EQUILIBRIA, REACTION RATES, AND YIELDS IN UNIT PROCESSES RAYMOND H. EWELL Purdue University, Lafayette, Ind.
Yields in chemical reactions are determined by two factors : the equilibrium state and the rates of all the reactions concerned. The necessity of clearly defining and separating the effect of these two factors is emphasized. Thermodynamics is able to state definitely what will be the final equilibrium state, but it can say nothing about the rate of attaining equilibrium. The knowledge of reaction rates and catalysis attempts to supply the answer to this question. The thermodynamics of some industrially important unit processes are discussed, including halogenation, hydration, hydrogenation, nitration, and isomerization. Important new sources of useful thermodynamic data and some methods of estimating approximately the heat, entropy, and free energy changes of reactions when exact data are not available are discussed. The importance of reaction rates to yields, and the relation of the activation energy and reaction mechanism to the rate are emphasized, including a discussion of the mechanisms of hydrogenation, addition and substitution chlorination, vapor-phase hydrolysis, and other unit processes.
Thermodynamics gives a definite and indubitable answer to the question: What is the final equilibrium state of any given system? Thermodynamics says clearly what will not happen and what may happen in a chemical reaction, but it can say nothing about what actually will happen. However, this last question is really the important one, and it is the knowledge of reaction rates and catalysis that attempts t o supply the answer.
Thermodynamics and the Equilibrium State Although the principles and methods of thermodynamics have been well known for a long time, it seems worth while to restate a few of the principles and give a few examples. For the prediction of the equilibrium state of any chemical system the following relations are of importance: AF = A H 1 F ” = AH‘
K,,,,,
There
I
MARCH, 1939
= e-AF’/RT
(1) (2) (3)
AF = free energy change of reaction AH = heat content change of reaction (negative of
ordinary heat of reaction) A S = entropy change of the reaction T = absolute temperature
K,,,,,, R e
= equilibrium constant = gas constant = 1.987 cal./ role-degree = base of natural logarithms = 2.718
Thc superscript zeros refer t o changes in the free energy, heat content, or entropy when all reactants and products are in the standard state of unit activity (approximately one atmosphere partial pressure for a gas). Complete discussions of the activity concept are available in any book on chemical thermodynamics (26). The equilibrium constant is a number characteristic of a given reaction at a given temperature. For the general reaction, aA
ICDUSTRIAL chemists and chemical engineers are primarily interested in obtaining the best possible yields from chemical reactions. Yields are determined by two factors: the equilibrium state and the rates of all the reactions concerned. These two factors, equilibrium and reaction rate, are unrelated; in order to have a clear understanding of the possibilities of a reaction, it is necessary to define clearly and to separate sharply the effect of each of these two factors. The term “reactivity” and the phrases “ease of reaction,” “reaction goes readily,” etc., are ambiguous in that they do not distinguish between these two factors. Instead, phrases such as “the equilibrium is in favor of such and such a product” and “the reaction giving such and such a product goes rapidly” should be employed. It is the purpose of this paper to emphasize the necessity of separating the equilibrium and rate factors and to point out the usefulness of and the differences between the three thermal quantities: (a) heat of reaction, (6) free energy of the reaction, and (c) activation energy of the reaction.
- TLS - TAS”
+ bB + . . . . . . = cC + d D + . . . . . .
(4)
the equilibrium constant has the following meaning:
......,. . a: a; . . . . . . . . .
= a: a :
KeQuil,
where a
=
(5)
the activity
The activity of a gas at moderate pressures is approximately equal to its pressure in atmospheres, so that we can write: Kequil.
=
p: pi p & p;l
...... ... .. ...,,..
Equations 3 and 5 (or 6) are quantitative expressions of Le Chatelier’s principle. The sign and magnitude of the free energy change determine the equilibrium constant and the course of a, chemical reaction. (In spite of the general understanding of thermodynamics, it seems that there is still a widespread misapprehension that it is the heat of reaction which determines the
INDUSTRIAL AND ENGMEERING CHEMISTRY
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course of a reaction.) The T A S term constitutes the difference between AH and AF (Equations 1 and 2 ) . The entropy, S,is therefore a very important property of substances, but i t is a rather tenuous concept to grasp. Entropy can be thought of as the quantitative measure of (a) the randomness of distribution and orient'ation and (b) the freedom of motion of the molecules in a system. The more formal, mathematical definitions of entropy can be found in any book on thermodynamics (25). A mixture of two gases has a higher entropy than the two gases separately (at the same pressure) because the distribution of the two species of molecules is more random. Gases have a much higher entropy than liquids or solids because there is a greater freedom of motion of the molecules, and for the same reason liquids have a higher entropy than solids. Entropy of motion is composed of entropy due to the translation, rotation, and vibration of the molecules, which are in this order of decreasing importance in their contribution to the total entropy of a gas. A dissociated gas has a higher entropy than the undissociated gas (e. g., C1, = 2C1) because there is a large increase in the entropy of translation, which overbalances the loss in entropy of rotation and vibration of the chlorine molecule. The entropy of any substance or mixture has a definite numerical value a t a given temperature and pressure, just like any other extensive property such as the weight or volume. For example, the entropy of a mole of methane gas a t 25" C. and 1 atmosphere pressure is 44.46 E. U. (entropy units or calories per mole-degree). Methods of computing entropies are given in any book on thermodynamics (16). AH and AS are usually of the same sign so that they oppose each other in Equations 1 and 2 ; AF will usually be smaller in numerical magnitude than AH and sometimes will even have the opposite sign. This opposing effect of the heat of reaction and the entropy change of the reaction are characteristic of these functions. A large value of -AH means the formation of more bonds during the reaction, while a large value of AS means a n increase in the freedom of motion of the molecules in the system; these are obviously opposing tendencies. I n other words, a chemical system likes t o have as many and as st'rong bonds as possible, and also as much entropy (freedom of motion) as possible, but i t cannot have a maximum of both a t the same time. The free energy change of the reaction, AF, represents a compromise between these two opposing effects. From Equation 1 it is obvious that the entropy term becomes increasingly important, the higher the temperature, which corresponds to the fact that vapor pressure, gaseous dissociation, etc., increase with bemperature. Of the available methods of determining the equilibrium constant of a reaction, the two most applicable to industrially important reactions are (a) direct experimental determination and (b) calculation by means of Equations 2 and 3. Many data are available in the literature for such calculations; notable new and more precise data have appeared during the last few years. Following is a list of new data or new collections of data with which industrial chemists should be familiar: Bichowsky and Rossini (1) give heats of formation, heats of fusion, vaporization, and transition, heats of solution and dilution, critically evaluated for inorganic substances and for organic substances up to two carbon atoms. Heats of combustion of normal paraffins, normal olefins, normal alcohols, and some branched-chain paraffins were recently determined precisely by Rossini at the National Bureau of Standards (69-35).
Heats of hydrogenation of forty-nine compounds containing carbon-carbon or carbon-oxygen double bonds and heats of halogenation of some olefins were recently determined precisely by Kistiakowsky and co-workers (2, 4, 6, 20-65) A series of eight bulletins of the U. S. Bureau of Mines by Kelley gives extensive thermodynamic data on many inorganic substances and a few simple organic substances (11-18). 268
Justi's book (10) includes 43 graphs and 116 tables of thermodynamic data on hydrogen, nitrogen, oxygen, carbon monoxide and dioxide, nitric oxide, nitrous oxide, water, sulfur dioxide, hydrogen sulfide, ammonia, methane, ethylene, acetylene, and air. Parks and Huffman give useful collection of thermodynamic data, even though it is old and sometimes not reliable ( 2 7 ) . Many of the data on entropy, heat capacity, and other thermodynamic data, calculated by means of the new methods of statistical mechanics, are given in recent volumes of the LandoltBornstein tables (24), in a Bureau of Mines bulletin (16), and in the book by Justi ( I O ) . Glasstone (8) gives references to all such data through 1935. New data from this source are constantly appearing-for example, those on hydrocarbons by Pitzer (28), on nitrogen dioxide and nitrogen tetroxide by Giauque and Kemp ( 7 ) , on the chloro- and bromomethanes, formaldehyde, and phosgene by Stevenson and Beach (35), and on cyclopropane by Linnett (66). Not only are the data calculated by these methods usually more accurate than can be determined experimentally, but it is often possible to calculate such data in cases where experimental determination is impossible.
Equilibrium Constants Following are examples of the calculation of equilibrium constants for some industrially important reactions:
+
+
EXAMPLE 1: CHd(g) Clz(g) = CH3Cl(g) HCl(g) 0 20,100 22,060 Qfk 18,290 55.99 44.66 44.46 53.31 Si98 AH& = -23,870 cal. A 9 8 = 2.84 E. U. AF& = -23,870 - 298 X 2.84 = -24,720 cal. where Qf& = standard heat of formation per mole from the elements at 25" c., and SL& = entropy of a mole of the gas at 25" C. and 1 atmosphere pressure. In this reaction AH is large and AS is small so that AH and A F are approximately equal, and the equilibrium is almost exclusively in favor of the chlorinated product at all temperatures. Approximately the same would be true for any other substitution chlorination. The entropy change is nearly always negligibly small when there is no change in the number of molecules, as in this case. When there is a change in the number of molecules, A S is usually between 25 and 40 for each change of 1 in the number of molecules. When no data arp available, an estimate of 30 for each change of 1 in the number of molecules will usually give A S closely enough for many purposes. EXAMPLE 2: NHs(g)
Q.f&
+ CO(g) = HCN(g) + HzO(g)
26,840 47.32 10,740 cal.
-30,700 57,800 48.25 45.17 AH& = ASig80.05 = E. U. A F o E AH" at all temperatures 11,000
46.05
S2"98
This is the same type as example 1, but here AS is actually almost zero, and AF and AH are almost the same.
+
EXAMPLE 3: CSHd(g) Hl(g) = CzH,(g) Qjzb -11,700 0 21,040 Si98 52.75 31.23 54.8 = -29.18 E.U. AH& = -32,740cal. A j & = -32,740 298 X 29.18 = -24,290 Cd.
+
This illustrates a case where AH is large and A S is also large, so that AF will be much smaller numerically than AH, and increasingly so as the temperature is raised. But TAS does not equal AH, and A F become zero, until 850" c.
+
EXAMPLE 4: CzH4(g) HzO(g) = CsHjOH(g) Qf& -11,700 57,800 57,070 %98 52.75 45.17 66.4 AS,",, = -31.5 E. U. AH,",, = -10,970 cal. 10
c.
T ASoa
AFoa
Iiequil.a
-2,370 79.4 -8,600 14.4 -1,580 -9,390 3.43 -786 -10,180 1.00 0 -10,970 0.347 f785 -11,750 0.062 2,340 -13,310 150 0.0152 3,930 -14,900 200 0.0020 7,070 -18.040 300 a Assuming t h a t AHo and A S o do not change with temperature. 0 25 50 75 100
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This reaction illustrates a case where AH is not very large and aS is fairly large, so that AF is negative at low temperatures, but at higher temperatures TAS becomes greater than A H , and AF becomes positive. These data show that good yields of ethylene can be obtained by dehydration of ethyl alcohol only at temperatures well above 75" C., and that at room temperature the proper catalyst should hydrate ethylene t o ethyl alcohol in good yields. This reaction is a fine example of the importance of using accurate data in making thermodynamic calculations. Parks and Huffman (27), using older data, calculated AF& = +4,060 calories; the above calculation, by means of the new precise data of Rossini, gives A F & = -1,580 calories, actually reversing the sign of AF and changing the equilibrium constant by a factor of 10,000. EXAMPLE 5.4: n-CaHlo(8) = iso-CaHlo(g) 687,940 686,310 Qcfs 75.1 70.8 8298 AH;ss = -1,630 cal. A s & = -4.3 E. C. AF& = -340 cal. AF" = 0 at 106" c. EXAMPLE 5B: n-CiH,z ( g ) = neo-C5HlZ(g) 845,270 840,580 &dm 85.0 73.7 s,",, AH& = -4,960 cal. As,",, -11.3 E. U. .AF& = -1,310 cal. A F " = 0 at 141" C.
where Qc,ze8= standard heat of combustion per mole at 25" C. The data on these two hydrocarbon isomerization reactions show some interesting facts: Branched-chain hydrocarbons have higher bond energy, but the straight chains have higher entropy; branched chains are more stable at lower temperatures and straight chains at high temperatures; and the temperature of neutral equilibrium ( A F o = 0) is of the order of 100-150" C. Therefore, straight-chain hydrocarbons in petroleum should isomerize spontaneously to the desirable branched chains at low temperatures. Unfortunately the rate of isomerization at low temperatures would be exceedingly slow, and if the temperature is raised to increase the rate of isomerization, the branched chains are no longer stable. EXAMPLE 6. HNOs(g) f CHa(g) 18,290 &figs 34,400 44.46 ? 87% AH& = -23,710 cal.
=
CH3NOz(g) 18,600
+ HzO(g) 57,800 45.17
? A s & = (estd.)
In studying vapor-phase nitration of methane a t Purdue, it was desired to know the equilibrium constant of the reaction, but the entropies of nitric acid and nitromethane were not known. The entropy change would certainly be small since there is no change in the number of molecules, but a comparison of the two structures shows that A S is probably even less than 1. The heavy atom skeletons of both nitric acid and nitromethane are planar and have nearly the same structure and distribution of mass in the molecule:
These data show that ethylene dichloride is more stable than vinyl chloride plus hydrogen chloride beloJv 270' C, and vice versa above 270 . Therefore, the chlorination of ethylene should lead to addition a t low temperatures and to substitution at high temperatures, which is in accord with experimental evidence.
Two Methods of Estimating Heats of Reaction When thermal data are not available, heats of reaction can sometimes be estimated by the difference in energy of the bonds broken and formed during the reaction. The bond energies in Table I were calculated from the heat of formation data of Bichowsky and Rossini ( I ) . The heat of sublimation of carbon was taken as 125,000 calories per gram atom. These bond energies are only approximate because the energy of a bond depends upon the nature of all the substituents on each of the two atoms forming the bond. For example, the carboncarbon bond in ethane would be considerably stronger than TABLEI. F-F CI-C1
APPROXIMATEBONDEXERGIES (IN KILOCALORIES) 64 58 46
C=C CZC
1-1
36
II-F H-CI
C-C
56
H-Br
Br-Br
72
H-I H-H C-H
104 88 104 67
C-F
C-C1
C-Br C-I 0-H
N-H
50 32
110 84
the one in hexachloroethane, and the energies of isomers often differ by several thousand calories. Nevertheless, useful approximations can be obtained from Table I. The following reactions show fair agreement between estimated and observed values of AH: Reaction
+
++
AH(estd.), kcal.
CZHB Clz = CzHaCl HC1 Br2 = CzHaBrz CzHa CZHB CzHa = n-CaHlo H2 = n-CsHn n-CaHla
++
AH(obsvd.), kcal.
-23 -20
-27 -22 -22 -30
-17
-32
The thermal data a t present are insufficient to determine the energy of bonds between carbon, and oxygen and nitrogen. Heats of reaction can often be estimated in such cases by comparing different reactions in which the same bonds are made and broken. For example, the following reactions, 2CH&HO = CHXHOHCHXHO (aldol condensation) CHacHO CH3N02 = C H ~ C H O H C H ~ N O ~ HCHO CHaCOOH = CHPOHCH2COOH
+
C
96
125 148 I03 87
+
CHI
HCHO
+ CH~CHZCOOC~H~ = 'CHCOOC2H6 cL2 I
OH
N - 0 distance
-
all have the same bonds made and broken: l.26A.
N - 0 distance- 1.21 A. N-C distance -1.46 A.
C=Ol .~ C-H
jbroken
,-,> 0-H
c-0
formed
so that the rotations and vibrations, and therefore the entropies,
of the two molecules will be nearly identical. EXAMPLE 7. CHzC1-CHzCl(g) = CH,=CHCl(g) Qf0m 29,300 -9,000 sibs
+ HCl(g)
22,060 44 66 A H & = 16,240 cal. = 30 E. U. (assumed) AF&8 = 16,240 - 298 X 30 = 7,300 cal. AF" = 0 at 270" C.
MARCH, 1939
?
?
From heats of combustion the first reaction has AH equals -15,000 calories, and it would be safe to assume that the others would be within 3,000 calories of this figure. The equilibrium constant of a reaction is a function of the temperature only; and if the equilibrium constant is known for a given temperature, thermodynamics can predict the final equilibrium state in any phase, solid, liquid, or gaseous, of any mixture of substances a t any pressure provided suffi-
INDUSTRIAL AND ENGINEERING CHEMISTRY
269
substituent group on the benzene ring. From the available data it is impossible to say which is the more stable among any group of ortho-meta-para isomers, but probably the free energy is not overwhelmingly in favor of any one isomer. The observed facts regarding the orientation of a second substituent, as summarized by the well-known Crum Brown and Gibson rule, are probably determined by + relative reaction rates in many cases, rather than by equilibrium. o\ t Another example is the sulfonation of hydrocarbons. Aromatic hydrocarbons sulfonate readily, and aliphatic hydrocarbons with difficulty. Since the free-energy change is fat absolute z e r o ) nearly the same in either case the observed difference in behavior must be due t o a difference in reaction rate. A similar state of affairs is true for the Friedel-Crafts reaction REACTION FIGURE 1. ENERGYPROFILEOF THE PATHOF A CHEMICAL on aromatic and aliphatic hydrocarbons. The yield obtained from any chemical process is the resultant of the relative rates of all the cient data are available. Such calculations are often quite diffireactions occurring. The yield of a desired product may cult and are too involved t o discuss here, but the methods are described in a number of books, such as that of Lewis and be reduced by two types of side reactions: (a) competitive reactions, such as the propylene and chlorine reRandall (65). The importance of such predictions cannot be actions considered above, and (b) successive reactions, which overstated; but when thermodynamics has said what the ultiuse up the desired product. For instance in the reaction mate equilibrium state will be, its work is finished, and we scheme must then consider the rate of attaining the equilibrium statei. e., the rates of chemical reactions.
I
s
Reaction Rates and Activation Energy I n order to show the importance of the rates of reactions, let us consider first the reaction between chlorine and propylene. There are seven possible reactions of which the net result is the reaction of one molecule of propylene with one molecule of chlorine :
the yield of product C will depend on the relative rates of the four reactions indicated. According t o the mass law the rate of a chemical reaction is proportional to the concentrations of the reacting' substances. Thus the rate of a unimolecular reaction, A+B
is given by
Thermodynamics shows that there is a large negative free-energy change for each of these seven reactions, and that the four dichloropropanes are all more stable than the three monochloropropylenes (plus hydrogen chloride) a t low temperatures and vice versa at high temperatures (compare example 7 above). There are not sufficient data to establish exact equilibrium relations among either of the two groups of isomers, but the equilibrium is probably not particularly in favor of any one isomer. At low temperatures, then, the product will be some mixture of the four diReacfanfs Ac fivuted comp/ex PMdUCfJ chloropropanes, and the composition of the mixture will be determined by the relative OF THE REACTION HZ IZ -,2HI FIGURE 2. MECHANISM rates of the four possible reactions. As a matter of fact the product is principally proand the rate of a bimolecular reaction pylene dichloride, showing that its rate of formation is much faster than that of the other dichloropropanes, although A+B+C thermodynamically there may be little to choose between them. Similarly at high temperatures, allyl chloride is the is given by principal product because its rate of formation is the greatest. Another instance where reaction rate is an important and _ _ d(A) _ = _ _ _ - - k(A)(B) (8) dt dt possibly the deciding factor is the introduction of a second
+
dm
270
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VOL. 31, NO. 3
Termolecular reactions-i. e., reactions caused by the simultaneous collision of three atoms or molecules-are very rare and need not be considered. The effect of concentration or pressure on the rate is thus seen to be relatively limited, and the principal factor determining the rate is the magnitude of k , the specific reaction rate constant. The specific reaction rate constant, k , is a function of the temperature as follows:
reaction. The r a t e of some reactions can be accelerated by passing a n appropriate wave length of light into the reaction mixture, thereby increasing the number of energy-rich molecules. Since the yield in a chemical reaction is determined by the - -A relative rates of all the reack = se R T (9) tions involved, the yield may obwhere A = activation energy viously be increased if a catalyst can be found that will s = a constant selectively increase the rate of the reaction giving the desired product. A good example is the homogeneous perThe constant s is sometimes called the “orientation” or oxide catalysis discovered by Kharasch (19). Hydrogen ‘[steric” factor, and it is a function of the probability that bromide a d d s slowly t o a l l y l b r o m i d e , giving the “normal” product, propylene dibromide, but in the presence of certain peroxides, the hydrogen bromide adds ‘(abnormally” t o give trimethylene bromide much more rapidly. Both reactions have a certain velocity in both cases, but in the absence of peroxides, the reaction giving propylene dibromide is much faster, even though it is still comparatively slom. I n the presence of peroxides, this latter reaction is unchanged (or even possibly accelerated a little), while the rate of the competing reaction giving trimethylene bromide is increased enormously and the relative rates are reversed. It is evident that the activation energy is the crux of the reaction rate problem, and this brings u p the question of reaction mechanism, for a particular activation energy is associated with a particular mechanism. Now POTESTIAL ENERGY CONTOUR MAP FOR A c 1 ATOM APFIGURE 3. a reaction will proceed by all possible mechaPROACHIKG AS Ha MOLECULE (9A) nisms, but chiefly by the one with the lowest activation e n e r g y 4 e., the lowest potential energy barrier to surmount-just as a railroad seeks out the two colliding molecules will be properly oriented with respect lowest pass over the mountains. to each other to cause reaction. Constant s is of secondary importance compared to the activation energy, A , in determining the rate of a reaction.
Activation Energy The rragnitude of the activation energy, A , is the principal factor in determining the rate of a reaction. It is the minimum energy which two colliding molecules (in the case of a bimolecular reaction) must have in order to react. Such a collision is called an “activated collision,” and the molecules are called “activated molecules.” As the temperature is raised, more and more pairs of colliding molecules will have this requisite energy, and the rate of reaction will be increased. The relation of the two energy quantities, heat of reaction and activation energy, is shown in Figure 1. It is obvious between Ithe hat activation for an exothermic energy and reaction either there the is heat no connection of reaction or the free energy of the reaction. Figure 1 also shows that the activation energy of an endothermic reaction is equal to the activation energy of the reverse exothermic reaction plus the heat of the reaction. For two colliding molecules to react, they must have a t least enough energy to surmount the potential energy hill whose height is A calories. The effect of a catalyst, either a homogeneous or a heterogeneous one, is to reduce the height of the barrier so that more molecules will have the energy to surmount it a t a given temperature. Usually, if not always, this is accomplished by a change in the mechanism of the MARCH, 1939
0;ta
&
“.‘$;;“
+@
FIGURE 4. MECHAXISMOF C1 Hz + HC1
+
Reoctmts
Producfs
+ HREACTION THE
Many reactions do not proceed by simple unimolecular or bimolecular reactions but by means of a series of uniniolecular and bimolecular reactions, called a “chain” reaction. Free radicals and free atoms in very small concentrations often take part in these reaction chains. As contrasting examples,
INDUSTRIAL AND ENGINEERING CHEMISTRY
271
Mechanism 6
Mechanism A
Reactants
Actiuoted complex
J.
+ Followed
by :
cus + c/,
-
C&Cl
u+
THE
REACTION Clz
+ CH,
while the analogous reaction between hydrogen and chlorine is a chain reaction,
reactants
c1, e 2c1 C1 + Hz +HCI + H
e. g.:
+ Clz +HCI + C1, etc.
__
212
+ HCl
Hz
activated complex +products H-H Iz t\+2HI
+ -4
The activated complex is considered t o have many of the characteristics of an ordinary molecule, except that it is very energy-rich so that it flies to pieces soon after its formation. The activated complex exists in very small concentration in equilibrium with the reactants. The heat of the reaction, reactants
=
activated complex
is the activation energy of the reaction. The energy and structure of the activated complex is thus of great importance. The activated complex in the reaction between Hzand Izis planar and tetratomic, as shown in Figure 2. I n determining the structure of activated complexes, use is
+ Re a c fan f s
CH&l
The most general treatment of reaction rates involves the concept of crossing a potential energy barrier by way of an activated complex (3, 6), thus:
+ Is +2HI
Either reaction will proceed by both mechanisms, but the bimolecular mechanism is faster in the case of iodine and the chain mechanism is faster in the case of chlorine (9). Each elementary reaction in a chain reaction has its own activation energy, A , and the rate of each elementary reaction is determined by this activation energy and by the concentrations of the reactants concerned. Frequently one of the elementary reactions is much slower than the others and is the rate-determining step in the chain, and its activation energy will be approximately the activation energy of the over-all chain reaction.
-
C/,+HC/+C/
Activated Complex Theory
consider the reaction between hydrogen and iodine, which has a simple bimolecular mechanism
H
by:
foffowed
+ Cl
FIGURE 5. Two POSSIBLE MECHANISMS FOR
Hz
0
+
Producfs
FIGURE 6. MECHANISM
OF THE REACTION CzH4 Clz + CHzC1-CHzC1
A c t i v a fed complex
+
Product
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VOL. 31, NO. 3
made of potential energy diagranis (or energy contour maps). Figure 3 illustrates this for the simple reaction between C1 and Hz(one of the elementary reactions in the chain reaction between Clz and H2). The contour map shows the potential energy surfaces about a hydrogen molecule as a chlorine atom approaches. The chlorine atom can approach the hydrogen molecule most easily along the line of the two hydrogens, which is a surprising but important result. The activated complex is then linear and triatomic, in which the hydrogenchlorine and the hydrogen-hydrogen bonds are only partially as strong as their normal single-bond strengths (Figure 4). I n general it may be considered that in the activated complex the bonds of the reactants are partially broken and the bonds of the products partially formed. The activated complex may seem far removed from practical reality, but such detailed studies of the mechanism of reactions provide the only scientific approach to an understanding of the effect of temperature and catalysts on reaction rates and yields. If the search for catalysts is to be anything but a rule-of-thumb approach, the reaction mechanisms must be better understood. The reaction between chlorine and parafins probably proceeds by means of a reaction chain analogous to the hydrogen-chlorine reaction. Taking methane as an example, there are two possibilities:
*
+* +
Mechanism A Clz Clz 2c1 C1 C1 CHI +HC1 CHI CHs + C12+CH3C1 C1, etc. H
+
+ +
Literature Cited Bichowsky and Rossini, "Thermochemistry of Chemical Substances," New York, Reinhold Publishing Corp., 1936. Conn, Kistiakowsky, and Smith, J . Am. Chem. Soc., 60, 2764 (1938).
Daniels, "Chemical Kinetics," Ithaca, Cornel1 Univ. Press, 1938.
Dollirer, Gresham, Kistiakowsky, and Vaughan, J . Am. Chem. SOC.,59, 831 (1937).
Mechanism B 2c1 CHd +CH&l H Clz --f HC1 C1, etc.
+
+
Calculations show that mechanism A would have the lower activation energy and is the more likely possibility, but experimental evidence to date does not rule out mechanism B definitely. If the colliding chlorine atom can approach the methane most easily collinear with a carbon-hydrogen bond, mechanism A will result (Figure 5 ) . But if the chlorine atom can approach the methane most easily perpendicular to the plane of three hydrogens, mechanism B will result (Figure 5 ) . Substitution chlorination of any other hydrocarbon will probably proceed by an analogous mechanism.
(SA) (10)
Dolliver, Kistiakowsky, Gresham, Smith, and Vaughan, Ibid., 60, 440 (1938). Eyring, Chem. Rev., 17, 65 (1935). Giauaue and Kemn J . Chem. Phvs.. 6.. 40 (1938). Glasstone, Ann. R i p t s . Chem. SOC.(Londonj, 32,' 66 (1935). Groggins, "Unit Processes in Organic Chemistry," pp. 159-65, New York, McGraw-Hill Book Co., 1938. Hirschfelder, Eyring, and Topley, J . Chem. Phys., 4 , 170 (1936). Justi, "Spezifische SVlirme,Enthalpie, Entropie, und Dissoziation technischer Gase," Berlin, Julius Springer, 1938. Kelley, U. S. Bur. Mines, BzdZ. 350 (1932). Ibid., 371 (1934). Ibid.. 383 (1935).
(11) (12) (13) (14) Ibid.,384 (1935). (15) I b i d . , 393 (1936). (16) Ibid.. 394 (1936'1.
n
n
+
The mechanism of more complicated reactions can only be surmised. Figure 7 illustrates a possible mechanism of the vaporphase hydrolysis of a halide. A chain reaction involving free radicals seems unlikely for this reaction since i t would involve the fission of either an oxygenhydrogen bond (110,000 calories) or a carbon-chlorine bond (67,000 calories), leading to a very high activation energy in either case. Nitration, sulfonation, amination by ammonia, and alkylation by various alkylating agents may proceed by similar mechanisms.
(19) Kharasch, J . Am. Chem. SOC.,55, 2468 (1933). (20) Kistiakowsky, Romeyn, Ruhoff, Smith, and Vaughan, Ibid., 57, 65 (1935). (21) Kistiakowsky, Ruhoff, Smith, and Vaughan, Ibid., 57, 876 (1935). (22) Ibid., 58, 137 (1936). (23) Ibid., 58, 146 (1936). (24) Landolt-Bornstein, Physikalisch-
Chemischen Tabellen, 5th ed., Erganxungsband 111,Sect. 265B, Berlin, Julius Springer, 1936. (25) Lewis and Randall, "Chemical Thermodynamics," New York, McGraw-Hill Book Co., 1923. (26) Linnett, J. Chem. Phys., 6 , 692
+
t
(1938). (27) Parks and Huffman, "Free Ener-
gies of Some Organic Compounds," A. C. S.Monograph 60, New York, Chemical Catalog Co., 1932. (28) Pitzer, J . Chem. Phys., 5, 473
Reoc f o n t s
A c fiva fed
MECHANISM OF THE FIGURE7 . POSSIBLE CH,OH
+ HCl
The addition of chlorine to a double bond may proceed either by a chain mechanism or by a simple bimolecular mechanism, of which the latter has the lower activation enerm -" ( 3 4 . Figure 6 depicts the mechanism of the reaction between ethylene and chlorine. ~
(1937). (29) Rossini, Ibid., 6, 168 (1938). (30) Rossini, J . Research Natl. Bur. Standards, 13, 21 (1934). (31) Ibid., 13, 189 (1934). 132) Ibid.. 15. 357 (1935). (33) Rossini and Knowlton,'Ibid., 19, 339 (1937). (34) Sherman, Quimby, and Gutherland, J . Chem. Phys., 4, 732 (1936). (35) Stevenson and Beach, Ibid., 6, 2j (1938),
complex Produc f s VAPOR-PHASEREACTION CHIC1 H@+
+
. I
MARCH, 1939
RECEIVED November 2, 1938.
INDtJSTRIAL AND ENGINEERING CHEMISTRY
213
(Center) Pilot-plant arrangement for a stainless-steel resin kettle Courtesy, B l a w - K n o x Company
(Left) S e l e c t i v e catalytic polymerization unit for the production of isooctane Universal Oil products C o m p a n y
Courtesy
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INDUSTRIAL AND ENGINEERING CHEMISTRY
VOL. 31, NO. 3
