % ethanol at a similar pH the photodeflection was substantially increased (-290 mv). Increasing the ethanol content to 42.2 wt yo ethanol produced no further increase in photodeflection. The addition of ethanol was observed to accelerate decomposition even in semidarkness. Ethanol-containing solutions could not be stored longer than 24 hr, whereas aqueous solutions were stable for several days. The decomposition of a solution upon standing would be evidenced first by low photodeflections, and within a short time colored decomposition products could be observed visually. For 1-naphthylamine, aqueous solutions gave a -50mv photodeflection; addition of 1 wt yo ethanol increased the photodeflection t o -210 mv, and 4 wt % ethanol caused a further increase to -240 mv. The effect of ethanol on photoreductions in aqueous solutions was extensively studied by Imamura and I
