I n d . Eng. C h e m . Res. 1987, 26, 1818-1822
1818
Kinetics of the Oxygen-Sulfite Reaction at Waterflood Concentrations: Effect of Catalysts and Seawater Medium Dennis B. Hobson, Peter J. Richardson,+and Peter J. Robinson* Department of Chemistry, Manchester Polytechnic, Manchester
MI
5GD, England
E. Alan Hewitt and Ian Smith Nalfloc Ltd., Northwich, Cheshire SW8 4 D X , England
Kinetic aspects of the sulfite-oxygen reaction have been studied at very low concentrations (ca. 50 and 11 pmol dmT3of SOS2-and 02,respectively). Neither sodium sulfite nor ammonium bisulfite reacts appreciably in the absence of a catalyst. In the presence of cobalt(I1) (19-48 nmol dm-3), the ammonium salt reacts ca. 2.5 times more slowly than the sodium salt. T h e strong variation of rate with p H was a n interesting and unexplained effect. Natural seawater was found to exhibit variable b u t substantial catalytic activity. Studies in synthetic seawater showed that (in order of reactivity) Co(I1) > Cu(I1) >> Ce(1V) > Mn(I1) were active catalysts, while Ni(II), Fe(II), and Fe(II1) were inactive. T h e effect on rate of the high ionic strength of seawater was semiquantitatively consistent with the Brmsted-Bjerrum theory. The reaction of sulfite ion with oxygen is widely used as a basis for deoxygenation treatments, an important application being in the secondary oil recovery waterflood process (Patton, 1977; Ogden, 1983; Mitchell et al.. 1980; Grogarty, 1983). In subsea oil fields, seawater is pumped into peripheral bore holes to drive oil through the porous rocks toward the collection point. Untreated seawater is highly corrosive toward the steel pipes, the main reason being reduction by dissolved oxygen of the cathodic potential which has to be overcome in order to dissolve iron at the anodic sites (Shreir, 1976). Reduction of the oxygen content thus decreases the corrosion rate drastically (Wheeler, 1975; Mitchell and Finch, 1981) and is usually achieved by preliminary "mechanical" deaeration (gas stripping) followed by chemical scavenging, normally using a sulfite formulation. The oxygen concentrations of interest are very low indeed, an acceptable final level being around 0.3 pmol dm-, O2 [7 ppb (parts per lo9) by weight]. Although the sulfite-oxygen reaction has been studied by many workers at concentrations in the range 10-1000 pmol dm-,, the first detailed study at very low concentrations was reported by us (Hobson et al., 1986). The system involved was sodium sulfite with a cobalt(I1) catalyst in distilled water, and the general rate eq 1 was derived. k [S032-]1[ Co2+105 [ 0 2 ] " -d[S03'-]
-
dt
1 + h'/[CO2'1
(1)
This equation correctly described the variation of the order of reaction with respect to cobalt, from 0.5 as reported previously at the higher concentrations (Chen and Barron, 1972; Bengtsson and Bjerle, 1975; Mishra and Srivastava, 1975, 1976) to 1.5 at very low concentrations. The orders with respect to sulfite and oxygen remained constant at 1.5 and 0, respectively, within experimental error. The existence of induction periods was clearly demonstrated, and their dependence on reactant concentrations was quantified. The results were interpreted in terms of a complex reaction mechanism composed of eq 2-9, and steady-state treatment of this scheme lead to the rate eq 10 in the low-concentration regime. 'Present address: Herd and Mundy Ltd., The Edgeley Institute, Stockport SK3 9JH, England.
0888-5885/87/2626-1818$01.50/0
initiation
+ + + + + + -
Co2++ S032- C0S03 CoSO, propagation
O2
02C0S03 SO,2-
[03S-02-CoS03]2- Co2+
02C0S0,
S032-
[O,S-O-COSO,]~termination
(4)
C0S03 + 02C0S03
(5)
S042-+ OCoSO,
=
(!&)( 2
+ Co2+ (6)
[O,S-O-COSO~]~-
SO:-
O2
-
202C0S03
-d[S032-1 dt
(3)
[03S-02-CoS03]2-
[03S-02-C~S03]2- Co2+ OCoSO,
(2)
+ 02C0S03
products
)~~[so32-l~,5[co2+ll,5
(7) (8)
(9) (10)
In the present work, we have used the low-concentration techniques previously developed to investigate related systems of more technological significance. The behavior of ammonium bisulfite rather than sodium sulfite scavenger has been studied, since this is an important chemical used commercially, and it has been suggested that it can be used without any added catalyst (Snavely, 1971; Dunlop, 1973; Scranton, 1979). The effect of using a seawater medium rather than distilled water is reported, and the effect of pH variations is considered. Finally, some alternative catalyst systems have been investigated in synthetic seawater for comparison with the previous results using cobalt(I1).
Experimental Section The apparatus and technique has been described in detail elsewhere (Hobson et al., 1986; Richardson, 1983). In essence, a 2-dm3 reaction flask was charged with the reacting solutions and samples run off at suitable intervals (e.g., every minute) into an acid-quenching solution containing fuschin-formaldehyde reagent by means of which the sulfite concentration was measured. A t pH below 2, the sulfite present is converted almost instantaneously into 0 1987 American Chemical Society
Ind. Eng. Chem. Res., Vol. 26, No. 9, 1987 1819 I
‘
Reaction Timelmin
I
I
6
12
18
I
24 1
I
4 f
1
I
5
IO
15
I
t
20
25
J I
Reaction Timelmin
3
Figure 1. Reaction profiles for 49.3 pmol dm-3 sodium sulfite (filled symbols) or ammonium bisulfite (open symbols) with 10.9 pmol dm-3 oxygen. Reactions in distilled water with 19.3 nmol dm-3 cobalt catalyst ( 0 , O ) ;reactions in Ekofisk seawater (A,A); reaction at pH 9 (m).
bisulfite and sulfurous acid, which are completely unreactive with oxygen (Penkett et al., 1979; Bengtsson and Bjerle, 1975; Templeton et al., 1963; Beilke and Gravenhorst, 1978). All runs were a t 295 K, and each concentration reading was a t least duplicated. As in most of our previous work (Hobson et al., 1986), the reactants were added to the flask in the order oxygen solution (in seawater where appropriate), sulfite solution, catalyst solution. The solutions were prepared as before, using BDH analytical grades of sodium sulfite and the relevant metal sulfates. Ammonium bisulfite was provided as the technical grade solution supplied by Nalfloc Ltd., with adjustment of the pH to 7.7 by addition of sodium hydroxide solution. Synthetic seawater was prepared to the ASTM specification (ASTM, 1980; Richardson, 1983). The basic composition includes sodium, potassium, magnesium, calcium, strontium, chloride, bromide, sulfate, bicarbonate, fluoride, and borate ions. This was used as such in some experiments and in others was augmented by the specified heavy metal salts, namely, barium, manganese(II), copper, zinc, lead(II), and silver nitrates.
Results and Discussion As found in our previous studies, the reaction between sulfite and oxygen in all the present work showed an initial induction period during which little or no reaction occurred, followed by an almost linear decay in sulfite concentration and finally a constant concentration when all the oxygen had reacted. Examples are shown in Figures 1 and 2. 1. Comparison of Reactions with Sodium Sulfite and Ammonium Bisulfite in Distilled Water. Ammonium bisulfite is normally used in North Sea waterflood applications without any added catalyst. However, on-site investigations have shown that the addition of catalyst is sometimes required to initiate the reaction, after which time, using a recirculation system, this addition can be stopped. This suggests that adventitious catalysis is usual, from steel surfaces and/or metals already present in natural waters, and these make the use of additional catalyst unnecessary in some field applications. In order to confirm that a catalyst from some source is needed with ammonium bisulfite, the uncatalyzed reaction in distilled water was investigated first. This gave very similar results to those with sodium sulfite, only a negligible reduction in sulfite concentration taking place (approximately 1 pmol dm-3 taking place in 40 min from an original concentration of 56 @moldm-3). These results
6
9
12
I
Reaction Timelmin
Figure 2. Reaction profiles for sulfite and oxygen (50.0 and 9.4 Kmol dm-3, respectively) in synthetic seawater with heavy metal solution (0) or with 0.05 ppm cobalt. (A), C.05 ppm copper (v),0.2 ppm manganese (o), or 0.2 ppm cerium (0). Table I. Cobalt-Catalyzed Reaction with Ammonium Bisulfite [Co2+]/nmol induction ratea/nmol sulfite dm-3 time/s dm-3 s-l 22.6, 27.1 NH.4 19.3 482, 613 406 63.5 NH4 29.0 NH4 38.6 200, 223 97.2, 97.8 83, 78 150, 158 NH, 48.3 llOb 104b Na 19.3 50b 550b Na 48.3 a
-d[SO$-] /dt.
Averages from previous work, for comparison.
confirm that the apparent reactivity of “uncatalyzed” ammonium bisulfite in natural waters is spurious and due to the presence of catalysts in the water source or the presence of metal. Previous investigators have also concluded that the apparent reactivity of uncatalyzed scavengers is due to the presence of natural catalysts (Miron, 1981; Templeton et al., 1963; Snavely, 1971; Snavely and Blount, 1969). Comparison was next made of the cobalt(I1)-catalyzed reactions of ammonium bisulfite and sodium sulfite in distilled water, using sulfite and oxygen concentrations of 49.3 and 10.9 pmol dm-3, respectively, with cobalt concentrations in the range 19.3-48.3 nmol dm-3. Some typical reaction plots are shown in Figure 1, and the extracted induction times and reaction rates are given in Table I. The results show that the induction periods were longer and the subsequent reaction rates lower (by a factor of approximately 2.5) when ammonium bisulfite was used instead of sodium sulfite, for equivalent cobalt concentrations. These results are consistent with those of previous investigators, who found the rate of the ammonium bisulfite reaction to be one-third (Mishra and Srivastava, 1975) and one-tenth (Matsuura et al., 1969) of the sodium sulfite-oxygen reaction rate (although the conditions used were not strictly comparable). We have proposed (Hobson et al., 1986) that initiation of the reaction is by formation of cobalt/sulfite complexes, but if free ammonia is present, coordination to produce cobalt ammines may also be important. It can be shown from standard dissociation and stability constants (Perrin, 1982; Sillen and Martell, 1964,1971) that the concentration of free (unprotonated) ammonia in these solutions was only 1.4 pmol dm-3, and the ratio of cobalt(I1) ammine complexes to hexaquocobalt(I1) is therefore very small (ca. However, cobalt(II1) ammine complexes have much greater stability constants, and the corresponding ratio for
1820 Ind. Eng. Chem. Res., Vol. 26, No. 9, 1987 Table 11. "Uncatalyzed"Reaction in Natural Seawaters induction rate/nmol medium" sulfite time/s dm-3 s-' A Na 156 23.2 A NH, 150 11.4 E Na 298 59.3 E 342, 426 13.4, 23.6 NH, s1 >1800 NH, s2 NH, 63, 75 80.1, 83.3 D Na >2400 2100 " A = seawater from Amlwych, E = seawater from Ekofisk, S l = synthetic seawater without heavy metals, S2 = synthetic seawater with heavy metals, D = distilled water for comparison.
cobalt(II1) is calculated to be about 16. Thus, the relative concentrations of cobalt(I1) and cobalt(II1) species can be drastically altered, even by these very low ammonia concentrations. The observed reduction of sulfite-oxygen reaction rate on using the ammonium salt can therefore be reasonably attributed to a disturbance of the chainpropagating cycle by selective complexation of the higher valence-state species involved. 2. Scavenging in Natural and Synthetic Seawater Media. When seawater is used for waterflood operations, it is likely that the many dissolved materials present will affect the scavenging reaction. To investigate these effects, sulfite/oxygen reactions were first performed in natural seawater samples. One of these was obtained from the area of the Ekofisk oil field in the North Sea, whilst the other was taken off the Welsh Coast at Amlych. Both sodium sulfite and ammonium bisulfite were used as oxygen scavengers, uncatalyzed reactions being investigated initially with sulfite and oxygen concentrations of 49.3 and 10.9 pmol dm-3, respectively. Typical results are plotted in Figure 1,and the rates and induction periods are given in Table 11. The results in seawater were more erratic than those obtained in distilled water and illustrate clearly that both natural seawaters exhibit catalytic qualities. The variability of results may arise from the release of catalytic and/or inhibiting materials from the plastic containers during storage (a source of previous difficulty with the distilled water experiments). The waters were analyzed by atomic absorption spectroscopy for the presence of transition metals but none of the metals investigated could be detected: however, trace amounts below the detection limits could still be catalytically active. These results indicated that there would always be doubts associated with rate measurements in natural seawaters. Since we wished to investigate the suitability of a range of metal ions as catalysts in seawater, oxygen scavenging was next performed in a synthetic seawater prepared to the ASTM (1980) specification. The complete formulation for this medium included the addition of a stock solution containing barium, manganese, copper, zinc, lead, and silver nitrates. Ammonium bisulfite was used as the scavenger, sulfite and oxygen concentrations being 50 and 9.4 pmol dm-3, respectively. Reactions (Table 11) without the addition of the heavy metal stock solution showed no significant reduction in sulfite concentration up to 30 min. However, when the stock solution of heavy metals was added, considerable catalytic action was observed, all the oxygen being consumed within the first 6 min. The most reactive metal in the synthetic seawater is probably the copper, which is present at 0.16 pmol dm-3; for comparison, 0.84 pmol dm-3 of copper caused complete scavenging in 1.3 min (see section 4). In view of the other metals present
Table 111. Effect of pH on the Induction Time and Rate induction time/s DH rate/nmol dm-3 4.0-7.7 >2400 no reaction 7.8 660, 756 2.81, 3.55 8.0 184, 116 25.3, 23.4 8.2 192, 265 115, 123 8.4 196 217 8.5 222 205 9.0 135 237 9.5 14, 4 421, 371 10.0 21 658 10.5 700
4
8.0
8.5
9.0
9.5
100
10 5
PH
Figure 3. log rate vs. pH for uncatalyzed reactions.
and possible synergistic effects, it is not surprising that the heavy metal stock solution caused a significant acceleration of the reaction. It can be concluded from the results of this section that many natural seawaters probably contain enough natural catalyst for efficient scavenging. Owing to its catalytic action, the heavy metal solution was omitted from the synthetic seawater recipe when investigating the suitability of various metal ions as catalysts in section 4. 3. Effect of pH on the Reaction Rate. The effect of pH was investigated in a series of reactions at constant sodium sulfite and oxygen concentrations (49.3 and 10.9 pmol dm-3, respectively). The reaction was very fast at high pH, and in order to obtain rates in the measureable range, the uncatalyzed reaction was studied in this series. Some typical results are included in Figure 1, and the induction periods and rates are given in Table 111. Figure 3 shows a plot of log rate vs. pH, which is seen to exhibit two separate linear regions. A t the higher pH (from 10.5 down to 8.3), the relevant dissociation constants (Perrin, 1982) show that the sulfite is almost completely in the form of S032-. The slope of ca. 0.3 under these conditions is suggestive of a possible additional reaction route involving the hydroxide ion as a catalyst or a complexing agent. At pH below 8.2, the slope is much greater (ca. 3), and although the dissociation of H2S03is being progressively suppressed in this region, the reduction in the concentration of SO3" can only account for a small part of the change in rate (a log-log plot of rate vs. [SO,2-] gives an "order" of 14!). Other workers have also commented on the high pH dependence of the rate of these reactions, although their results referred to a heterogeneous system and are therefore not directly comparable with ours (Linek
Ind. Eng. Chem. Res., Vol. 26, No. 9, 1987 1821 Table IV. Scavenging Times and Induction Periods Using Various Catalysts
catalyst Co(I1) Co(I1) Cu(I1) Cu(I1) Ce(1V) Mn(I1) Ni(I1) Fe(I1) Fe(II1)
[catalyst]/ wmol ppm dm-3 0.2 3.62 0.05 0.91 0.2 3.38 0.05 0.84 0.2 1.55 0.2 3.91 0.2 3.65 0.2 3.86 0.2 3.86
induction time/s 1600 >1600
rate/nmol dm-3s-l >640 535, 583 >600 345, 278, 297 19.5, 15.3 13.1, 17.1 no reaction no reaction no reaction
scavenging time/ min > Ce(1V) > Mn(I1) >> Ni(II), Fe(II), Fe(II1). At the 0.2 ppm catalyst level, both Co(I1) and Cu(I1) were effective in producing complete scavenging in less than half a minute. The greater catalytic activity of Co(I1) over Cu(I1) was illustrated when the concentration of the catalyst was reduced to 0.05 ppm, the Co(I1) giving complete reaction in 0.7 min compared to about 1.3 min for Cu(I1). All the other catalysts were investigated at the 0.2 ppm concentration only. Ce(1V) and Mn(I1) both gave scavenging times on the order of 20-25 min, while the remaining three catalysts, Ni(II), Fe(II), and Fe(III), failed to promote any detectable reaction. As scavenging times of less than 2 min are normally required in actual waterflood operations, it can be seen that the only suitable catalysts among those studied are cobalt and copper, the latter producing slightly longer reaction times. The selection of possible alternative catalysts for investigation was based partly on a consideration of the type of complexes that would be formed in the mechanism. In the case of cobalt(II), the formation of a complex between the CoSO, species and O2 was proposed as a possible second step in the mechanism, partial oxidation of the catalyst to its higher (+III) oxidation state occurring in the formation of this complex. Other systems were therefore sought, having two valency states between which similar one-electron-transfer processes might be possible. From half-cell potential data (Parsons, 1959; Johnson, 1982; Weast, 1986), the relative stabilities of the +I11 and +I1 states of Mn are similar to those for Co (E‘ = 1.60 and 1.93 V, respectively). It might therefore be expected that Mn(I1) would be an effective catalyst in the sulfite oxidation reaction, and in fact it does act as a catalyst, but not to the same extent as Co(I1). A detailed explanation for the difference would be extremely complex, particularly in view of the presence of ammonia in the systems, but it
seems likely that the lesser effectiveness of Mn is associated with a less favorable balance of the stabilities of various species in the propagating cycle. Similar comments would apply to the Ce(IV)/Ce(III) system (Eo= 1.61 V), which is slightly more effective than Mn(I1) on a molefor-mole basis. The lack of reactivity of Ni(I1) is consistent with the relative inaccessibility of the Ni(II1) state (@ = 4.2 V). The reactivity of Cu(I1) can be attributed to the formation of Cu(1) complexes due to reduction by sulfite. The Cu(II)/Cu(I) potential is heavily dependent on the nature of the specific species involved, but a propagating cycle based on these states certainly appears feasible. The lack of catalytic activity of Fe(I1) and Fe(II1) can be attributed at least in part to their tendency to be hydrolyzed and precipitated out of solution. An increasing turbidity was observed during the reaction when either Fe(I1) or Fe(II1) was employed as catalyst: it was thought that this was due to oxidation of Fe(I1) and precipitation of hydrated oxide species (Flynn, 1984). On the other hand, these results suggest that the adventitious catalysis found in on-site work does not arise from iron, either in the form of solid surfaces or as soluble corrosion products. 5. Effect of Ionic Strength on the Cobalt-Catalyzed Reaction Rate. The reaction rate obtained for the cobalt-catalyzed reaction in synthetic seawater (Table IV) is approximately one-tenth of the rate predicted for the same reactant concentrations in distilled water, using the general rate equation derived by Hobson et al. (1986). This difference can tentatively be attributed to the effect of the high ionic strength of the seawater medium (0.42 mol dm-3) on the activity coefficients of the ions involved. Application of the Brmsted-Bjerrum equation (Clark and Wayne, 1969) to the proposed mechanism (eq 2-9) leads (Richardson, 1983), in the low-concentration region, to eq 11 where k is the rate constant a t ionic strength I (mol log ( k / k , ) = -6.1I1I2
(11)
dm-3) and k , is that a t infinite dilution. This equation correctly predicts a decrease in rate with increasing ionic stength, but the rate ratio at I = 0.42 is calculated as compared with the observed 0.1. This difference is not surprising, since the ionic strength of 0.42 is well outside the range of validity of the limiting law (up to ca. I = 0.04). An approximate application of the extended theory (Robinson and Stokes, 1959), using the Debye-Huckel B term with a typical ionic radius and ignoring the C terms, leads to a rate factor k / h , of approximately 0.01 for I = 0.42, much closer to the observed value. Inclusion of the C term is expected to raise the predicted ratio even further, and the interpretation of the ionic strength effect in terms of ionic activity coefficients certainly seems to be semiquantitatively consistent with the mechanism proposed by Hobson et al. (1986).
Acknowledgment This work was supported by an SERC Studentship (to P J Richardson) and by Nalfloc Ltd. Registry No. Na2S03,7757-83-7; NH4S03H,10192-30-0; Co, 7440-48-4; Cu, 7440-50-8; Ce, 7440-45-1; Mn, 7439-96-5.
Literature Cited ASTM “Standard Specification for Substitute Ocean Water” Report D1141-75, 1980. Bengtsson, S.; Bjerle, I. Chem. Eng. Sci 1975, 30, 1429. Beilke, S.; Gravenhorst, G. Atmos. Enoiron. 1978, 12, 231. Chen, T. I.; Barron, C. H. Ind. Eng. Chem. Fundam. 1972,11(4), 466. Clark, I. D.; Wayne, R. P. “Theory of Elementary Reactions in Solution” In Comprehensiue Chemical Kinetics; Bamford, C. H.,
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Tipper, C. F. H., Eds.; Elsevier: Amsterdam, 1969; Vol. 2, Chapter 4. Dunlop, A. K. Corrosion Inhibitors; Nathen, C. C., Ed.; NACE: Houston, 1973; p 76. Flynn, C. M. Chem. Reu. 1984, 84, 31. Grogarty, W. B. J . Pet. Technol. 1983, 35, 1581, 1767. Hobson, D. B.; Richardson, P. J.; Robinson, P. J.; Hewitt, E. A.; Smith, I. J . Chem. Soc., Faraday Trans. 1 1986,82, 869. Johnson, D. A. Some Thermodynamic Aspects of Inorganic Chemistry, 2nd ed.; Cambridge University Press: Cambridge, 1982; Chapter 4. Linek, V.; Tvrdik, J. Biotechnol. Bioeng. 1971, 13, 353. Matsuura, A,; Harada, J.; Akehata, T.; Shirai, T. J . Chem. Eng. Jpn. 1969, 2(2), 199. Miron, R. L. Mater. Perform. 1981, 20(6), 45. Mishra, G. C.; Srivastava, R. D. Chem. Eng. Sci. 1975, 30, 1387. Mishra, G. C.; Srivastava, R. D. Chem. Eng. Sei. 1976, 31, 969. Mitchell, R. W.; Grist, D. M.; Boyle, M. J. J . Pet. Technol. 1980, 32, 904. Mitchell, R. W.; Finch, E. M. J . Pet. Technol. 1981, 33, 1141. Ogden, P. H. Chemicals i n the Oil Industry; Royal Society of Chemistry: London, 1983. Parsons, R. Handbook of Electrochemical Constants; Butterworths: London, 1959; p 69. Patton, C. C. Oilfield Water Systems; Campbell Petroleum Series;
Campbell Petroleum: New York, 1977. Penkett, S. A,; Jones, B. M. R.; Brice, K. A,; Eggleton, A. E. J. Atmos. Enuiron. 1979, 13, 123. Perrin, D. D. Ionisation Constants of Inorganic Acids and Bases in Aqueous Solution, 2nd ed.; IUPAC Chemical Data Series 29; Pergamon: Oxford, 1982; p 101. Richardson, P. J. PhD Thesis, CNAA (Manchester Polytechnic), 1983. Robinson, R. A,; Stokes, R. H. Electrolyte Solutions, 2nd ed.; Butterworths: London, 1959; Chapter 4. Scranton, D. C., Jr. Mater. Perform. 1979, 18(9), 46. Shreir, L. L. Corrosion;Newnes-Butterworth: London, 1976; Vol. 1. Sillen, L. G.; Martell, A. E. Stability Constants of Metal-Ion Complexes, 2nd ed.; Special Publication 17; Chemical Society: London, 1964; p 151 et seq. Snavely, E. S.,Jr. J . Pet. Technol. 1971, 23 (April), 443. Snavely, E. S.; Blount, F. E. Corros. NACE 1969, 25(10), 397. Templeton, C. C.; Rushing, S.S.; Rodgers, J. C. Mater. Perform. 1963, 2(8), 42.
Weast, R. C., Ed. CRC Handbook of Chemistry and Physics, 67th ed.; CRC: West Palm Beach, 1986; p D-151. Wheeler, D. Pet. Eng. 1975, Nou, 68.
Received for review June 19, 1986 Accepted May 26, 1987
Detailed Gas Chromatography/Mass Spectrometric Structural Determination of Olefin Oligomerization Products Alan L. Chaffee* CSIRO, Division of Energy Chemistry, Menai, NSW 2234, Australia
Kingsley J. Cavell,+Anthony F. Masters,’ and Robert J. Western CSIRO, Division of Materials Science, Clayton, Victoria 3168, Australia
Reaction gas chromatography/mass spectrometry methods have been applied in determining the molecular structure of individual C7 olefins present in a complex mixture of isomers formed by the cooligomerization of C3 and C4 olefins. The catalytic oligomerization of low molecular weight olefins to higher molecular weight products, useful in the production of plasticizers, lubricants, detergents, fuels, etc., has been practiced commercially for many years (Hobson and Pohl, 1973). There are several processes which employ a range of catalytic types and operating conditions (Sittig, 1978; Chauvin et al., 1974; Freitas and Gum, 1979; Bogdanovic, 1979). These processes lead to different product distributions, as a result of the operation of different reaction pathways and their relative importance under particular conditions. For many applications, high product selectivity is required (e.g., linear a-olefins or branched internal olefins); hence, a detailed definition and understanding of these reaction pathways is required. A novel example of propylene oligomerization over a highly specific nickel catalyst has recently been reported (Masters and Cavell, 1985). In this case, the product distribution was relatively simple. By using gas chromatography (GC) and coinjection of standard compounds, it was demonstrated that eight C, olefins were produced as Current address: Department of Chemistry, University of Tasmania, Hobart, Tas. 7001, Australia. Current address: Department of Inorganic Chemistry, University of Sydney, Sydney, NSW 2006, Australia.
*
0888-5885/87/2626-1822$01.50/0
“primary” products and that four more resulted from subsequent double bond isomerization. In this case, since the mechanism specifically excludes highly branched isomers (such as the 2,2-dimethylbutenes), it was possible to obtain all the necessary standard compounds. Unfortunately, for olefins CnHPn,the number of possible isomers increases rapidly with n so that oligomerization and cooligomerization of even simple olefins can produce complex mixtures which vary in both carbon number and isomer distribution. As isomer complexity and carbon number increase, it becomes increasingly difficult or impossible to obtain all of the necessary standard compounds to facilitate the definitive assignment of GC peaks by coinjection. Conventional electron impact gas chromatography/mass spectrometry (EI-GC/MS) is also of limited value since the E1 mass spectra of many isomeric olefins are nearly identical. The determination of reaction pathways necessarily starts with a detailed and exact characterization of individual components in the product mixture. To help overcome the difficulties outlined above, we have employed a technique of postcolumn reaction (hydrogenation) GC/MS recently described in the literature (Chaffee and Liepa, 1985). This procedure is an ideal aid in assigning definitive molecular structures to individual 0 1987 American Chemical Society
