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ENERGIES OF AMIDES IN BENZENE SOLUTIONS
where the temperature is much lower and the concentration of oxygen is a few tenths of one per cent., the identity of the mass 30 is not certain and is probably due to both a trace of formaldehyde and NO which have diffused inward. The effect of formaldehyde on the composition of the diffusion flame was obtained by placing pellets of paraformaldehyde on the grid of the burner port. Under these conditions, as may be seen in Table I a considerable concentration of formaldehyde results in higher concentrations of Hz and of CO than in the normal flame. Also, the concentration of 0 2 is lower than usual, presumably because of reaction with the formaldehyde vapor. The formaldehyde concentration in the flame decreases sharply near the edge of the flame where higher temperatures and higher concentrations of oxygen are encountered.
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Inside the luminous cone where practically a1 the methane disappears, it is most unlikely that, formaldehyde plays any role in the chemistry of the methane disappearance. The primary reaction of methane is probably by oxygen and other free radical induced pyrolysis. At the edge of the flame, the importance of the role of formaldehyde in the combustion cannot be resolved by our experiments. Acknowledgments.-The authors wish to acknowledge the assistance of Dr. Richard H. Knipe for many valuable discussions.. Mr. Maynard Hunt gave considerable help in the temperature study. Mr. William Golyer aided in the experimental work and assisted in the operation of the mass spectrometer. Mr. Andreas V. Jensen also assisted in the operation of the mass spectrometer, and Mrs. Helen R. Young aided in the reduction of the mass spectral data.
ENERGIES AND ENTROPIES OF ASSOCIATION FOR AMIDES IN BENZENE SOLUTIONS. PART I BY MANSELDAVIESAND D. K. THOMAS The Edward Davies Chemical Laboratories, University College of Wales, Aberystwyth, England Received September 7,1966
A precise isopiestic method has been used to study the association of trichloroacetamide and some anilides in benzene a t temperatures 25-45". The assumptions involved and the method of analyzing the simultaneously occurring association equilibria are discussed and illustrated for the solutes measured.
The structural character and the energy involved in the molecular association of the amides are of particular interest: the former has frequently been considered in relation to their infrared absorption spectra and has even been advanced as a criterion for deciding the geometrical ("cis" or "trans") structure of the monomeric molecules1; the energy involved is a prime factor in the Hbridge interaction which is BO important in the various configurations of peptide and protein chains. Even so, there are very few adequate determinations of these factors, much of the data being of an essentially qualitative character over a limited range of conditions (concentration and temperature), or involving distinctly arbitrary elements in their interpretation.2 The present account presents the results and analysis of systematic vapor pressure studies of amide solutions in benzene. As data have been obtained by two experimental methods, these have been segregated into Parts I and 11: most of the results are in Part 11, but the general assumptions and method of handling the data are given in the present Part I. Experimental A precise volumetric form of the isopiestic method has already been d e ~ c r i b e d . ~ The more rapid distillation and better draining of benzene compared with water improved both the speed and accuracy of the determinations. Thus the uncertainty of the apparent molecular weights of the solute was usually less than *0.5'% and reaults could be relied upon down to concentrations of 0.02m. (1) M. Tsuboi, et al., Bull. Cham. SOC.Japan, BB, 215 (1949). (2) Muoh of the earlier data is referred t o in: W. E. 9. Turner, "Molecular Association," Longmans Green and Go., London, 1915. (3) M. Davies and D. K. Thomas, THIS JOURNAL,60, 767 (1956).
The standard solute chosen for benzene solutions wai biphenyl. Of the many careful studies of this closely-ideal system we will quote only some of Tompa's results at 2504: no departures from ideality could be established up to 0.25 m , whilst at 1.4 m the deviation from unity of the factor vapor pressure lowering/(v.p. solvent) (mole fraction solvent) was only 1.004. Accordingly, the mean apparent degree of asspciation (f) of any non-volatile solute in benzene is closely approximated t o by the ratio of its stoichiometric mole fraction to that of the biphenyl iu the solutions which are isopiestic. Materials. Benzene.-AnaIar solvent was well-dried (CaS04) and redistilled in an all-glass apparatus. When procurable Analar reagents were used in all the following preparations. Biphenyl.-On distillation the middle fraction, b.p. 217" at 300 mm., was collected. This melted sharply at 71.0" (Heilbronn). Acetanilide: twice recrystallized from benzene; m.p. 113.0". Formanilide: repared from redistilled aniline and formic acid, excess of tge product was treated with boiling petrol ether (60-80O) and after cooling the solid was filtered off and dried; m.p. 49". Trichloroacetamide.-The ethyl ester was first prepared and treated with excess of aqueous a y o n i a and the product recrystallized from benzene; m.p. 141 Trichloroacetanilide: twice recrystallized from benzene : m.p. 95'.
.
Results A typical set of observations is shown in Table I. The temperature read on a thermometer calibrated a t the N.P.L. to f0.02" was 35.01°, the maximum fluctuations being f0.002' (Beckmann) . I n this table, concentrations are in molalities, volumes in ml. The final volumes are those for which no change occurred on further equilibration for 4 to 6 hours and so are taken to give the isopiestic concentrations of the last column. (4)
H. Tompa, J . Chem. Phya.. 16, 292 (1948).
MANSELDAVIESAND D. K. THOMAS
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TABLE I BENZENE SOLUTIONS OF TRICHLOROACETAMIDE-BENZENE SOLUTIONS, BIPHENYL AT 35.01 10.002" Run no.
Biphenyl Trichloroaaetamide Biphenyl Trichloroacetamide Biphenyl Trichloroacetamide Biphenyl Trichloroacetamide Biphenyl Trichloroacetamide Biphenyl Trichloroacetamide
1
2 3 4 5 6
Initial conan.
Initial, vol.
Final vol.
Vol. change
Mean change
Final concn.
0.0172 .0165 .0265 .0298 .0320 .0386 .0393 .0487 .0450 .0529 .0480 .0583
5.812 5.900 5.869 5.577 5.711 5.818 5.456 5.631 5.887 5.749 6.231 5.861
6.122 5.586 5.894 5.552 5.606 5.917 5.372 5.713 6.004 5.626 6.341 5.749
+0.310 -0.314 +O. 025 -0.025 -0.105 +o. 099 -0.084 $0.082 +0.117 -0.123 + O . 110 -0.112
f0.312
0.0164 .0175 .0264 .0300 .0326 .0380 .0399 .0480 .0441 .0541 .0472 .0594
1 0 .025
*o.
102
fO .083 10.120
AO.111
TAEILE I1 lo* X N
f KIP
(f - 1)/f
10'
xN
I Kiz
(f - I)/; 103 X N
f Kir
(f - 1)/N
108 X N
f 108 X N
f los X N
103 X N
2.48 1.13 78 59
(1) Formanilide in benzene a t 24.98' 3.69 4.85 6.47 1.19 1.24s 1.32 94 110 141 61 63 65
0.88 1.038 48 44.6
(2) Acetanilide in benzene A t 24.98' 2.22 2.39 3.54 1.09 1.092 1.14 54 51 61 44.0 41.8 46.6
2.18 1.078 42 36.0
2.38 1.08* 45 37.8
At 35.01' 3.46 1-11, 46 36.8
4.31 1.140 50 37.1
8.35 1.38 162 63
4.28 1.164 64 44.6
5.78 1.218 75 45.0
5.72 1.186 58 38.s
2.38 1.19
(3) Trichloroacetamide in benzene At 24.98' 2.98 3.71 4.63 1.23 1.27 1.326
5.06 1.35
1.42 1.11s
2.29 1.136
At 35.01O 2.93 1.17
3.72 1.21
4.26 1.236
4.57 1.25
1.46 1,058
2.11 1.0%
At 45.05' 2.66 1.11
3.36 1.14
4.06 1.17
4.68 1.195
9.04 1.08'
11.72 1.084
15.02 1.085
8.95 1.06
11.58 1.07
14.95 1,072
f
0.88 1.021
108 X N
0.81
(4) Trichloroacetanilide in benzene At 24-68' 7.34 1.65 4.93 1.044 1.07, 1.084 At 35.02' 4.78 7.20 1.56 1.054 1.07 1.02
1.01, f Table I1 provides a summary of the isopiestic observations. N is the stoichiometric mole fraction of t,he solute in the solution. The mean degree of association for the solute (f) is obtained by dividing N by the effective mole fraction of the solute-in practice by the mole fraction of biphenyl in the isopiestic solution. This involves the major assumption that departures from ideality of the solute are due entirely to its association : the adequacy of this assumption is discussed later. A greater
range of N values was precluded in some cases by cojresponds to limited solubility. N = 10 X 0.13 M . The factors Klzand (f - 1 ) / N are in mole fraction units. Discussion In 1943 Kreuaerb made a number of important general deductions by applying the Law of Mass Action to associative equilibria. Some similar (6) J. Kreuzer. 2. phyaik. Chem., B53, 213 (1943).
results have since been deduced by statistical methods, Firstly, we define Kreuzer’s factors in terms of mole fractions-as he used the thermodynamically less acceptable molar concentration unit-and then quote those of his relations we shall use. Nnis the mole fraction of a n-mer in solution. (1) (2)
(3) (4)
(5)
(6)
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ENERGIES OF AMIDESIN BENZENE SOLUTIONS
June, 1956
N = C n N , = total concn. of solute in terms of mole n fraction as monomer 8 = E N . = total effective mole fraction of all n n-mer8 f = N/N = mean degree of association LY = NI/N = stoichiometric fraction of solute present as monomer E = N1/$ = fraction of total independent molecular unit formed by monomer K,, = l/lcnl = N,/Nzn = association constant for formation of n-mer from monomer: (knlis the dissociation constant).
Micro pipets Loading chamber
Level of water in the thermostat
Rubber bung Insulating material
Pair of matched thermistors
w
Without making any assumptions as to the number or nature of the associative equilibria involved, apart from their individual conformity to the Law of Mass Action in mole fraction units of the species directly involved, Kreuzer’s general relations become
Dewar vessel Well of solvent
Fig. 1.
have studied the appearance of “association” bands in the infrared absorption, the anomalous variation of the molecular polarizations* and the = JON‘ (1/a)dN1 (7) departures from osmotic ideality are already so marked a t low concentrations (
