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Ion-molecule reactions in ethyl chloride
29 International
Journal of Mass
Spectromtry
and Ion Physics,
I4 (1974)
29-44
,Q
ION-MOLECULE
2. ~UCZYl?SKI Department
REACTIONS
IN ETHYL
CHLORIDE
AND I-i. WiNCEL
of Radiation
Chemistry,
Institute
of Nuclear
Research,
03-19.5 Warsrawa
(Poland)
(Received 13 October 1973)
ABSTRACT
The ion-molecule reactions in gaseous ethyl chloride have been investigated in a mass spectrometer at pressures from 0.01 to 2 torr and at temperatures between 190 and 400 I(. The major primary ions, CzH,Cl+ and CZHSL, react with ethyl chloride to form CZH,CLH+ which further produces C,H,,CI’. At low temperatures, in the range 190-270 K, the C4H,eCl+ ion undergoes a stepwise solvation C2H,0Cl’(C2HsCl),_,+C2H,Cl + C4H,eCl+(C2H,C1),, where n = 1, 2, 3 and 4. Thermodynamic values for the 0,l and I,‘>-equilibria have been obtained from the equilibrium constants and their temperature coefficients determined in pure &H&I and in the H+Z,H,Cl system.
INTRODUCTION
The ion-molecule reactions in ethyl chloride have been invest,gated utilizing a variety of experimental techniques [I, 23. The results obtained in these studies seem to be in general agreement as to the identification of the rr,actions occurring in the system. The majority of reactions lead to the forma&n of the C,H,&li ion which remains the main ionic product at high pressire and relatively high (250 “C) [1 ] as well as moderate [Z] temperatures. It has been demonstrated by Field et al. [3-91 that at high pressures in the reaction chamber of the mass spectrometer the lifetime of the “unstable” ions increased markedly as the temperature was lowered and many interesting association ions could be observed at subambient temperatures. The present study was undertaken to examine the ionic reactions of ethyl chloride- occurring in a mass spectrometer under conditions of low temperature and high pressure in C,H&l and HZ-&H,Cl systems. The ethyl chloride in the. C&Cl-H, mixture is ionized by chemical ioa-
30 ization H3+ + C,H,Cl
4
C,H,ClH*
i- H2
(1)
The H, + ion is an easily accessible source of protons and has been used previously by several investigators [lO-141. H3+ formed by the reaction H,++H2
--+ Hs++H
(2)
is initially in a highly excited vibrational state (about 2 eV) 115, 163 which is rapidly moderated Sy subsequent collisions with Hz. It is reasonable to assume that at high pressures, as in this case, Hs+ is in the ground vibrational state.
EXPERIMENTAL
The results reported here were obtained with the magnetic mass spectrometer described in detail elsewhere [17J A modified version of the ion source has been used in these studies. The main aim of the source modification was to control and measure the gas temperature inside the reaction chamber. The new arrangement is schematically presented in Fig. 1. Cooling is accomplished by flowing nitrogen
SECTION
G E N
2
z
a
84
9
,o
G
A-A
RF
II II-
TH-
W
Fig. 1. Schematic drawing of ion source. RC = reaction chamber; EG = e!ectron gun; R = repeller e!ectrode; ET = electron trap; F, PI, P2 = ion focl;sing electrodes; G = grounded electrode; S = screening mesh; TH-1 and TN-II = thermocouples I and II, respectively; B = reaction chamber block; CT = copper tubing.
gas or other suitable coolants through copper tubing CT to the reaction chamber block B. Coolant is circulated through this tubing continuously to maintain a constant temperature while data are taken. The temperature of the reaction chamber walk. can be controlled at selected values between 120 R and 400 K. By. using thermocouples located in different places in the chamber, in
31
separate experiments we found that quite large temperature gradients occur on the chamber walls, due to absorption of the radiant energy of the hot hlament in the region of the electron entrance slit and nonuniform heat losses. In such a situation measurement of the gas temperature by means of thermocouplek placed in the chamber walls, as is usually done, seems to be very problematical
and there-
fore we paid special attention to measuring the temperature of the reactant was. In order to obtain the reaction temperature the special thermocouple TH-I placed directly inside the reaction chamber, as shown in Fig. 1, continuously indicated the temperature near the reaction region. The thermocotiple is calibrated after mounting in the ion source, with a copy of the reaction chamber specially manufactured for this purpose. The second thermocouple TI-I-II is placed outside the reaction chamber at the position shown in Fig. 1. The temperature was measured with an average precision of 0.2 “C. Temperature values obtained with both thermocouples at different reaction chamber pressures for several gases are given in Fig. 2 as an example. As can be seen, at constant wall temperature (T,) the temperature measured inside the chamber (Ti) varies with pressure up to about 0.5 torr and depends on the sort of gas introduced. Furthermore, when
350 1 h&t-Q-@--e-+-t-+l@qJv-v-v-v-v-o-
-oz-
VQ
Fig. 2. Dependence of the temperature measured by thermocouple TH-I inside the reaction chamber (open points) and by thermocouple TH-II in the chamber walls (full points) on the pressure of different gases.
32 the temperature rises, the difference T,- T; decreases, in this temperature range. Similar results were obtained with a thermistor temperature sensor. The observed behaviour of Ti with pressure is the result of a combination of the heat exchange processes: (1) conductivity through the gas, (2) conductivity by electric contact of the thermocouple, and (3) heat exchange by radiation. The gas temperature, Tg, can be evaluated on the basis of the following expression;vhich fortows from theenergetic balanceof the processes mentioned j181
*ET= Ti -I-a(Ti - 7”) -t-b(T: - T$)
(3)
where 7” is the temperature of the electric contact of the thermocouple; a and b are coefficients* depending on the instrumental parameters, on the gas used and its temperature and on the gas pressure for the low pressure range (< 0.4 torr). At higher pressures (above some 0.5 torr), with viscous gas flow, the coefficients a and b should be pressure independent. Indeed, we found (Fig. 2) that under these conditions the measured value of Ti is relatively pressure independent_ 4t
20
200 5,
K
*’
300
Fig. 3. Variation of the temperature values meascred inside the reaction chamber (thermocouple TH-I) with the gas temperature for different gases at a pressure of 1 torr. The insert of the figure shows the variation of the temperature differences (T,- Ti) and (Tp-TV) with TI and TW, respectiveljr, for several gases (pressure = 1 torr). * The coefficients a and f~can be expressed as a = q/Z - a and b = Q/L - a where 1 and CYare the coefficients, respectively: cl and cz are coefficients depending on instrumental parameters.
thermalconductivityand accommodation
33
With the use of various TS and T,,, values measured for several gases with different thermal conductivity coefficients [19] (the gases Hz, N, and Xe were used here) and taking for simplicity the accommodation coefficients of the gases used as 1 and with use of eqn. (3) we can evaluate the Q, b and T, values for a given temperature*. Using these values and the measured values .of ri and T, the true gas temperature can be evaluated by means of eqn. (3). Figure 3 shows the plot of Ti versus Ts,- in the insert of the figure the dependence of the differences (r’ - Ti) and (Tg - T,,.) on Y’, and T,, respectively, are presEnted, and it can be seen that with the exception of Hz the measured temperatures (Ti and T,) differ markedly from the gas temperature and these differences increase as the thermal conductivity of the gases decreases. This fact, if more common, may lead to some discrepancies between the thermodynamic values obtained in different laboratories and it should perhaps be taken into account, especially in the case when a hot filament located close to the reaction chamber is used to create the primary ions. In this work T’ values calculated in the manner described above were used to obtain the thermodynamic values. The pressure in the reaction chamber was measured with a capacitance MCT-Atlas Werke micromanometer connected directly to the chamber through a 5 mm-i.d. pipe. The temperature of the pressure sensor was maintained at 55 “C. The gauge was calibrated with the static pressures measured by mercury manometers. In the studies on C2H,CI-Hz mixture the typical C,H,Cl concentration was 10 “/ However, the true concentration of C2H,Cl in the reaction region may differ from this due to adsorption-desorption on the walls of the chamber, especially in the experiments performed at different temperatures. In order to see how this effect depends upon temperature, the temperature dependence measurements of CzHsCl concentration in the C2H,Cl - H, mixture after leaving the chamber were performed in separate experiments with the reaction chambercopy (as used in our mass spectrometer). The mixture composition being at approximately atmospheric piessure wa s analyzed by means of a gas chromatography method. The data are presented in Fig. 4 and, as may be expected, the &H,Ci concentration in the mixture decreases somewhat as the temperature falls. Assuming that this effect is also quantitatively valid for the ion-source pressure range investigated, the correction in the pressure values used for the determination of the thermodynamic quantities has been made. The ionizing electrons enterin g the reaction chamber had an ener,gy of 1000 eV. Some experiments with different electron energies were also made. The ion accelerating voltage, U, was 2 kV. In most of the experiments the electric fields in the reaction region were * The values obtained for the apparatus coefficients cl and c2 4.25 X lo-” m K WV-‘, respectively.
are equal to 1.02X lo-”
and
Fig_ 4. Dependence chamber
of the CIHsCl
concentration
in the HZ-C~H&l
mixture Ieaving the reaction
on the gas temperature.
reduced by operating the electron trap (U,) and repeller (I;rr) at the ionization chamber potential (U). Some experiments were also performed at different repeller fields and at different trap voltages. Space-charge effects were minimized by operating at ionizing currents of about 0.1 ,uA (the filament current was kept constant during the experiments). In the subambient temperature range the total ion current measured is significantly reduced as the temperatur, a of the chamber is lowered. However, it seems that these changes resulted in no pronounced changes in the spectra. This effect is probably due to the collection of some charge in the reaction chamber (perhaps by the adsorbed gas). ‘rhe sensitivity of our ion detection system is quite high and no pushing force has been applied to overcome this effect. In order to reduce penetration of the ion acceleration voltage into the reastion chamber region, the distance between the ion exit slit and the first electrode of the ion focusing system was changed from 4 to 13 mm and the voltage apphed to this electrode was reduced to minimum operating potential (4 V) for most of the experiments. Besides, the region between the reaction chamber and the focus electrode was screened by a highly transparent wire mesh maintained at the reaction chamber potential. Thus, this modification ptoduced a weak electric field around the reaction chamber and created apprcximately field-free conditions in the chamber. Data - presented in this paper were obtained w.ith two ion-focusing
’
systems; Optics I: electrode F at - 20 V, half-plates P, and P, at -40 and -21 V, respectively; Optics II: electrode F ac -4 V, half-plates-P, and F2 at -34 and - 17 V, respectively.
35 In order to improve the operating conditions of our mass spectrometer for low temperature and high pressure studies the examination of the ionic processes of C,I-I,Ci was done on a binary mixture of Hz (reactant gas) and QH,Cl (material under investigation).
PRESSURE,
Torr
Fig. 5. Pressure-dependent variations of the relative ktersities of the prominent ions in pure CzH5CI (open points, To = 257 K, optics I) and in the H2-CJi5Cl system (full points, To = 233 K, optics II).
We give in Fig. 5 a plot of thz relative abundances of the ion currents determined using each of the two optics systems with. the apparatus, against the pressure_ It can be seen that the ions exhibit different behaviour for the cases mentioned; particularly, in the studies before modification, at pressures above 1.2 torr, the intensities of both solvates, C,H1&lfC2H,Cl and C4H,&1* (CZH5Cl)z, decrease sharply while those of C,H,,Clf and CaHS+ rise with pressure. In addition, some nonintegral mass numbers corresponding to the loss of C2H,CI molecules from C,JYX,~CI’ and their solvates were observed. These data indicate that collision-induced dissociation outside the reaction chamber occurred very eaciently before the modification. The results obtained after the modijkation (Fig. 5) seem to prove that this process is significantly reduced. Materials The ethyl chloride products of BDH Ltd. (English product) and “Frkgata” (pro narcosi; Polish product) were used without further purification. The hydrogen
36 used was Polish product prepurified grade. The C2H,C1 before being introduced into the mass spectrometer was degassed by the freezing and thawing technique at liquid nitrogen temperature.
RE9JL-S General
AND
DISCLJSSION
pressure and remperature
dependences
The observed variations with pressure of the intensities of the main ions in ethyl chloride are shown in Fig. 6 (for experiments performed at 285 and 403 K). As can be seen, the relative intensities of the C2HS+, C2H5+, C,H,Cl’ and C,H,CIHi ions decrease sharply and that of C,H,,Cl’ rises with pressure
m4 6-e-o-e Q
& CIj
1.02
Cl
0.025
qoso
0,075
0,loo
0
as
1.5
:
Fig. 6. Pressure-dependent variations of the relative intensities of the major ions in pure C2H5CI at indicated temperatures; optics I. The points given in graph (b) refer to the same ions as indicated in graph (a).
which is in close agreement Hughes [l ] and by Eeauchamp
with trends previously reported by Tiernan and et al. [2]. The reactions which lead to the formation
of C4H10Cl’ have been documented in earlier studies [I, 21 and were observed here and are summarized in Table 1. As is seen in Fig. 6a, the C,H, + ion is formed in considerable amounts, particularly at pressures above 0.3 torr. The ion-cyclotron study of ethyl chloride demonstrated [2] that this ion is formed in the two reactions: C2H,’
+C,H,Cl
-+ C,H,+ +Cl
(4)
37 TABLE
1
ION--WOLECULE
OBSERVED
REACTIONS
IN
ETHYL
CHLORIDE
Reaction
Reaction ivo.
C2H5C!l+ +C2HsCl
-+ -+ -+ +
C~H~ClH’i(C,H,Cl) C,H,Cl+ i- (CzHSCl, C,H, + f (HCl, Cl) C,Hs + + (2 HCI)
Hz)
tw
hHr (kcai/nzole)a
Re$
-11
1,2
(171
f12 -3-(13) -2O-(-335)
1,2 2 2
(16)
(5)
C2HS + f &H&l
+ CLH&lH+ f (CzHs) -+ CsHs’i(CH4, HCl)
(1s) WI
-14 - 12-(-22)
1, 2 1
&I&+
+ -+ + -+ -+
&H&l + + (C,H,) C2H&.Zl+ i (C,H,) CaHs + -i- (HCI 1 C,H, + + (Cl) C2H, + i (CzHoCI)
(20) (21) (22)
-11 f12 -27-(-42) -lO-(-20) -i-12
1, 2 1, 2 1, 2 2 3
-+ C2H5ClHi -!- (C~HZ) + CzHs + t (&H&l)
(23) (24)
-28 -21
1, 2 1
+ C,H, + -!- (C&Cl) + c,I-X&+ i (&H&I)
(25) (26)
-53
1 1
+ CzHs + i (CHzCIz)
(27)
-21-(-41)
2
4
(98)
+&H&l
C2Hs + f&H&I C&
•,-i-
CH2Cli
GHXI fC2HsCI
C,H,CiH+
+CzH,CI
C,HI,,Cl’+HCl
(44) (6)
1, 2
n The heats of the reactions were calculated on the basis of the following thermochemical data: a F. H. Field and J. L. Frank!in, EZectron hpacr Phenomena and Properries of Gaseozzs Ions, Academic Press, New York, 1957. b Tizermai-Constants of Organic Substances, (in Russian), AN SSSR, Moscow, 1970. c M. H. Karapet’yanc and M. L. Karapet’yanc, FzzndamentaJThermodq-namic Co.-rants of Inorganic and Organic Substances, (in Russian), Khimija, MOSCOW, 1968. d V. I. Vedeneev, L. Gurvich, V. N. Kondrat’yew,.V. A. Medvedev and Ve. L. Frankevich, Bend Energies, Ionization Potentials and Electron Aflnities, (in Russian), AN SSSR, MOSCOW, 1962, p. 139.
and C2H,CI’+C2H,C1
+ C,H,++HCI+CI
OS the other hand, the experiments of Tiernan and Hughes [I ] performed. with tandem mass spectrometer show that no C4HPi is present among the ionic products from the reactions of CzH4+ and C,H,Ci* with C,H,Ci. These investigators found, however, that instead of reaction (4) the following reactionoccurs the
C,H,+ + C&Cl
+ C,H, * + (C2H,Cl)
(6)
The high-pressure studies El] have shown that the CJ&,’ ion is formed in great amount (10 %), but no suggestion concerning its origins has been given by the authors. The previous [l ] as well as the present high-temper-ature experiments* * From the point of view of ion-residence tims the conditions under which the high-pressure experiments of Tieman and Hughes [l] (at field strength of 10.6 V/cm) have been carried out are approximately equivalent to those under which our data, presented in Fig. 6a, were obtained-
38 seem to indicate that the C,H, + ion is essentially unreactive with the C2H,Cl. At high pressures and relatively low temperature, however, this ion disappears, as shown by the results given in Fig. 6b. The figure also shows that as the pressure increases the intensity of C&I,&1 decreases somewhat and at the same time the CtH,’ and C,H,Cl+ ions are produced. It is most probable that these ions are formed as a result of the collisional dissociation of the CSH,&l ions in the regions outside the source by the following reactions*: C,H, ,Cl’
+ M + C,H,
+ + C2H,Cl
+ C2H,Cl’
+C,H,
(7) (8)
The overall trends observed (Fig. 7) for C,H,,Clf, CzHsi and C2H,Cli with increasing temperature are. similar to those found with pressure changes (Fig. 6b) but the behaviour of the C,H, + ion is different (the production of C,H,+ rises with the temperature). The effect of temperature on the observed changes may be explained in terms of an increased internal excitation of the C4H1&1+ ions and as a result of an increased probability of their decomposition (reactions (7) and (S)) msl _ -de and outside the reaction chamber as the temperature is increased. The fact that the collisional decomposition of CSHl,,Cl+ occurs by loss of C,H, and &H,Cl additionally supports the previous investigators’ suggestions [I, 21 that the C,H,,Cli ions are of diethylchloronium type structure C2H,Cl*C,H,. The observed change in intensity of C,H,’ with pressure (Fig. 6) and temperature (Fig. 7) may be the result of the influence of the internal energy of the reaction complexes formed in the collisionsC,H,Cl’-C,H,Cl and&H*+-C2H,Cl on the relative probabilities of the competing reaction channels. However, the relative yields of the ionic products (excepting that of C2H,+) of these reactions are substantiaxy independent of pressure above 0.3 torr (Fig. 6) and no adduct ions resulting from the mentioned collisions have been observed in the previous [I, 23 or the present experiments even at low temperatures (up to 193 K). In view of this, the above explanation dealing with the behaviour of C4HQt with pressure and temperature seems to be unlikely** and we suspect that in high-pressure experiments all the C,Hg f ions do not necessarily originate from reactions (4) and (5). The additional step for production of this ion may be CSH&l+*
+ C4H9+ +HCl
(9)
* The ratio of the intensities of the peaks occurring at nonintegral masses and corresponding to the reactions (7) and (S), to the parent ion intensity, C4HIoC1+, generaliy increased with increasing pressure. ** in contradistinction to reaction (4), reaction (6) is endothermic by about 12 kcal/mole, and it
w&d be difficult to understand the increase of product yield through reaction (6) when the internal energy of the collision complex CzHo+-C2H5Ci, as may be reasonably assumed, decreases with pressure (due to collisional deactivation of the CzH4+ ion).
39 1PC
I.075
0.7: !j
L C Ij
c'j
Q5l
IPSO
3.025
0.251 /
./=
i
CL%@d
/
oio/Q4 270
/--
290
280
300
310
0
Fig. 7. Relative intensities of the C,H,,C!I +, C4Hg+. C2HsCli and C2Hsi as a function of gas temperatuie. Pressure = 0.96 torr, optics 1.
180
210
3,
K
X0
ions in pure CzHsCi
270
Fig. 8. Dependences of the relative intensities of the major ions in pure C2HsCI perature. Pressure = 0.6 torr, optics I.
on the gas tern-
40 Then, the observed pressure dependence is readily explained in terms of the C,H, + ion production by decomposition of excited C4H,,Ci’*. The decomposition rate might decrease as the latter ion is stabilized and, of course, it increases when the temperature rises.
Salvation processes At the high pressures and relatively low temperatures used in the present experiments, the C,HroCl* ion undergoes a stepwise salvation. A typical plot of the fractional intensities of observed ions in pure ethyl chloride at P = 0.6 torr as a function of temperature is shown in Fig. 8. It can be seen that as the temperature decreases the intensity of C,H,,Cl’ decreases accompanied by an increase in the intensities of the solvated ions, C,H,&l~(C2H,C.l),. In pure C&H,Cl, solvates with n up to 3 have been usually observed, and only at temperatures below 193 K has the solvate with 12 = 4 sometimes been seen with a very low yield. Observation of C,H,.Clt(CzH,C1), was complicated by condensation of the ethyl chloride at a temperature of about 190 K. Data obtained for the H,+H,C!l system wer,p similar except for the solvates with n > 2, which were below our detection limit. The intensity changes with pressure and temperature (Figs. 5 and 8) suggest that the observed solvates are formed by reversible reactions, so that we can write C,H,&i’
(C2H,Cl),_
The equilibrium constants,
IL-
1.n =
1 i- CaH,Cl
it C,H,,,Cl+
K, _ , , “, of these reactions
k4HloCI+(C2H5CI), k.,HloCI
+ (C2H5CI),_
1 - pCLHsCl
(&H&l),
(10)
may be expressed by
(11)
are the intensities of the C4H10Cl+ (C,H,Ci), and C,H,<(C2H,C1),,_, ions, produced in the reaction chamber, respectively, while P C2H5C~is the pressure of &H,Cl. Figure 9 shows the plots of K,, r and K,, Z as a function of pressure for pure C,H,C!l and the H,-C2H5CI mixture, and one can see that rather different behaviour of the equilibrium constants is observed for both these systems: in pure C2H,C1 the I&, r value decreases over the whole pressure range investigated, while K,. Z is constant; in the mixture H,-C,H,Cl both I&,, 1 and KISz initially decrease (up to about 0.9 torr) but afterwards remain constant_ This suggests that in the latter system both the 0,l and 1,Zequilibria approach or are in an equilibrium state. Most probably the difference in behaviour of K,, r and I&, z in the C2H,CI and H,+H,CI systems arises from the experimental conditions, which in the former case were more drastic and therefore collisional dissociation outside the reaction chamber was more efkient than in the latter cage. The temperature dependence of the measured equilibrium constants for both systems
where
~C4HloC1+(C2H,CI),
and
~C4H&I’(C2H5CI),-l
41
5.x+*
0.5
PRESSURE
1.0
1.5
2.0 *5-103
, Torr
Fig. 9. Variation of the equilibrium constants Tg = 233 K, optics II; (0, B) pure C2H&I,
KO. I and Kl . z with pressure. T, = 253 K, optics I.
(g,
v ) Hz-CzH5CI,
was investigated. Since the pressure in the reaction chamber varied as the temperature was changed, it was always corrected after every change of temperature in order to maintain its value constant. Duplicate experiments were made on other days and good agreement between the results was found. Values of the equilibrium constant are plotted as a function of inverse temperature in a conventional van ‘t Hoff plot, as for example shown in Gig_ 10. The experiments with the H,C,H,Cl mixture yielded straight line van ‘t HOE plots, which gave an enthalpy value reproducible at two different pressures (0.95 and 1.4 torr); the plot for the 0,1-equilibrium in pure CzH,Cl is linear at higher temperatures and nonlinear at low temperatures (below 227 K). The enthalpy values obtained from the van ‘t Hoff plot for both systems* under study are given in Table 2. We see that a significant difference between the respective AHnc_ In values of these systems exists; values found for pure C2H,C1 presumably are too high due to perturbation of the measured ion intensities by the dissociation outside the reaction chamber. If reaction of desolvation in this region occurs via an intermediate the dissociation of which depends upon the internal ener,gy, then one may expect that this process will be temperature-deptndent 1201.This effect may introduce a systematic error in the temperature-dependence measurements * The ANonS1 value for pure &H&l plot (Fig. 10).
was obtained
from the linear portion
of the van ‘t Hoff
42 1~
IO~T.K
a.5
o
10~
Kn-I,n
!:
atm-i
/ 10
/
:J
"&/
2 °,
t ~0
, 4.5
v 5.0
5.5102
10')T, K
Fig. 10. V a n "t H o f f plots for C4HIoCI+(CzHsCI),,_x--kCzHsC1 ~- C 4 H l o C I + ( C 2 H s C I ) , u n d e r different c o n d i t i o n s . F u l l p o i n t s ; p u r e C2H~CI, PC2H~CI = 0.9 ton., optics I; o p e n p o i n t s : H z - C z H s C1, PtC_,HsCI+H2) -----0.95 t o n ' , optics IL
o f the equilibrium constants, a n d as a c o n s e q u e n c e o f this the e n t h a l p y values o b t a i n e d f r o m van 't H o f f plots m a y be in error. E l e m e n t a r y algebra shows t h a t this effect can raise as well as r e d u c e the true e n t h a l p y values. This d e p e n d s on interference o f various f a c t o r s , such as the relative c o n c e n t r a t i o n s o f p r o d u c t a n d r e a c t a n t ions involved in p a r t i c u l a r equilibria o f the reaction sequence, a n d the m e c h a n i s m a n d kinetics o f the desolvation outside the c h a m b e r . If, f o r example, it is a s s u m e d t h a t dissociation o f solvated ions XMn +, where n = I a n d 2, occurs by p a t h s involving the r a t e - d e t e r m i n i n g steps o f the reactions outside the c h a m b e r (occurring with the same probabilities) XM2 + + M --+ X M + + 2 M
(12)
XM++M
(13)
--+ X + + 2 M
a n d a n equilibrium o f r e a c t i o n XM2_ 1 + M ~ - XMn +
(14)
43 TABLE
2
‘r=RblODmAMIC V.km_rE!j
I-EIE
tC+HJZI K
n--l,
3ooa
(Lzarm-‘)
03
52
I,2
64
J For the H2-C2H5CI
b For pure C2H5Cl.
I&WI-mNS:
+
4.GC300= (kcaljmole)
4W’
4s 21
4EP
(kcallmole)
(eu)
(kcaljmole)
-2.6iO.15 -2.5&0.15
-5.210.3 -4X&0.3
-&7&1_5 - 7.7i1.5
-1313 - 9&3
system; the pressure fraction of C2HsCI = 0.1.
is reached under conditions when Ix >> Ixhi ~2 IxM+ (where Ix, IXM and IXMl are the intensities of the X’, XMf and XMzl ions, respectively), then the measured equilibrium constants for both the O,l and 1,Zequilibria would decrease more rapidly with increasing temperature than in the absence of dissociation outside the chamber. This is for the following case*: we found K,- l,n values to be lower for pure C2H,C1 and the slopes of the van ‘t Hoff plots are steeper than for the case of the Hz--C2H,Ci mixture (Fig. 10) when dissociation outside the chamber seems to be insignificant. On the other hand, if under equilibrium conditions, the ion intensities are as & << kM +z kMz, then, in the case when dissociation outside the chamber is noticeable, the decrease with increasing temperature should be slower for Ke, r and faster for K,, 2 when compared with the case without contribution of the effect mentioned. Of course, the influence of dissociation outside the chamber on the measured enthalpy values may be negligible under certain conditions (when the concentrations of the ions involved in the equilibria do not differ very much), especially for the higher equilibria of the reaction sequence. This seems to explain why satisfactory agreement between thermodynamic quantities has been found in different laboratories [3, 4, 211 for higher equilibria, while discrepancies for the lower ones are very large. We believe that the thermodynamic quantities determined in the present study in the system Hz-CZH5C1, when dissociation outside the chamber seems to be insignificant, are close to the true ones. As can be seen from Table 2, the difference in enthalpy values for 0,l and 1,2-equilibria is insignificant and lies within the limit of the experimental error. * Equation (11) involving the contribution outside the chamber may be rewritten K’,_
%.a =
of dissociation (both collisional and spontaneous) (11’)
k~H~oCi+(C~H~CI),---Ic &HIOCI+
(C2H5CI),_
l +&)kzH,CI
where I, is the intensity of ions arising from dissociation
outside
the chamber.
44 Our data indicate that the association energy of C,H,Cl with the C,H1,CI” ion is relatively low, and it would appear that the interaction between the ligands in the solvate C4H,&1’(C,H,C1), is very weak. It should be noted that the binding strength between C,H,Cl and C,H,,Ci’ is close to that found by Bennett and Field [6] for the association complex CF3 - CEYI,~ (AH” = -4.55 kcal mole-‘) which e IS f ormed at subambient temperatures in the CF4-CH4 -mixture. ACKNOWLEDGEMENi-
The authors are indebted to Mr. K. Grabowski
for technical
assistance.
REFERENCES 1 T. 0. Tieman and B. M. Hughes, A&zn_ Clrem. Se,-_, 82 (1968) 412. 2 J. L. Beauchamp, D. Hoitz, S. D. Woodgate and S. L. Patt, J. Amer_
Chem.
Sot.,
94 (1972)
2798.
3 45 6
7 8 9
10 1I 12 I3
14 15 16 17 18
19 20 21
D. D. F. S.
P. Beggs and F. 11. Field, J. Amer. Chem. Sot., 93 (1971) 1567. P. Eeggs and F. H. Field, J. Amer. Chem. Sue., 93 (1971) 1576. H. Field and D. P. Beggs, J. Amer. Chem. SOL, 93 (1971) 158.5. L. Bennett and F. H. Field, J. Amer. Chem. Sot., 94 (1972) 5188. S. L. Bennett and F. H. Field, J. Amer_ Chem. Sot., 94 (1972) 6305. S. L. Bennett and F. H. Field, J. Amer. C/tern. Sot., 94 (1972) 5186. S. L. Bennett, S. G. Lias and F. H. Field, J. Phys. Chem., 76 (1972) 3919. V. Aquilanti and G. G. Volpi, J. Chem. Phys., 44 (1966) 3174. V. Aquilanti, A. GaBi and G. G. Volpi, J. Chem. Phys., 47 (1967) 831. V. Aquilanti, A. Galli, A. Giardini-Guidoni and G. G. Volpi, J. Chem. Ph_v.s., 48 (1968) 43 10. M. T. Bowers and D. D. Elleman, J. Amer. Chem. Sue., 92 (1970) 1847. M. T. Bowers and P. R. Kemper, J. Amer. Chem. Sot., 93 (1971) 5352. J. 3. Leventhal and L. Friedman, J. Chem. Phys., 50 (1969) 2928. M. T. Bowers and D. D. Elleman, J_ Amer. Chem. SLX., 92 (1970) 7258. H. Wince], ht. J. iWass Spectrom. Ion Phys., 9 (1972) 267. J. Groszkowski, Technika wysokiej pr&ni, Wydawnictwa Naukowo-Techniczne, Warszawa, 1972, pp. 49-52 and 300-305. K. RaZnjeviE, Tab&e ciepbze z wykresami, Wydawnictwa Naukowo-Technicme, Warszawa, 1966. H. Wince1 and J. A. Herman, JCS Faraday Trans., 69 (1973) 1797. A. J. Cunningham, J. D. Payzant and P. Kebarle, J. Amer. Chem. Sot., 94 (1972) 7627.
Journal of Mass
Spectromtry
and Ion Physics,
I4 (1974)
29-44
,Q
ION-MOLECULE
2. ~UCZYl?SKI Department
REACTIONS
IN ETHYL
CHLORIDE
AND I-i. WiNCEL
of Radiation
Chemistry,
Institute
of Nuclear
Research,
03-19.5 Warsrawa
(Poland)
(Received 13 October 1973)
ABSTRACT
The ion-molecule reactions in gaseous ethyl chloride have been investigated in a mass spectrometer at pressures from 0.01 to 2 torr and at temperatures between 190 and 400 I(. The major primary ions, CzH,Cl+ and CZHSL, react with ethyl chloride to form CZH,CLH+ which further produces C,H,,CI’. At low temperatures, in the range 190-270 K, the C4H,eCl+ ion undergoes a stepwise solvation C2H,0Cl’(C2HsCl),_,+C2H,Cl + C4H,eCl+(C2H,C1),, where n = 1, 2, 3 and 4. Thermodynamic values for the 0,l and I,‘>-equilibria have been obtained from the equilibrium constants and their temperature coefficients determined in pure &H&I and in the H+Z,H,Cl system.
INTRODUCTION
The ion-molecule reactions in ethyl chloride have been invest,gated utilizing a variety of experimental techniques [I, 23. The results obtained in these studies seem to be in general agreement as to the identification of the rr,actions occurring in the system. The majority of reactions lead to the forma&n of the C,H,&li ion which remains the main ionic product at high pressire and relatively high (250 “C) [1 ] as well as moderate [Z] temperatures. It has been demonstrated by Field et al. [3-91 that at high pressures in the reaction chamber of the mass spectrometer the lifetime of the “unstable” ions increased markedly as the temperature was lowered and many interesting association ions could be observed at subambient temperatures. The present study was undertaken to examine the ionic reactions of ethyl chloride- occurring in a mass spectrometer under conditions of low temperature and high pressure in C,H&l and HZ-&H,Cl systems. The ethyl chloride in the. C&Cl-H, mixture is ionized by chemical ioa-
30 ization H3+ + C,H,Cl
4
C,H,ClH*
i- H2
(1)
The H, + ion is an easily accessible source of protons and has been used previously by several investigators [lO-141. H3+ formed by the reaction H,++H2
--+ Hs++H
(2)
is initially in a highly excited vibrational state (about 2 eV) 115, 163 which is rapidly moderated Sy subsequent collisions with Hz. It is reasonable to assume that at high pressures, as in this case, Hs+ is in the ground vibrational state.
EXPERIMENTAL
The results reported here were obtained with the magnetic mass spectrometer described in detail elsewhere [17J A modified version of the ion source has been used in these studies. The main aim of the source modification was to control and measure the gas temperature inside the reaction chamber. The new arrangement is schematically presented in Fig. 1. Cooling is accomplished by flowing nitrogen
SECTION
G E N
2
z
a
84
9
,o
G
A-A
RF
II II-
TH-
W
Fig. 1. Schematic drawing of ion source. RC = reaction chamber; EG = e!ectron gun; R = repeller e!ectrode; ET = electron trap; F, PI, P2 = ion focl;sing electrodes; G = grounded electrode; S = screening mesh; TH-1 and TN-II = thermocouples I and II, respectively; B = reaction chamber block; CT = copper tubing.
gas or other suitable coolants through copper tubing CT to the reaction chamber block B. Coolant is circulated through this tubing continuously to maintain a constant temperature while data are taken. The temperature of the reaction chamber walk. can be controlled at selected values between 120 R and 400 K. By. using thermocouples located in different places in the chamber, in
31
separate experiments we found that quite large temperature gradients occur on the chamber walls, due to absorption of the radiant energy of the hot hlament in the region of the electron entrance slit and nonuniform heat losses. In such a situation measurement of the gas temperature by means of thermocouplek placed in the chamber walls, as is usually done, seems to be very problematical
and there-
fore we paid special attention to measuring the temperature of the reactant was. In order to obtain the reaction temperature the special thermocouple TH-I placed directly inside the reaction chamber, as shown in Fig. 1, continuously indicated the temperature near the reaction region. The thermocotiple is calibrated after mounting in the ion source, with a copy of the reaction chamber specially manufactured for this purpose. The second thermocouple TI-I-II is placed outside the reaction chamber at the position shown in Fig. 1. The temperature was measured with an average precision of 0.2 “C. Temperature values obtained with both thermocouples at different reaction chamber pressures for several gases are given in Fig. 2 as an example. As can be seen, at constant wall temperature (T,) the temperature measured inside the chamber (Ti) varies with pressure up to about 0.5 torr and depends on the sort of gas introduced. Furthermore, when
350 1 h&t-Q-@--e-+-t-+l@qJv-v-v-v-v-o-
-oz-
VQ
Fig. 2. Dependence of the temperature measured by thermocouple TH-I inside the reaction chamber (open points) and by thermocouple TH-II in the chamber walls (full points) on the pressure of different gases.
32 the temperature rises, the difference T,- T; decreases, in this temperature range. Similar results were obtained with a thermistor temperature sensor. The observed behaviour of Ti with pressure is the result of a combination of the heat exchange processes: (1) conductivity through the gas, (2) conductivity by electric contact of the thermocouple, and (3) heat exchange by radiation. The gas temperature, Tg, can be evaluated on the basis of the following expression;vhich fortows from theenergetic balanceof the processes mentioned j181
*ET= Ti -I-a(Ti - 7”) -t-b(T: - T$)
(3)
where 7” is the temperature of the electric contact of the thermocouple; a and b are coefficients* depending on the instrumental parameters, on the gas used and its temperature and on the gas pressure for the low pressure range (< 0.4 torr). At higher pressures (above some 0.5 torr), with viscous gas flow, the coefficients a and b should be pressure independent. Indeed, we found (Fig. 2) that under these conditions the measured value of Ti is relatively pressure independent_ 4t
20
200 5,
K
*’
300
Fig. 3. Variation of the temperature values meascred inside the reaction chamber (thermocouple TH-I) with the gas temperature for different gases at a pressure of 1 torr. The insert of the figure shows the variation of the temperature differences (T,- Ti) and (Tp-TV) with TI and TW, respectiveljr, for several gases (pressure = 1 torr). * The coefficients a and f~can be expressed as a = q/Z - a and b = Q/L - a where 1 and CYare the coefficients, respectively: cl and cz are coefficients depending on instrumental parameters.
thermalconductivityand accommodation
33
With the use of various TS and T,,, values measured for several gases with different thermal conductivity coefficients [19] (the gases Hz, N, and Xe were used here) and taking for simplicity the accommodation coefficients of the gases used as 1 and with use of eqn. (3) we can evaluate the Q, b and T, values for a given temperature*. Using these values and the measured values .of ri and T, the true gas temperature can be evaluated by means of eqn. (3). Figure 3 shows the plot of Ti versus Ts,- in the insert of the figure the dependence of the differences (r’ - Ti) and (Tg - T,,.) on Y’, and T,, respectively, are presEnted, and it can be seen that with the exception of Hz the measured temperatures (Ti and T,) differ markedly from the gas temperature and these differences increase as the thermal conductivity of the gases decreases. This fact, if more common, may lead to some discrepancies between the thermodynamic values obtained in different laboratories and it should perhaps be taken into account, especially in the case when a hot filament located close to the reaction chamber is used to create the primary ions. In this work T’ values calculated in the manner described above were used to obtain the thermodynamic values. The pressure in the reaction chamber was measured with a capacitance MCT-Atlas Werke micromanometer connected directly to the chamber through a 5 mm-i.d. pipe. The temperature of the pressure sensor was maintained at 55 “C. The gauge was calibrated with the static pressures measured by mercury manometers. In the studies on C2H,CI-Hz mixture the typical C,H,Cl concentration was 10 “/ However, the true concentration of C2H,Cl in the reaction region may differ from this due to adsorption-desorption on the walls of the chamber, especially in the experiments performed at different temperatures. In order to see how this effect depends upon temperature, the temperature dependence measurements of CzHsCl concentration in the C2H,Cl - H, mixture after leaving the chamber were performed in separate experiments with the reaction chambercopy (as used in our mass spectrometer). The mixture composition being at approximately atmospheric piessure wa s analyzed by means of a gas chromatography method. The data are presented in Fig. 4 and, as may be expected, the &H,Ci concentration in the mixture decreases somewhat as the temperature falls. Assuming that this effect is also quantitatively valid for the ion-source pressure range investigated, the correction in the pressure values used for the determination of the thermodynamic quantities has been made. The ionizing electrons enterin g the reaction chamber had an ener,gy of 1000 eV. Some experiments with different electron energies were also made. The ion accelerating voltage, U, was 2 kV. In most of the experiments the electric fields in the reaction region were * The values obtained for the apparatus coefficients cl and c2 4.25 X lo-” m K WV-‘, respectively.
are equal to 1.02X lo-”
and
Fig_ 4. Dependence chamber
of the CIHsCl
concentration
in the HZ-C~H&l
mixture Ieaving the reaction
on the gas temperature.
reduced by operating the electron trap (U,) and repeller (I;rr) at the ionization chamber potential (U). Some experiments were also performed at different repeller fields and at different trap voltages. Space-charge effects were minimized by operating at ionizing currents of about 0.1 ,uA (the filament current was kept constant during the experiments). In the subambient temperature range the total ion current measured is significantly reduced as the temperatur, a of the chamber is lowered. However, it seems that these changes resulted in no pronounced changes in the spectra. This effect is probably due to the collection of some charge in the reaction chamber (perhaps by the adsorbed gas). ‘rhe sensitivity of our ion detection system is quite high and no pushing force has been applied to overcome this effect. In order to reduce penetration of the ion acceleration voltage into the reastion chamber region, the distance between the ion exit slit and the first electrode of the ion focusing system was changed from 4 to 13 mm and the voltage apphed to this electrode was reduced to minimum operating potential (4 V) for most of the experiments. Besides, the region between the reaction chamber and the focus electrode was screened by a highly transparent wire mesh maintained at the reaction chamber potential. Thus, this modification ptoduced a weak electric field around the reaction chamber and created apprcximately field-free conditions in the chamber. Data - presented in this paper were obtained w.ith two ion-focusing
’
systems; Optics I: electrode F at - 20 V, half-plates P, and P, at -40 and -21 V, respectively; Optics II: electrode F ac -4 V, half-plates-P, and F2 at -34 and - 17 V, respectively.
35 In order to improve the operating conditions of our mass spectrometer for low temperature and high pressure studies the examination of the ionic processes of C,I-I,Ci was done on a binary mixture of Hz (reactant gas) and QH,Cl (material under investigation).
PRESSURE,
Torr
Fig. 5. Pressure-dependent variations of the relative ktersities of the prominent ions in pure CzH5CI (open points, To = 257 K, optics I) and in the H2-CJi5Cl system (full points, To = 233 K, optics II).
We give in Fig. 5 a plot of thz relative abundances of the ion currents determined using each of the two optics systems with. the apparatus, against the pressure_ It can be seen that the ions exhibit different behaviour for the cases mentioned; particularly, in the studies before modification, at pressures above 1.2 torr, the intensities of both solvates, C,H1&lfC2H,Cl and C4H,&1* (CZH5Cl)z, decrease sharply while those of C,H,,Clf and CaHS+ rise with pressure. In addition, some nonintegral mass numbers corresponding to the loss of C2H,CI molecules from C,JYX,~CI’ and their solvates were observed. These data indicate that collision-induced dissociation outside the reaction chamber occurred very eaciently before the modification. The results obtained after the modijkation (Fig. 5) seem to prove that this process is significantly reduced. Materials The ethyl chloride products of BDH Ltd. (English product) and “Frkgata” (pro narcosi; Polish product) were used without further purification. The hydrogen
36 used was Polish product prepurified grade. The C2H,C1 before being introduced into the mass spectrometer was degassed by the freezing and thawing technique at liquid nitrogen temperature.
RE9JL-S General
AND
DISCLJSSION
pressure and remperature
dependences
The observed variations with pressure of the intensities of the main ions in ethyl chloride are shown in Fig. 6 (for experiments performed at 285 and 403 K). As can be seen, the relative intensities of the C2HS+, C2H5+, C,H,Cl’ and C,H,CIHi ions decrease sharply and that of C,H,,Cl’ rises with pressure
m4 6-e-o-e Q
& CIj
1.02
Cl
0.025
qoso
0,075
0,loo
0
as
1.5
:
Fig. 6. Pressure-dependent variations of the relative intensities of the major ions in pure C2H5CI at indicated temperatures; optics I. The points given in graph (b) refer to the same ions as indicated in graph (a).
which is in close agreement Hughes [l ] and by Eeauchamp
with trends previously reported by Tiernan and et al. [2]. The reactions which lead to the formation
of C4H10Cl’ have been documented in earlier studies [I, 21 and were observed here and are summarized in Table 1. As is seen in Fig. 6a, the C,H, + ion is formed in considerable amounts, particularly at pressures above 0.3 torr. The ion-cyclotron study of ethyl chloride demonstrated [2] that this ion is formed in the two reactions: C2H,’
+C,H,Cl
-+ C,H,+ +Cl
(4)
37 TABLE
1
ION--WOLECULE
OBSERVED
REACTIONS
IN
ETHYL
CHLORIDE
Reaction
Reaction ivo.
C2H5C!l+ +C2HsCl
-+ -+ -+ +
C~H~ClH’i(C,H,Cl) C,H,Cl+ i- (CzHSCl, C,H, + f (HCl, Cl) C,Hs + + (2 HCI)
Hz)
tw
hHr (kcai/nzole)a
Re$
-11
1,2
(171
f12 -3-(13) -2O-(-335)
1,2 2 2
(16)
(5)
C2HS + f &H&l
+ CLH&lH+ f (CzHs) -+ CsHs’i(CH4, HCl)
(1s) WI
-14 - 12-(-22)
1, 2 1
&I&+
+ -+ + -+ -+
&H&l + + (C,H,) C2H&.Zl+ i (C,H,) CaHs + -i- (HCI 1 C,H, + + (Cl) C2H, + i (CzHoCI)
(20) (21) (22)
-11 f12 -27-(-42) -lO-(-20) -i-12
1, 2 1, 2 1, 2 2 3
-+ C2H5ClHi -!- (C~HZ) + CzHs + t (&H&l)
(23) (24)
-28 -21
1, 2 1
+ C,H, + -!- (C&Cl) + c,I-X&+ i (&H&I)
(25) (26)
-53
1 1
+ CzHs + i (CHzCIz)
(27)
-21-(-41)
2
4
(98)
+&H&l
C2Hs + f&H&I C&
•,-i-
CH2Cli
GHXI fC2HsCI
C,H,CiH+
+CzH,CI
C,HI,,Cl’+HCl
(44) (6)
1, 2
n The heats of the reactions were calculated on the basis of the following thermochemical data: a F. H. Field and J. L. Frank!in, EZectron hpacr Phenomena and Properries of Gaseozzs Ions, Academic Press, New York, 1957. b Tizermai-Constants of Organic Substances, (in Russian), AN SSSR, Moscow, 1970. c M. H. Karapet’yanc and M. L. Karapet’yanc, FzzndamentaJThermodq-namic Co.-rants of Inorganic and Organic Substances, (in Russian), Khimija, MOSCOW, 1968. d V. I. Vedeneev, L. Gurvich, V. N. Kondrat’yew,.V. A. Medvedev and Ve. L. Frankevich, Bend Energies, Ionization Potentials and Electron Aflnities, (in Russian), AN SSSR, MOSCOW, 1962, p. 139.
and C2H,CI’+C2H,C1
+ C,H,++HCI+CI
OS the other hand, the experiments of Tiernan and Hughes [I ] performed. with tandem mass spectrometer show that no C4HPi is present among the ionic products from the reactions of CzH4+ and C,H,Ci* with C,H,Ci. These investigators found, however, that instead of reaction (4) the following reactionoccurs the
C,H,+ + C&Cl
+ C,H, * + (C2H,Cl)
(6)
The high-pressure studies El] have shown that the CJ&,’ ion is formed in great amount (10 %), but no suggestion concerning its origins has been given by the authors. The previous [l ] as well as the present high-temper-ature experiments* * From the point of view of ion-residence tims the conditions under which the high-pressure experiments of Tieman and Hughes [l] (at field strength of 10.6 V/cm) have been carried out are approximately equivalent to those under which our data, presented in Fig. 6a, were obtained-
38 seem to indicate that the C,H, + ion is essentially unreactive with the C2H,Cl. At high pressures and relatively low temperature, however, this ion disappears, as shown by the results given in Fig. 6b. The figure also shows that as the pressure increases the intensity of C&I,&1 decreases somewhat and at the same time the CtH,’ and C,H,Cl+ ions are produced. It is most probable that these ions are formed as a result of the collisional dissociation of the CSH,&l ions in the regions outside the source by the following reactions*: C,H, ,Cl’
+ M + C,H,
+ + C2H,Cl
+ C2H,Cl’
+C,H,
(7) (8)
The overall trends observed (Fig. 7) for C,H,,Clf, CzHsi and C2H,Cli with increasing temperature are. similar to those found with pressure changes (Fig. 6b) but the behaviour of the C,H, + ion is different (the production of C,H,+ rises with the temperature). The effect of temperature on the observed changes may be explained in terms of an increased internal excitation of the C4H1&1+ ions and as a result of an increased probability of their decomposition (reactions (7) and (S)) msl _ -de and outside the reaction chamber as the temperature is increased. The fact that the collisional decomposition of CSHl,,Cl+ occurs by loss of C,H, and &H,Cl additionally supports the previous investigators’ suggestions [I, 21 that the C,H,,Cli ions are of diethylchloronium type structure C2H,Cl*C,H,. The observed change in intensity of C,H,’ with pressure (Fig. 6) and temperature (Fig. 7) may be the result of the influence of the internal energy of the reaction complexes formed in the collisionsC,H,Cl’-C,H,Cl and&H*+-C2H,Cl on the relative probabilities of the competing reaction channels. However, the relative yields of the ionic products (excepting that of C2H,+) of these reactions are substantiaxy independent of pressure above 0.3 torr (Fig. 6) and no adduct ions resulting from the mentioned collisions have been observed in the previous [I, 23 or the present experiments even at low temperatures (up to 193 K). In view of this, the above explanation dealing with the behaviour of C4HQt with pressure and temperature seems to be unlikely** and we suspect that in high-pressure experiments all the C,Hg f ions do not necessarily originate from reactions (4) and (5). The additional step for production of this ion may be CSH&l+*
+ C4H9+ +HCl
(9)
* The ratio of the intensities of the peaks occurring at nonintegral masses and corresponding to the reactions (7) and (S), to the parent ion intensity, C4HIoC1+, generaliy increased with increasing pressure. ** in contradistinction to reaction (4), reaction (6) is endothermic by about 12 kcal/mole, and it
w&d be difficult to understand the increase of product yield through reaction (6) when the internal energy of the collision complex CzHo+-C2H5Ci, as may be reasonably assumed, decreases with pressure (due to collisional deactivation of the CzH4+ ion).
39 1PC
I.075
0.7: !j
L C Ij
c'j
Q5l
IPSO
3.025
0.251 /
./=
i
CL%@d
/
oio/Q4 270
/--
290
280
300
310
0
Fig. 7. Relative intensities of the C,H,,C!I +, C4Hg+. C2HsCli and C2Hsi as a function of gas temperatuie. Pressure = 0.96 torr, optics 1.
180
210
3,
K
X0
ions in pure CzHsCi
270
Fig. 8. Dependences of the relative intensities of the major ions in pure C2HsCI perature. Pressure = 0.6 torr, optics I.
on the gas tern-
40 Then, the observed pressure dependence is readily explained in terms of the C,H, + ion production by decomposition of excited C4H,,Ci’*. The decomposition rate might decrease as the latter ion is stabilized and, of course, it increases when the temperature rises.
Salvation processes At the high pressures and relatively low temperatures used in the present experiments, the C,HroCl* ion undergoes a stepwise salvation. A typical plot of the fractional intensities of observed ions in pure ethyl chloride at P = 0.6 torr as a function of temperature is shown in Fig. 8. It can be seen that as the temperature decreases the intensity of C,H,,Cl’ decreases accompanied by an increase in the intensities of the solvated ions, C,H,&l~(C2H,C.l),. In pure C&H,Cl, solvates with n up to 3 have been usually observed, and only at temperatures below 193 K has the solvate with 12 = 4 sometimes been seen with a very low yield. Observation of C,H,.Clt(CzH,C1), was complicated by condensation of the ethyl chloride at a temperature of about 190 K. Data obtained for the H,+H,C!l system wer,p similar except for the solvates with n > 2, which were below our detection limit. The intensity changes with pressure and temperature (Figs. 5 and 8) suggest that the observed solvates are formed by reversible reactions, so that we can write C,H,&i’
(C2H,Cl),_
The equilibrium constants,
IL-
1.n =
1 i- CaH,Cl
it C,H,,,Cl+
K, _ , , “, of these reactions
k4HloCI+(C2H5CI), k.,HloCI
+ (C2H5CI),_
1 - pCLHsCl
(&H&l),
(10)
may be expressed by
(11)
are the intensities of the C4H10Cl+ (C,H,Ci), and C,H,<(C2H,C1),,_, ions, produced in the reaction chamber, respectively, while P C2H5C~is the pressure of &H,Cl. Figure 9 shows the plots of K,, r and K,, Z as a function of pressure for pure C,H,C!l and the H,-C2H5CI mixture, and one can see that rather different behaviour of the equilibrium constants is observed for both these systems: in pure C2H,C1 the I&, r value decreases over the whole pressure range investigated, while K,. Z is constant; in the mixture H,-C,H,Cl both I&,, 1 and KISz initially decrease (up to about 0.9 torr) but afterwards remain constant_ This suggests that in the latter system both the 0,l and 1,Zequilibria approach or are in an equilibrium state. Most probably the difference in behaviour of K,, r and I&, z in the C2H,CI and H,+H,CI systems arises from the experimental conditions, which in the former case were more drastic and therefore collisional dissociation outside the reaction chamber was more efkient than in the latter cage. The temperature dependence of the measured equilibrium constants for both systems
where
~C4HloC1+(C2H,CI),
and
~C4H&I’(C2H5CI),-l
41
5.x+*
0.5
PRESSURE
1.0
1.5
2.0 *5-103
, Torr
Fig. 9. Variation of the equilibrium constants Tg = 233 K, optics II; (0, B) pure C2H&I,
KO. I and Kl . z with pressure. T, = 253 K, optics I.
(g,
v ) Hz-CzH5CI,
was investigated. Since the pressure in the reaction chamber varied as the temperature was changed, it was always corrected after every change of temperature in order to maintain its value constant. Duplicate experiments were made on other days and good agreement between the results was found. Values of the equilibrium constant are plotted as a function of inverse temperature in a conventional van ‘t Hoff plot, as for example shown in Gig_ 10. The experiments with the H,C,H,Cl mixture yielded straight line van ‘t HOE plots, which gave an enthalpy value reproducible at two different pressures (0.95 and 1.4 torr); the plot for the 0,1-equilibrium in pure CzH,Cl is linear at higher temperatures and nonlinear at low temperatures (below 227 K). The enthalpy values obtained from the van ‘t Hoff plot for both systems* under study are given in Table 2. We see that a significant difference between the respective AHnc_ In values of these systems exists; values found for pure C2H,C1 presumably are too high due to perturbation of the measured ion intensities by the dissociation outside the reaction chamber. If reaction of desolvation in this region occurs via an intermediate the dissociation of which depends upon the internal ener,gy, then one may expect that this process will be temperature-deptndent 1201.This effect may introduce a systematic error in the temperature-dependence measurements * The ANonS1 value for pure &H&l plot (Fig. 10).
was obtained
from the linear portion
of the van ‘t Hoff
42 1~
IO~T.K
a.5
o
10~
Kn-I,n
!:
atm-i
/ 10
/
:J
"&/
2 °,
t ~0
, 4.5
v 5.0
5.5102
10')T, K
Fig. 10. V a n "t H o f f plots for C4HIoCI+(CzHsCI),,_x--kCzHsC1 ~- C 4 H l o C I + ( C 2 H s C I ) , u n d e r different c o n d i t i o n s . F u l l p o i n t s ; p u r e C2H~CI, PC2H~CI = 0.9 ton., optics I; o p e n p o i n t s : H z - C z H s C1, PtC_,HsCI+H2) -----0.95 t o n ' , optics IL
o f the equilibrium constants, a n d as a c o n s e q u e n c e o f this the e n t h a l p y values o b t a i n e d f r o m van 't H o f f plots m a y be in error. E l e m e n t a r y algebra shows t h a t this effect can raise as well as r e d u c e the true e n t h a l p y values. This d e p e n d s on interference o f various f a c t o r s , such as the relative c o n c e n t r a t i o n s o f p r o d u c t a n d r e a c t a n t ions involved in p a r t i c u l a r equilibria o f the reaction sequence, a n d the m e c h a n i s m a n d kinetics o f the desolvation outside the c h a m b e r . If, f o r example, it is a s s u m e d t h a t dissociation o f solvated ions XMn +, where n = I a n d 2, occurs by p a t h s involving the r a t e - d e t e r m i n i n g steps o f the reactions outside the c h a m b e r (occurring with the same probabilities) XM2 + + M --+ X M + + 2 M
(12)
XM++M
(13)
--+ X + + 2 M
a n d a n equilibrium o f r e a c t i o n XM2_ 1 + M ~ - XMn +
(14)
43 TABLE
2
‘r=RblODmAMIC V.km_rE!j
I-EIE
tC+HJZI K
n--l,
3ooa
(Lzarm-‘)
03
52
I,2
64
J For the H2-C2H5CI
b For pure C2H5Cl.
I&WI-mNS:
+
4.GC300= (kcaljmole)
4W’
4s 21
4EP
(kcallmole)
(eu)
(kcaljmole)
-2.6iO.15 -2.5&0.15
-5.210.3 -4X&0.3
-&7&1_5 - 7.7i1.5
-1313 - 9&3
system; the pressure fraction of C2HsCI = 0.1.
is reached under conditions when Ix >> Ixhi ~2 IxM+ (where Ix, IXM and IXMl are the intensities of the X’, XMf and XMzl ions, respectively), then the measured equilibrium constants for both the O,l and 1,Zequilibria would decrease more rapidly with increasing temperature than in the absence of dissociation outside the chamber. This is for the following case*: we found K,- l,n values to be lower for pure C2H,C1 and the slopes of the van ‘t Hoff plots are steeper than for the case of the Hz--C2H,Ci mixture (Fig. 10) when dissociation outside the chamber seems to be insignificant. On the other hand, if under equilibrium conditions, the ion intensities are as & << kM +z kMz, then, in the case when dissociation outside the chamber is noticeable, the decrease with increasing temperature should be slower for Ke, r and faster for K,, 2 when compared with the case without contribution of the effect mentioned. Of course, the influence of dissociation outside the chamber on the measured enthalpy values may be negligible under certain conditions (when the concentrations of the ions involved in the equilibria do not differ very much), especially for the higher equilibria of the reaction sequence. This seems to explain why satisfactory agreement between thermodynamic quantities has been found in different laboratories [3, 4, 211 for higher equilibria, while discrepancies for the lower ones are very large. We believe that the thermodynamic quantities determined in the present study in the system Hz-CZH5C1, when dissociation outside the chamber seems to be insignificant, are close to the true ones. As can be seen from Table 2, the difference in enthalpy values for 0,l and 1,2-equilibria is insignificant and lies within the limit of the experimental error. * Equation (11) involving the contribution outside the chamber may be rewritten K’,_
%.a =
of dissociation (both collisional and spontaneous) (11’)
k~H~oCi+(C~H~CI),---Ic &HIOCI+
(C2H5CI),_
l +&)kzH,CI
where I, is the intensity of ions arising from dissociation
outside
the chamber.
44 Our data indicate that the association energy of C,H,Cl with the C,H1,CI” ion is relatively low, and it would appear that the interaction between the ligands in the solvate C4H,&1’(C,H,C1), is very weak. It should be noted that the binding strength between C,H,Cl and C,H,,Ci’ is close to that found by Bennett and Field [6] for the association complex CF3 - CEYI,~ (AH” = -4.55 kcal mole-‘) which e IS f ormed at subambient temperatures in the CF4-CH4 -mixture. ACKNOWLEDGEMENi-
The authors are indebted to Mr. K. Grabowski
for technical
assistance.
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