D) Exchange Catalysis in Alkanes - ACS

Review pubs.acs.org/acscatalysis

Cite This: ACS Catal. 2018, 8, 2296−2312

Hydrogen/Deuterium (H/D) Exchange Catalysis in Alkanes Aaron Sattler* Corporate Strategic Research, ExxonMobil Research & Engineering Company, 1545 Route 22 East, Annandale, New Jersey 08801, United States ABSTRACT: The catalytic exchange of hydrogen and deuterium (H/D exchange) in light alkanes has been studied for almost a century. While alkanes and their C−H bonds are relatively inert, H/D exchange studies have shown that a large number of materials can catalytically activate these bonds. These studies helped elucidate the mechanisms by which alkane C−H bonds interact with catalytic materials. This Review serves to highlight this area of research, focusing on two main classes of heterogeneous materials, metals and metal oxides, and their trends in reactivities and selectivities are described in detail. Furthermore, the ability of these materials to carry out C−H bond activation and H/D exchange catalysis is compared with that of molecular organometallic complexes, and the mechanistic relationships and similarities in these processes are proposed. KEYWORDS: H/D exchange, mechanism, isotopes, deuterium, C−H activation, alkanes

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INTRODUCTION The use and value of isotopes, especially deuterium (D, heavy hydrogen 2H), to study and gain insight about the mechanisms of chemical systems through kinetics and labeling studies is well-known and has been studied extensively.1−4 Over 80 years ago, research concerning the exchange of hydrogen and deuterium, which is replacing a H atom with a D atom or a D atom with a H atom, now commonly called H/D exchange, in light alkanes began, demonstrating that transition metals were capable of catalytically exchanging their relatively inert C−H and C−D bonds.5−8 This work started a new field of research, and immense investigation ensued on catalytic H/D exchange in light alkanes. Many transition metal surfaces were found to be capable of this transformation, but did so with different efficacies and selectivities. This area has been the subject of several reviews.9−15 Other systems were also discovered that could carry out H/D exchange, including metal oxides (especially dehydrated aluminas, detailed below), superacids,16,17 and molecular organometallic complexes.18−28 This research has implications for hydrocarbon conversion to more useful products and is therefore of interest to the catalysis community. In addition to fundamental studies of catalysis, other research areas have exploited H/D exchange, and some examples are labeling studies for polymer characterization by neutron scattering techniques,29−34 fundamental studies on equilibrium isotope effects in organometallic compounds,35−44 kinetic and equilibrium studies on the geochemical isotopic distribution of natural occurring hydrocarbons,45−53 and applications of H/D exchange for small molecule synthesis and pharmaceuticals.54−58 This article serves to review the literature on H/D exchange in alkanes, starting from its initial discovery in 1935. A brief background on isotope effects is followed by the H/D exchange studies and is broken down into two classes: metal catalysts and metal oxide catalysts. © XXXX American Chemical Society

Additionally, heterogeneous H/D exchange catalysis is compared with organometallic chemistry, indicating the correlations and mechanistic similarities in these systems.

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ISOTOPE EFFECTS Origins of Isotope Effects. Although this Review does not focus on isotope effects, it is useful to describe and understand why changing isotopes affect chemical systems, as many of the literature reports undoubtedly encounter these circumstances. Isotope effects arise from the differing vibrational energies of isotopologues (e.g., H−Cl and D−Cl), which is expressed in eqs 1 and 2. Isotopologues are molecules that differ only by their isotopic composition (e.g., CH4 and CH3D). Equation 1 is the harmonic approximation for a molecular vibration (where Ev = energy of the vibration of vibrational state v; v = vibrational quantum number: 0, 1, 2···; k = force constant; μ = reduced mass) and eq 2 defines the reduced mass (μ). ⎛ 1⎞ 1 Ev = h⎜v + ⎟ ⎝ 2 ⎠ 2π

μ=

k μ

m1m2 m1 + m2

(1)

(2)

Thus, by changing from H to D (atomic masses of ∼1.008 and ∼2.014, respectively), the reduced mass increases thereby decreasing the vibrational energy (Ev). The force constant (k) is nearly constant for isotopologues. The difference in vibrational energy is commonly observed in vibrational (IR) spectroscopy; e.g., a C−H stretch is ∼3000 cm−1, while a C−D stretch is Received: December 7, 2017 Revised: January 15, 2018

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DOI: 10.1021/acscatal.7b04201 ACS Catal. 2018, 8, 2296−2312

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ACS Catalysis ∼2200 cm−1. There are many different vibrations, stretches and bends, symmetric and asymmetric, etc., but the reader can just focus on one as an example. The zero-point energy (ZPE) is the energy of the bond when v = 0, and is nonzero because of 1 the v + 2 term (eq 1). As such, the ZPEs of isotopologues are different, thereby making their free energies different. This is exemplified in a Morse potential energy diagram shown in Figure 1; a molecule with deuterium has a lower free energy and therefore a greater bond dissociation energy (BDE) than a molecule with hydrogen.

should be noted that in some reactions the energies of the transition states become quite important, because the activation energy is the difference in energy of the transition and ground states. Thus, a greater difference in the transition state energy for X−D vs X−H as compared with the ground state energy difference leads to a KIE in which kH < kD (kH/kD < 1), which is known as an “inverse” KIE.60,61 Common examples where inverse isotope effects are observed are in organic reactions where hybridization changes occur.62 It is worth noting that the observations of inverse isotope effects (and KIEs in general) are often a combination of several KIEs (and equilibrium isotope effects, see below); the above information focuses on understanding KIEs for a single elementary step. Equilibrium Isotope Effects (EIEs). Equilibrium isotope effects (EIEs), which are variations in equilibrium constants (K) due to isotopes (e.g., KH/KD), also exist because isotopologues, or isotopomers, have different free energies. Isotopomers are molecules that have the same isotopic composition but differ in isotope position (e.g., CH3−CHD-CH3 and CH2D-CH2−CH3). Unlike kinetic isotope effects (KIEs) which are typically normal (kH > kD) as described above, EIEs can either be normal (i.e., KH > KD) or inverse (i.e., KH < KD) depending on the assignment of reactants vs products. In fact, the reverse of a chemical reaction has the reciprocal EIE in the forward direction. For example, consider the EIEs shown in Figure 3, which express the

Figure 1. Morse potential energy diagram of a diatomic X−H/D, highlighting the difference in bond dissociation energies between X−H and X−D.

Kinetic Isotope Effects (KIEs). Kinetic isotope effects (KIEs), which are variations in rate constants (k) due to isotopes (e.g., kH/kD), arise from the difference in zero-point energies. In the ground state, X−D is lower in energy than X−H, and thus KIEs are typically “normal”, meaning that the rate constant of a reaction that involves breaking a bond with H (kH) is greater than that with D (kD) (i.e., kH > kD; kH/kD > 1) because the activation energy Ea(D) is greater than Ea(H), making kH > kD as shown in Figure 2. The example shown ignores the transition state energies to simplify the effect;59 it

Figure 3. Example of an equilibrium isotope effect (EIE) for the preference of D to reside on the center or terminal position of propane.

equilibria for the reaction exchanging H/D on the center and terminal positions of propane (the experimental value of which is currently not known). It can also be measured by determining two separate equilibrium isotope effects, Kc and Kt , and taking the ratio of the two (Figure 3). In this example, if Kc/Kt > 1, that indicates that D prefers to be in the center position over the terminal, and if Kc/Kt < 1, it indicates the terminal position is favored. The origins of EIEs are complicated and have many contributing factors. There have been many theoretical, but few experimental, studies to try to better understand EIEs in alkanes.63−70 One reason why limited studies exist in alkane systems may be (i) EIEs are often close to unity, making their measurement difficult, and (ii) finding appropriate conditions and catalysts to equilibrate these systems is often nontrivial. The normal or inverse nature can sometimes be predicted based on the fact that D prefers to reside in the higher frequency vibrational oscillator (i.e., the higher vibrational energy) compared with H, as depicted in Figure 4, but this is certainly not a rule. In this regard, alkane EIEs can be more complicated because of similarity of the isotopologues and isotopomers (i.e., all C−H/D bonds), such that simple approximations are often insufficient and experimental or theoretical data is needed.

Figure 2. Energy diagram depicting the origin of a normal kinetic isotope effect (KIE) for a reaction cleaving an X−(H/D) bond. The activation energy (Ea) for X−D is greater than X−H, such that kH > kD (i.e., normal KIE, kH/kD > 1). 2297


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Table 1. Previously Reported Apparent Activation Energies (Ea) for H/D Exchange Catalysis in Alkanes on Metal Surfaces catalyst Ni/kieselguhr Ni/kieselguhr Ni film Ni film Pd film Pd film Pt film Pt film Rh film Rh film W film W film W film Mo film Ta film Zr film Cr film V film Ni film Ni film Pt film Pd film Rh film W film Rh film Ni film Pd film Pd pumice Pt pumice Rh pumice Ir pumice Pt W film Ni film Pt Pt Pd film W film Pd film Rh film Pt Pt Pd film Pd film Pt pumice Pd film Pd film Rh film Pt film W film Ni film

Figure 4. Energy diagram depicting the origin of an equilibrium isotope effect (EIE) for a reaction exchanging an X−H/D with Y to give Y−H/D and X. The larger vibrational frequency (right) has a larger energy difference between H and D (ΔE(Y−H*)) compared to ΔE(X−H*) (left), indicating KD > KH.

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H/D EXCHANGE ON METAL CATALYSTS Initial Studies. Research on H/D exchange in alkanes first started in 1935 by Morikawa, Benedict, and Taylor,5 when they demonstrated the production of deuterated methanes from CH4 and D2 by using a reduced nickel-kieselguhr catalyst at temperatures between 184−305 °C.71 They showed by an Arrhenius type analysis that the apparent activation energy (Ea) for the exchange between CH4 and D2 was 28 kcal mol−1 (for conversion to kJ mol−1, 1 kcal mol−1 = 4.184 kJ mol−1), while H/D exchange was found to be only 19 kcal mol−1 for CH4 and CD4.72 The experimental Ea values are tabulated in Table 1. The increase in Ea when using D2 has been shown to be an effect due to competitive adsorption of D2 on the catalyst surface, thereby limiting sites capable of carrying out the C−H/D exchange. Another explanation is that D2 cleavage is rate determining, but this has been shown to not be the case, as H2/D2 exchange is rapid12 in comparison with C−H/D exchange. Morikawa, Benedict, and Taylor also demonstrated that Nikieselguhr could catalyze H/D exchange in ethane73 and propane.74 H/D exchange occurred in ethane at temperatures above 110 °C, while exchange in propane occurred at temperatures above 80 °C. At higher temperatures, cracking (hydrogenolysis) was observed; ethane methanation occurred at 150 °C and cracking of propane occurred at 138 °C (their product analysis indicated that the initial products were ethane and methane). It is readily apparent from these studies that over a Ni catalyst, the rate of exchange follows the order C3H8 > C2H6 > CH4, based on the temperatures required to observe catalysis (80 °C, 110 °C, and 184 °C, respectively). Furthermore, at elevated temperatures, other reactions occur such that H/D exchange is not the only reaction that affects product distributions. Farkas and Farkas 75 also worked in this area and demonstrated H/D exchange could be catalyzed by Pt for a variety of organic molecules.76−78 They showed that Pt was proficient for H/D exchange in alkanes like propane and butane,79 hexane and cyclohexane.80,81 Butane exchange was 4 times faster than propane exchange, and propane exchange was 36 times faster than ethane. Remarkably, they found very mild conditions (temperatures as low as 26 °C) could catalyze H/D

alkanea,f

temp range (°C)

Ea (kcal mol−1)

ref

CH4 (CD4) CH4 CH4b CH4c CH4b CH4c CH4b CH4c CH4b CH4c CH4b CH4c C2H6 C2H6 C2H6 C2H6 C2H6 C2H6 C2H6 C2H6 C2H6 C2H6 C2H6 C3H8 C3H8 C3H8d C3H8 C3H8 C3H8 C3H8 C3H8 C3H8 i-C4H10 i-C4H10e n-C4H10 n-C4H10 n-C5H12 neo-C5H12 neo-C5H12 neo-C5H12 n-C6H14 n-C6H14 n-C6H14 n-C6D14 (H2) c-C3H6 c-C5H10 c-C6H12 c-C6H12 c-C6H12 c-C6H12 c-C6H12

138−218 138−218 206−255 206−255 243−308 243−308 159−275 159−275 138−217 138−217 92−174 92−174 −80 to −29 −50 to 0 −44 to 0 158−192 149−215 102−160 162−195 0−75 134−192 145−207 0−70 −82 to −24 −25 to −16 −47 to 0 130−200 100−200 100−250 50−200 100−200 20−40 −80 to −27 −47 to 0 26−41 77−95 60−130 0−45 112−174 0−76 55−124 31−55 60−140 60−140 −18 to 52 0−37 18−82 −48 to 0 0−31 −69 to −48 −35 to 0

19 28 23.8 29.4 22.0 34.2 20.8 24.5 20.0 24.4 8.9 11.3 8.2 7.0 7.8 15.4 13.9 20.7 18.0 6.5 12.5 21.4 11.7 9.0 13.3 10.4 23.7 17.2 18.2 17.3 17.7 25 7.9 9.0 26.3 11.0 14.5 11.2 33 15.5 9 17 19.4 21.6 8.0 14.2 13.0 10.4 12.0 11 10.8

72 72 87 87 87 87 87 87 87 87 87 87 102 102 102 102 102 102 102 102 102 102 102 108 108 108 121 112 114 113 113 82 108 108 79 79 121 117 117 117 81 81 116 116 134 128 128 128 128 128 128

a

H/D exchange was with D2 unless otherwise noted in parentheses. Rates determined for production of CH3D (i.e., monodeuteration). c Rates determined for production of CD4 (i.e., perdeuteration). d Exchange of secondary hydrogens (CH2) only. eExchange of tertiary hydrogen (CH) only. fAbbreviations for alkanes are n- normal, i- iso, ccyclic. b

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because a stepwise process would be expected to give largely monodeuteration at low conversion. Stepwise is defined here by the gas-phase products observed (not the intermediates that may be on the catalyst surface); more detail about the mechanisms that are expected to produce gas-phase product distributions from stepwise or multiple exchange pathways are described below. Figure 7 shows the simulated distribution of methane isotopologues expected from stepwise H/D exchange of CH4 (Figure 6). Several assumptions in this simulation should be noted which are (i) the back (reverse) reactions under these conditions are negligible, which is reasonable when a large excess of D is present, such that there is near-zero H concentration, and (ii) the four rate constants (per hydrogen atom) are equal (k1 in Figure 6) indicating negligible secondary KIEs (i.e., the rates of deuteration of CH4 and CH3D are equal on a per hydrogen basis). It is clear from Figure 7 that at low conversion (e.g., 10% at time = ∼ 0.12) in a stepwise process, the major product is CH3D (>90%), with negligible CD4 ( Rh > Pt > Ni > Pd.87 The high activity of W was attributed to the low activation energy of methane adsorption, as methane was readily adsorbed even at 0 °C.87 It is important to note that upon heating the methane treated W surface to 70 °C, the pressure increased more than expected for simple methane desorption, indicating unsaturated carbonaceous species were formed on the surface with liberation of H2. Arrhenius analysis was carried out to find the apparent activation energies (Ea) of these processes (Table 1).94 It was shown that two principle mechanisms occur giving disparate distributions of deuterated methanes; one gave CH3D as the major product, while the other gave CD4, consistent with the results on Co.85,95−97 At this point, it is instructive to identify the differences in these two proposed mechanisms and why they affect the product distributions. Metal surfaces appear to carry out H/D exchange in alkanes by activating the C−H/D bonds to give alkyl hydride intermediates on the metal surface (Figure 8, right). In the literature, this is often referred to as adsorption or dissociative adsorption (also described as forming alkyl radicals which should not be confused with gas-phase radicals) of the alkane, and the microscopic reverse is called desorption. In organometallic chemistry, these reversible transformations are analogous to oxidative addition (adsorption) and reductive elimination (desorption). As such, adsorption is a chemical reaction (chemisorption) in this case, occurring on the metal surface, and not “physisorption” for which no substrate bonds are fully broken; physisorption likely precedes chemisorption. This is specifically noted because the nomenclature in the literature can be confusing, especially since it is almost a century old. On metal catalysts the experimental results (e.g., time scales for reactions to occur or temperatures required to observe exchange) show that activities of various alkanes trend with the bond dissociation energy (BDE) of the C−H bonds (see Table 2). These results indicate that an initial step involves

exchange in hexane and cyclohexane. As in the case of Ni catalysis, catalysis on Pt followed similar trends, namely, the rate of exchange follows the order C4H10 > C3H8 > C2H6 > CH4. It will become apparent that this is a common feature of H/D exchange catalysis on metals. Trends and Selectivities. About a decade later in 1951, an unprecedented study by Taylor and co-workers used mass spectrometry analysis to assess the regiochemical distribution of deuterium in the H/D exchange of propane on a Pt catalyst.82 Their studies indicated for the first time that the rate of exchange was faster at the center (secondary) position of propane than at the terminal (primary) position. They revealed that there are two apparent trends that affected the observed rates: the size of the hydrocarbon and the position on hydrocarbons containing chemically inequivalent hydrogen atoms. These relative rates follow the same general trends in homolytic bond dissociation energies (BDEs). The BDEs of various alkanes are tabulated in Table 2 and trends in BDE can be generalized by assessing radical stability as depicted in Figure 5. Table 2. Bond Dissociation Energies (BDEs) of C−H Bonds in Alkanesa alkane methane ethane propane propane n-butane n-butane i-butane i-butane neo-pentane cyclopropane cyclobutane cyclopentane cyclohexane

position

bond dissociation energy (kcal mol−1)

− − CH3 CH2 CH3 CH2 CH3 CH − − − − −

105.0 100.5 100.9 98.1 100.7 98.3 100.2 95.7 100.3 106.3 (109.0)b 97.8 (99.9)b 95.6 (96.9)b 99.5 (100.0)b

a From ref 83. bThe BDEs for cyclic alkanes have also been determined by thermodynamic cycles; see ref 84.

Figure 5. Relative radical stabilities for hydrocarbons (R is an alkyl group).

Methane and Mechanistic Implications. In the same year (1951), a new metal catalyst, a cobalt-thoria Fischer− Tropsch catalyst was studied and found to promote H/D exchange in methane, ethane, propane, and butanes at 183 °C.85 Most interestingly, a detailed analysis of the gas-phase product distribution by mass spectrometry demonstrated that multiple H/D exchange and perdeuteration (full deuteration) occurred, even at low conversion. This indicated that a simple stepwise exchange process (Figure 6) was not occurring

Figure 6. Stepwise exchange of CH4 with deuterium (CH4 + D → CH3D + H, etc.). 2299


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Figure 7. Simulated distribution of methane isotopologues from stepwise CH4/D exchange. d-Methanes plotted to show that the sum of the products is equal to [CD4]final − [CH4]t. The x-axis, time, has arbitrary units.

(e.g., in propane: primary C−H BDE: 100.9 kcal mol−1, secondary C−H BDE: 98.1 kcal mol−1), and the observed rate can be a function of both bond cleavages. It is important to note here that on the metal surfaces studied, the exchange of H and D on the surface of a metal is rapid compared to the activation of the C−H/D bond. This has been demonstrated by studies measuring the rates of H2/D2 exchange to give mixtures of H2, D2, and HD.12 Thus, after the formation of an alkyl hydride on a metal surface, the mobile hydrides can more rapidly exchange metal sites with other Hs or Ds (H/D scrambling), and if a desorption event occurs, H/D exchange has occurred. If only these three events occur (i.e., adsorption, H/D scrambling on the metal surface, desorption), this describes a stepwise process (shown above in Figure 6 and Figure 7), and at low conversions, principally CH3D would be expected. This implies that other elementary

Figure 8. Dissociative adsorption of H2 (left) or a C−H bond (right) over a metal (top) or two metals (bottom).

the cleavage of a C−H/D bond, and that this step is rate determining for H/D exchange. Methane is the slowest alkane to undergo H/D exchange because it has the highest BDE (105.0 kcal mol−1), which is why it requires higher reaction temperatures, compared with other alkanes that have lower BDEs. Furthermore, alkanes like propane that have chemically inequivalent hydrogens, have multiple rates of exchange due to the difference in BDEs for the different types of C−H bonds

Figure 9. Energy diagram and mechanism for stepwise exchange between CH4 and excess D2 on a metal catalyst, which is expected to give CH3D as major product at low conversions. 2300


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Figure 10. Energy diagram and mechanism for multiple exchange between CH4 and excess D2 on a metal catalyst, which is expected to give CD4 as major product at low conversions.

Figure 11. Coordination modes of hydrocarbons on metal surfaces.

Table 3. Temperatures Required to Achieve Equal Conversions (1%) for Ethane/D2 Exchangea metal temperature (°C) a

W

Ta

Mo

Rh

V

Cr

Pt

Zr

Pd

Ni

−54

−14

−12

26

125

136

142

145

177

192

From ref 102.

tantalum, zirconium, chromium, vanadium, nickel, platinum, palladium, and rhodium were active catalysts.102 Manganese and silver were inactive for H/D exchange, and iron and cobalt were more active for cracking. Interestingly, cracking of ethane on Fe with D2 produces CD4, but Fe does not catalyze H/D exchange between CH4 and D2 under the same conditions.103 By using the temperature required to achieve equal conversion (1% conversion) as a metric for activity, they reported the order of activity as shown in Table 3, with W being the most active again.102 Their reported apparent activation energies are listed in Table 1. More detailed analyses on the kinetics of these reactions has been conducted previously, describing the relationship between rates and the Arrhenius parameters,104 and compensation effects.15,105,106 As with methane, two main mechanisms appear to be occurring for ethane, accounting for two distinct distributions; one gives d1-ethane (C2H5D, monodeuteration) and one gives d6-ethane (C2D6, perdeuteration),107 and the distributions were metal dependent. The process that gives these two distributions in ethane is similar to that of methane; the main difference is that the second C−H activation after the ethyl hydride has formed also occurs on the β-carbon, thereby forming an

steps are slow compared to adsorption and desorption. If, however, after the adsorption event, a second C−H activation occurs,98 multiple exchange would be expected at low conversions, because more than one bond is being activated per adsorption event.99 Contrasting stepwise exchange, multiple exchange requires that the C−H activations subsequent to adsorption are faster than the desorption step. Example energy diagrams for stepwise and multiple exchange pathways for CH4 on a metal surface with excess D are shown in Figures 9 and 10, respectively. The aspect to focus on in these diagrams is the ΔEa of a methyl hydride (i.e., the adsorption product of methane on a metal surface) to either (i) desorb to produce methane or (ii) activate a second C−H bond to give a methylidene species (αα adsorption, Figure 11). For stepwise exchange, the latter Ea is greater but for multiple exchange, the former Ea is greater.100,101 The same is true for other alkanes, except that the second C−H activation likely occurs on a different carbon (analogous to an αβ interaction for ethane adsorption, see below Figure 11). Higher Alkanes and Mechanistic Implications. Kemball investigated H/D exchange of ethane and D2 over a large series of metal films, demonstrating that tungsten, molybdenum, 2301


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Figure 12. Representation of multiple exchange pathway, showing plausible C−H bond activations (H and D atoms on metal (M) surface are not shown for clarity).

exchange; the αβ intermediate was also significant for propane, and the only observed pathway for ethane. For longer alkanes, αβγ (i.e., allyl) intermediates (Figure 11) were also shown to be accessible exchange pathways.121 Additional longer chain alkanes were studied by Burwell Jr. and co-workers, looking at H/D exchange of heptane, (+)3methylhexane, 3,3-dimethylhexane, and 2,2,3-trimethylbutane with D2 on a Ni-kieselguhr catalyst.122−124 These molecules were chosen to address (i) multiple exchange pathways in the presence of a quaternary carbon and (ii) the racemization of a chiral center during H/D exchange. In line with the neopentane results of Kemball above,117 they found that multiple exchange could not propagate on a metal surface past a quaternary carbon at lower temperatures, based on the observation that for 3,3-dimethylhexane and 2,2,3-trimethylbutane exchange, a maximum of d7-species were formed (n-propyl and i-propyl branches, respectively). Additionally, they found that enantiopure (+)3-methylhexane racemized under the experimental conditions, at similar rates to H/D exchange, leading the authors to conclude that racemization occurs during the isotopic exchange.122 Burwell Jr. also did similar studies using Pd catalysts, finding similar results to that of Ni noting that at lower temperatures Pd is more likely to promote multiple exchange than Ni.125 Racemization most likely occurs through an olefin-like intermediate, but additional steps are required for the hydrogen atom on the chiral carbon to epimerize. The olefin intermediate is prochiral but requires either attack from the opposite face or desorption and readsorption. Desorption seems unlikely because at the reaction temperatures dehydrogenation is significantly endergonic. A more likely scenario involves an allylic-like intermediate, requiring additional C−H bond cleavages, thereby facilitating attack from the opposite side by allowing the olefin to flip while the alkyl fragment is still covalently bound to the metal surface. This is related to the distribution of deuterated products for cyclic alkanes during H/D exchange catalysis (see below), and is akin to the observed cis/trans selectivities when o-xylene is hydrogenated.126,127 Cyclic Alkanes and Mechanistic Implications. Kemball and Anderson studied the reactivity of cyclic alkanes with D2 on various metal films.128−133 They found that the H/D exchange activity trends for cyclohexane with transition metals was similar to other acyclic alkanes, being W > Rh > Pt > Pd (i.e., required temperatures of −66 °C, −45 °C, 0 °C, 19 °C, respectively, to reach 1% conversion).128 Interestingly, the order of reactivity for the series of cycloalkanes studied was cyclopropane > cyclopentane > cyclohexane.128,134 This is unexpected based on the BDEs (Table 2), indicating other factors can contribute, such as sterics, accessibility and ring strain (and possibly a change in the rate-limiting step, vide inf ra). At low temperature, all metals showed monodeuteration (stepwise H/D exchange), such that stronger adsorbed species were not being formed. At higher temperature, Rh and Pd

ethylene-like intermediate referred to as αβ coordination (Figure 11). Kemball postulated that a possible reason for the different mechanisms and distributions have to do the with crystal structures of the metal films.102 For example, W, Mo, Ta, Cr, and V all have body-centered cubic (BCC) structures and give d1-ethane as the primary product at low conversions.102 In contrast, Rh, Pd, and Pt have face-centered cubic (FCC) structures, and give d6-ethane as the primary product at low conversions, such that secondary C−H activations on these metals are more facile than ethane desorption. Propane and isobutane H/D exchange over W, Rh, and Ni were later studied by Kemball.108,109 Results were similar to that of ethane: namely, W was the most active catalyst at low temperature and showed monodeuteration, whereas Rh brought about multiple exchange to give highly deuterated products.110,111 Activity is again measured here by the temperature required to have equal disappearance of starting alkane (not D2), as the multiple exchange process requires more D.108 This type of measurement is therefore indicative of the first C−H bond activation and ignores secondary processes. A representation of the multiple C−H bond activations needed for multiple exchange in propane is shown in Figure 12. The third C−H activation to form the allyl intermediate is not strictly required but is shown as it is a likely intermediate. Furthermore, it was found that secondary C−H bonds in propane reacted faster than primary C−H bonds, and tertiary C−H bonds in isobutane reacted faster than the primary C−H bonds.108 Addy and Bond also contributed to this research, studying propane, and cyclopropane, exchange on Pd,112 Rh and Ir,113 and Pt.114,115 Exchange of n-hexane on Pd was studied and showed perdeuteration, demonstrating that this was a common feature on Pd.116 Additionally, they measured an EIE between H2 and n-hexane, and they found it to be 2.18 at 137 °C, measured as [(KD/KH)hexane]/[(KD/KH)H2].116 Neopentane exchange was also studied on W, Ni, Pd, and Rh films, with the order of reactivity being W > Rh > Pd based on the temperature required to achieve equal conversion.117 Several notable differences between neopentane and the other alkanes were that (i) monodeuteration (to form d1neopentane, C5H11D) was the major product on all metals at low conversions, (ii) Rh and W showed more multiple exchange at higher temperature, proposed to occur via an αγ intermediate (see Figure 11), and (iii) Rh showed multideuteration on one methyl group, which is reminiscent of methane exchange to make CD4. These results indicated that the presence of a quaternary carbon inhibited the multiple exchange process and differentiated neopentane from all other alkanes that have been mentioned above. A common observation of H/D exchange catalysis with Rh was multiple exchange, while for W stepwise exchange was typical.118−120 Furthermore, with improved NMR spectroscopic techniques, Kemball and co-workers showed that the αγ intermediate proposed for neopentane was also an intermediate for propane 2302


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Table 4. Previously Reported Apparent Activation Energies (Ea) for H/D Exchange in Alkanes over Metal Oxide Catalysts catalyst SiO2/Al2O3 SiO2/Al2O3 Al2O3b Al2O3c Al2O3 CaO CeO2 Cr2O3 Fe3O4 Ga2O3 Gd2O3 La2O3 MgO MnO Nd2O3 Pr6O11 Sm2O3 TiO2 V2O3 Yb2O3 ZnO ZrO2 Al2O3c Al2O3d Al2O3d CaO MgO SiO2/Al2O3 Al2O3c Al2O3c Al2O3c Al2O3c Al2O3c

activation temp (°C) e

825 515 515 627 540 540 540 540 540 540 540 540 540 540 540 540 540 540 540 540 540 540 627 550 550 570 570 570 627 627 627 627 627

alkanea,f

temp range (°C)

Ea (kcal mol−1)

ref

CH4 (CD4) CH4 (CD4) CH4 (CD4) CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 CH4 C3H8 C3H8 C3D8 (Al2O3) C3H8 C3H8 C3H8 i-C4H10 n-C4H10 c-C3H6 c-C4H8 c-C5H10

345−384 460−540 0−62 11−87 450−520g 450−520g 480−550g 370−450g 480−550g 290−300g 420−520g 420−500g 420−480g 460−520g 440−520g 420−520g 450−520g 490−540g 440−520g 450−550g 350−400g 460−540g 19−90 20−55 135−240 300−440 300−440 300−440 0−97 0−62 −23 to 16 0−44 16−66

13.0 33.4 5.7 4.1 34.7 27.7 54.3 30.4 46.6 39.0 29.6 29.6 46.1 39.9 30.6 40.2 28.0 47.6 31.8 44.9 32.0 31.8 8.6 8.7 9.2 15.8 18.2 16.2 7.9 8.4 6.2 10.8 10.0

149 151 151 173 252 252 252 252 252 252 252 252 252 252 252 252 252 252 252 252 252 252 173 156 156 156 156 156 173 173 173 173 173

H/D exchange was with D2 unless otherwise noted in parentheses. bAl2O3 formed by neutral hydrolysis of Al(OiPr)3, and dehydrated at 515 °C. Boehmite dehydrated at 627 °C. dGibbsite dehydrated at 550 °C. eThe silica−alumina cracking catalyst is called C-825-C, so it is assumed the activation temperature is 825 °C, but the activation temperature was not reported. fAbbreviations for alkanes are n: normal; i: iso; and c: cyclic. g Temperature range inferred from catalytic performance results and Arrhenius plot from ref 252. a c

proposed as intermediates, similar to studies above with the chiral alkanes, and the mechanism by which multiple exchange occurs has been studied in depth.138−145 More detailed studies on the exchange between cyclopentane and D2 on Pd/alumina exhibited five product distributions, which the authors suggested indicate five distinguishable processes.146,147 They assigned this distribution to different sets of surface sites; these results show that multideuteration, and thus the propensity to have a second C−H activation occur, is dependent not only on temperature, metal, and crystallite size but also on the facet and crystal termination (e.g., terrace, face, edge, step, corner) at which exchange occurs. This is not surprising because many catalysis studies show dependency of crystallite size and facet.148 General Conclusions for H/D Exchange Metal Catalysts. To summarize this section, it is clear that metal surfaces catalyze H/D exchange in acyclic and cyclic alkanes. In general, the activity of alkanes correlates with the C−H BDE, although there are some exceptions (e.g., cyclopropane). Furthermore, multiple pathways exist for exchange after the first C−H bond has been activated on a metal surface. One major pathway gives monodeuterated products, while others give multideuterated

showed multiple exchange like with acyclic alkanes, indicating the subsequent C−H activations become competitive with alkane desorption. Interestingly, at these temperatures, selectivities showed maxima at d5-cyclopentane (i.e., C5H5D5) and d6-cyclohexane (i.e., C6H6D6). These results indicated that propagation of H/D exchange occurs facilely on one side of the ring, but flipping the ring on the metal surface is slow compared with desorption.135 Burwell Jr. and co-workers studied cycloalkanes on Ni and found contrasting results; namely, maxima at half deuterated species were not observed, but this was at elevated temperatures.136 They conducted additional studies at lower temperature, which did display maxima at C5H5D5 and C6H6D6, clearly showing the importance of the reaction conditions.137 Thus, higher temperature conditions allowed ring-flippage without desorption, which was not observed at lower temperatures, consistent with more multiple exchange. In addition to finding that multiple-exchange was more likely at high temperatures (consistent with Kemball’s work above), the authors also found that lower D2 pressures and larger crystallites promoted multiple exchange; this was ascribed to having more open sites to activate additional C−H bonds.137 Allylic and pi-bonded intermediates have been 2303


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methane.151 H/D exchange of propane over alumina was stepwise, with C3H7D being the initial product. The work of Flockhart et al.156 and Larson and Hall151 both indicated that only a minority of the surface sites were active for H/D exchange catalysis. In this regard, Larson and Hall showed that the alumina catalyst had different surface sites that carried out H/D exchange at different rates, such that all surface sites were not the same and the most active sites were only a small portion (∼1%) of the surface Al−OH sites. Moreover, it was shown that if CD4 was allowed to exchange with alumina to produce CD3H (and presumably Al−OD), the D atoms were quantitatively recovered after the system was evacuated and exposed to CH4.151 This result indicated that the sites that were carrying out the rapid H/D exchange in methane did not exchange with other less active −OH sites of the alumina surface on the time scale of the experiment; if they did, the −OD from CD4 would be expected to convert to −OH sites, resulting in a net “loss” of D in the subsequent exchange with CH4. These same sites were found to be poisoned by CO2 and olefins. To help us better understand why some sites are more reactive than others, it is useful to describe the properties of alumina in more detail. Aluminas. The synthesis, production, structure, and surface properties of aluminas are complicated and have received considerable attention. In general, alumina (often shown as Al2O3) is produced by the thermal dehydration of alumina hydrates. The Bayer process produces α-alumina trihydrate, which is also the naturally occurring mineral gibbsite (Al(OH)3).157−160 A second trihydrate, β-alumina trihydrate, also known as bayerite, is another source of alumina. These hydrates can be heated to form alumina, but prior to the formation of alumina, two intermediate phases are typically formed, namely monohydrates, which are known as α-alumina monohydrate (boehmite) and β-alumina monohydrate (diaspore). Importantly, the phase of the alumina formed is dependent on the structure of the hydrate and the thermal conditions that are used during the transformation. For example, γ-alumina161 can be produced by heating boehmite at ∼550 °C (γ- is specifically noted here because it appears to be the most active phase for H/D exchange), but a different phase (e.g., η-alumina) may be formed if a different precursor is used. There is a wide range of transition aluminas known, including phases χ, γ, η, κ, and θ, which differ in their bulk structure (as determined by X-ray diffraction, for example) and surface chemistry (e.g., −OH vibrational stretches). Furthermore, the surface chemistry and spectroscopic properties of each phase is also dependent on the temperature at which the final dehydration occurs (e.g., γ-aluminas are not all the same). If any of the transition aluminas are heated at elevated temperature (ca. 1200 °C), the thermodynamic phase, αalumina (corundum), is formed. Peri and co-workers studied the properties of the γ-alumina surface that could be formed when dehydrated at different temperatures and proposed that there were five likely types of surface −OH groups (called A, B, C, D, and E sites).162,163 These −OH surface hydroxyls were classified on the basis of the number of next nearest neighbors that lost water after dehydration, ranging from zero to four. Importantly, a surface −OH site formed in which there were zero next nearest neighbors (C-site), produced an −OH that was adjacent to a strong Al Lewis acid, because the Al is bound to less oxygen atoms; the C-site was determined experimentally by using IR spectroscopy, and this site was found to have the lowest

and perdeuterated products. These distributions depend heavily on (i) the metal, (ii) the temperature of the reaction, and (iii) the alkane, because they change the relative rates of the first and subsequent C−H bond activations. At higher temperatures, secondary reactions are also expected to occur (e.g., hydrogenolysis); these reactions typically emanate from an initial C−H bond cleavage, thereby allowing H/D exchange to also provide information about these reactions.

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H/D EXCHANGE ON ALUMINA AND METAL OXIDE CATALYSTS Initial Studies. In 1948, Taylor and co-workers reported for the first time in a short communication that silica−alumina catalysts were capable of inducing H/D exchange between CH4 and CD4 at temperatures above 340 °C.149 About two decades later, Larson and Hall reported that on an alumina catalyst,150 H/D exchange occurred between methane and surface hydroxyls of alumina at room temperature.151−155 An Arrhenius analysis of the alumina system demonstrated that the reaction had a low apparent activation energy of 5.7 kcal mol−1 (Table 4). Their studies showed that there was a direct reaction between the alumina and methane, and that there was a normal primary KIE of ∼2, when comparing the rates of disappearance of CD4 (on protio alumina) and CH4 (on deuterated alumina) in separate experiments, indicating that the C−H/D bond breaking step occurs prior to or during the rate-determining step. Furthermore, like catalysis on metal surfaces, H2/D2 exchange was rapid (experimentally instantaneous) at −78 °C. However, contrasting reactivity of some metal surfaces, only stepwise exchange occurred, meaning that at low conversions, monoexchange (i.e., CH4 to CH3D) was the exclusive product (Figure 7). Other oxides were tested in addition to alumina.151 Silica was studied and was found to be 1000 times less active than silica−alumina, and silica−alumina was significantly less active than alumina. For example, aluminacatalyzed H/D exchange at 25 °C while silica−alumina catalysts required temperatures of ∼450 °C. Thus, surface OH sites were not the only functional group necessary to carry out H/D exchange at low temperature. Following up on the initial work of Larson and Hall,151 Flockhart, Uppal and Pink reported that alumina catalysts (gibbsite dehydrated at ∼550 °C) were capable of carrying out H/D exchange between propane and D2.156 Contrary to later reports (see below), they found evidence that the secondary position (methylene hydrogens, CH2) exchanged faster than the primary position (methyl hydrogens, CH3), based on line shape analysis of 1H NMR spectra. Furthermore, they claimed that without D2 (or H2) cofeeds, exchange did not occur and proposed that “the hydroxyl groups in alumina are not involved as vehicles for the deuterium and that the exchange of hydrogen between the alumina hydroxyl groups and the hydrocarbon is not an intermediate step in the exchange reaction.” However, based on other literature, this analysis seems to be incorrect, and Flockhart et al. most likely did not observe the exchange because the number of active sites on the alumina are minimal compared to the quantity of propane added. In this regard, Larson and Hall quantified the reaction sites, stating only ∼1% of the −OH groups are active.151 Nevertheless, dehydrated alumina was shown to be an active catalyst for H/D exchange of propane. Furthermore, mixtures of silica/alumina were much less active for propane H/D exchange and required higher temperatures, in accord with Larson and Hall’s results with 2304


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ACS Catalysis frequency O−H stretch (3700 cm−1). Most importantly, Peri observed experimentally that the C-site most rapidly exchanged with D2 to form surface −OD groups. It should be noted that Ratnasmay et al. pointed out that the Peri studies focused only on the {100} facet as the surface plane exposed, and proposed that other planes can contribute to the surface chemistry and spectroscopic properties, giving a different set of five surface −OH groups (called 1a, 1b, IIa, IIb and III sites).164 Regardless of the exact nature of the site, it was apparent that the rate of exchange of a surface −OH with D2 was site-dependent. Additionally, the exchange was dependent on the phase of alumina, and the lowest frequency stretch was not always found to be the most active for H/D exchange. For example, Sinfelt and co-workers found that on η-alumina, the highest frequency −OH most readily exchanged with D2.165 Other reactions besides H/D exchange were studied; the C-sites observed by Peri were found to irreversibly react with NH3 by dissociative adsorption, to form Al−NH2 and −OH species.162 This is specifically noted because Larson and Hall proposed that these are the same sites that carry out the rapid H/D exchange in methane.151 Thus, on dehydrated γ-alumina, it appears that the surface −OH groups adjacent to Lewis acidic Al sites were the most active sites for H/D exchange. Why or how these sites carry out the catalytic exchange is, however, still unclear, and is discussed in more detail below. Trends and Mechanistic Implications. Kemball and coworkers studied H/D exchange of n-butane and isobutane on alumina catalysts and found that the rate of n-butane exchange was more rapid (2.3 times faster) than that of isobutane.166 Furthermore, after extensive exchange of isobutane, they found that almost no perdeutero product, C4D10, had formed, but instead consistent primarily of C4D9H, which still had one H atom. These results indicated that only the methyl sites were exchanging and doing so at least 100 times more rapidly than the methine tertiary hydrogen, thereby forming (CD3)3CH. Additionally, the methyl groups on n-butane exchanged 55 times faster than the methylene secondary hydrogens. Notably, these trends are opposite to those for BDEs (Table 2) and catalysis on metal surfaces. These data led them to propose that exchange on alumina occurred via carbanionic intermediates, based on the stability of carbanions (Figure 13) (i.e., the acidity

methyl anion is more stable than a t-butyl anion. Thus, the acidity (pKa) of an alkane may be used to estimate the conjugate base (carbanion) stability. Opposite trends would be expected for radical (Figure 5) or carbocationic intermediates (Figure 13), thus ruling against these types of intermediates in a rate limiting step. Kemball then did additional studies on a larger series of alkanes and cycloalkanes for H/D exchange on dehydrated γalumina,173 demonstrating the generality of the H/D exchange catalysis. For propane H/D exchange, the methyl:methylene rate constant ratio was found to be 170:1 at 82 °C, again consistent with carbanion stability and opposite what is expected based on a homolytic bond energy (BDE).174 Furthermore, they found that the tertiary hydrogen in isobutane did not exchange under the reaction conditions. The cyclic alkanes showed similar trends correlating with their pKa, such that cyclopropane was the fastest to exchange, followed by cyclobutane and then cyclopentane (pKa’s are ∼38, ∼43, ∼44, respectively).173 These trends are also opposite of BDE (Table 2); interestingly, metal-catalyzed H/D exchange has similar trends for cyclic alkanes. The apparent activation energies have been previously determined and are listed in Table 4. In a follow up study,175 Kemball and co-workers showed that the activation (pretreatment) temperature for the γ-alumina that gave the highest rates of H/D exchange was 567 °C, demonstrating the sensitivity of treatment conditions for aluminas. Furthermore, they found that for a series of linear alkanes, the rates of exchange were heptane > hexane > pentane, which was rationalized again based on the stability of the carbanions that would be formed; although maybe counterintuitive, longer chain alkyl groups have similar electron-donating ability based on inductive arguments, but can better stabilize a negative charge via delocalization. In line with these trends, higher acidity hydrocarbons like olefins exchanged at faster rates, which is expected for sp2 carbons as they are stronger acids, but were also catalyst poisons. Copéret, Sautet, and co-workers further studied the effects of temperature for γ-alumina catalyst activation, and they found that at a dehydration temperature of 700 °C, the chemisorption of methane was maximized and equated this with catalytic activity.176,177 Their density functional theory (DFT) calculations indicated the C−H bond of methane is cleaved across a dehydrated Al−O species to give fragments containing Al−CH3 and O−H, and they proposed that these sites are responsible for catalysis. This is in line with what Larson and Hall proposed in their initial report,151 namely, that CH4 is dissociatively adsorbed across an Al−O fragment, a process they state is analogous to NH3 dissociative adsorption observed by Peri (see above).162 Specifically, Larson and Hall write “The driving force for this is so high that analogous processes occur with other molecules, e.g., with ammonia... Since methane is readily deuterated over alumina, it may be supposed that the molecule is cleaved as it is chemisorbed. When reacting with the dual

Figure 13. Trends in carbanion (top) and carbocation (bottom) stability.

of hydrocarbons).167−172 In general, electron-donating alkyls destabilize a carbanion relative to hydrogens, which is why a

Figure 14. Depiction of the sigma-bond metathesis mechanism on γ-alumina proposed to be responsible for H/D exchange in alkanes (CD4 shown). 2305


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Figure 15. Previously proposed (ref 180) sigma-bond metathesis (SBM) pathway for exchanging C−H and C−D bonds of methane by a Cp*2ScMe catalyst.

acid−base surface sites in a manner analogous to NH3 or H2O, then Al−CH3 and adjacent OH groups would be formed.”151 Comparisons with Organometallics. Although the mechanisms may be similar for these transformations of NH3 and CH4 with dehydrated alumina, one major difference is the lack of a lone pair on methane. While ammonia (or water) can first bind to the Lewis-acidic Al site, and then transfer a proton to the Lewis-basic O moiety, methane cannot. In this regard, the reactivity of methane is reminiscent of a mechanism proposed in organometallic chemistry, namely, sigma-bond metathesis (SBM), and is described in detail below. On dehydrated alumina, a SBM mechanism for H/D exchange can be described as is shown in Figure 14,178,179 and this is similar to the DFT structures calculated previously.177 It is important to note here that while the driving force for ammonia or water dissociative adsorption may be high, the energy surface for methane adsorption must be relatively flat, as this is a requirement based on the rapid H/D exchange catalysis observed. If this was not the case, methane adsorption would be too exothermic, thereby implying that the microscopic reverse (methane desorption) would be slow, which it is not. The trends described above for alkane C−H activation over dehydrated alumina may suggest that the transition state species that occur in these reactions have carbanionic character. As briefly mentioned above, there is an entirely different chemical system, namely, a family of organometallic scandium complexes studied by Bercaw and co-workers, that is capable of exchanging H/D on alkanes, and other hydrocarbons (e.g., olefins and alkynes),180 that show similar reactivity trends as compared to γ-alumina. Specifically, in the Cp*2ScR (Cp* = pentamethylcyclopentadienyl, C5Me5) system studied, it was found that rates of C−H/D exchange in hydrocarbons followed the same acidity trends found for exchange on alumina (i.e., sp > sp2 > sp3), in addition to the propensity to exchange primary C−H bonds more rapidly than secondary C−H bonds.181 They proposed the SBM mechanism to account for this reactivity; this mechanism describes a pathway in which two sigma-bonds approach each other to form a four-membered square-like transition state, which then allow the two former sigma-bonds to break while two new sigma-bonds form. This pathway is depicted in Figure 15 for the original Cp*2ScMe system described by Bercaw and co-workers,180 with the center structure emphasizing the four-membered transition state. The trends in reactivity (sp > sp2 > sp3) made the researchers question whether the transition states in these reactions gave carbanionic character (Figure 16),180 similar to what is proposed for H/D exchange on alumina (a proposed transition state showing the possible similarities on alumina is shown in Figure 16). However, they found that solvent effects were minimal, with nonpolar solvents having high rates; this is unexpected for polar transition states.180 Furthermore, exchange of substituted arenes with electron-donating and electron-withdrawing groups did not show trends in selectivity

Figure 16. Proposed transition states for sigma-bond metathesis mechanisms for the reported organometallic Sc system (left) and dehydrated alumina (right), indicating the partial charges that may develop.

that would be expected for charged transition states or intermediates, like what is commonly observed in electrophilic aromatic substitution reactions. Another explanation proposed by Bercaw and co-workers for these reactivity trends is based on an orbital overlap argument in the four-membered transition state for SBM,180,182 with the higher s-character hydrocarbons having better overlap with Sc−R, lowering the transition state energy and giving faster rates (i.e., sp > sp2 > sp3).183 Additionally, the electron-deficient Sc center has a higher affinity for the more electron-rich hydrocarbons, as does the Lewis-acidic and electron-deficient Al center. Figure 16 shows that similar transition states can be drawn for SBM on both the Sc−R system and the alumina system, and suggests that H/D exchange on Al2O3 may occur via a SBM type mechanism (Figure 14). Other organometallic systems like the Cp*2ScMe system have been realized to carry out H/D exchange in alkanes, most notably a [Cp*Ir(PMe3)Me]+ system discovered by Bergman and co-workers,18,19 and others referenced in the introduction.184 In the Ir systems, a SBM pathway was also proposed as a likely pathway for H/D exchange, but a discrete oxidative addition and reductive elimination sequence is also plausible with the increased dn count on Ir.185 The trends for sigma-bond metathesis are described above: namely that stronger C−H bonds (based on BDE) and more acidic C−H bonds are more reactive. Indeed, this is generally true for C−H activation in organometallic systems by oxidative addition, and systems that demonstrate this have been studied comprehensively by Jones,186−200 Bergman,18,19,201−210 and others,211−217 and this area has been reviewed.218−223 One of the most important systems to show these trends was described by Shilov, which could activate the C−H bond of methane and subsequently oxidize it to produce methyl chloride and methanol.224 Research on analogous systems to better understand the mechanism of Shilov chemistry have been studied in detail by Bercaw and Labinger.225−237 Clearly, there is an abundance of literature on C−H activation in organometallics now, but, at the time of this H/D exchange work (1950s−1970s), there was not! Indeed, the first clear example of C−H activation by oxidative addition was shown by Chatt and Davidson in 1965,238 and the following research working out these trends and mechanistic implications did not happen until many years later. 2306


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research on H/D exchange catalysis, and is at least one way to bridge the two material classes. As an interesting case, cyclic alkanes like cyclopropane, exchange rapidly on both metals and alumina; as described above, this is not expected to be the case for metal-catalyzed H/D exchange because of the high BDE of cyclopropane (Table 2). Thus, maybe in the case of cyclic alkanes, the ratelimiting step has changed for metal catalysis, or the first coordination step also involves a larger change/distortion in the ring structure; this may be expected due to ring strain, which may be relieved upon coordination. The second step (k2) is then either no longer limiting, or the relevant bond strength has been significantly modified after coordination with the metal surface, and no longer trends with the BDE of the free alkane.

In general, the reason that reactivity trends do not correlate with BDE is because the rate-limiting step is coordination of the alkane to the metal center, prior to oxidative cleavage, and not the cleavage step itself (Figure 17). As such, studies on σ-

Figure 17. Two step process of oxidative addition: alkane C−H bond coordination, followed by oxidative cleavage.

complexes and agostic239 interactions were conducted, and proposed as important intermediates in C−H bond activation, and the focus of significant research.240−244 For completion purposes, it is worthwhile noting that a unique system has been established by Wayland and co-workers, in which dinuclear complexes can homolytically cleave an alkane C−H bond,245−247 a mechanism disparate from more traditional oxidative addition. Metal- vs Alumina-Catalyzed H/D Exchange. Seeing as the activity of alumina-catalyzed H/D exchange can be related to organometallic chemistry, can the same be done with metalcatalyzed H/D exchange? When the stepwise mechanism of oxidative addition (Figure 17) is compared with what is proposed from alumina-catalyzed H/D exchange (Figure 18, top), it can be seen that there is a similar kinetics description. For alumina catalysis, the intermediate is drawn as a transition state (the four-center SBM-like transition state), but this is not required. Nevertheless, in both cases, the kinetics can be simplified if the steady-state approximation is made, such that the intermediate (or transition state) is either short-lived or its concentration is not changing; therefore, the rate equation for product formation can be written as shown in eq 3. Two limiting situations can be described: k2 ≫ k−1 or k−1 ≫ k2. For the former (k2 ≫ k−1), the rate equation reduces to be proportional to k1: this indicates that coordination of the alkane is rate limiting, and is consistent with both organometallic oxidative addition, sigma-bond metathesis, and aluminacatalyzed H/D exchange. The latter case (k−1 ≫ k2) gives a rate equation that reduces to be proportional to K*k2 (where K is the equilibrium constant for alkane coordination, k1/k−1): in this situation, the cleavage step is expected to be slow and rate limiting, therefore, it can be postulated that this step is more likely to trend with BDE. Of course, this is just a suggestion, but this simple approximation does appear to fit the large body of

rate =

k1k 2 [catalyst][alkane] k −1 + k 2

(3)

Other Metal Oxides. It is important to note that research on other metal oxides has been carried out. These other systems are not described in detail in this Review, but are mentioned if the reader is interested. CrOx systems were reported early on by Burwell Jr. to catalyze H/D exchange between hexane and D2 at temperatures above 300 °C, also by a stepwise pathway.248,249 Other oxide catalysts have been studied and display H/D exchange activity indicative of carbanionic intermediates (active: MgO, CaO, SrO, BaO, La2O3, CeO2, Pr2O3; inactive: Tb2O3, TiO2, ZrO2, ThO2, ZnO, WO2),250 and the mechanism over MgO was examined in greater detail, again showing that carbanion stability correlated with activity.251 Ga2O3 showed catalysis for both H/D exchange in alkanes and olefin isomerization.252 In a more recent study, a large range of metal oxides was studied (including Al2O3 and Ga2O3) showing that many of these metal oxides were active for H/D exchange in alkanes.253−255 Their study concludes that Ga2O3 is the most active catalyst, after normalizing by surface area, and comparing rates at 500 °C. The apparent activation energies of these catalysts were determined and are tabulated in Table 4. General Conclusions for H/D Exchange on Alumina and Metal Oxide Catalysts. To conclude this section, the activation of C−H/D bonds on the surfaces of dehydrated Al2O3 (and other metal oxides) to promote H/D exchange occurs via a mechanism that correlates well with the stability of the carbanion that would be formed if the alkane was deprotonated (i.e., pKa of the alkane). Importantly, these trends are opposite of what is expected based on the homolytic bond strength (BDE) of the C−H bond in the alkane. Sigma-

Figure 18. Comparison on kinetics of R−H activation over alumina and metal catalysts. 2307


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ACKNOWLEDGMENTS I wish to thank John Bercaw for reviewing this manuscript prior to submission and providing valuable feedback and advice. In addition, Jay Winkler, Wesley Sattler, Gerard Parkin, and my colleagues at ExxonMobil, namely, Pedro Serna, Micaela Taborga Claure, Michele Paccagnini, Randall Meyer, and Stu Soled, are thanked for helpful discussions.

bond metathesis pathways are proposed here and can account for the same trends in reactivity. Regardless of the mechanism, dehydrated aluminas are proficient in exchanging the C−H bonds in light alkanes, especially methane and ethane, and the methyl groups of propane and other higher alkanes, remarkably at room temperature.

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CONCLUSIONS Many catalysts have been discovered over the past century that are capable of carrying out H/D exchange in light alkanes. This Review focuses on two classes of catalyst, namely, (i) metals and (ii) metal oxides (specifically alumina), which demonstrates several features and trends in each class of catalyst. For metal catalysts, reactivity of the alkane trends with the BDE of the alkane. Thus, methane is the slowest alkane to exchange and therefore requires the highest temperature reaction conditions. In an alkane with chemically inequivalent hydrogens, reactivity increases in the order: primary < secondary < tertiary (in-line with BDEs). At lower temperature, stepwise (monoexchange) is most common, and at higher temperature, multiple-exchange becomes more common and is dependent on the metal (e.g., Rh promotes multiple-exchange while W promotes monoexchange). Different metals require different temperatures to carry out the exchange, and they give different product distributions and secondary reactions. Multipleexchange (and perdeuteration) is strongly inhibited by the presence of a quaternary carbon, giving strong indication to the importance of olefin and allylic like intermediates. Furthermore, at higher temperatures, secondary reactions become more prevalent (e.g., hydrogenolysis) as does catalyst poisoning. For alumina catalysis, reactivity of the alkane trends with the pKa (stability of the carbanion that would be formed) of the alkane. Thus, methane is the fastest acyclic alkane to exchange and exchanges at high rates at room temperature. These reactivity trends are the opposite of what is found for transition metal surfaces and BDEs. In an alkane with chemically inequivalent hydrogens, reactivity increases in the order: tertiary < secondary < primary (opposite of BDEs). Stepwise (monoexchange) is the only reported exchange distribution shown in the literature thus far. Lastly, both metal and alumina-catalyzed H/D exchange have been compared with organometallic chemistry regarding C−H activation at a molecular transition metal center. Interestingly, the selectivities and rates of alumina H/D exchange catalysis have very similar trends with organometallic systems that carry out C−H activation. While metal-catalyzed H/D exchange displays the opposite trends, it is possible that similar mechanisms can be envisioned, but with different kinetics and rate-limiting steps. With this large variety of catalysts, it is clear that many materials can activate the relatively inert C−H bonds of alkanes, and H/D exchange catalysis is a prime tool to better understand how these activations occur.

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Review

REFERENCES

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*E-mail: [email protected]. ORCID

Aaron Sattler: 0000-0001-5871-8646 Notes

The author declares no competing financial interest. 2308


Review

ACS Catalysis

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ACS Catalysis (98) After the second C−H activation, additional bond cleavages to give further activated hydrocarbons can also occur, but are not required to give the same isotopic distributions. (99) For a related example in organometallic catalysis, see: Yung, C. M.; Skaddan, M. B.; Bergman, R. G. J. Am. Chem. Soc. 2004, 126, 13033−13043. (100) Differences in the nature of the metal sites have been proposed to account for multiple exchange. See: Frennet, A.; Lienard, G.; Crucq, A.; Degols, L. J. Catal. 1978, 53, 150−163. (101) Bond has also described these differences in observed distributions, and the different interpretations in the literature. For more details, see reference 15. (102) Anderson, J. R.; Kemball, C. Proc. R. Soc. London, Ser. A 1954, 223, 361−377. (103) Additional studies were carried out on Fe films, showing that at temperatures of 113−229 °C, very strong adsorption of alkanes occurred, causing catalyst deactivation. When exchange was observed, like other metals, both monodeuteration and perdeuteration occurred, and the order of reactivity was butane > propane > ethane. In line with strong adsorption, the principle product of Fe hydrogenolysis is methane, while over Pt only a single C−C bond is cleaved, indicating weaker adsorption. See: Dowie, R. S.; Kemball, C.; Whan, D. A. J. Phys. Chem. 1976, 80, 2900−2903. (104) Kemball described in more detail the observed relationships between the frequency factors and apparent activation energies. See references 88, 89, and 102. (105) Bond, G. C.; Keane, M. A.; Kral, H.; Lercher, J. A. Catal. Rev.: Sci. Eng. 2000, 42, 323−383. (106) For a review on compensation effects and related topics, see: Liu, L.; Guo, Q. X. Chem. Rev. 2001, 101, 673−695. (107) d2-Ethane (C2H4D2) is also formed, indicating that another minor pathway may also be present. Analysis of the fragmentation patterns in the mass spectra indicate that the regiochemistry is 1,2-D2C2H4D2 (i.e., CH2D−CH2D), which is expected for an ethylene bound intermediate, rather than an ethylidene. (108) Kemball, C. Proc. R. Soc. London, Ser. A 1954, 223, 377−392. (109) For exchange between propane and D2 over a light activated Pt/TiO2 catalyst, see: Herrmann, J. M.; Courbon, H.; Pichat, P. J. Catal. 1987, 108, 426−432. (110) Studies on the facets of Ni indicate that ordered {111} and {100} surfaces are responsible for multiple exchange pathways, while random polycrystalline shows mono exchange for ethane and propane H/D exchange. For methane, no correlation was observed. See: Anderson, J. R.; Macdonald, R. J. J. Catal. 1969, 13, 345−359. (111) For additional studies on site selectivities for exchange of propane on Ni/alumina and hydrogenolysis activity, see: Oz, H.; Gaumann, T. J. Catal. 1990, 126, 115−125. (112) Addy, J.; Bond, G. C. Trans. Faraday Soc. 1957, 53, 368−376. (113) Addy, J.; Bond, G. C. Trans. Faraday Soc. 1957, 53, 383−387. (114) Addy, J.; Bond, G. C. Trans. Faraday Soc. 1957, 53, 388−392. (115) Maegawa, T.; Fujiwara, Y.; Inagaki, Y.; Esaki, H.; Monguchi, Y.; Sajiki, H. Angew. Chem., Int. Ed. 2008, 47, 5394−5397. (116) Meyer, E. F.; Kemball, C. J. Catal. 1965, 4, 711−716. (117) Kemball, C. Trans. Faraday Soc. 1954, 50, 1344−1351. (118) Brown, R.; Kemball, C.; Oliver, J. A.; Sadler, I. H. J. Chem. Res. (S) 1985, 274−275. (119) For additional studies on Rh and Pt H/D exchange catalysis, which has demonstrated the affinity of Rh for multi-exchange, see: McKee, D. W.; Norton, F. J. J. Catal. 1965, 4, 510−517. (120) Studies on the Pt{111} surface of a single crystal demonstrate the propensity for mono v. perdeutero exchange and show that two mechanisms occur, which have activation energies of 17.0 kcal mol−1 and 6.5 kcal mol−1, respectively. See: Zaera, F. Catal. Lett. 1991, 11, 95−104. (121) Gault, F. G.; Kemball, C. Trans. Faraday Soc. 1961, 57, 1781− 1794. (122) Burwell, R. L., Jr.; Briggs, W. S. J. Am. Chem. Soc. 1952, 74, 5096−5102.

(123) For related a related study, see: Long, M. A.; Moyes, R. B.; Wells, P. B.; Garnett, J. L. J. Catal. 1978, 52, 206−217. (124) For related work by Burwell Jr. on unsaturated hydrocarbons, see: Meyer, E. F.; Burwell, R. L., Jr. J. Am. Chem. Soc. 1963, 85, 2881− 2887. (125) Burwell, R. L., Jr.; Shim, B. K. C.; Rowlinson, H. C. J. Am. Chem. Soc. 1957, 79, 5142−5148. (126) Neyestanaki, A. K.; Maki-Arvela, P.; Backman, H.; Karhu, H.; Salmi, T.; Vayrynen, J.; Murzin, D. Y. J. Catal. 2003, 218, 267−279. (127) Neyestanaki, A. K.; Maki-Arvela, P.; Backman, H.; Karhu, H.; Salmi, T.; Vayrynen, J.; Murzin, D. Y. J. Mol. Catal. A: Chem. 2003, 193, 237−250. (128) Anderson, J. R.; Kemball, C. Proc. R. Soc. London, Ser. A 1954, 226, 472−489. (129) For related work on H/D exchange of cyclopentane on Pd/ alumina, see: Lomot, D.; Karpiński, Z. Catal. Lett. 2000, 69, 133−138. (130) For related work on H/D exchange and dehydrogenation of cyclohexylamine on Pt, see: Moss, R. L.; Kemball, C. Nature 1956, 178, 1069−1069. (131) For related work on H/D exchange and cyclization on Pt, see: Shi, B. C.; Davis, B. H. J. Catal. 1994, 147, 38−47. (132) Cyclobutane has also been studied. See Hilaire, L.; Maire, G. Bull. Soc. Chim. Fr. 1967, 886−894. (133) Early studies were done by Farkas and Farkas. See reference 81. (134) For studies on cyclopropane hydrogenolysis with D2 over Pt, see: Bond, G. C.; Turkevich, J. Trans. Faraday Soc. 1954, 50, 1335− 1343. (135) For use of 2H NMR spectroscopy for cyclopentane exchange assignments, see: Faro, A. C., Jr.; Kemball, C.; Brown, R.; Sadler, I. H. J. Chem. Res. (S) 1982, 342−343. (136) Rowlinson, H. C.; Burwell, R. L., Jr.; Tuxworth, R. H. J. Phys. Chem. 1955, 59, 225−231. (137) Burwell, R. L., Jr.; Tuxworth, R. H. J. Phys. Chem. 1956, 60, 1043−1049. (138) Rooney, J. J. J. Catal. 1963, 2, 53−57. (139) Ledoux, M. J.; Gault, F. G.; Bouchy, A.; Roussy, G. J. Chem. Soc., Faraday Trans. 1 1978, 74, 2652−2666. (140) Rooney, J. J.; Gault, F. G.; Kemball, C. Proc. Chem. Soc. 1960, 407−408. (141) Gault, F. G.; Rooney, J. J.; Kemball, C. J. Catal. 1962, 1, 255− 274. (142) Gault, F. G.; Rooney, J. J. J. Chem. Soc., Faraday Trans. 1 1979, 75, 1320−1329. (143) Clarke, J. K. A.; Hegarty, B. F.; Rooney, J. J. J. Mol. Catal. 1990, 62, L39−L43. (144) For related work on cycloolefins, see: Hilaire, L.; Gault, F. G. J. Catal. 1971, 20, 267−287. (145) Bond has commented on the different interpretations for cyclic alkane exchange reactions; see reference 15. (146) Schrage, K.; Burwell, R. L., Jr. J. Am. Chem. Soc. 1966, 88, 4549−4555. (147) Schrage, K.; Burwell, R. L., Jr. J. Am. Chem. Soc. 1966, 88, 4555−4560. (148) For related support effects on Pt-catalyzed H/D exchange in cyclopentane, see: Oudenhuijzen, M. K.; van Bokhoven, J. A.; Koningsberger, D. C. J. Catal. 2003, 219, 134−145. (149) Parravano, G.; Hammel, E. F.; Taylor, H. S. J. Am. Chem. Soc. 1948, 70, 2269−2270. (150) The alumina was dehydrated at 650 °C, and prepared from neutral hydrolysis of aluminum isopropoxide, thereby forming mixtures of η- and γ-alumina. (151) Larson, J. G.; Hall, W. K. J. Phys. Chem. 1965, 69, 3080−3089. (152) For related studies of D2 exchange on silica, silica-alumina, and alumina, see: Hall, W. K.; Leftin, H. P.; Cheselske, F. J.; O’Reilly, D. E. J. Catal. 1963, 2, 506−517. (153) For related studies on D2 exchange on Pt/alumina, see: Hall, W. K.; Lutinski, F. E. J. Catal. 1963, 2, 518−527. 2310


Review

ACS Catalysis (154) For related studies on catalytic sites of alumina and silicaalumina, see: Hall, W. K.; Lutinski, F. E.; Gerberich, H. R. J. Catal. 1964, 3, 512−527. (155) For related studies on ortho-para H2 conversion and H2/D2 exchange on alumina, see: Van Cauwelaert, F. H.; Hall, W. K. Trans. Faraday Soc. 1970, 66, 454−468. (156) Flockhart, B. D.; Uppal, S. S.; Pink, R. C. Trans. Faraday Soc. 1971, 67, 513−525. (157) Stumpf, H. C.; Russell, A. S.; Newsome, J. W.; Tucker, C. M. Ind. Eng. Chem. 1950, 42, 1398−1403. (158) Brown, J. F.; Clark, D.; Elliott, W. W. J. Chem. Soc. 1953, 84− 88. (159) Souza Santos, P.; Souza Santos, H.; Toledo, S. P. Mater. Res. 2000, 3, 104−114. (160) Stacey, M. H. Langmuir 1987, 3, 681−686. (161) For a detailed description of the structure of γ-alumina (viewed as a defect oxyhydroxide), see: Soled, S. J. Catal. 1983, 81, 252−257. (162) Peri, J. B.; Hannan, R. B. J. Phys. Chem. 1960, 64, 1526−1530. (163) Peri, J. B. J. Phys. Chem. 1965, 69, 220−230. (164) Ratnasamy, P.; Sivasanker, S. Catal. Rev.: Sci. Eng. 1980, 22, 401−429. (165) Carter, J. L.; Lucchesi, P. J.; Corneil, P.; Yates, D. J. C.; Sinfelt, J. H. J. Phys. Chem. 1965, 69, 3070−3074. (166) Robertson, P. J.; Scurrell, M. S.; Kemball, C. J. Chem. Soc., Chem. Commun. 1973, 799. (167) Dessy, R. E.; Kitching, W.; Psarras, T.; Salinger, R.; Chen, A.; Chivers, T. J. Am. Chem. Soc. 1966, 88, 460−467. (168) Streitwieser, A.; Keevil, T. A.; Taylor, D. R.; Dart, E. C. J. Am. Chem. Soc. 2005, 127, 9290−9297. (169) Streitwieser, A.; Caldwell, R. A.; Lawler, R. G.; Ziegler, G. R. J. Am. Chem. Soc. 1965, 87, 5399−5402. (170) Streitwieser, A.; Young, W. R.; Caldwell, R. A. J. Am. Chem. Soc. 1969, 91, 527−528. (171) Streitwieser, A.; Caldwell, R. A.; Young, W. R. J. Am. Chem. Soc. 1969, 91, 529. (172) Streitwieser, A.; Young, W. R. J. Am. Chem. Soc. 1969, 91, 529− 530. (173) Robertson, P. J.; Scurrell, M. S.; Kemball, C. J. Chem. Soc., Faraday Trans. 1 1975, 71, 903−912. (174) It should be noted that this trend is the opposite of what was found in reference 156, but is consistent with the other literature in the field. (175) John, C. S.; Kemball, C.; Pearce, E. A.; Pearman, A. J. J. Chem. Res. (S) 1979, 400−401. (176) Joubert, J.; Salameh, A.; Krakoviack, V.; Delbecq, F.; Sautet, P.; Copéret, C.; Basset, J. M. J. Phys. Chem. B 2006, 110, 23944−23950. (177) Wischert, R.; Copéret, C.; Delbecq, F.; Sautet, P. Angew. Chem., Int. Ed. 2011, 50, 3202−3205. (178) The coordination number and geometry around Al and O is not meant to be specific, but just simplified to exemplify the SBM mechanism (e.g., it is not implied that the Al starts as a two coordinate species). (179) It should be noted that there is a lone pair on the oxygen atom in alumina, which is not the case in the organometallic scandium system. (180) Thompson, M. E.; Baxter, S. M.; Bulls, A. R.; Burger, B. J.; Nolan, M. C.; Santarsiero, B. D.; Schaefer, W. P.; Bercaw, J. E. J. Am. Chem. Soc. 1987, 109, 203−219. (181) Also in accord with reactivity trends on alumina, cyclopropane H/D exchange was rapid in the Bercaw study (reference 180). (182) For a review on sigma-bond metathesis reactions, see: Lin, Z. Y. Coord. Chem. Rev. 2007, 251, 2280−2291. (183) In this regard, it is interesting to note that in both systems that undergo H/D exchange, C−C bond formation (with subsequent H2 formation) does not occur, and was suggested in reference 180 to not occur because of the same orbital overlap argument (i.e., sp3 orbital on C has less s character than the lone s orbital on H). (184) A surface organometallic species was synthesized previously and carried out H/D exchange catalysis, and was proposed to do so by

a SBM mechanism. See: Casty, G. L.; Matturro, M. G.; Myers, G. R.; Reynolds, R. P.; Hall, R. B. Organometallics 2001, 20, 2246−2249. (185) Strout, D. L.; Zaric, S.; Niu, S. Q.; Hall, M. B. J. Am. Chem. Soc. 1996, 118, 6068−6069. (186) Jones, W. D.; Feher, F. J. J. Am. Chem. Soc. 1982, 104, 4240− 4242. (187) Jones, W. D.; Feher, F. J. J. Am. Chem. Soc. 1984, 106, 1650− 1663. (188) Jones, W. D.; Feher, F. J. J. Am. Chem. Soc. 1985, 107, 620− 631. (189) Jones, W. D.; Feher, F. J. Acc. Chem. Res. 1989, 22, 91−100. (190) Jones, W. D.; Feher, F. J. J. Am. Chem. Soc. 1986, 108, 4814− 4819. (191) Chin, R. M.; Dong, L. Z.; Duckett, S. B.; Partridge, M. G.; Jones, W. D.; Perutz, R. N. J. Am. Chem. Soc. 1993, 115, 7685−7695. (192) Atesin, T. A.; Li, T.; Lachaize, S.; Brennessel, W. W.; Garcia, J. J.; Jones, W. D. J. Am. Chem. Soc. 2007, 129, 7562−7569. (193) Vetter, A. J.; Rieth, R. D.; Brennessel, W. W.; Jones, W. D. J. Am. Chem. Soc. 2009, 131, 10742−10752. (194) Evans, M. E.; Li, T.; Jones, W. D. J. Am. Chem. Soc. 2010, 132, 16278−16284. (195) Choi, G.; Morris, J.; Brennessel, W. W.; Jones, W. D. J. Am. Chem. Soc. 2012, 134, 9276−9284. (196) Jiao, Y. Z.; Evans, M. E.; Morris, J.; Brennessel, W. W.; Jones, W. D. J. Am. Chem. Soc. 2013, 135, 6994−7004. (197) Jiao, Y. Z.; Morris, J.; Brennessel, W. W.; Jones, W. D. J. Am. Chem. Soc. 2013, 135, 16198−16212. (198) Evans, M. E.; Burke, C. L.; Yaibuathes, S.; Clot, E.; Eisenstein, O.; Jones, W. D. J. Am. Chem. Soc. 2009, 131, 13464−13473. (199) Jones, W. D. Acc. Chem. Res. 2003, 36, 140−146. (200) Jones, W. D. Inorg. Chem. 2005, 44, 4475−4484. (201) Wasserman, E. P.; Moore, C. B.; Bergman, R. G. Science 1992, 255, 315−318. (202) Hoyt, H. M.; Michael, F. E.; Bergman, R. G. J. Am. Chem. Soc. 2004, 126, 1018−1019. (203) Janowicz, A. H.; Bergman, R. G. J. Am. Chem. Soc. 1983, 105, 3929−3939. (204) Janowicz, A. H.; Bergman, R. G. J. Am. Chem. Soc. 1982, 104, 352−354. (205) Buchanan, J. M.; Stryker, J. M.; Bergman, R. G. J. Am. Chem. Soc. 1986, 108, 1537−1550. (206) Bergman, R. G. Science 1984, 223, 902−908. (207) Alaimo, P. J.; Bergman, R. G. Organometallics 1999, 18, 2707− 2717. (208) Klei, S. R.; Golden, J. T.; Burger, P.; Bergman, R. G. J. Mol. Catal. A: Chem. 2002, 189, 79−94. (209) Arndtsen, B. A.; Bergman, R. G.; Mobley, T. A.; Peterson, T. H. Acc. Chem. Res. 1995, 28, 154−162. (210) Fiedler, D.; Leung, D. H.; Bergman, R. G.; Raymond, K. N. Acc. Chem. Res. 2005, 38, 349−358. (211) Rybtchinski, B.; Cohen, R.; Ben-David, Y.; Martin, J. M. L.; Milstein, D. J. Am. Chem. Soc. 2003, 125, 11041−11050. (212) Bolig, A. D.; Brookhart, M. J. Am. Chem. Soc. 2007, 129, 14544−14545. (213) Hoyano, J. K.; Graham, W. A. G. J. Am. Chem. Soc. 1982, 104, 3723−3725. (214) Bullock, R. M.; Headford, C. E. L.; Hennessy, K. M.; Kegley, S. E.; Norton, J. R. J. Am. Chem. Soc. 1989, 111, 3897−3908. (215) Ghosh, C. K.; Graham, W. A. G. J. Am. Chem. Soc. 1987, 109, 4726−4727. (216) Wick, D. D.; Goldberg, K. I. J. Am. Chem. Soc. 1997, 119, 10235−10236. (217) Fu, R.; O’Reilly, M. E.; Nielsen, R. J.; Goddard, W. A., III; Gunnoe, T. B. Chem. - Eur. J. 2015, 21, 1286−1293. (218) Balcells, D.; Clot, E.; Eisenstein, O. Chem. Rev. 2010, 110, 749−823. (219) Lersch, M.; Tilset, M. Chem. Rev. 2005, 105, 2471−2526. (220) Shilov, A. E.; Shul’pin, G. B. Chem. Rev. 1997, 97, 2879−2932. (221) Crabtree, R. H. Chem. Rev. 1985, 85, 245−269. 2311


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ACS Catalysis (222) Halpern, J. Inorg. Chim. Acta 1985, 100, 41−48. (223) Crabtree, R. H. Chem. Rev. 1995, 95, 987−1007. (224) Goldshleger, N. F.; Shteinman, A. A.; Shilov, A. E.; Eskova, V. V. Zh. Fiz. Khim. 1972, 46, 1353−1354. (225) Driver, T. G.; Day, M. W.; Labinger, J. A.; Bercaw, J. E. Organometallics 2005, 24, 3644−3654. (226) Stahl, S. S.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 1996, 118, 5961−5976. (227) Owen, J. S.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2006, 128, 2005−2016. (228) Bercaw, J. E.; Hazari, N.; Labinger, J. A.; Oblad, P. F. Angew. Chem., Int. Ed. 2008, 47, 9941−9943. (229) Heyduk, A. F.; Driver, T. G.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2004, 126, 15034−15035. (230) Holtcamp, M. W.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 1997, 119, 848−849. (231) Iverson, C. N.; Carter, C. A. G.; Baker, R. T.; Scollard, J. D.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2003, 125, 12674− 12675. (232) Johansson, L.; Tilset, M.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2000, 122, 10846−10855. (233) Luinstra, G. A.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 1993, 115, 3004−3005. (234) Stahl, S. S.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 1995, 117, 9371−9372. (235) Williams, T. J.; Caffyn, A. J. M.; Hazari, N.; Oblad, P. F.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2008, 130, 2418−2419. (236) Zhong, H. A.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2002, 124, 1378−1399. (237) Driver, T. G.; Williams, T. J.; Labinger, J. A.; Bercaw, J. E. Organometallics 2007, 26, 294−301. (238) Chatt, J.; Davidson, J. M. J. Chem. Soc. 1965, 843−855. (239) Brookhart, M.; Green, M. L. H.; Parkin, G. Proc. Natl. Acad. Sci. U. S. A. 2007, 104, 6908−6914. (240) Brookhart, M.; Green, M. L. H. J. Organomet. Chem. 1983, 250, 395−408. (241) Wick, D. D.; Reynolds, K. A.; Jones, W. D. J. Am. Chem. Soc. 1999, 121, 3974−3983. (242) Northcutt, T. O.; Wick, D. D.; Vetter, A. J.; Jones, W. D. J. Am. Chem. Soc. 2001, 123, 7257−7270. (243) Vetter, A. J.; Flaschenriem, C.; Jones, W. D. J. Am. Chem. Soc. 2005, 127, 12315−12322. (244) Hall, C.; Perutz, R. N. Chem. Rev. 1996, 96, 3125−3146. (245) Wayland, B. B.; Ba, S.; Sherry, A. E. J. Am. Chem. Soc. 1991, 113, 5305−5311. (246) Sherry, A. E.; Wayland, B. B. J. Am. Chem. Soc. 1990, 112, 1259−1261. (247) Cui, W. H.; Wayland, B. B. J. Am. Chem. Soc. 2004, 126, 8266− 8274. (248) Burwell, R. L., Jr.; Littlewood, A. B. J. Am. Chem. Soc. 1956, 78, 4170−4171. (249) Burwell, R. L., Jr.; Littlewood, A. B.; Cardew, M.; Pass, G.; Stoddart, C. T. H. J. Am. Chem. Soc. 1960, 82, 6272−6280. (250) Utiyama, M.; Hattori, H.; Tanabe, K. J. Catal. 1978, 53, 237− 242. (251) Hoq, M. F.; Nieves, I.; Klabunde, K. J. J. Catal. 1990, 123, 349−363. (252) Carleton, F. B.; Quinn, H. A.; Rooney, J. J. J. Chem. Soc., Chem. Commun. 1973, 231−232. (253) Hargreaves, J. S. J.; Hutchings, G. J.; Joyner, R. W.; Taylor, S. H. Appl. Catal., A 2002, 227, 191−200. (254) For related methane exchange studies on metal oxides, see: Quanzhi, L.; Amenomiya, Y. Appl. Catal. 1986, 23, 173−182. (255) For related n-butane exchange on ZSM-5, see: Narbeshuber, T. F.; Stockenhuber, M.; Brait, A.; Seshan, K.; Lercher, J. J. Catal. 1996, 160, 183−189.

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D) Exchange Catalysis in Alkanes - ACS

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