PULSE RADIOLYSIS OF METHYLIODIDE AND METHYL BROMIDE
1919
Pulse Radiolysis of Aqueous Solutions of Methyl Iodide and Methyl Bromide. The Reactims of Iodine Atoms and Methyl Radicals in Water‘
by J. K. Thomas Chembtry Division, Argonne National Labmatmy, Argonne, IUinok
(Received December 19, 1966)
The pulse radiolysis of aqueous solutions of methyl iodide and methyl bromide has been studied on the ANL linear accelerator. In the case of methyl iodide, methyl radicals and iodide atoms are produced: eaqCHaI + CHa I- and OH C H a I T CHaOH I(CHJ). With methyl bromide, methyl radicals are produced but no Br atoms could be detected. The absolute rates of several methyl radical reactions in water have been measured and these are discussed in terms of the corresponding H and OH radical reactions.
+
Introduction I n the past 5 years, the technique of pulsed radiolysis has provided extensive data on the nature of radiolytic processes in liquids and has elucidated many kinetic patterns of the radiolysis fragments. I n water reliable rate data are available for reactions of H atoms, OH radicals, and hydrated electrons, eaq-.28 Previous work2b has shown that hydrated electrons react with rates that approach diffusion control with methyl iodide and with other alkyl halides. These reactions probably lead t o bond breakage,8 e.g., eaq- R(ha1ide) +R halide. Thus, these systems may be used as a convenient source of alkyl radicals in liquids where solvated electrons can be generated. The first part of the work uses the pulse radiolysis technique to study in detail the processes occurring in the radiolysis of aqueous solutions of methyl iodide. The second part describes the measurement of the rates of reaction of methyl radicals in water.
+
+
Experimental Section Preparation of Solutions. Reagent grade methyl iodide from the Baker Chemical Co. is washed by shaking with triply distilled water. The mixture is allowed to settle and the supernatant water is poured away; the process is repeated three times. This treatment gives a water-saturated methyl iodide which is free of hydrogen iodide. A small quantity (-2 ml) of this purified methyl iodide is placed with 50 ml of triply distilled water in a 100-ml syringe, with no gas space. Pure argon gas (50 ml) is introduced into the
+
+
+
syringe and the mixture is shaken vigorously for 3 min. The gas phase of air, methyl iodide, and argon is expelled and the whole procedure repeated four times. This gives a saturated solution of methyl iodide (0.1 M ) in water, with less than 2 X 10-7 M oxygen. Dilute solutions of methyl iodide are then prepared by the syringe dilution te~hnique.~Solutions of methyl bromide are prepared by bubbling CH,Br through water for 15 min. This produces a saturated solution of CHaBr in water which is free of oxygen. Dilute solutions are prepared by mixing the saturated solution with degassed water via the syringe technique. Analysis. Iodine is determined by mixing the irradiated sample with an equal volume of 0.2 M potassium iodide and measuring the 11- at 350 mp on a Gary spectrophotometer, the extinction coefficient of 13at 350 mp being 26,500. Gases such as methane and oxygen are determined by stirring the irradiated solution in a Van Slyke apparatus and injecting the liberated gases into a gas chromatograph.6 ~~
(1) Based on work performed under the auspices of the U.S. Atomic
Energy Commission. (2) (a) M. Anbar and P. Neta, Intern. J . Appl. Radicrtion Isotopes, 16, 227 (1965); (b) A. Szutka, J. K. Thomas, S. Gordon, and E. J. Hart, J. Phys. Chem., 69,289 (1986). (3) D. N. Skelly, R. G. Hayes, and W. H. Hamill, J . Chem. Phys., 43, 2795 (1965). (4) E. J. Hart, S. Gordon, and J. K. Thomas, J. Phys. Chem., 68, 1271 (1964). (5) J. K. Thomas and E. J. Hart, Radktion Res., 17, 408 (1962).
Volume 71, Number 6 May 1967
J. K. THOMAS
1920
Pulse Radiolysis Apparatus. This is described in detail elsewhere6 and here it suffices to outline the procedure. The solutions are irradiated in a 4-cm quartz cell by 0.4- or 1.0-psec pulses of 15-Mev electrons from an Arc0 linear accelerator. An analyzing light beam passes through the cell twice, giving a path length of 8 cm, and then after passing through a Bausch and Lomb monochromator, it is monitored with a 1P28 photomultiplier tube. The output of the tube is amplified and displayed on a Tektronix 545 A oscilloscope. The rise or response time of the whole apparatus is 80 nsec. Dosimetry is carried out by direct observation of the ea,- in degassed water at 400 mp where c is 27207 or by direct observation of the 12-spectrum in loR3M potassium iodide where c is 14,000 at 385 mp.* The primary yields used in conjunction with the above dosimetry are G(eil,-) = 2.40, G(0H) = 2.40, and G(H) = 0.60.
I5
14
t
Amp
-- -
1.3 1.2
OH
+
I.!
1
0.3 0.2
+
removal of H OH to give unreactive prodCi" +C2H,OH ucts
The kinetics of decay of eaq-, measured at 600 mp, and formation of I-, measured at 230 mp, in a solution of 25 p M CH31-10-3 M C2H4 are shown in Figure 2. The kinetics are apparently first order with k(,,, + C H ~ I ) = 1.65 f 0.25 X 1O'O sec-I. The plot of log [e,,-] vs. log ([I-IO - [I-]*) is linear with a slope of unity showing that ea,- immediately produces I-. Formation of Iodine Atoms. Figure 3 shows the spectrum obtained in the pulsed radiolysis of solutions of 2 X M N2O and M CH3I. The intensity of the spectrum is reduced by approximately half in the absence of N20 and is completely removed by large concentrations of OH radical scavengers such as methanol and ethanol. Assuming that only OH radicals produce the spectrum, an absorption coefficient E 5000 is measured at 310 mp. The intensity of the spectrum is unaffected by oxygen and this spectrum is attributed to a complex of an iodine atom with a methyl iodide molecule, I (CHJ). The Journd of Physical Chemistry
-
1.0 -
+ CH3I +CH3 + I-
+ C2H4 +C P H ~
-
Figure 1. Spectrum of I- in a solution of 50 p M CHJ-10-* M ethylene: - , literature data for I-.
Rate of Appearance of I-. Figure 1 shows the spectrum produced in a deoxygenated solution of 50 p M CH31-10-a M c2H4. The spectrum grows in after the pulse to give a permanently absorbing product. The similarity of the spectrum to that given in the literature for iodide ion suggests that the observed product is indeed the stable iodide ion.
H
I
8
€1-
Results
eaq-
,
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.0 0.9 1.0 tag [ m o - l I - l , ] -
1.1
1.2
Figure 2. Kinetic plot of results in Figure 1, concentrations of eaQ-and I- expressed in arbitrary units.
Decay of the I ( C H J ) Complex. In the presence of iodide ion the spectrum decays by a first-order process proportional to [I-] and produces directly the spectrum of 12-. The mean of ten measurements gives ~ I ( c H , I ) + I - = 6.0 f 0.5 X log sec-'. The rate of decay of the spectrum is also increased by methyl and ethyl alcohols but not by sodium acetate and acetone. The rate of decay in the presence of alcohol is first order and directly proportional to [alcohol]: ~ I ( C H , I ) + C H ~ O H= 1.7 i 0.4 X 109sec-1. The decay of the I(CHa1) complex in the absence of an added scavenger is first order at low intensities and (6) J. K. Thomas, S. Gordon, and E. J. Hart, J. Phys. Chem.,68, 1524 (1964).
(7) E. J. Hart, private communication. (8) J. K. Thomas, Trans. Faraday Soc., 61, 702 (1965).
1921
PULSERADIOLYSIS OF METHYL IODIDE AND METHYL BROMIDE
,
0.181
,
,
,
,
,
,
,
,
,
,
,
, 0.2
-
t A SLOPE
?L----l 0.0270 I 280
290
310
300
320 330 340 350 360 570 360 390 400 0 IO 20 30 40 50 60 70 80 90 10
WAVELENGTH Imp)-
Figure 3. Spectrum of I(CHJ) in a solution of 2 X 10-* M NnO-lO-' M CHJ.
2.9 2.8 2.7 2.6 2.5 2.4 2.3 2.2 2. I
t
2
2.0
-g
1.8 1.7
1
1
,
1
&
1
1
,
,
1
Figure 5. Plot of A S vs. (I(CHII))~ from Figure 4; AS = (slope of curve slope of c w e at low intensity); I(CHs1)Z expressed in moles per liter.
-
1
1
1
X
a30 0.28
-
1
I
I
1
0.22 0.20 o.18 -
0.26
0.24
t
1.9
-
81)
0.16
c
0.14-
2
0.12-
:f//--;
I.6
0.04
I.5 I.4
0.02
I .3
0
I .2
I
NO. PULSES
I. I
I .o
0.9 ~
O*'O
4 8 12 16 20 24 28 32 36 40 44 48 52 56 Time 1psec)-
Figure 4. Decay of I(CHs1) expressed as -log (optical density I(CHg.1) a t 310 mp) us. time, p s e c . The point a t time zero is moved on the -log OD axis to a convenient position. The initial [I(CHJ)], expressed as optical density at 310 mp, and the intensity, ev/l./pulse X lolo, respectively, are, for low intensity, 0.008-0.041 and 0.25-1.29; O, 0.148 and 4.65; A, 0.208 and 6.55; and X, 0.38 and 12.0.
independent of intensity. At high intensities the rate of decay increases with intensity. The results are shown in Figures 4 and 5. Iodine Yields. Figure 6 shows the yield of iodine Iz (measured as 13-) os. dose expressed as number of pulses in the pulse radiolysis of solutions of CHJ with and without NzO. In the CH3I solutions alone the
-
Figure 6. Yield of 12 as optical density of 11- us. dose expressed as number of pulses. The intensity, ev/l./pulse X lolo, is for A, 1.32 (2 X M N z O - ~ O -CHaI) ~ (1); 0, 1.82 (10-3M C&I) (2); and 0,2.80 (2 X 10-2 M Nz0-10-8 M CHsI-10-8 M methanol (3). (Because of the higher intensity the abscissa is plotted as number of pulses X 0.65.)
initial yield of iodine G(I2) = 1.10 while in N20 saturated solutions G(Iz) = 2.50, i.e., I/4G(eaq-) G(0H)). There is no effect of millimolar concentrations of methanol and ethanol on these 12 yields in spite of the fact that the decay of the I(CH3I) is more rapid in these solutions. Radiolysis of Methyl Bromide. In many instances methyl bromide is a more suitable source of methyl radicals than methyl iodide (a) because the OH radical does not give a strongly absorbing species analogous to I(CH31)and @I) the reaction product Br- is more transparent than I-.
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Volume 71, Number 6 Mag 1967
J. K. THOMAS
1922
The pulse radiolysis of deoxygenated solutions of CH3Br shows a weak transient below 280 mp with an absorption that increases with decreasing wavelength. I n the presence of N20 the intensity of the species is nearly doubled. In this system the OH radical abstracts an H atom from the CH3 group of CH3Br to give the radical CHzBr, the yield of this radical being doubled in N20. There is no evidence that Br atoms are produced. In the presence of oxygen a fairly strong absorption is observed in the ultraviolet region with a peak at 260 mp (Figure 7). This spectrum is similar to the spectra of HOz and 0 2 - and is attributed to the species CH3O2 and OZCH2Br. When nitrous oxide is added to the solution the electrons are captured by the NzO yielding OH and subsequently CH2Br is the only radical formed. The spectrum of CH302 may then be obtained as the difference between the spectra with and without N20 in the solution. Another approach used to obtain the spectrum of CH302is to scavenge the OH radicals with potassium thiocyanate to give the radical CKS. The net result is that CNS is formed rapidly while CH302 is formed slowly. The difference in the rate of formation of CXS and CH302 is sufficiently great so that the spectrum of CH302 can be measured. The agreement between the two methods is good and is shown in Figure 7. The oscilloscope tracing of the rate of production of CH3O2 at 260 mp in a solution containing potassium thiocyanate, methyl bromide, and oxygen is shown in Figure 8. The first instantaneous rise of the trace is
1 I5 %
I-
1
'
g
0.3 0.8
-.--as -f I
"
I
'
'
'
I
250 260 270 280 230 300 310 320 330 A (mpl-
Figure 7. Spectra of 02CH2Brand CH~OZ:0 , M CHaBr--10-4 M 0210-2 M KCNS; 0 , 10-3 M CH3Br-10-4 M 02; x, 10-3 M CH3Br-10-' M Orsaturated NzO; - - - -, solution x divided by 1.92, Le., spectrum of 02CH2Br; -.-., difference between solution 0 and solution x /1.92, Le., CHsOz.
The Journal
of
Physical Chemistry
F
4
due to the CNS radical, while the portion which grows to a plateau over many microseconds is due to CH302. This picture is read on the ANL reader CHLOE and the digitalized information is fed into the ANL, CDC 3600 computer. This machine calculates the optical density a t time t and plots out the results in the desired form. Figure 9 shows the first-order plot from Figure 8. The computer calculates the least-square slope of the line and hence the rate constant of the reaction CH3
+
0 2
--+ CH30z
(1)
The mean of ten results gives ICl = 4.7 f 0.7 X lo9 sec-l a t 23" ; from the effect of temperature on the rate constant an activation energy E = 3.5 i 0.3 kcal is measured for reaction 1. Production of Methane. I n the low intensity Cow y radiolysis of aqueous solutions of M CHJ and alcohols, methane is produced with G(CH4) = 2.45 f 0.15
+ RHOH +CH4 + ROH
(2)
The same result is obtained with solutions of methyl bromide and alcohol. The yield of methane is independent of alcohol concentration from 0.1 to 1.0 M and of CH3Br from to lon3 M . With CHd, however, G(CH4) increases with CH31 concentrations above 3 X M reaching 10.9 at [CHJ] = 5 X M . This is due to a chain reaction ROH
t
0.1 0 240
8
I
Figure 8. Oscilloscope trace of solution 10-3 M M 02. CHaBr-10-2 M KCNS-5.25 X
CH3
I
+
+-
+ CH3I +RIOH + CH3
Oxygen lowers the yield of methane due to reaction 1 competing with reaction 2. This expression now holds Icz[RHOHl G(CH4) G"(CH4) - G(CH4) ki [OzI where G"(CH4)is the yield of methane in the absence of oxygen. A typical plot of the above equation is shown in Figure 10 for the competition between isopropyl alcohol and oxygen for CH3 radicals. The slope
PULSE RADIOLYSIS OF METHYL IODIDE AND METHYL BROMIDE
-0.30
- 0.70 [02] = 5.25 x I O - 5 ~ SLOPE = 0.105x IO' k =4.6~ IO9
-1.10
-8
-1.50
1923
of I- places a limit of less than 80 nsec on the lifetime of any negative ion intermediate. The fate of the H atom is not illustrated in these experiments; it probably reacts rapidly with CH31 to produce CH3 and HI.9 The presence of the methyl radical is demonstrated by the production of methane gas in the lowintensity y radiolysis of aqueous solutions of M CHaI and 1 M isopropyl alcohol. The methane yield G(CH4) = 2.45 0.151°comes from the reaction
*
CH3
0.00 2.00
4.W
6.03
8.00 10.00 12.00 14.00 16.00 18.00 20.00 TIME (MICROSECONDS)
Figure 9. Firsborder plot of the results of Figure 8.
+ (CH&CHOH
--t
(CH&COH
The third short-lived species, Figure 3, is attributed to the complex of the iodine atom with CH31. This is in accord with the doubling of the intensity of the spectrum on addition of N2O via reaction 3
N20
+ eaq- 5 OH + N2
(3) and by the elimination of the spectrum by OH radical scavengers. The OH radical could react wia reactions 4-6. The absence of any effect of oxygen on the spec-
P CH21 + H20 OH + CH3I +CH3 + HOI CH30H
[ISOPROPANOL] [021
K
IO'
- IO-' [BUTADIENE]
or
x
a 0.3-
[METHANOL]
Figure 10. Competition plots for the reaction 02: CHa isopropyl alcohol vs. CHI 0, [isopropyl alcohol] = 1.0 M; and CHa methanol us. CH, butadiene: A, [methanol] = 0.95 M .
+
+
+
+
of the line gives the ratio k 2 / h and as kl = 4.7 X log sec-l, k2 can be calculated. Some typical results are given in Table I together with the rates of addition of CHI radicals to unsaturated compounds, which are measured relative to k2. A typical plot for the competition between methanol and butadiene is shown in Figure 10.
Discussion Methyl Iodide Solutions. The results clearly demonstrate that three products are produced in the radiolysis of oxygen-free solutions of methyl iodide. The stable species in Figure 1is the iodide ion produced from the reaction eaqCHd. The correlation between the rate of disappearance of esq- and the appearance
+
+ CH,
+I
(4) (5) (6)
trum rules out reaction 4 as the radical CH2I would give O&H21 with 02 and alter the spectrum. Reaction 5 is unlikely as the spectrum of HOI is not that shown in Figure 4. This leaves reaction 6 as the most likely reaction path for the OH radical reaction with CH31. The yield of 12, G(L) = 2.5 in N20-CH31 solutions, and the direct decay of the spectrum with I- to give Izalso support reaction 6. The exact nature of the spectrum is, however, less certain. In the gas phase iodine atoms show an absorption below 200 mp;" no absorption has been reported above 300 mp. However, many spectra of I atom complexes have been reported above 300 mp.I2 In the present work a complex between I and water or CH31is possible. A search was made for an absorption due to the IHzO complex in the pulse radiolysis of potassium iodide solutions. The conditions of the experiments are such that the OH radical concentration is in excess of the I- concentration, e#., [OH] = 10 p M , [I-] = 5 p M . Here the OH radical produces I via OH
+
(9)T.J. Hardwick, J . Phya. Chem., 66,2246 (1962). (10) J. K. Thomas, 3rd International Congress of Radiation R e search, Cortina, Italy, 1966. (11) R. J. Donovan and D. Husain, Nature, 206, 171 (1965). (12) T. H. Glover and G. Porter, Proc. Roy. SOC.(London), 262,476 (1961).
Volume 71, Number 6 Mag 1967
J. K. THOMAS
1924
Table I : Rat,esof Reaction of Radicals and Atoms in the Gas Phase and in Water a t 25' Compound
kl X 1 0 g a * b
Ethylene Propylene 1-Butene Isobutylene Butadiene Methanol Ethanol Isopropyl alcohol Oxygen Iodine
4.1 6.5 6.5 5.0 6.5 0.47 0.72 1.74
...
...
kz X lo@%"
3.2 3.9 5.1 10.0 10.0 1.9 x 10-3 1.6 x 5 . 0 x 10+ 20 -10
ka X 10"
4.9 5.3 30.0 39.0 1.25 X 105 0.22 0.59 3.4 4.7 x 106 6 . 0 X lo6
k4
X 10"
2.0 4.0
... ... 1.18 X lo2 0.05 0.14 0.45 Diffusion controlled Diffusion controlled
' kl, OH radical rate; kz, H atom rate; and ka, CHI radical rate, all in water phase; k4, CHI radical rate in gas phase. by I- method, J. K. Thomas, Trans. Faraday SOC.,61,702 (1965). ' Measured via competition with methanol, ref 15.
+
I- + I OH- and an I-H,O complex is formed. However, the only absorption observed is a weak Izband. It is thus left to postulate that the observed spectrum is due to the I(CH31) complex. The nature of iodine complexes has been discussed by Ratzin,l3 who proposes that complexes showing an absorption with two peaks can be considered as perturbed iodide ion spectra. I n the present case if a partial electron transfer occurred to the iodine atoms, then a complex of the form 16-(CH3)6+16-is formed. I n the interpretation of ref 13 this should exhibit a perturbed iodide ion spectrum similar to that shown in Figure 3. The first-order decay of I(CH31) at low intensjties in water and with alcohols is interpreted as a reaction of these solutes with the I(CH31) to give complexes, which like the I(H20) complex have no measurable absorption. The fact that G(12) is not affected by the alcohols shows that the I atom does not undergo an irr 3versible chemical reaction in these systems. At hig'i radiation intensities the I atom complexes recombine and inci-ease the decay rate. This is il1ustrLted in Figure 4; the increasa in the reaction rate over that at low intensities is expressed by k8(I(CHd))2.
+ HYO +I-H2O + CHJ I(CH31) + I(CH31)--+Iz + 2CHJ I(CH3I)
(7) (8)
Figure 5 shows that the above relationship holds and gives 2k8 = 3 X 109 1. M-' sec-'. The value for kl is given by the first-order decay a t low intensity, i.e., kl = 3.0 X lo4sec-'. The Reaction of Methyl Radicals. I n the gas phase the rate of reaction of methyl radicals with oxygen and iodine is very rapid and in the case of oxygen the activation energy is zero. I n water the rates are again rapid and the activation energy for the oxygen reaction is 3.5 kcal/mole. In order for a species such as the The Journal of Physical Chemistry
Measured
CH, radical to diffuse, an adjacent water molecule has to move to create a vacant site. On this basis the activation energy for a diffusion-controlled reaction in water is the activation energy for viscosity change, which is 3.7 kcal/mole. As the measured activation energy of reaction 1 is in agreement with this value, the reaction of a methyl radical with oxygen is diffusion controlled in water as in the gas phase. The comparison of the rates of CH3 radical reactions in water and the gas phase shows that the same order of reactivity occurs in both phases. However, the rates in water are on the whole faster than those in the gas, a similar effect being noted with H a t ~ m s . ' ~ JThis ~ may be due to the solvation of the solute by the water, i.e., a change in reaction energy, or it may be due t o the caging in of the reactants by the solvent. This would promote more encounters between the reactants resulting in an increased rate of reaction. It is instructive to compare the rates of reaction of CH3radical reactions in water with the rates of H atoms and OH radicals. For abstraction reactions the same general increase in reactivity is noted for all radicals in going from methyl alcohol to isopropyl alcchol and shows that the same reaction mechanism holds for abstraction by all these radicals. For addition reictions to unsaturated hydrocarbons the CH3 radical behaves in a similar fashion to the H atom, the rates folloming the atom-localization energy of the hydrocarbon.16 The OH radical reactions show no dependence on the atom-localization energy, however, all the rates being similar and approaching diffusion control. The ab(13) L. I. Katsin, J . Chem. Phys., 23, 2055 (1955); L. 1. Katsin and R. L. McBeth, J . Phys. Chem., 62, 253 (1958). (14) H.Schwars, ibid., 67, 2827 (1963). (15) J. P. Sweet and J. K. Thomas, ibid., 68, 1363 (1964). (16) 5. Sato and R. J. Cvetanovib, J . Am. Chem. Soc., 81, 3223 (1959).
NOTE^
1925
solute value of the CH3 radical rates is significantly less than the H atom rates. This is due to two factors: (a) the methyl radical is planar and an activation energy is required to reach the transition state which has the character of the tetrahedral methane molecule and (b) the preexponential or A factors for the methyl radical reactions are much smaller than the A factors for the H atom reactions. This is due to the significant amount of entropy lost by the CHI and the solute in forming the transition state of the reaction.’? The above data demonstrate the feasibility of using methyl iodide as a convenient source of methyl radicals in water and Other in which are stabilized. For the particular case of water the methyl
iodide is also a convenient source of iodine atoms and should provide valuable data on the various I atom complexes that may be formed.
I wish to thank Mr. B. E, clift and Mr. E. Backstrom, who operated the Argonne linear accelerator. I also wish to extend special praise and thanks to Mr. A. Lent and Mr. J. Butler of the Applied Mathematics Division, who carried out all the programming on CHLOE and the 3600 computer, thereby enabling me to greatly improve the processing of my data. (17) 8. Beneon, ‘$TheFoundations of Chemical Kinetics,” McGraw~ i l Book i CO.,Inc., New York, N. Y., 1960, p 286.
NOTES
Phosphorus-31 Chemical Shifts
of Phosphonate Anions by Jean G. Riess,’ John and John H. Letcher
R.Van Wazer,
Central Research Department, Monsanto Company, St. Louis, Mieaouri (Received August 16, 1966)
In a publication from this laboratory2 10 years ago it was stated that the chemical shifts of “the phosphonic acids and the phosphonates offer a direct way to measure the relative electron-donating ability of organic radicals . . . [with] the stronger electrondonating groups causeling] the lesser shielding of the phosphorus nucleus.” Recently the quantum mechanical theory of 31Pchemical shifts has been elucidated8*‘ for the entire range of phosphorus compounds and it is shown that the chemical shift of compounds in which phosphorus has four substituents is primarily sensitive to the polarity of the u bonds and the total occupation of the d, orbitals of the phosphorus. In view of this theoretical work, it seemed desirable to obtain the 31Pchemical shifts of a number of phosphonates under carefully controlled conditions. These data and the recently published6 chemical shifts of quaternary triphenylphosphonium salts are compared with each
other in light of the general theory and the problem of measuring electron-donating ability. The $lP chemical-shift measurements on the phosphonates were carried out at 40.5 Mc using a Varian HR-100 spectrograph on aqueous solutions in the range of 0.1-0.5 M in phosphorus, employing 85% HsP04 in a capillary tube as the reference standard. In order to avoid complications due to hydrogen association and hydrogen bonding and to make the data comparable, solutions of the phosphonic acids were brought to pH 14 by addition of tetramethylammonium hydroxide. As long as the two hydrogens of the phosphonic acid are neutralized, it was found that the 31P chemical shift is insensitive to rather large variations in alkalinity or dilution.6 In the region of measurement, a severalfold change in either phosphorus or alkalinity concentration caused less than 0.2-ppm change in the chemical shift. The experimental data are presented in Table I. (1)
On leave from the University of Strasbourg, 1964-1966. R. Van Wazer, C. F. Callis, J. N. Shoolery, and R. C. Jones,
(2) J.
J . Am. Chem. Soc., 78, 5715 (1956). (3) J. H. Letoher and J. R. Van Wazer, J. Chem. Phys., 44, 815 (1966). (4) J. H. (5)
Letcher and J. R. Van Wazer, ibid., 45, 2916, 2926 (1966). 9. 0. Grim, W. McFarlane, E. F. Davidoff, and T. J. Marks,
J . Phys. Chem., 70, 581 (1966). Also see Nature, 208, 995 (1965). (6) Also see K. Moedritzer, submitted for publication.
Volume 71, Number 6 May 1967
