S46
PETER J. BERKELEY, JR.,AND MELVINTV. HANXA
electronic considerations must also be responsible to some extent for the observed differences between nickel and platinum for this reaction. Both of these metals belong to group VI113 and therefore possess partially filled d-bands, but it is not possible a t this stage to
Vol. 67
assess the relative importance of morphological and electronic factors in determining the catalytic acti7-ity. Acknowledgments.-The author wishes to thank Dr. E. J. Cairns and W. T. Grubb for many helpful discussions during the course of this investigation.
S.N.R. STUDIES OF HYDROGEN BOXDING. I. BISARY MIXTURES OF CHLOROFORRI ASD SITROGEN BASES BY PETER J. BERKELEY, JR.,ATD RIELVIN W. HANNA Department of Chemistry, University of Colorado, Boulder, Colorado Received October 6 , 1968
A discussion is given of the theory of the dilution-shift technique for examining 1: 1 hydrogen-bonded complexes by n.m.r. It is shown that a t infinite dilution of the proton donor in an acceptor solvent, the limiting slope and the intercept (the limiting shift) of a plot of donor concentration vs. shift of the donated proton is sufficient to determine uniquely the equilibrium constant for hydrogen-bond formation and the shift of the pure hydrogenbonded species. A method for taking solvent effects into account is included. Data for chloroform plus pyridine, ethylideneisopropylimine, acetonitrile, and E-methylpyrrolidine are reported. These data are analyzed by the above theory and the importance of solvent effects is evaluated. A qualitative discussion is given of the calculated equilibrium constants and shifts upon complex formation.
Introduction
It is now well known that the position of a hydrogenbonded proton in an n.m.r. spectrum is shifted downfield with respect to the position of a non-liydrogenbonded proton.2 This shift has become known as the hydrogen-bond shift. In recent years an abundance of experimental information has appeared on many different systems in an attempt to correlate the hydrogenbond shift with hydrogen-bond strength, infrared stretching frequencies, and other physical properties of the hydrogen-bonds3 There has been considerable variation in the means of interpreting this data, however, with the result that it is difficult to get a clear picture of the relation between n.m.r. hydrogen-bond shifts and properties of the hydrogen-bond itself. This recent growth of the literature has also introduced certain nomenclature ambiguities. The situation is aggravated by the fact that tm-o distinct types of dilution-shift experiments may be carried out. In the case of hydroxylic materials, one compound can act as both proton donor and acceptor, and here it is usual to observe the shift of the hydroxyl proton as such a compound is diluted with inert solvent. This paper concerns itself, however, with the situation where a proton donor is diluted with acceptor and the shift of the donated proton observed. It is common to utilize, in both cases, the shift at infinite dilution of the acid-base in the first case, or of the acid in the second case. These two kinds of infinite dilution shifts, however, have quite a different significance. Unless otherwise stated, all that follows will be concerned with the second case. A further nomenclature difficulty involves the meaning of the phrase “hydrogen-bond shift.” I n both kinds of experiment mentioned above, a distinction should be made between the observed dilution shift and a (1) Supported in part by Grant GM 091&7-01S1 from the National Institutes of Health, Public Health Service. (2) See for instance: J. A. Pople, W. G. Schneider, and H. J. Bernsteina “High Resolution Nuclear Magnetic Resonance,” Chapter 15, hIcGraw Hill Book Co., Ken York. N. Y., 1959. (3) G . C. Pimentel and A. L. McClellan, “The Hydrogen Bond,” W.H. Freeman and Co., San Francisco, Calif., 1960. Chapter 4.
-
quantity which can be calculated from this observed shift, the hydrogen-bond shift of pure dimer. I n this paper the former quantity will be distinguished from the latter by being called the observed H-bond shift. It is the purpose of this paper to present a simplified approach that is applicable to the study of 1: 1 hydrogen-bonded complexes by n.m.r. and to point out the significance of ‘[observedH-bond shifts” as a criterion of hydrogen-bond strength. Specifically, it will be shown that for the case of a 1: 1 complex the data relating the chemical shift of the donated proton to apparent concentration of this proton are sufficient to determine uniquely the association constant of the complex without resorting to iterative procedures, to curve fifitting, or to information from other sources (such as infrared). This theory is then applied to data for the association of chloroform with pyridine, acetonitrile, et hylideneisopropylimine, and N-methylpyrroljdine. Theory Consider the equilibrium
A
+D
AD
where A is the proton acceptor and D, the donor. If the usual4 assumption is made that the system is an ideal mixture of species (after the above equilibrium has been taken into account), concentrations can be substituted for activities and the equilibrium expression takes the form
(1)
where the X’s are the mole fractions, the No’s are the initial number of moles, and X is the moles of complex a t equilibrium. Applying equation 1 to cases where only an average spectral line is observed for the hydrogeii(4) J. E. Hildehrand and R. L. Scott, “The Solubility of Non-electrolytes,” 3rd Edition, Chapter XI, Reinhold Publ. Corp., New York, N. P.,
1950.
N.M.R.STUDIES OF HYDROGEN BOXDING
April, 1963
bonded and non-hydrogen-bonded protons, Huggins, et al.,5 have shown that
K
60
= - AD
l + K
+
- 6 ~ ) 6~
(2)
where 60, BAD and 6~ are the chemical shifts of the donated proton in infinitely dilute donor, of the pure complex, and of the pure donor, respectively. For compactness in what follows, the symbol A,D will be introduced to represent the hydrogen bond shift of pure donor, i.e., AAD = AD - 6 ~ .It is necessary to realize here that the above applies oiily to binary systems with the acceptor being the second component. When one component KS both acceptor and donor and the other is inert solvent, the shift a t infinite dilution equals that of the moi~omer.~ An important point to be re-emphasized6 in these equations is that the observed hydrogen-bond shift (6, SD) depends o n both the equilzbrium constant, K , and the H-bond shift of ths pure dimer, AAD. Thus, when a correlation between n.m.r. data and hydrogen-bond properties is attempted, it is extremely important to state precisely what is being correlated. As shall be seen below, the values of K and AAD sometimes go in opposite directions for a series of acceptors with the same donor, and a decision must be made as to which quantity is to be used as a criterion of hydrogen-bond strength. In some cases, however, the observed hydrogen-bond shift is a good criterion of the H-bond shift of the pure dimer. Thus if I< is very large AAD = (60 SD), but this condition does not hold in many instances. Additional information can still be obtained from the data, however. The use of limiting slope in n.m.r. studies of hydrogen-bonding was introduced in connection with self-ass’ociationof phenols.’ Applying standard mathematical procedures to a plot of observed chemical shift us. apparent mole fraction of acceptor, the slope a t high X A O is given by
Therefore the 1im.itingslope, 80, is =
K (IfK12
(4)
Now if AAD is eliminated between (2) and (4)
847
where 6o - 6~ is small (say less than 1.0 p.p.m.) because it is known that the chemical shift is solvent That is to say, part of the observed Hbond shift, (ao - a ~ ) ,will be dependent upon the solvent used, unless the shifts are referred to the gas phase. Bothner-By9 has recently provided an empirical means of taking this into account. He has found that the unexpected downfield shift of a proton in a material “i” dissolved in a solvent “j” can be expressed as a product of two numbers: x i , characteristic of the compound containing the proton and g,, characteristic of the solvent. Thus he writes for a cylindrical sample tube
(7) where 6i0 is the shift of the proton in the gas phase and xj is the volume diamagnetic susceptibility of the solvent. Experimentally, 6 has to be observed with respect to some reference compound, either external or internal. If the absolute shift of an internal reference is 6,, the ) then be expressed as measured shift ( 6 ~ can 631 =
6
- 6,
If equations 5 and 6 are now rederived with the inclusion of solvent effects, there results
and
where (6& is the measured shift a t infinite dilution, Y A is the “Vj” of the acceptor (base), and A, is the difference in x value between donor and reference. The term Ao, which is the gas phase difference in shifts between pure donor and reference, can be obtained by an independent measurement in the gas phase or it can be obtained from liquid phase measurements using equation 7 if the appropriate x and y values and susceptibilities are known. The use of equation 8 eliminates the influence of differential solvent effects on donor and reference. It automatically takes into account the small self-association of chloroform,1°for example.
Experimental
Thus 8 0 (t.he limiting slope) and 6o (the shift at infinite dilution) of a plot of observed shift vs. proton donor coiicentration are suficieiit information to yield AAD and K uniquely. Equation 6 is n.ot satisfact’ory, however, in the cases
The reagent grade chloroform used was extracted with water to remove alcohol, dried over anhydrous potassium carbonate and distilled immediately before use. Reagent grade cyclohexane was used without further purification. Reagent grade pyridine r a s dried over anhydrous K&Oj and distilled. N-Methplpyrrolidine, obtained from the Aldrich Chemical Companp, was placed in contact with metallic sodium overnight and then distilled from over metallic sodium. Reagent grade acetonitrile wits extracted with a saturated solution of potassium hydrouide, dried over anhydrous KPCCl, and then distilled over K?COd. The ethylideneisopropplimine was prepared in the manner ontlined hv Campbell, et al.”
( 5 ) C. M. Huggins. G. C . Pimentel, and J. N. Shoolery, J. Chem. PhUs., 23,1244 (1955). (6) This point was made sometime ago by Hungins, et al., ref. 8. (7) C. M. Huggins, GI. C . Pimentel, and J . N. Shoolery, J. Phys. Chem., 60, 1311 (1966); see also E. 11. Becker, U. Liddel, and J. N. Shoolery, J. Mol. Spectry.. 2 , 1 (1958).
(8) Reference 1, Chapter 16. (9) A. A. Bothner-By, J. Mo2. Spectry, 6 , 52 (1960). (10) C. F.Jumper, hl. T . Emerson, and B. B. Howard, J. Chem. Phus.. 35, 1911 (1961). (11) K. N. Campbell, A. H. Sommers. and B. K. Campbell, J. A m . Chem. Soc., 66,82 (1944).
and
PETER J. BERKELEY, JR.,A N D
84fi I
I
I C - -
370
.-.
-
dol
3600
-
-
Ob2
,b3
0!05
0.b4
Apparent Mole Frociion CHCI,,
- XA WAX, ( 6 M h - A' f 0.p.s yaAx. C.P.S. Proton acceptor x 106 II Pyridine 0 600" - ,078 -4 3 -95.6 -0 0 -840 N-Methylpyrrolidine 710b 0 037 2 0 -86 2 509" CHaCH=NCH(CHa)* 0 059 3 2 -28 7 514d Acetonitrile a Reference 2 after correction to 37.5'. Estimated from that of pyrrolidine found in A. Pacault, Ann. Chzm. [12], 1,527 (1946). Measured using the A-60 spectrometer, see equation (4-2), H. Francois and J . Hooran, Compt. page 79, in reference 2. rend., 240, 1220 (1955). TABLE
solveiit effect parameters needed for these corrections. (Xote that Ao was obtained from a gas phase measurement of W. G. S~hneider.'~)Significant contributions from solvent corrections (>lo%) are not obtained until ( 6 ~ ) falls o below 1.0 p.p.in., which is the case onlywith acetonitrile. It is of interest that the values of the solvent eff ect-corrected shifts of pyridine and acetonitrile are in good agreement with those found by AIartinI5 for what she calls A, (shift of pure liquid chloroform minus shift of infinitely dilute chloroform). I n Table I11 are listed the values of K and AAD calcuTABLE I11 CALCULATED PROPERTIES O F CHLOROFORM-PROTON ACCEPTOR
SOLUTIONS AAU
Proton acceptor
Pyridine N-Methylpyrrolidme CH3CH=XCH(CH3)2 Acetonitrile
K
0 2 5 3
69 2 2 2
C.P.8.
-234 - 123 -119 - 37.6
P P m.
-3.90 -2 05 -1.98 -0 63
lated using equations 8 and 9, neglecting the derivative in the denominator of equation 9. It would be tempting to spend some discussion on the relative order of magnitudes of the derived K's. However, it is felt at this time that such discussion would of necessity be highly speculative. Experiments are currently under way to determine the full significance of these K's. AHo values for such weak interactions will be sufficiently small that the ASo values will be of major (13) The z's for cyclohexane and chloroform were found to be 3 1 and 7.0, respectively ?j of cyclohexane IS glven by Bothner-Byg as 0 033. All other y's used are gii en In Table 11. (14) Private communication' -337 2 0.p.s (at 60 blc.) after coi~oction from a gaseous neopentnne refeience t o a gaseous cyclohexane ieference. (15) Lf.Martin, Ann. P k y s , 7, 35 (1962).
April, 1963
REACTIOSOF BYDROGEK ATOMSWITH
significance in determining K . However, these values are of the correct magnitude as can be seen by comparison with data of Huggin@for triethylamine-chloroform and with that of Jumperl8 for acetonitrile- and pyridine-chlorofori~i. Classical physical chemistry techniques also give similar results. For example, for the chloroform-dioxane system,17 K , = 1.1. It should not be expected, however, that these K’s will follow the same order as found for the KB’s of the same hases because of the lesser importance of entropy effects. The calculated AAD’S are of interest, however, bepause they do appear to be related to the basicity of the acceptor. At the same time, they demonstrate the care which must be exercised when n.m.r. data are compared with other properties of hydrogen-bonded systems. On the basis that the base strength ( K B )is what determines the strength of the hydrogen-bond shift and hopefully the value of AAD, one would expect the A l ~ ’ s to decrease in the order sp3 > sp2 > sp.’* I n fact, with the exception of pyridine this is the order that is observed even though the sp2 and sp3 uiishared pairs give AAD’S which are very close to each other. This correlation, however, must take magnetic anisotropy effects into account. Thus, there is an additional downfield shift in the pyridine-chloroform complex because of the aromatic ring in the base. The Ahn in pyridine is, therefore, not a measure of a property of the hydrogen-bond until the effect of the aromatic ring currents has been subtractpd out. A similar argument can be applied to acetonitrile except that in this case paramagnetic contributions to the nuclear screening of the hydrogen-bonded proton produce an upfield shift. The “true” hydrogen-bond A ~ is D therefore larger than the AAD in Table 111. d schematic representation of these effects is shown in Fig. 2. It is also highly possible that a larger difference between sp3 and sp2 unshared pair shifts would be found if the paramagnetic contribution to the shift due to the carbon-nitrogen (16) C F Jumper, “A Study of Hydrogen Bonding by Dielectric hIethods and b y Nuclear Magnetic Resonance ” Thesis, Florida State IJniversity, L C Card No Rfic 61-1285 (19611 (17) M . L McGlashnn and R P Rastogi Trans Fa,aday S o c . 64, 498 ( 1 950). (18) C R Nollei, “Chemistry of Organic Compounds,” W. B Saundeis Co , Philadelphia, Pa 1951 (19) J. A. Pople, Pro6 Roy Soc (London), A239, 541, 560 (1957).
LIQUID
OZONE
849
Fig. 2.-Schematic representation of ring currents in pyridine and acetonitrile. The small arrows in the center of the ring represent the direction of the field induced by H,, the applied field. The dashed lines represent the lines of force of the induced field; and their direction at the nitrogen lone-pair orbitals indicate whether the induced field adds to or subtracts from H Ofrom the point of view of the chloroform proton.
double bond were subtracted. Since theoretical estimations of these effects are about as uncertain as the magnitudes of AAD’S themselves, it is difficult to take these contributions into account in any satisfying quantitative way. The basic point is that it would be uiisound to try to correlate these n.ni.r. shifts in any more than a general way with the properties of the hydrogen bond formed in these systems, but it does seem that the general order of expected H-bond strength sp3> sp2> sp is confirmed for the non-aromatic compounds since it is likely that the correction to the acetonitrile would be of the order of 0.5 p.p.ni.20 It appears on the basis of the above work, therefore, that in cases where the magnetic anisotropy does not vary greatly, AAD might be a useful criterion of H-bond and could be correlated with other properties such as infrared hydrogen-bond shifts. Acknowledgments.-The authors wish to thank Professor E. King for many helpful discussions, and Dr. W. G. Schneider for communicating his gas phase chloroform shift. (20) I n reference 2, page 179, a calculated value of 10 p.p m. is given for the upfield shift of the acetylene proton. Because this shift varies a s the reciprocal of the bond length t o the third power, the 10 p.p.m. must be mul0.05. tiplied by about 13/(--2.7)3
THE REACTION OF HYDROGEN ATOMS JTITH LIQUID OZOKE~ BY J. A. WOJTOWICZ, F. ~IARTIKEZ, AKD J. A. ZASLOTTSKY~ Olin Mathieson Chemical Corporation, Organics Division, New Haven, Connecticut Received October 6, 1968 The reaction of atomic hydrogen with liquid ozone a t -196“ gives a product which upon warming evolves molecular oxygen and leaves a residue of aqueous hydrogen peroxide. It is shown that under certain reproducible conditions the ratio of evolved oxygen to residual hydrogen peroxide is unity. The absence of hydrogen peroxide as a primary reaction product (under these conditions) was confirmed by infrared examination.
Introduction Many invest,igators haye reported the format,ion of hydrogen peroxide in frozen col-&llsates obtained by the reactioli of hydrogel1 and oxygen contai11illg radi-
c a k 3 Despite the general accord regarding the forniation of hydrogen peroxide, t.here is disagreement as t,o the explanation of the evolution of oxygen which occurs when the condensates are warmed from -196’ to
(1) Presented in part a t the Fifth International Symposium Radicals, Uppsala, Sweden. July, 1981. 12) Address inquiries to J. A. Zaslowsky.
(3) (a) -\. M. Bass and N. P. Broida, “Formation and Trapping of Free Radicals,” Academic Press, N e n P o r k , N. Y., 1960: (b) G. J. iifinkoff, “Frozen Free Radicals,” Int.ersoience, New York, N. Y., 1960.
011
Free
