NOTES
4416
Fz
results from an alteration in the paramagnetic term of the Ramsey shielding equation. A linear relation between the ortho-proton shifts for the aromatic halides was found with an empirical parameter, &. This parameter was defined as P/Irs, where P is the bond polarizability, I is the first ionization potential, and r is the bond length. A plot of the chemical shifts for the various ortho protons in Table I against Q is shown in Figure 4. Benzene is also shown in the plot. As can be seen, the fit to the plot is excellent, and it is presumed that this mode of deshielding of the ortho protons dominates in these compounds. Reasonably, the additivity rule of Martin and Dailey would not be expected to hold in this situation.
12
+I
0
-1
-a
Figure 3. The energy level diagram for l-chloro2-bromobenzene showing the line assignments and frequencies (leveb not to scale; add 400 C.P.S. to each frequency).
2.8
i
2.6
-
Experimental Section The compounds used in this study were all commercially available. Spectra were determined with a Varian A-60 at the ambient temperature of the probe (ca. 40"). Solutions of each compound were made up 5 1 0 % by volume in carbon tetrachloride and outgassed. One drop of tetramethylsilane was added as an internal standard. The instrument was calibrated against chloroform. Each spectrum was recorded several times (4-lo), and the averaged line positions were used in the calculations.
Acknowledgment. We wish to express our gratitude to the Robert A. Welch Foundation for its generous support of this work.
.
w% .
Ff: o
b
-
2.4
.
.Y
a 2
'
(9) P. Hruska, H. M. Hutton, and T. Schaefer, Can. J . Chem., 43, 2392 (1966).
Dissociation Constants for Some Nitrophenols 2.2
and Salicylic Acid in Deuterium Oxide
.
by Paul K. Glasoe L
I 3
2
4
8.
wittenberg University, Springfield, Ohio (Received August 11, 1966)
Figure 4. Plot of the chemical shift data of the o-hydrogens us. the Q values of Hruska, Hutton, and Schaefer.
range van der Waals type interaction between the halogen and the adjacent proton. No quantitative calculations were attempted, and this explanation has been subsequently discarded in favor of a more reasonable explanation. Hruska, Hutton, and Schaefer9 have presented evidence that the low-field shift The J O U T ~of U ~Phvsical Chemistry
I n a report on the dissociation constants of some acids in deuterium oxide, Bell1 noted the poor agreement between his result for 2,4-dinitrophenol and that reported earlier by McDougall and Long.2 He also -~ ~
(1) R. P. Bell and A. T. Kuhn, Trans. Faradall Spc., 59, 1789 (1963). (2) A. C. McDougall and E". A. Long, J . Phys. Chem., 66, 429 (1962).
4417
NOTES
Table I: p K Values for Some Nitrophenols and for Salicylic Acid in Water and in Deuterium Oxide Compound
o-Nitrophenol p-Nitrophenol 2,PDinitrophenol Salicylic acid
Long and McDougall PKH APK
7.19 7.26 4.12 2.94
0.75 0.48 0.70 0.75
Bell and Kuhn PKH AP K
...
...
... ...
4.07
0.52
...
...
Martin and Butler PKH APK
PKH
APK
7.25 7.24 4.02 3.00
7.22 7.14" 4.12 3.01
0.60 0.58 0.56 0.56
a R. Robinson and A. Peiperl give the value 7.156: J. Phus. Chern., 67, 1723 (1963). 624 (1936).
pointed out similar large discrepancies in the values reported by McDougall and Long and those of Martin and Butler3 for o-nitrophenol and p-nitrophenol. A summary of these values i;3 given in Table I. Noting that the results of McDougall and Long were obtained using the glass electrode to measure pH or pD whereas his results and those of Martin and Butler were made using spectrophotometry, Bell suggests the possibility that the glass electrode is not wholly reliable in deuterium oxide solutions. Since we have published several results for dissociation constants of acids in deuterium oxide using the glass electrode, we undertook the measurement of pKH and PKDfor the nitrophenols in question using the procedure described in an earlier notes4 The solutions were in general very dilute, usually less than 0.01 M , owing to the low solubility of the nitrophenols in water. The ionic strength was kept around 0.05 by addition of KCl. It was found that the glass electrode was much more stable in such a solution than it was in the nitrophenol alone. I n all cases, the pK values are corrected for the ionic strength using the Debye limiting law. The values we have obtained are shown in Table I. I n each case, our value checks reasonably well with Bell or with Martin and Elutler. I n view of these results, we conclude that in these cases the glass electrode gives proper results in deuterium oxide solutions. Long and McDougal12 report a value for ~ K .-D ~ K (or E ApK) for salicylic acid of 0.75. This is much higher than that given by La Mer and Korman,6 0.61. We have measured the PKD and the ~ K Efor I salicylic acid and obtain 0.56. We conclude that the deuterium isotope effect in salicylic acid is only slightly greater than that of "normal" carboxylic acids.'s6 (3) D. C. Martin and J. A. V. Butler, J . Chem. SOC.,1366 (1939). (4) P. K.Glasoe and L. Eberson, J. Phys. Chem., 68, 1560 (1964). (5) See footnote b of Table I. (6) C. K. Rule and V. K. La Mer, J. Am. Chem. Soo., 60, 1974 (1938).
0.57 0.56 0.52 O.6lb
This work
V. K. La Mer and S. Korman, Science, 83,
The Proton Magnetic Resonance Spectrum of Phenanthrene
by Robert C. Fahey and Gary C. Graham Department of Chemistry, University of California at S a n Diego, L a Jolla, California (Received August 18, 1966)
The proton n.m.r. spectrum of phenanthrene (Figure 1) contains two complex multiplet patterns. Because inter-ring coupling constants are usually negligibly small in aromatic hydrocarbons, the 9,lO-protons should give rise to a single line. The intense line a t
the center of the high-field multiplet is reasonably assigned to these pr0tons.l The remaining lines in the spectrum comprise an ABCD pattern, the lowfield multiplet of which has been assigned to the 4,5protons based on studies of phenanthrene-9-dl and 4methylphenanthrene.2 An approximate analysis of the spectrum has been reported previou~ly.~We report here a complete analysis of the spectrum and compare our results with the earlier findings. Spectra were measured at several concentrations (8.9 to 17.4%, w./v.) in CDC1, on a Varian HR-60 spectrometer. Peak positions were determined relative to tetramethylsilane as internal standard using the side-band technique. The spectra were analyzed using (1) N. Jonathan, S. Gordon, and B. P. Dailey, J . Chem. Phys., 36, 2443 (1962). (2) H. J. Bernstein, W. G. Schneider, and J. A. Pople, Proc. Roy. SOC. (London), A236, 515 (1956). (3) T. J. Batterham, L. Tsai, and H. Ziffer, Australian J . Chem., 17, 163 (1964).
Volume 69, Number 12 December 1966
