COMMUNICATIONS TO THE EDITOR
3518 oxygen of the P=O group, which may have little influence on the relaxation time of the phosphorus spin.'la I n a similar manner triethylphosphite (V) is easily decoupled in a Ni(AA)2-CDCl3 solution, while triethylphosphate (IV) is not decoupled. I n order to examine the effect of the chelate concentrations, the nmr spectra of 5,5-dimethyl-Zphenoxy-1,3,2-dioxaphosphorinane(VI) and 5,5-dimethyl2-t-buthoxy-l,b,2-dioxaphosphorinane (VII) l 2have been examined at various concentrations of Ni(AA)2 and CO(AA)~.As is illustrated in Figure 3a, the signal of the methylene protons in the 1,3,2-dioxaphosphorinane ring of VI may be analyzed approximately as an AB part of an ABX system.12 The axial protons (HA) give a pair of doublets (JAP = 2.2; JAB = 10.5 Hz) at 3.40 ppm13and the signal of the equatorial protons (HB)
OR
8 . 1
perturbed trip1et ( J B p = J A B = appears as a 10.5 He) at 6 3.40,13 As the concentration of Ni(AA)2 increases, the lower signal (HA)is decoupled first due to a smaller coupling constant with phosphorus, While the triplet of HB remains coupled (7.5 X M ; Figure 3b), At 1 x 10-3 M , the HBtriplet changes into a very broad doublet (Figure 3c) and a sharp doublet appears
at 2 X M (Figure 3d). A completely decoupled spectrum, obtained at 1 X lov2M , is simply analyzed as an AB system (Figure 3e), which affords 6 4.30 (HA), 3.42 (H)B,and JAB = 10.6 Ha. A very similar result was obtained in VI1 using cobalt(I1) acetylacetonate [CO(AA)~] instead of Ni(AA)2, and a small long-range coupling between phosphorus and protons of the t-butoxy group (JPOCCH = 0.86 Hz) vanished also. This technique may afford a very convenient tool for phosphorus-proton spin decoupling in various phosphites because unfavorable line broadening and conor Ad tact shift are negligibly small for Ni(AA)2 and CO(AA)~. (9) The upfield shift of B methyl protons may be partly due to the diamagnetic shielding effect of the chelate ring which may still be attached to the coordinated nickel. I n that case, however, this assignment is unaffected. (10) F. A. L. Anet and A. J. R. Bourn, J . Amer. Chem. Soc., 87,5250 (1965). (11) W. D. Horrocks, Jr., R. C. Taylor, and G. N. LaMar, ibid., 86, 3031 (1964); L. N. Pignolet and W. D. Horrocks, Jr., ibid., 90, 972 (1968). ( l l a ) NOTEADDEDIN PROOF.After the completion of this article, Frankel's paper attracted our attention: (L. S. Frankel, J . Chem. Phys., 50,943 (1969)). He showed that the phosphorus-proton spin coupling is vanished also in the Co2+ complex of 111. (12) Nmr spectra of the six-membered cyclic phosphites have been discussed, J. H. Hargis and W. G. Bentrude, Tetrahedron Lett., 5365 19689 and references given. (13) A preferred conformation of six-membered cyclic phosphites has been considered to be the following [D. W. White, G. K. IMcEwen, and J. G. Verkade, ibid., 5369 (1968). (cf. ref 12) I.
C O M M U N I C A T I O N S T O THE E D I T O R
The Near-Infrared Spectrum of Liquid Hydrogen Sulfide'
Sir: There are reported herein what are believed to be the first measurements of the near-infrared spectrum of liquid H2S. I n the dilute gas state its vibrational spectrum is very similar to that of H2O and the rotation-vibration bands have been analyzed in some detail.2-4 Though its molecular weight is almost twice that of HzO, its low freezing point (-82.9') and critical point (100.4') are indicative of relatively weak intermolecular forces. A comparison of the effect of temperature upon the parameters of the vibrational band cluster which includes the (111) transition in liquid HzS a t 1.6 p with the same parameters of the much studied (111) band cluster in liquid H20 near 1.2 p should differentiate the role of H bonding in affecting the latter band. The Journal of Physical Chemistry
The spectra in Figure 1 were taken with a standard Cary Model 1431 recording spectrophotometer equipped with a special cell assembly previously described.6 Liquid HzS,'j handled in an all-Inconel system, was dried in contact with silica gel, fractionally distilled several times, and finally distilled directly into the 2.54-cm sample cell. Measurements at each temperature were made vs. CC1, at 25', and the ordinates of each scan were normalized to zero at 6700 cm-l. (1) Research sponsored by the U. S. Atomic Energy Commission under contract with the Union Carbide Corporation. (2) H. C. Allen, Jr., and E. K. Plyler, J . Chem. Phys., 22, 1104 (1954) : J. Res. Nat. Bur. Stand., 52, 205 (1954). (3) G. L. Ordway, P. C. Cross, and E. J. Blair, J . Chem. Phys., 23, 541 (1955). (4) H. C. Allen, Jr., and E. K. Plyler, ibicl., 25, 1132 (1956). (5) W. C. Waggener, A. J. Weinberger, and R. W. Stoughton, J. Phys. Chem,., 71,4320 (1967). (6) Matheson hydrogen sulfide, CP grade; minimum purity by volume, 99.70%.
COMMUNICATIONS TO THE EDITOR
3519 I
i\
0 0
0 3 0 r
c
0 4500
DEG
5000
(cell
0.29 ATM
0.944 G / M L
2f.03
0.776
59.48
0.622
91158
0.423
WAVENUMBER
(1201
I
I -
A
(040)
I
(200)
L I0501 (1301
I1011
(0311
loo?)
I
1
I 7500
7000 (cm-')
1060)
12101 11311
I
1'431 IG41)
10121
8500
8000
12201 1 2 0 l \ ~ 3 C O I I121 I 1102, 10221 1303!
Figure 1. Absorption spectra of liquid hydrogen sulfide from near the freezing point to near the critical point. The vertical lines below the curves indicate the positions and relative intensities of vibrational transitions in the region for the dilute gas state. Quantum numbers, VI, VZ, vat for the upper state are in parentheses; the lower state in each case is (000).
Table I : Parameters of Corresponding Absorption Bands for Orthobaric Liquid H28 (75') and HzO (250')
Position, cmL1 Band width, cm-1 Maximum intensity (1. mol-' cm-l) X lo3 Integegrated intensity (1. mol-' cm-2 a
Change (cm-1) from the freezing point.
-Band
1-
HsS
Hz0
5200 7200 291 300 39.0 425 12.7
164
7 -
Band
--Band
2--
HzS
H10
6250 ($44)" 8673 ($425)" 130 (+92)* 450 (-44)b 18.2 (-37)* 1 9 . 2 (+137)* 3 . 2 ( ~ 0 ) 9~. 0 (+25)*
HzS
7609
3Ha0
10,400 440 1.3 8.4 . . . 4.2
...
--Band
4-7
HIS
HzO
9650
12,000
...
1.0
... ..
0.7
Per cent change from the freezing point.
Figure 1 shows the spectrum of orthobaric liquid HzSa t 74.8" in the region between 1.1 and 2.2 p. The four slightly structured and well separated bands arise from the corresponding clusters of the possible vibrational transitions which are diagrammed below the curves. The harmonic and combination frequencies for the dilute gas state were calculated from the vibrational constants for HzS,4and the vertical heights were drawn to represent a decrease in intensity with the number of terms in the combination. For HzO, these clusters of transitions occur between 0.8 and 1.6 p, and the locations of the corresponding bands in the spectrum of orthobaric liquid HzO at 250" are given in Table I along with measured band parameters. The HzO bands are characterized by their extreme band widths and temperature sensitivities. The widths
of these bands decrease in going from 0 to 250" (viz. band 2 in Table I) reflecting the strong intermolecular H bonding and the concomitant bond breaking which occurs with increasing temperature. The HzS spectrum differs from that of HzO in that the bands are narrower, are in general less intense, and broaden with increasing temperature from the freezing point. This broadening, which must arise from kinetic processes connected with weak intermolecular association, is similar to that which we have observed in pure Lorentzian bands of liquid C02,5except that the band broadening and shifting for the same range of reduced temperature are in liquid COZ 8 times smaller and indicative of still weaker intermolecular forces. Interestingly, a least-squares fit of our 75" data for the 6250-cm-l HzS band required five highly non-Lorentzian components Volume 73, Number 10 October :860
3520
COMMUNICATIONS TO THE EDITOR
just as in the case of the 8675-cm-’ HzO band at 25OO.7 The component HzS bands ranged from 0.78 to 1.00 in fraction Gauss profile with half-widths averaging nearly 150 cm-I. A complete discussion of our results pertaining to HzO and H2Swill be published soon.
a 6
w 0
(7) W. C. ’Waggener, A. J. Weinberger, and R. W. Stoughton, Abstract of Paper 75, Division of Physical Chemistry, 156th National Meeting of the American Chemical Society, Atlantic City, N. J., Sept 9-13,1968,
Q
W. C. WAGGENER
CHEMISTRY DIVISION RIDGE NATIONAL L.4BOR.4TORY OAK RIDGE, TENNESSEE 37830
0
A. J. WEINBERGER
OAK
R. W. STOUGHTON
RECEIVED JULY 17, 1969
6
Evidence of Chain Association of Benzoic Acid in Benzene?
Sir: Harris and Dunlopl published the results of careful isopiestic measurements of the osmotic coefficient of the solutions of benzoic acid in benzene at 25’ for the molality range 0.36 to 0.84. The variation of the osmotic coefficient cannot be interpreted in terms of dimerization of benzoic acid and the authors were not able to calculate the association constant Klz = x2/x12 (xz and x1 are the mole fractions of dimer and monomer), The data by Harris and Dunlop can be interpreted in n/a, where terms of the mean association number2 x n = xin, is the nominal number of moles of the solute (mass divided by the molecular weight of the monomer), n, being the number of moles of i-mer and ii = x n i . x and the practical osmotic coefficient, 4, are interrelated as follows: cp is defined by the equation ps = p 2
- cpRT(n/nd
(1)
where p, is the chemical potential of the solvent and n, is the number of moles of the solvent; on the other hand, in an ideal associated solution ps = p 2
- RT
In (1
+ ri/ns)
(2)
0,4
- l)/r 5 $
(3)
muo, where m is molality and wo is one thousandth of the solvent molecular weight (notation of Harris and Dunlop). For dilute solutions, (2) and (3) simplify to ps = p2
- RT(ri/ns)
(2%)
0,8
1
2 100
:
3
Figure 1. a, Mean association number, x, of benzoic acid in benzene against molality; b, y (2 - x)[%/(x l ) ]l i Z as a function of f : 0, x = I/+; 0, x = l/$.
-
=
Table I x
A
B
I/$ l/+
0.01390
0.5072 0.7433
0,01541
KI?;
5 2 0 0 2 ~600 4200 =t420
K
372C4 48& 4
+
2/(1 nl/n). Therefore, if deviations from ideality are attributed exclusively to the association of benzoic acid, the presence of complexes formed by three or more monomolecules of benzoic acid must be admitted. The dimerization constant, Klz, and multimerization constant, K = xt/xt-lxl (for i > 2 ) , can be evaluated easily by the linear equation3
(2 - x)[E/(x - l)]’”= A
From (1) and (2) it follows
l/x = (exp cpr the mole ratio r = n/n, =
0,B
molality
+
- BZ
+
(4)
whereZ=ii/(ii n,) = 1/(1 x/r); 1/A2 = Klz and B/A = K . I n a previous paper,4 I proved the equation
(2
- C / E ) ( C / E - 1)
= ( E K ~ ~ * ) - ”* ( eK*) ~
(4a)
where c, E, KIZ*, and K* are expressed in molalities. As C / E = x, c;Klz* = 5Klz and EK* = ZK, eq 4 is equivalent to eq 4a. For K = 0 eq 4 reduces to the
and 1/x = cp
(34
(1) K. R. Harris and P. J. Dunlop, J . Phys. Chem., 71, 483 (1967).
I n Figure la, x is plotted against molality. x is greater than 2 for more concentrated solutions, but cannot exceed 2 for monomer-dimer equilibria as x =
(2) 13. Buchowski, Rocz. Chem., 42, 165 (1968). (3) H. Buchowski, ibid., 42, 167 (1968). (4) H. Buchowski and R. Lewandowski, J. Chim. Phys., 64, 1345 (1967); see eq 11.
The Journal of Physical Chemistry
