6194
J. Phys. Chem. 1995,99, 6194-6198
Thermochemistry of the Deacon Process Mohamed W. M. Hisham? and Sidney W. Benson" Donald P. and Katherine B. Laker Hydrocarbon Research Institute, Department of Chemistry, University of Southem Califomia, University Park, Las Angeles, Califomia 90089-1661 Received: July 12, 1994; In Final Form: December 1, 1994@
The Deacon process is composed of a number of discrete steps whose thermochemistry (AH, AS) is examined for a large number of main groups and transition metal oxides which may show catalytic behavior. The overall process comprises a catalytic cycle which can be examined in terms of two independent steps: (1) HC1 absorption by a metal oxide to form metal chloride (or oxychloride) plus water and (2) oxidation of chloride by 0 2 to regenerate metal oxide and free C12. It is shown that in the temperature range 150-450 "C the exothermic chlorination step (l), if too exothermic, will be followed by a very slow endothermic oxidation step since AH1 A H 2 = -13.6 kcal, the overall enthalpy change in the process. An examination of the thermochemistry of a large number of main group and transition metals shows that only the latter have values of AHl and A H 2 that are sufficiently small as not to render the entire process too slow. Some experimental observations are reported on the behavior of six transition metal oxides and two main group oxides in steps 1 and 2. Step 1, the chlorination, is fast and exothermic for all metal oxides. Step 2, the oxidation, is very slow for all the metal chlorides except for Cu, below 450 "C. For the two main group metals examined, A1203 and MgO, no evolution of chlorine from the chlorides is observed below 750 "C in the presence of 1 atm of 02(g).
+
Introduction The industrial use of elemental chlorine (C12) in chlorination processes produces an equimolar byproduct stream of HCl according to the general stoichiometry
RH
+ C1, -.RC1+ HC1
This HCl stream is environmentallyundesirable and has a very restricted market. Many metallurgical processes also produce large HC1 byproduct streams. A recently patented process for the direct conversion of CH4 into C2H4, C2H2, and higher hydrocarbons uses Clz as a recyclable reagent.''2 Consequently, there has been growing interest in finding efficient methods for converting HC1 back into Cl2. In 1868, Deacon developed a process by which Cl2 is produced by direct oxidation of HC1 with 0 2 in the presence of CuC12 as a catalyst. The process is described by the following net chemical reaction: HCl(g)
+ 'I4O2(g)= 1/2H20(g)+ '/$12(g) + 6.8 kcal
(1)
Reaction 1 is an exothermic process. AH1"(300)= -6.8 kcal. In the presence of the catalyst it is expected to reach equilibrium at a normal operating temperature of 650-750 K. However, there are a number of practical problems associated with this process. Higher temperatures increase the rate of the reaction, but simultaneously reduce the HC1 equilibrium conversion. Furthermore, at an elevated temperature (2675 K) the problems associated with corrosion become more critical, and the catalyst activity rapidly decreases mainly due to volatilization. Since the early 1900s, various efforts have been made to improve the Deacon process. Several modifications of the catalyst's composition have been suggested, such as addition of less volatile rare earth metal salts with CuC12 and the addition of chlorides or oxides of a number of other metals such as V, Present address: Occidental Chemical Corporation, Technology Center, 2801 Long Rd., Grand Island, NY 14072. Abstract published in Advance ACS Abstracts, March 15, 1995. @
0022-365419512099-6194$09.0010
Be, Mg, Bi, and Sb, as promoters. A number of researchers have proposed the addition of NaCl and KC1, which form double salts with both CuCl and CuCl2 to reduce their volatility. Cr2O3 and V203 have also been shown to be efficient catalysts for the p r o c e ~ s .Numerous ~~~ patents claim improved catalytic activity, and there are presumably efficient industrial processes which have found limited commercial application including the AirCo, Kellog, and Shell However, none of these processes can be characterized as completely successful, since the conversions obtained under normal operating conditions are less than those claimed in the literature, and practical problems such as corrosion, decline in catalytic activity, losses due to volatilization, and higher temperature operation still remain and contribute significantly to the final product cost. The main reason that would explain the limited success of all the Deacon-type processes is the lack of a more detailed understanding of the catalytic reaction mechanism associated with the process. Although a large number of studies have been published on the process i t ~ e l f , ~ -the ' ~ great majority are semiempirical in nature,l3.l4 and in most instances catalytic selectivity and catalytic reactor design have been based on a limited amount of laboratory data and incomplete knowledge of the detailed thermochemistry and kinetics of the process. There seems to be a general agreement in the literature's.'6 that the reaction mechanism for the Deacon process (when CuC12 is used as the catalyst) can be described by the following overall reaction scheme.
+ '/2c12(g) 2CuCl(s) + '/,02(g) = CuO(s) + CUCl,(S) CuO(s) + 2HCl(g) * CuCl,(s) + H,O(g) CuCl,(s)
CuCl(s)
(2) (3) (4)
Although CuCl is the active species for the oxidation step, reaction 2 is thennodynamically unfavorable so that CuCl is never more than a few percent of the total C~C12.l~ Equations 0 1995 American Chemical Society
J. Phys. Chem., Vol. 99, No. 16, 1995 6195
Thermochemistry of the Deacon Process
TABLE 1: Thermochemical Parameters Obtained for Reaction Schemes 1 and 2" HC1 absorption metal
AW
oxidation
ASo
-32.9 -55.6 -68.0 -70.4 -71.4 -33.5 7.5 -11.6 -26.0 -35.1 -43.3 -47.3 -14.8 -22.8 -14.3 -22.8 -15.7 - 14.7 2.0 -8.6 -21.2 -17.8 -34.0 -31.1 -32.6 -29.7 -26.0 -25.7 -23.3 -16.9 -20.0 -21.2 - 17.9 -17.1 -18.9 -9.4 -17.8 -4.1 -6.3 -12.3 -4.7 -16.8 -9.8 -9.3 -5.6 -15.1 -31.3 -1.2 -12.8 -8.7 -14.5 1.4 -6.5 -0.5 -6.9
AGO423 -27.7 -49.7 -62.0 -64.0 -64.8 -27.7 13.4 -5.4 - 19.9 -28.8 -36.6 -40.2 -8.8 -16.5 -8.2 -17.2 -9.1 -8.3 9.3 -1.3 -13.8 - 10.6 -26.7 -23.9 -25.4 -22.5 -18.7 - 18.7 -16.1 -9.7 -12.8 - 14.0 - 10.8 - 10.0 -11.7 -2.1 -10.7 2.9 0.9 -5.1 2.6 -9.6 -3.2 3.2 6.8 -2.2 -12.7 11.6 -0.2 4.1 -1.7 14.2 -0.2 5.1 8.0
-12.4 -13.9 -14.3 -15.1 -15.5 -13.6 -14.0 -14.7 -14.4 -15.0 -15.8 - 16.7 -14.1 -14.9 -14.4 -13.2 -15.5 -15.1 -17.2 -17.2 -17.5 -17.1 - 17.3 -17.1 -17.1 -17.1 - 17.2 - 16.6 -17.1 -17.1 -17.0 -17.0 -16.7 -16.8 -17.0 -17.3 -17.0 -16.6 -17.1 -17.1 -17.3 -17.0 -15.5 -29.6 -29.3 -30.4 -44.0 -30.3 -29.8 -30.3 -30.3 -30.3 - 14.9 -13.2 -35.2
AH" 26.1 48.8 61.2 63.6 64.6 26.7 -14.3 4.8 19.2 28.3 36.5 40.5 8.0 16.0 7.5 16.0 8.9 7.9 -8.8 1.8 14.4 11.0 27.2 24.3 25.8 22.9 19.2 18.9 16.5 10.1 13.2 14.4 11.1 10.3 12.1 2.6 11.0 -2.7 -0.5 5.5 2.1 10.0 3.0 -4.4 -8.0 -1.5 17.7 -14.8 -0.8 4.9 0.9 -15.0 -0.3 -7.3 -6.7
ASo 4.7 6.2 6.6 7.4 7.8 5.9 6.2 6.9 6.6 7.2 8.0 8.9 6.3 7.1 6.6 5.4 7.7 6.5 9.5 9.5 9.8 9.4 9.6 9.4 9.4 9.4 9.5 8.9 9.4 9.4 9.3 9.3 9.0 9.1 9.3 9.6 9.3 8.9 9.4 9.4 9.6 9.4 7.8 14.0 13.7 14.8 28.4 14.7 14.2 14.7 14.7 14.7 7.1 5.4 19.6
AGO673 22.9 44.6 56.8 58.6 59.4 22.7 -18.5 -0.2 14.8 23.5 31.1 34.5 3.8 11.2 3.1 12.3 14.0 3.5 -15.2 -4.6 7.8 4.5 20.7 18.0 19.5 16.6 12.8 12.9 10.2 3.8 6.9 8.1 5.0 4.2 5.8 3.9 4.7 -8.7 -6.8 -0.8 -4.3 3.7 -2.2 -13.8 - 17.2 -11.4 -1.4 -24.7 -10.4 -5.0 -9.0 -24.9 -5.1 -10.9 -19.9
a AWr and AGOr are in kcal/mol of metal, while So are in cal/(mol K). No corrections have been made for CPin estimating AGO423 or AG0673. AH" and ASa are values at 298 K. Values in parentheses are the initial valence state of the metal in the oxide form.
2 and 3 can be combined to give the following net reaction ( 5 ) for the oxidation step: CuCIz(s)
+
1/z02(g)== CuO(s)
+
CI2(g)
AHo = +I50 kcal/mol A S o = 13.4cal/mol
(5)
In 1962, AllenI5 suggested that for a metal chloride to be an efficient catalyst the free energy change associated with reactions such as (4)and ( 5 ) must be on the same order as RT and showed that only a limited number of metals satisfy this criteria. We have recently developedI7 an empirical technique for the quantitative estimation (k2 kcal) of the thermodynamic properties of various types of inorganic solids. As a result, it is now
possible to explore in more detail the thermochemistry of the Deacon process. Unless stated otherwise, thermochemical data used here are taken from NBS tables'* and from ref 17. The values given are in kilocalories/mole of metal atom at 300 K. For each step to be examined the change in heat capacity is generally small so that while corrections can become significant, they are below the uncertainties in the estimates of AfW and So for the metal salts, and so we neglect them. Before proceeding, a caveat is in order. The thermochemistry we shall present is for pure bulk compounds. Catalytic reactions, however, take place at surfaces where the catalytic species is itself dispersed over the surface of a high-area metal
Hisham and Benson
6196 J. Phys. Chem., Vol. 99, No. 16, 1995 oxide such as aluminum, silica, or bentonite. The energies and enthalpies of surface atoms and molecules are significantly more positive than in the bulk phases of pure materials. Entropies of surface species are only slightly increased relative to the bulk phases so that, in general, the free energies are more positive on the surface. In the absence of a tested, general method for correcting for these changes, the use of bulk properties for surface species implies a compensation by the more positive free energies of products for those of the reactants. The error involved in such an assumption is uncertain so that conclusions reached, although presented quantitatively, must be considered rather to be semiquantitative, at best. Conclusions will always have to be compared with observations before assigning them quantitative certainty. Ameliorating somewhat these considerations is the fact that the observed overall equilibrium (eq 1) is entirely in the gas phase, for which the thermochemistry is very accurate. The sums of changes in AHi" and AS? for each step must add up to AH,and AS1, respectively:
i
CI
+
CI
=Cu-OH@) CI
I
AHso = -21 kcal/mol AS,' = -32.8cal/mol
CI
(6)
+
HnO(g)
(7)
A W = 13.4 kcal/mol A W = 33.2cal/mol
I
SCHEME 1: Chloridation Reaction Monovalent oxides
+ HCl(g) - MCl(s) + 1/2H20(g)
+ HCl(g) - M(OH)(Cl)(s) 2M(OH)(Cl)(s) - M20C12(s) + H20(g) overall: MO(s) + HCl(g) - '/,M,OCl(s) + '/,H20(g) MO(s)
Trivalent oxides
M202(OH)(C1) M,O(OH),(Cl), overall:
M203(s)
+ HC1-
-
M,O(OH),(Cl),
+ H20(g) 2MOCl(s) + H20(g)
2MOCl(s)
+ 2HCl(g)
+ 2HCl(g) -M(OH),Cl,(s) M(OH),Cl,(s) - MOC12(s) + H20(g) overall: M02(s) + 2HCl(g) - MOCl,(s) + H20(g) M02(s)
CI
I I =Cu-0-Cu(s)
Cu-0 v H ( s ) + i u ( s )
CI
(7') Eqs 6' and 7' emphasize the changing, nonmetallic composition of the surface over the adsorption (chlorination) part of the catalytic cycle. The role of the metal, important though it is, is assigned a passive role. In a similar vein, eq 8 implies a phase separation on the surface of the oxide and chloride species. Reaction of HCl with Metal Oxides. Table 1 summarizes the thermochemical data for HC1 reactions with metal oxides for monovalent to hexavalent metals. Because of the exothermicity of the reaction, we have selected 423 K (150 "C) for the reaction temperature, and the reaction is given in Scheme 1. =r
Tetravalent oxides
I
HCI(g)
(6')
Divalent oxides
Note also that in the Deacon process the metal content of the surface is constant over the catalytic cycle. This is true of the promoter, additive, and support. Only the chlorine/oxygen/ hydrogen content of the surface changes during the cycle. CuCh as a Catalyst. Before proceeding to the mechanism for different metal chlorides, a few words on the mechanism with CuC12 are in order. The mechanism shown in reactions 2-4, involves three main stages: HCl absorption by CuO given by reaction 4, formation of Cl2 via valence changes (reaction 2), and regeneration of catalyst by reaction with 0 2 (reaction 3). On the basis of our knowledge of the thermochemical properties of intermediate species, and after testing thermochemically the validity of a number of alternative mechanisms,the HCl absorption by CuO can be described by reactions 6-8.'' Some remarks should be made at this point. Since the
CuO(s)
+ HCl(g) == Cl(s) + OH(s) 20H(s) H20(g) + O(s)
O(s)
MO,,,(s)
CAH; = AH,';CAS,' = ASl i
OH(s):
CI
I
Cu-O-Cu(s)
+CuO(s) + CuClp(s)
AH," = -0.2 kcal/mol AS$ = 4.0 cal/mol
(8)
reaction mechanism of the process consists of both exothermic (reaction 6) and endothermic (reaction 2) steps, one expects there to be an optimal temperature range for the process, which has indeed been found to be the case experimentally. Further, the true catalyst for the process is a mixture of the oxide and chloride forms of Cu,the oxide form being necessary for HC1 absorption, while the chlorine form participates in the C12 release step by 0 2 uptake. Therefore, for any metal oxidekhloride to be an active catalyst in the process, initially oxide must be able to absorb HC1 and subsequently the intermediate formed must be able to liberate Cl2 by reaction with 0 2 . The constancy of the Cu species on the surface during the catalytic cycle permits us to present eqs 6 and 7 in an alternate form involving only the nonmetallic species Cl(s), O(s), and
M203(OH),C12(s) overall:
M20,(s)
-
+ 2HCl(g) -
Hexavalent oxides MO,(s)
overall:
+ 2HCl(g)
M03(s)
+ H,O(g) 2M02Cl(s) + H,O(g)
2MO,Cl(s)
-
MO(OH),(Cl),(s)
+ 2HCl(g)
-
MO,Cl,(s)
+ H20(g)
Liberation of Clz with Oxygen (Oxidation Step). Table 1 also summarizes thermochemical data for the species involved
Thermochemistry of the Deacon Process
J. Phys. Chem., Vol. 99,No. 16, 1995 6197
in C12 liberation reactions. The oxidation reaction is given in Scheme 2, and because of endothermicity of these oxidations, a higher reaction temperature of 673 K (400 "C) was selected.
SCHEME 2: Oxidation Reaction Monovalent 2MCl(s)
+ '/,02(g)
Divalent M,OCl,(s)
M,O(s)
+ Cl,(g)
+ '1202(g)
2MO(s)
+ Cl,(g)
+ '/,02(g)
M20,(s)
+ C12(g)
Trivalent 2MOCl(s)
-
TABLE 2: Thermochemical Comparison of Oxychloride Decomposition (AI&) and HCl Absorption ( A H d metal
decomposition reaction
cu
cu2oc12 c u z o Cl2 HgzOClz HgzO Clz FeOCl FeO '/2C12 COOCl c o o + 1/2c12 NiOCl NiO '/2C12 CrOCl CrO '/zC12 UOCl2 uo c12 MnOClz MnOCl l/2C12 MnOC12 MnO '/2C12 VOZCl v02 '/2C12 NbOzCl NbO2 '/2C12 M002C12 Moo2 C12 M002C12-M002C1+ '/2C12
Hg Fe co Ni Cr
u
Mn Mn
v
Nb MO MO
-.
+ + +
-
+ + + +
-28.6 -45.6 -4.1 -6.3 -12.3 -9.8 -9.3 -3 1.3 -31.3 -6.5 -0.5 -6.5 -6.5
catalyst
reaction with HCl
oxidation (Cl2 liberation)
CuO
At 150 "C, HCl was absorbed and H20 liberated.
At 375 "C, C12 was liberated slowly. At 420 "C, Cl2 was liberated rapidly but catalyst vaporized. At 560 "C, C12 was liberated. At 625 OC, C11 was liberated. At 650 "C, C12 was liberated. At 200 OC, dark brown liquid was liberated. At 250 OC catalyst completely vaporized. No C12 was evolved even at 780 "C. No Cl2 was evolved even at 750 "C.
+ 1/202(g)-,M205+ C12(g) At 150 "C, HCl was absorbed and C12 HzO liberated. At 150 "C, HCl was absorbed NiO with H20 liberation. H20 was liberated when HCI COO was absorbed at 150 "C. V205 At 150 "C, HC1 was absorbed and only H20 was liberated. Moo3 At 150 "C, HC1 was absorbed and only H20 was liberated. A1203 At 150 "C, HC1 was absorbed and H20 was liberated. At 150 "C, HCl was absorbed MgO and H20 was liberated. MnO2
Hexavalent
The following comments can be made concerning the tables. (a) Except for three metals (Be, Al, and Zr), all other metals show negative values for AH and AG in HCl absorption steps. In order to favor the Cl2 liberation steps, the value should be between 0 and 10 kcal. This is because the total reaction is 2HC1 '/202 == H20 Cl2, which has "3= -13.6 kcal, hS0298 = -15.4 cal, and AGO298 = -9.0 kcal/(mol K). More negative values for AGO in HC1 absorption steps will result in more positive values for AGO in the Cl2 liberation steps, which will slow this latter process. Except for Mg and Ga, no other main group metals satisfy t h i s criterion. Moreover, the values of AGO mainly depend on the values of AHO of corresponding oxides and chlorides (oxychloride values can be related to values for chlorides and oxides"). It is interesting to note that when the difference between AfHO of chlorides and corresponding oxides is about 20 kcal or less, favorable AGO values are obtained for both stages. Predominantly, transition metal compounds satisfy these conditions. (b) In the HC1 absorption steps, the reaction is exothermic. The heat liberated by the reaction may be sufficient to decompose the intermediate (oxychloride) formed. We compared such decomposition with the heat liberated by the corresponding HCl absorption step, and the results are summarized in Table 2. The decomposition of MnOC12 to MnOCl '/2C12 is only 28 kcal endothermic, whereas HCl absorption is exothermic reaction by 31.3 kcal, so that the combined two steps are exothermic. (c) Because of the exothermic nature of HCl absorption steps, catalytic bed temperatures may get sufficiently hot to vaporize the low-boiling oxyhalide intermediate formed. (d) In the oxidation steps, Cl2 can be produced at various stages. High temperature can cause the decomposition of the intermediate halide, resulting in a lower valence compound. These low-valency compounds can vaporize (e.g., CuCl from CuC12) or can form lower valence oxides. In subsequent reactions these oxides can react with HC1, but liberating Cl2 from this intermediate may be difficult.
+
--.
+
49.6 81.3 32.1 21.1 67.0 50.0 84.0 28.0 49.9 14.0 29.4 30.6 16.4
See reaction in Scheme 1.
Pentavalent
+
----
+ +
AHdkcal"
TABLE 3: Reactions of HCl and 0 2 with Metal Oxide Catalysts
Tetravalent
2MO,Cl(s)
--
rvldfltcal
+
+
Although the thermodynamic approach is very useful, it can only provide constraints on possible reactions. In practice, kinetic parameters and physical properties will dominate. We have carried out a number of experiments with selected catalysts which were chosen on the basis of the values of thermodynamic parameters discussed, and the results are summarized in Table 3. MgO or A1203 did not liberate Cl2 even at 750 "C. As we expect, MnO2 liberated Cl2 in the HCl absorption step. Catalysts such as Moo3 and V2O5 vaporized at low temperatures after treatment with HC1. The experiments were performed by adding HC1 and N2 carrier to an all-glass system containing about 10 g of the metal oxide supported on either a-Al203 or silica substrate. An equimolar amount of NaCl or KCl was also used as cocatalyst. The catalyst was put into a U-tube immersed in an electric furnace, where the temperature was monitored by two chromelalumel thermocouples. An all-glass, magnetically operated glass piston circulated the gases through the system,I9 and a U-trap immersed in dry ice-acetone mixtures permitted us to withdraw H20 as it formed. Total pressure (and thereby total HC1 absorption) was monitored by a Validyne metal diaphragm manometer used as a null instrument. A mercury manometer on the other side was used to measure absolute pressure. In the oxidation experiments, no trap was needed and the 0 2 uptake was monitored by pressure increase. At the end of these oxidation experiments, the Cl2 gas was frozen in a liquid nitrogen trap, the residual 0 2 measured and pumped away, and then the C12 allowed to vaporize and its pressure measured.
Hisham and Benson
6198 J. Phys. Chem., Vol. 99, No. 16, 1995
Summary It has been shown for some time that the exothermic gas phase oxidation of HC1 to Cl2 plus H2O is too slow to proceed in the gas phase below 800 K. Instead it requires a solid metal-based catalyst. The venerable Deacon process, which has used such a catalyst, is known to proceed in two stages, a chlorination step in which a metallic oxide, spontaneously and exothermically absorbs HCI to form a chloride plus water and an endothermic oxidation step in which the metallic chloride is restored to its initial oxide state by reaction with 0 2 at higher temperature. Examination of the thermochemistry of these two stages for metals of the periodic tables shows that main group 1 and 2 metals generally have AHoxthat are too exothermic to permit use in a cycle. Only transition metals and very few rare earths and some actinides have AHoxand AHClwhich are small enough to fit the kinetic requirements. Qualitative observations are reported on the kinetic behavior of some of the first-row transition metal oxides and on MgO and A1203. CuO, which has been the basis of most Deacon oxychlorination catalysts, was the only one of these to fit the requirements of a complete cycle below 700 K.
References and Notes (1) Benson, S. W. U. S. Patent 4,199,533, 1980. (2) Weissman, M.; Benson, S. W. Int. J . Chem. Kinet. 1984, 16, 307.
Allen, J. A.; Clark, A. J. Rev. Pure Appl. Chem. 1971, 21, 145. Engel, W. F.;Waale, M. J.; Muller, S. Chem. Ind. (London),1976, Dow Chemical Co. U.S. Patent 2,577,808. Shell Development Co. U S . Patent 2,542,961. Air Products Co. U S . Patent 2,204,172,2,204,733;2,271,056. van Dijk, C. P.; Schreiner, W. C. Chem. Eng. Prog. 1973,69, 57. Johnstone, H. F. Chem. Eng. Prog. 1948, 44, 657. Oblad, A. G. Ind. Eng. Chem. 1929, 7, 23. Amold, C. W.; Kobe, K. A. Chem. Eng. Prog. 1952, 48, 293. Ruthven, D. M.; Kenny, C. N. Chem. Eng. Sci. 1968,23, 981. Furusaki, S. A.1.Ch.E. J . 1973, 19, 1009. Kenny, C. N. Cutul. Rev. Sci. Eng. 1975, 11, 197. Allen, J. A. J. Appl. Chem. (London) 1962, 12, 406. Mallikarjunan, M. M.; Zahed Hussain, S . J. Sci. Ind. Res. 1981, (16) 42, 209. (17) Hisham, M. W. M.; Benson, S. W. The Estimation of the Enthalpies of Formation of Solid Salts. In From Atoms to Polymers; Eds.; VCH: New York, 1989; Chapter 10. (18) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schum, R. H.; Halow, I.; Bailey, S. M.; Chumey, K. L.; Nutall, R. L. J. Phys. Ref. Datu 1982,l I , Suppl. No. 2. (19) Hisham, M. W. M.; Benson, S. W. Rev. Sci. Instrum. 1989,60,7. JP94 174 1N
