Recent Advances and Advisable Applications of Bond Energetics in

Perspective Cite This: J. Am. Chem. Soc. 2018, 140, 8611−8623

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Recent Advances and Advisable Applications of Bond Energetics in Organic Chemistry Jin-Dong Yang,†,+ Pengju Ji,†,+ Xiao-Song Xue,‡ and Jin-Pei Cheng*,†,‡ †

Center of Basic Molecular Science (CBMS), Department of Chemistry, Tsinghua University, Beijing 100084, China State Key Laboratory of Elemento-organic Chemistry, Collaborative Innovation Centre of Chemical Science and Engineering, College of Chemistry, Nankai University, Tianjin 300071, China

‡

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S Supporting Information *

Although applications of these equations to analyzing the observed chemistry have achieved great success in the past, it should be noted that hasty use of these relatively simple guidelines to handle complicated situations may often encounter problems. This should not, however, be viewed as a collapse of the governing roles of bond energetics. Rather, it points out the need of a more inclusive consideration of the other factors at work in treating such cases, as will be illustrated herein. Indeed, this field seems to have met unexpected challenges over the past 10 to 20 years because the target problems in contemporary chemistry have become much more complicated. Two examples come to mind. First, in C−H activation, the rates of metal-oxo mediated hydrogen-atom transfers (HAT) are well observed to linearly correlate with C−H BDEs.8 In many other cases, however, a stronger C−H bond, instead of a weaker one, can be selectively functionalized,9 degrading a simple use of energetic criteria. Second, in contrast to abundant R−H BDE and pKa data,4 knowledge about many other important bonds, such as C−C, C−N, C−O, or C−M (metal), as well as some weak bondings (e.g., H-bond, halogen-bond), those are needed for handling the more complex situations encountered today, is lacking, yet hampering attempts to apply bond energetics in research. Nevertheless, while such new challenges are not insurmountable, they do pinpoint areas where practical bottlenecks are located and where further energetic research and application should be emphasized. It is true that, this field does not attract the attention as it once did; nevertheless, the motivation and research activities on bond energy are still enticing investigators into intriguing and puzzling experimental observations based on trial-and-error bench research. Advances in this broad area in recent years are very impressive, ranging from the hydricity scale,10 the pKa slide rule,11 and energetic study on enantioselectivity,12 to name a few, as well as “effective BDE” analysis for proton-coupled electron transfer (PCET),13 or “unorthodox non-covalent interactions”.14 We will describe studies that are most representative of the power of bond parameters in unraveling complex experimental situations. We will primarily focus on experimental not computational works, although many more could be found for the latter. We will start by showing the excellent potential of using the proper blends of basic energetic quantities, such as those demonstrated in developing PCETs by merging pKa and redox data, to explore or to understand organic chemistry. This

ABSTRACT: Most organic transformation involves cleavage and formation of various covalent bonds, and naturally, can be regarded as a process of bond reorganization, which should be intrinsically related to bond energies (e.g., pKa, BDE, etc.). However, in many cases such as in C−H bond activation/functionalization, direct correspondence between the bond energy and reaction rate or other relevant properties is only occasionally observed when applying the bond data by simple rules like the Linear Free-Energy Relationships (LFERs) in handling intricate reaction systems. In this Perspective, we present examples to argue that the abovementioned situation is not a consequence of a diminishing role of the bond energetics in research, but most likely, comes from an improper use of energetic strategy, or simply due to a faulty selection of the data from unsuitable sources. Some advisable applications of bond energies in unscrambling the problems in modern day chemistry are exemplified through representative recent advances of the researches in this connection. Some of the possible directions of future research endeavors in the field of bond energetics and its prudent applications are recommended.

1. INTRODUCTION Since the initial description of chemical bond by Lewis1 in 1916 and the groundwork laid afterward by London2 and Pauling3 on the nature of bonding, chemical bond research has become one of the core fields in chemistry and has prospered as such for nearly a century. This success can be attributed to the innate nature of bond energetics associated with an understanding of the mechanism and driving force of reaction. This is true because chemical transformations can be viewed as bond reorganization through bond breaking and reconstruction intrinsically governed by bond energies. As we now know, enormous bond energy data of different types, especially those associated with homolysis and heterolysis of covalent bonds, have been measured and accumulated in the past4 and have played a vital role in advancing traditional chemistry to become a more rational science. Many key principles in physical organic chemistry still widely used nowadays were in fact founded on the basis of various energetic parameters, taking the three landmarks of the Hammett equation,5 Brönsted catalysis law,6 and Marcus equation7 as examples. © 2018 American Chemical Society

Received: April 17, 2018 Published: June 19, 2018 8611

DOI: 10.1021/jacs.8b04104 J. Am. Chem. Soc. 2018, 140, 8611−8623

Perspective

Journal of the American Chemical Society

strong X−H bonds20 via PCET are along this line and very fruitful. However, an exhaustive summation of the new achievements in this area does not meet the aim of this paper and will not be attempted. Indeed, the most noteworthy feature of PCET that merits the special attention of this Perspective is its surprisingly broad manipulatable range of thermodynamic driving force “BDFE” (25−110 kcal/mol). This seemingly new energetic term, the bond dissociation free energy, or the “effective BDFE”, is achievable on the basis of properly blended pKa and redox potentials of the catalyst system, i.e., a Brönsted acid-reductant pair or a base-oxidant pair (Scheme 1), selected for a particular reaction. In contrast to the regular solution-BDE that is derived by joining the energetic quantities of the same species and also used for assessing the H•-donor capacity of the parent substrate itself, the BDFE of the PCET catalyst, though it is also based on the idea of blending, is obtained otherwise from the pKa and redox potential of the respective species that compose the catalyst pair. As shown in Scheme 1, in a PCET, the electron and proton could be transferred respectively from two independent sites of the catalyst pair. Consequently, the thermodynamic capacity of a given combination to donate a formal hydrogen atom can be identified by the BDFE value (eq 1).

will be followed by a description of the work related to complementary use of the bond-breaking and bond-forming energies to rationalize some baffling observations in C−H activation or enzymatic oxidation that are beyond one’s intuition. Then, in the third part, we present typical examples as a reminder for that basic bond energy data (e.g., pKa in ionic liquids) are irreplaceable for untangling many key issues in organic chemistry and for disclosing insights into observations behind the data. The outlook for future bond energetic research as well as its advisable applications will be addressed in relevant sections.

2. BLENDING ENERGETIC QUANTITIES FOR PREDICTING OR ANALYZING REACTIONS The energy required to homolytically break an R−H bond to generate R•/H• radicals is called bond-dissociation-energy (BDE) by convention, which, early on, relied entirely on gas phase chemistry, but in solution, was essentially inaccessible. However, Bordwell et al. developed a thermochemical approach,15 by taking advantage of abundant pKa and readily achievable redox data, to evaluate many BDEs in solution,4a the invaluable role of which has now been well recognized not only for providing the data for relatively large organic molecules that are unsuitable for measuring in the gas phase but also for its inimitable utility in promoting new synthetic strategies especially for reactions involving radical generation such as HAT, and more recently, for proton-coupled electron transfer (PCET) in solution. PCET, which occurs when an electron and a proton are removed in a concerted way as depicted in Scheme 1, is

BDFE = 1.37pK a(AH) + 23.06[E 0(Redn / n + 1) + E 0(H0/ +)]

(1)

This formalism has already been generally accepted and extensively explored.10,21 Experimental BDFE values of PCET reagents and excellent examples of their successful applications have been summarized by Mayer21 and Knowles13 in their elegant reviews. For a broader perspective of related studies, those are the two best recommended sources of previous literature. In principle, the attention of the present Perspective will primarily be directed toward some of the most representative applications of BDFEs and their potential in predicting and understanding the reactions. 2.1. Governing Role of BDFE in Developing Reductions of Polar π-Bonds via PCET. Generation of versatile Ccentered radicals via traditional HAT to organic π-bonds (e.g., CO) is normally energetically inaccessible, essentially because the newly formed OH bond is much weaker (∼30 kcal/mol, see Figure 1) compared to that of most HAT reagents. The broad hydrogen donating capability of PCET reagents may provide an alternative approach. In 2013, Knowles applied PCET to successfully discover the way for ketyl-olefin couplings.19a In their system, a ketyl radical was generated by the catalyst combination involving Ru complex and acid via reductive PCET (Figure 1). Subsequent intramolecular ketylolefin coupling and esterification released the cyclic products. Thermodynamic analysis illustrated that catalyst combinations with BDFE less than 30 kcal/mol could offer good to excellent yields (74−93%), while those with BDFE far larger than 33 kcal/ mol could not promote couplings. The reaction efficiency was

Scheme 1. Reductive and Oxidative PCET Processes, where H-A, B, Redn and Oxn Refer to Brönsted Acid, Base, Reductant and Oxidant, Respectively

ubiquitous in biological metabolism16 as well as in many chemical transformations.17 Compared to the conventional HAT, the multisite PCET affords distinct advantages for developing new chemistry. In the past decade, exploitation of PCET to generate key radical intermediates, especially combined with photochemical activation,13,18 has become a mainstream endeavor in organic chemistry. The recent seminal discoveries on reduction of polar π-bonds19 and homolysis of

Figure 1. Application of PCET in catalytic ketyl-olefin coupling. 8612


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100 kcal/mol. These are hardly ever activated even with the most powerful HAT acceptors. The adjustable hydrogen abstracting ability of PCET combinations may provide a possible avenue to mildly cleave such a strong X−H bond. The pioneering work in this regard was reported by Knowles in 2015.20b By varying the compositions of a dual catalyst system (Ir oxidant/base), they were able to broaden a BDFE window to 80−108 kcal/mol. It was found that catalyst combinations with BDFE considerably lower than the N−H BDFE of substrates (∼99 kcal/mol) were unable to efficiently promote the reaction (see Figure 3, yields < 6%). However, combinations with BDFE

found roughly proportional to the BDFEs of catalyst combinations. In light of this energetic trend, a similar PCET combination was adopted to achieve asymmetric aza-pinacol cyclizations.19b Another PCET catalyst, the Ir/amine pair, was recently found to be an efficient combination in the reductive aldehyde, ketone and imine homocouplings,19c as well as in aldimine/aldehydeaniline couplings.19d Notably, in this system developed by Rueping et al., the generated amine radical cation served as strong Brönsted acid and the reduced Ir complex as the reductant. Although not explicitly stated, the significance of BDFE can be realized upon a careful examination. As seen, several Ir/amine combinations with BDFE ranging from 25 to 35 kcal/mol could enable the homocouplings with various efficiency (44−78%), while when an Ir complex was replaced with a weak reductant Ru complex, the couplings could not occur (Figure 2). That is because Ru complex renders BDFE of

Figure 3. Catalytic alkene carboaminations enabled by PCET.

around or greater than 99 kcal/mol could successfully generate amidyl radicals and enable intramolecular carboaminations of olefins (yields 16−92%). Mechanistic studies showed that the reactions proceeded via a PCET pathway. The scope of N−H substrates described above was limited to N-aryl amides. Replacement of N-aryl by alkyl groups results in an average increase of 4 pKa units in DMSO and 6 kcal/mol in BDFE for N−H bonds. Attempts to activate stronger N−H bonds in N-alkyl amides (107−110 kcal/mol) by the previous catalyst combinations proved to be unsuccessful. Thus, Knowles20c and Rovis20d independently developed more powerful PCET combinations to enable remote C−H bond activation, and they revealed the dependence of catalyst efficiency on the energetic quantity BDFE (Figure 4). Through an intramolecular 1,5-HAT, amidyl radicals generated in situ by PCET were observed to selectively activate the remote δ-sp3C− H bond in the presence of many other seemingly indistinguishable bonds of similar kind. It was found that the product yield (10−82%) corresponded proportionally with the BDFE of the catalyst combination (92−103 kcal/mol), which could be modulated by varying the electronic effects of substituents on the dipyridyl ligands of the Ir complexes. The most effective combination selected in Knowles’ study20c to furnish the vital Ncentered radical is [Ir(dF(CF3)ppy)2(5,5′-(dCF3)bpy)]+ (E0* = 1.30 V vs Fc0/+ in AN) and Bu4NOP(O)(OBu)2 (pKa of its conjugate acid: ∼13 in AN). On the other hand, in Rovis’ trifluoroacetamide system,20d the weaker oxidant [Ir(dF(CF3)ppy)2(dtbbpy)]+ (E0* = 0.83 V) was employed jointly with a much stronger base K3PO4 (pKa of its conjugate acid > 20 in AN), making their BDFE (∼103 kcal/mol) comparable to Knowles’ catalyst combination. The equivalent capacities of the catalysts from these two groups were verified by the observation of nearly identical yields of similar substrates in two independent systems (Figure 4), again indicating the predictive value of the BDFEs. More recently, a similar PCET strategy has been extended to embrace more synthetic practices involving homolysis of strong N−H bonds, such as the intermolecular anti-Markovnikov hydroamination of unactivated alkenes,24 βamination,25 and so on. The tendency of reaction efficiencies in accordance with their thermodynamic BDFEs was further demonstrated.

Figure 2. Photoredox-catalyzed pinacol coupling of benzalmization studies.

catalyst combinations over 40 kcal/mol, far greater than that is needed to break CO π-bond to form a radical. When BDFEs of catalyst combinations were modulated from about 25 to 50 kcal/mol by introducing additives [pKa of (conjugate) acids: ∼1−20 in AN] into reaction systems, further investigations demonstrated that the product yields varied correspondingly. Accordingly, this combination has recently been extended to reductive arylation of carbonyl/iminyl derivatives, and similar BDFE control was obeyed.22 Recently, reductions of carbonyls and of electron-rich enamines were also realized by different groups via PCET using a SmI2/H2O combination.23 Based on a thermodynamic analysis, the BDFE of the SmI2/H2O pair to release hydrogen is estimated to be 26 kcal/mol, which is among the weakest XH bonds of stable reagents. Therefore, the SmI2/H2O combination can be regarded as one of the most powerful hydrogen donors, and hence, should be able to reduce the electron-rich CC bond of enamines, otherwise incapable of unimolecular hydrogen reduction. Although substantial achievements have been accumulated in the synthetic applications of PCET in the hydrogen reductive process, the substrate scope has, up to now, mainly been limited to aromatic carbonyl and imine compounds. In view of the wide range of BDFE (25−110 kcal/mol) for PCET reagents, it is reasonable to expect that many other unsaturated bonds could, at least theoretically, be hydrogenated, provided that a proper PCET reagent can be identified. That is to say, many PCET reductive systems remain to be explored with the anticipation of discovering new PCET combinations for valuable synthetic methodology. 2.2. Governing Role of BDFE in Homolysis of Strong X−H Bonds via Oxidative PCET. Polar X−H bond in amides and alcohols is a class of strong chemical bonds with BDFE over 8613


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Figure 4. Catalytic alkylation of remote C−H bonds via an intramolecular 1,5-HAT strategy.

strategy is likely also feasible to other analogous ET-motivated atom or group transfer, such as hydride, halogen, OH, CF3, or nitric oxide (NO, NO+, NO−) transfer, etc. Toward this end, substantial research as well as the necessary bond energy measurements are expected.

Besides the N−H systems, analogous thermodynamic guidelines also work well for the activation of strong O−H bonds (BDFE ∼ 102 kcal/mol). In 2016, isomerization of cyclic alcohols into linear ketones was achieved by using the PCET strategy with a photocatalytically generated arene radical cation as the actual oxidant (Figure 5).20e Inspection of various catalyst

3. COMPLEMENTARY PLAY OF BOND-FISSION AND BOND-FORMATION ENERGIES Normally, in a C−H activation study, one would first want to speculate the strength of the target C−H bond, paying much less or no attention to the BDEs of the bonds that are newly formed in reaction, such as C−M or O−H, etc. Neglecting this could result in a serious breach of judgment, although this common practice may be sometimes right for relatively simple reactions under certain circumstances, e.g., in most HAT reactions. However, as contemporary organic chemistry is burgeoning so fast and becoming much more complicated than 10 to 20 years ago, this would likely bring wrong images to one’s understanding of the reaction. As will be introduced in this section, the newly formed bonds often play significant or even more important roles in affecting the consequence of reaction. To exemplify why a single-parameter energetic analysis is often inadequate to untangle the puzzles observed in experiment and the hero behind has to be properly credited, we selected two typical cases in this regard. One is on catalytic bond transformation by transition metal complex to reveal the role of M−C bond. The other is on metal-oxo mediated biomimetic HAT reaction, where the effect of formation of an O−H bond on switching the chemo-selectivity for productive vs nonproductive pathways is identified. In both cases, the relatively greater increase of bond strengths of the newly formed ones upon structural modification vs the C−H bonds of concern is proven to be responsible for the experimental observations beyond intuition. 3.1. Linear Correspondence of M−C BDE with C−H BDE. In metal-mediated C−H bond activation processes, if the substrate bears both aromatic and aliphatic components, the functionalization is observed to occur mostly, if not entirely, at a stronger aromatic C−H bond rather than at a weaker aliphatic C−H site of the side chain.26 This is rationalized on the basis of the knowledge for the trend of the M−C bond strength variation against that of the C−H bond, a series work done primarily by Jones27 (see below), and immediately points out the crucial importance of M−C bond energy study. Determination of the gas phase M−C BDEs can be traced back to mid-20th century.28 However, most of the presently known M−C BDEs are referred to average values based on the

Figure 5. Catalytic ring-opening of cyclic alcohols by PCET activation of strong O−H bonds.

combinations (radical cations and bases) with BDFE ranging from 77 to 105 kcal/mol indicated that only the pairs with BDFE near or over the threshold value (102 kcal/mol) could homolytically cleave the O−H bond to generate the expected product. The energetic governing roles in the selection of catalyst or substrate addressed above suggest that the BDFE can serve as a very useful tool to predict the competence of PCET. However, activations of N−H bonds are now mainly restricted to aryl amides and sulfamides, leaving alkyl amides and amines untouched. Since the N−H BDFEs of the latter substrates are comparable to those of the former, employment of the PCET maneuvers may potentially be amenable to other N−H and O− H systems. One more problem to consider is the exploration of new PCET combinations with suitable BDFEs to allow the rate of radical generation to be attenuated to a moderate level in order to make intermolecular transformations more accessible (vs intramolecular) and manipulatable. Besides the PCET that demonstrates the power of a combinatorial use of energetic parameters from blended catalysts in analyzing organic reactions, a reasonable question may arise in one’s mind. That is, would this two-dimensional strategy be applicable for reaction systems analogous to but not the PCET as well? It appears to us no obvious reason to say NO. Indeed, as the PCET can actually be thought as electron-coupled proton-transfer (i.e., ECPT) due to a higher energetic competency of ET vs PT in most circumstances, it should be logical to anticipate that, by delicate mixing of the relevant energetic quantities of a properly blended catalyst system, this 8614


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correlation lines for relative Rh−C BDEs of Tp’Rh(CNneopentyl)RH with the corresponding C−H BDEs.35 The solid line with blue squares is the LFER for parent hydrocarbons and the dashed line with red triangles for the αsubstituted substrates. Both plots give an identical slope (ρ = 1.4) with a 7.5 kcal/mol vertical energetic gap in between. Similar trends in LFER analyses were exhibited also with the M− C complexes centered at other transition metal system.33 By virtue of the clues extracted from these LFERs, at least two simple but quite useful guiding principles for handling reactivity and selectivity issues concerned in C−H bond activation can be illustrated: 1) The ionicity (polarity) of a bond has dominant impact to M−C bond strength whereas repulsive (steric) effect of a C-centered ligand comes to be a secondary factor. This renders the overall influence of such ligand on the relative M−C BDE to generally follow an order as alkynyl > aryl > alkenyl > primary alkyl > other alkyl. 2) The LFER slopes of >1 indicate greater sensitivity of M− C bond to the changes of C-centered ligand than that of a C−H bond. This means that one can make use of such ρ value, jointly with the order above, for choosing metal complex that allows offsetting the energy loss in breaking the C−H bond of interest. From the above, it is obvious that the significance of solution M−C BDE is increasingly pronounced in modern organic chemistry. Despite the remarkable advances in relative M−C BDE investigations, direct comparison of the stability of M−C complexes that bear the same C-ligands but attaching to different core metals cannot be made using these relative bond values. This, unfortunately, prevents selecting a suitable metal across different transition metals for a more predictable rational design of synthesis and for acquiring the knowledge that is necessary to make wiser use of the intrinsic properties of different types of metals with adjustable ligand structure. However, all these necessities can be achieved if the absolute M−C BDEs could be precisely and adequately determined in solution. More importantly, the absolute BDEs can serve as the benchmarks in the design or calibration of standard computation models for more readily access of other absolute BDEs that are experimentally inaccessible. In the context of the rapidly growing demand of M−C bond energies for understanding the countless transformations involving pivotal roles of metal complexes, substantial endeavors need to be devoted to this area. Appropriate experimental protocols should be developed to enable reliable measurements of the absolute M−C BDEs. 3.2. Vital Role of O−H pK a in Understanding Preferential Oxidation of Strong C−H Bond by Cytochrome P450 While Remaining the Protein Itself Intact. Besides the transition metal systems addressed above, in highvalent metal-oxo mediated HAT reactions, the energy issue of bond formation should also be noted. A typical example is given below to elaborate the critical role of pKa of the newly formed O−H bond in Cytochrome P450 (i.e., P450) for keeping the fragile protein framework from damage during its oxidative cleavage of a strong C−H bond. It is well-known that thiolate ligated P450 can smoothly oxygenate inert hydrocarbons at physiological temperature, leaving its relatively fragile proteins intact.36 To understand this, qualitative explanations, such as by taking specific factors like the strong axial electron donation of thiolate ligation37 or a kinetic

total energies required to cleave all the similar M−ligand bonds in homoleptic metal complexes with simple ligands, e.g., CO, halogen. Hence, the average BDE is, of course, not the same as the energy of the very first bond dissociation, which should be denoted the exact intrinsic absolute M−C BDE. On the other hand, in the condensed phases, the early interests in determination of M−C BDEs were provoked from the Co−C bonds in coenzyme B 12 and its models by kinetic, 29 equilibrium30 or calorimetric31 methods, due to their biological importance. These works are quite remarkable, for they supplied reasonably good access to estimate the upper or lower limits of M−C BDEs. However, the confidence level of the derived bonding data was not high, e.g., experimental BDEs of a particular Co−C bond by different methods may span a range of 25 to 35 kcal/mol,29c−f substantially greater than the error level of ≤3 kcal/mol estimated by Bordwell15 for the solution R−H BDEs. Due to the difficulties in determining absolute M−C BDE in solutions, relative values from experiments and computations were utilized in normal applications of such data,32 and so as in this subsection. As conveyed previously, the C−H BDEs are usually observed to linearly correlate with reaction rates in metal-oxo mediated HAT reactions.8 Linear free-energy relationships (LFERs) of relative M−C vs C−H BDEs were also frequently found in experimental and computational thermodynamic studies of transition metal mediated C−H activation, often with slopes >1.32,33 This implied that the energy gain from M−C bond formation should usually be more than offset of the energy loss from C−H bond scission. Besides thermodynamic analyses, linear dependences of kinetic barriers with those BDEs were also disclosed.34 This, again, demonstrated a greater importance of the M−C bond vs C−H bond in C−H activation, further reminding of the implication that, in many cases, care must be taken while neglecting the energy of bond formation when applying the bond energetic criteria. Jones et al. have done a great deal of work on examining the LFERs between M−C and C−H BDEs for many metal complexes by kinetic methods. Figure 6 shows one of the typical observations of this kind, presenting two distinct

Figure 6. Plots of relative experimental M−C BDE vs C−H BDE. The solid line is fit to the hydrocarbons and aliphatic nitriles −(CH2)n−CN (n = 2−5) (blue ■), and the dashed line is fit to the −CH2X substrates and −CHF2 (red ▲). (Adapted with permission from ref 35. Copyright 2013 American Chemical Society.) 8615


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Journal of the American Chemical Society preference of P450 for C−H bond38 into considerations, were proposed. However, a quantitative understanding of the governing factors for this selectivity was only achieved until quite recently in 2013, attributed to the pKa determination of the Iron(IV)hydroxide in P450.39 Figure 7 illustrates the two competing reaction paths for P450 to possibly take. The one on the right is the productive path

bond functionalization and many other metal-oxo mediated processes are also complex systems, which may bear similar characters as that in heme analogues. The potential to utilize this strategy learnt from nature (P450) to design new syntheses remains to be explored.

4. ACIDITY SCALE AS A POWERFUL TOOL TO REVEAL INSIGHTS BEHIND THE DATA From the preceding sections and a recent review in this aspect on C−H activations,8 one may have gathered a fresher impression on the importance of these basic bond data, like the pKa addressed above, in promoting chemistry to become a rational science, although oftentimes a joint application with other key elements may be necessary. Needless to say, pKa is definitely the most significant and also the best accessible energetic parameter for understanding organic transformations in solution, not to mention derivation of other highly important bond data in organic frontiers, such as the BDEs (or BDFEs), the BDEs/pKas of radical cations/anions, etc., has to be originated from the precise experimental determinations of pKa. Despite the amazing facts above and that in the past, some frustrating problems hampering a healthier development of its research and application have to be better handled. The most noteworthy one among others is how to avoid choosing wrong pKa values from a database that is actually unsuitable for one’s particular research. This is, in fact, the most frequently seen mistake for many people who are not familiar with the pKa techniques and with reasonably unscrambling the solvation effect on pKa variation. Hence, in this section, we will first present a recommended guideline for organic chemists to properly use the existing acidity data. Then, we explicate our initial purpose to work on setting up the pKa scales in ionic liquid (IL) and also showcase the revived power of this traditional term to disclose the key insights behind the data. 4.1. About pKa Values Measured in Molecular Media and Advisable Use of the Data of Different Sources. The pKa, i.e., equilibrium acidity, describes the Gibbs free energy of heterolytic bond dissociation of a Brönsted acid (HA) in a solvent (Scheme 2).

Figure 7. Productive and nonproductive pathways for P450 oxidations.

which expresses a HAT between iron(IV)oxo (P450-I) and substrate RH (alkyl) to generate an R• radical and iron(IV) hydroxide (P450-II). Its energetic driving force (ΔGp) is reflected by the difference of the C−H and O−H (newly formed) BDEs. As seen, the new O−H bond does affect the productive process, whose energy can be evaluated from the aqueous pKa of P450-II and reduction potential of P450−I (E°I vs NHE) from eq 2 (a derivative of eq 1). On the other hand, the nonproductive path depicted on the left side of Figure 7 represents an electron transfer (ET) between P450-I and tyrosine or other enzymatic residues, whose thermodynamics (ΔGnp) is the difference of their redox potentials. Hence, the thermodynamic tendency for P450-I to undergo an HAT or ET (ΔGrel) can be reflected by the difference between ΔGp and ΔGnp as in eq 3 (E°tyr for redox potential of protein residues). ΔGp = BDE(C−H) − BDE(O−H) = BDE(C−H) − 1.37pK a(O−H) − 23.06 E I° − 57.6

(2)

ΔGrel = ΔGp − ΔGnp ° − 57.6 = BDE(C−H) − 1.37pK a(O−H) − 23.06 Etyr

(3)

From the existing information of hydrocarbon BDE(C−H) of ∼101 kcal/mol,4 P450-II pKa of 11.9 (or 16.3 kcal/mol) and E°tyr of ∼1.0 V (or 23.06 kcal/mol),37 a ΔGrel value of about 3 kcal/mol in eq 3 can be derived. Although by looking at this value alone the HAT path for P450-I seems slightly unfavorable, taking the synchronous kinetic effect into account, an overall trend for P450 in favor of the productive HAT path should be realized. This is due to a huge deduction to the ET rate of P450-I brought in by the thiolate ligation to replace a histidine ligation as in other type of peroxidases, as anticipated from a kineticthermodynamic relation shown in the Marcus equation.7 Taking the histidine ligated HRP-II (horseradish peroxidase, pKa ≤ 3.5)40 as an example, one can realize an over 104-fold reduction of the nonproductive ET rate for P450-I compared to that for HRP from Marcus equation on the basis of a pKa difference of ≥8.4 (or ≥11.5 kcal/mol) for these two heme proteins. This, acting as a final adjusting weight, eventually finishes up with an obvious superiority for P450 to be able to activate an inert hydrocarbon C−H bond and at meantime avoids an oxidative degradation of its surrounding protein frameworks. As demonstrated above, the iron(IV) hydroxide pKa can provide a simple criterion for assessing the capacity of heme analogues to selectively activate hydrocarbons in the presence of other electron reductants. Generally, in organic syntheses, C−H

Scheme 2. Acidic Dissociation in the Gas Phase and in Solution

About “Absolute Acidity” in Solution. Since most organic reactions are carried out in solution via polar transition states, numerous acidity scales have been established in various neat or mixed solvents in an attempt to quantitatively understand the insight behind. As a result, collection of the solution phase acidities has far exceeded that in the gas phase and is now the most available bond energetic database used in chemistry.4b In tradition, the gas phase heat of deprotonation can be called absolute acidity, for it reflects the intrinsic nature of the A−H bond (Scheme 2).41 However, in the solution acid−base chemistry, if the derived pKa is referenced to the particular solvent of measurement as the standard, then this value can be 8616


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course, most pKa data in DMSO, AN, and water (in a pH range of 2−12)42−44 can be trusted in this respect. General Methodology. The past century has witnessed a considerable ingenuity of devising various methods to study the acid−base equilibria in various media, mainly including spectrophotometric methods and electrochemical approaches, such as potentiometry, conductometry and voltammetry. Several occasionally used techniques by virtue of kinetics, calorimetry, NMR, etc., are also known. The methodology for pKa determination also evolves with the advance of technology. Some new approaches, such as with capillary electrophoresis (CE),48 high performance liquid chromatography (HPLC),49 etc., were developed which are helpful to streamline generation of pKa data under certain circumstances. However, it is not the goal of this Perspective to address the details of these methods. There are several (but not limited to) excellent reviews50 and monographs42,46,51 specifying on these. Acidity Scales Recommended for Organic Chemists. Although there are many pKa scales in various solvents and also a plethora of methods for acidity determination to choose from, how to find the appropriate pKa values suitable for one’s research from the literature is not so straightforward and often ends up in confusion. One of the aims of this subsection is to provide some information for readers to reasonably sort out the required data from sprawling sources. Most organic reactions are carried out in nonaqueous solvents, due primarily to the solubility issue and the need for manipulating the water-sensitive species/intermediates (e.g., carbanion) most commonly used in organic synthesis. The wellknown acidity scale in water is, however, not particularly suitable for studying the problems in organic solvents, especially when a C−H or C−C bond-breaking is involved, because the reliable C−H pKa values in water are surprisingly inadequate (no more than one to two hundreds, mostly for α-EWG substituted C−H acids) which are far less than the known C−H pKas in DMSO.2 Moreover, the former also corresponds poorly with the latter due to differential solvation by the uneven hydrogen bonds in water. Therefore, except for certain special cases, the aqueous C−H acidity values reported in early literature especially before the blossom of various nonaqueous acidity scales around mid- to late 20th century are not recommended for modern research. The most widely used equilibrium acidities in dipolar aprotic solvents come almost entirely from the pKa scales in DMSO43 and AN,44 with >3000 and 1500 values determined in these two organic solvents, respectively.4b In general, DMSO is more suitable for studying weak acids, such as carbon (C−H) acid and nitrogen (N−H) acid, whereas AN more suitable for stronger acids because its basicity is weaker. However, they may complement each other. In addition, the span of the measurable pKa ranges in DMSO and AN are very broad, about 2 to 32 and 2 to 30,52 respectively, both are much wider than the best measurable pH range of 2 to 12 in water. This is beneficial for organic chemists in both synthetic and theoretical study. Most pKa data in these two scales were determined by the indicatoroverlapping method (IOM) with UV−vis, as described in the milestone paper of Bordwell.53 Scheme 3 (vide supra) depicts the routine of this approach. The most remarkable advantage of the spectrophotometric method compared to most others is that the measurement can be conducted at very low concentrations (∼10−4−10−5 M) of the target acid and the indicator base, and hence effectively avoid the annoying interferences from ionpairing, homo/heteroassociation effects (Scheme 4), etc. to the pKa value and thus warrants its genuine identity. Bordwell’s pKa

regarded as absolute acidity by convention, and the scale built up by the data of this kind can be called an absolute acidity scale. The pKa scales in water,42 DMSO,43 acetonitrile (AN),44 etc., meet this criterion, and so, belong to this category. The prerequisite for a scale to be called absolute requires that all the data must have a single reference state; that is, the solvent itself acting as the only base. To illustrate, a schematic diagram is given below (Scheme 3), where HA′ refers to an anchor acid Scheme 3. Autoprotolysis of Anchor acid HA′ in Solvent S and the Principle of Indicator Overlapping Method (IOM)a

a

The activity coefficients in these equations are assumed as unity.

whose pK′HA′ is measured in an absolute sense against the solvent S as the reference base, HIn and AH represent an indicator and a target acid, respectively. It should be emphasized that the equations in this section are expressed entirely on the basis of thermodynamics but not meant to imply mechanism (for ion-association complexes formed during reaction process, please refer to the discussions in the literature45). Besides the spectroscopic determination of autoprotolysis pKa for anchor acid showed above, absolute acidity can also come electrochemically by measuring pH of the solution with a control of ionic strength and the use of various buffers (e.g., in water).42,46 However, due to a much slower electrode response in most aprotic organic media and a greater uncertainty to acidity therefrom,47 the latter is seldom used now in organic chemistry in general. It should be pointed out that, although the relative data are still usable in some cases, the absolute acidity in solution, especially with high precision, is far more valuable than just showing the acidity orders in the same solvent, which is perhaps the only thing that a relative scale can do. To exemplify: (1) an absolute scale allows a quantitative judgment on, among different media, which solvent is more powerful for deprotonation than others, a question usually hard to answer by other means (e.g., according to pKa values, IL is more powerful than AN, but solvent polarity tells the opposite); (2) it provides bond energy benchmarks that are vitally useful for rational design or calibration of computation models for solvation and bond transformation in a particular solvent; (3) it is critically important to identify the subtle factors for some key issues, such as the effects of ion-pairing, homo/hereto-association, etc., which are easily overlooked but crucial for rationalization of experimental results, either with respect to the solvent or to the substrate structure. Moreover, if an acidity is reported “absolute” and also is verified not being spoiled by ion-association (e.g., ionpairing), it represents a “true pKa” which should be the most recommended term for people to use in their related work rather than taking a relative or apparent acidity, whenever avoidable. Of 8617


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graphs42,51 and databases4b,63 have been compiled for the pKa values as well as the methods of pKa measurement in water. Due to the leveling effect, about 95% aqueous pKa values are between 2.0 to 12.0.50e The pKa values falling in this range were precisely determined and are trustable, although the majority of them are for oxyacids and the conjugated acids of nitrogen-containing bases, along with a small handful of thiols. The most direct and convenient approach to an acidity in water is to measure the pH of the solution and calculate the pKa from the Henderson− Hasselbach equation. However, for weak acids, such as N−H acids and carbon acids, as addressed above, the measurement could be difficult, and is not applicable for acids with a pKa > 12.0. Although the related information for very weak acids may be achieved kinetically (i.e., the so-called kinetic pKa),50a,64 a comment is not intended in this Perspective focusing mainly on equilibrium acidity. One last but important reminder needs to emphasize is that the pKa values obtained in water do not necessarily correlate well (though occasionally do) with the corresponding data in DMSO or other aprotic solvents.65 For one such example, phenol is found substantially less acidic than malononitrile in DMSO (pKa: 18.0 vs 11.1), but in water, it is more acidic than the latter (pKa: 9.99 vs 11.12). This is mainly caused by an uneven H-bonding solvation of water as compared to the aprotic DMSO. As a consequence, the aqueous acidity scale often cannot reflect exactly the intrinsic order of acidity for many compounds measured in this solvent.65 4.2. pKa Scales in Ionic Liquids and Implications Behind Energetic Data. In the above, the current status of the pKa studies in the molecular media was concisely summarized. In the following, we will brief the recent progresses in equilibrium acidity study in ionic media to further illustrate the necessity of these fundamental data in disentangling the problems in ionic liquid (IL) chemistry that would otherwise not be solved with the available information from molecular solvent. The room-temperature ionic liquids (RTILs, or abbreviated as ILs), known as a rising mainstream category of solvent system composed of all ions, are conceptually and structurally distinctive from those of molecular solvents.66 During the past 2 decades, ILs have been extensively employed in chemical, biological and industrial processes.67 However, compared to the tremendous attention paid to the aspects of practical importance, relatively little is known about the chemical respects of theoretical significance, such as how would the composition of the ions of ILs alters the energies of bond dissociation of the substrates so as to change their chemistry, etc. To answer fundamental questions like this, the pKa scale is naturally the first thing to consider according to the experiences learned from the chemistries in molecular media.

Scheme 4. Homo/heteroassociation Complication between Dissociated and Undissociated Species

scale in DMSO is most generally accepted in organic chemistry largely for this reason.43 The UV−vis indicator method has also been applied to building the acidity scale in AN by Leito.44c−e Significantly, in contrast to other existing means, this method allows the UV−vis sensitive species in a reaction, i.e., the acid and base interacted equilibrium, to be directly “seen” by monitoring the changes of their absorptions. This feature can immediately tell whether an unwanted specific ion-association, like the one shown in Scheme 4, is affecting the measurement or not, so that to remind one if a calibration is necessary.54 It should be pointed out that, though the electrochemical method (e.g., potentiometric titration) is also used to derive pKa values of weak acids in organic media especially at early times,47 it suffers significant decay of accuracy caused by much slower electrode response (vs in water) and/or a demand to use high concentrations (∼10−1−10−3 M) for weakly acidic substrate and the electrolyte that are inherently needed for detecting meaningful electrode signals. Nonetheless, for relatively strong acids, such limitation may be managed to an acceptable level and allows electrochemistry to be conditionally used to obtain the required information with high efficiency.46 We also should mention, with great respect, that there have been many well renowned scientists, to name a few: Wheland,55 McEven,56 Streitwieser,57 Breslow,58 Arnett,59 Ritchie,60 Kolthoff,50d,54b et al.,61 who established their “relative” and/or “ion-pair” C−H acidity scales in aprotic media in early years. Hammett also established an acidity function H_ to describe the relative basicity of weak acids in alcohol solution of DMSO, and then extended to almost pure DMSO.61,62 Despite that these scales are all relatively small and their predictive power is no longer as strong as before and therefore would not be highly recommended in this Perspective for in-depth theoretical analysis, these milestone works indeed greatly advanced chemistry to the current stage and are definitely worthy of our high tribute. About Acidity Scale in Water. The aqueous pKa scale is no doubt the largest database of equilibrium acidities until present, and hence, is very useful for people, including organic chemists, whose research is primarily conducted in aqueous phase and focused on heteroatom (especially O, N)-centered substrates. Over the past about 100 years, a tremendous amount of aqueous pKa has been measured and a wealth of mono-

Figure 8. Structure of the aprotic and protic ILs used for the acid−base equilibrium study. 8618


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Figure 9. Double-line Hammett plots for the pKa values of benzenethiols in 4 AILs. (Adapted with permission from ref 71e. Copyright 2014 American Chemical Society.)

the same line with other electron-withdrawing groups (EWGs) featured by having either π- or a pair of p-electrons. This is because that the IL cation (imidazolium or pyrrolidinium), while solvating the thianion at the sulfur site as conventionally seen in molecular solvent, can also strongly solvate the paraEWG by the so-called CSAR (i.e., Cation-solvation-assisted Resonance) effect,71e as depicted in eq 4. The CF3 group is unable to facilitate such an extended resonance in solvation with IL by this mechanism; therefore, it is biased from the EWG line. Obviously, this IL effect would not possibly be identified without the precisely determined pKa data.

Determination of the pKa values in ILs is more challenging than in molecular solvents mainly because of their high viscosity and hygroscopic nature. Despite this, a few pioneer studies on the absolute pKa values of some N−H acids in neat ILs were carried out by using the electrochemical method.68 However, the uncertainties of the derived data were reported large (0.4− 1.0 pK unit) due to the inherent limitations of the methods.68 In this regard, the indicator method (IOM) should be the best choice (vide supra). In recent years, we have carried out a systematic study on the absolute pKa values by the IOM in pure ILs based on our expertise in pKa measurement in DMSO.69 Several most commonly used aprotic (AIL) and protic (PIL) ILs were selected as “standard ILs” for the acidity study purpose (Figure 8) and were used under the “standard conditions” outlined in our initial work.70 Absolute acidity scales of a series of C−H, O−H, N−H, N+−H and S−H acids which cover a broad acidity span (∼20 pK units) and now include more than 500 pKa values for about 200 structurally different organic compounds were established in the AILs and a PIL with high precision (uncertainly ≤ ±0.05 pK units).70,71 These basic bond energetic data, serving as a powerful tool, enabled a number of key concerns in IL chemistry to be rationalized as exemplified below. “Ionic Liquid Effect”. One of the fundamental aspects regarding the nature of ILs is their microscopic solvation behaviors on respective entities in IL solutions. The relevant research in this connection is mainly restricted to theoretical approaches based on the knowledge from molecular media.72 ILs may show unique solvation effects under certain circumstances that are not likely to occur in molecular solvent, which is labeled as “IL effect”,73 and are normally subtle and difficult to detect. In this connection, the precisely measured pKa scale in AILs is an ideal tool to screen out these characteristic solvation behaviors of ILs. For examples, contrary to the scenarios observed in molecular solvents, the Hammett plots of the pKa values of para-substituted benzenethiols in all the four imidazolium and pyrrolidinium based AILs demonstrate a distinct double-line relationship with identical slope and excellent linearity (Figure 9), presenting a typical “ILs effect”. Note that, in all these four lines, the CF3 group does not fall in

Does Ion-Association of Solute Ions Occur in IL Solution? It is known that a salt dissolved in a solvent of low polarity can form various modes of ion-associations (e.g., ion pairs).74 As such, what would be the situation for a salt dissolved in an IL solvent which itself is a “liquid salt” consisting of the charged cations and anions? This should be of a necessary concern in a serious pKa study because, if ion-pairing occurs (Scheme 5), it would result Scheme 5. Ion-Pairing Complication in pKa Determination

in an “apparent pKa” (or called ion-pair pKa) rather than the desired true pKa (or “free-ion pKa”), which is much superior to the apparent one for profound analysis of the structure− property relationships. Studies for ILs revealed that their polarity is generally low to moderate according to their dielectric constants (ε = 10.0−15.0 8619


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Journal of the American Chemical Society for AILs and 20.0−25.0 for PILs),75 which are substantially lower than that of the aprotic dipolar molecular solvent DMSO (ε = 46.5)76 and AN (ε = 36.1),77 the favorite media for acidity studies of organic compounds. If the general rules for ionassociation found in molecular media in light of dielectric constant also stand in ionic liquid,78 it would suggest an occurrence of ion-association in ILs, which is deleterious for obtaining a “true pKa”. To ensure a high-fidelity absolute pKa to be measured, it is necessary to examine the ion-pair effect in ILs under the IOM conditions. The results of the experiments specially designed for this purpose are presented in Table 1,

BDE in ILs is essentially absent. This is mainly due to the unusual complexity in handling solvation quantities in ILs, and also, to a severe dearth of precisely determined bond parameters to serve as authentic benchmarks for modeling the methodologies. In this regard, our experimental pKa values in ILs71 should be beneficial to serving as the standard references for calibrating computational methods. By taking advantage of the IL pKa values in hand, we have developed a model named as “ion-biased cluster-continuum model (IB-CCM)” for computing the solvation free energies for neutral and ionic species in ILs, and then, have successfully applied it to the calculations of absolute pKa values.79 In this so-called IB-CCM model, certain numbers of the “explicit” cation or anion of IL were employed to modulate the inner solvation shell of the cluster-continuum model80 to quantum mechanically describe the specific solute− solvent interactions of the proton and the conjugate anion of the acid with the IL anions and cations, respectively. The rationality of designing the IB-CCM was rooted in experimental70,71c and theoretical findings,81 where the IL cation (e.g., [Bmim]+) and its counteranion (e.g., [NTf2]−) do not strongly associate with each other to form long-lived ion pairs, which allowed a separate use of the IL cation or anion (instead of using them as a whole) in modeling the inner sphere of the solvation model. Application of the IB-CCM model to compute experimental pKa values of a serial of common acids in [Bmim][NTf2] gave a mean unsigned error of ≤0.5 pK units,79 which is better than or at least equally good as that of the best pKa calculations in the conventional molecular solvents.51f,82 Notably, use of other methods originally designed for computing pKa values in molecular solvent all rendered quite poor modeling. Even Truhlar’s SMDGIL (generic ionic liquid solvation model),83 which is the best among all others, still led to reproduction of the experimental pKa values with marginally satisfactory errors (>2 pK units).68f This reflected an inherit limitation of the continuum solvation models for the calculation of bond energies in ionic media, despite that the SMD-GIL was very valuable for predicting solvation free energies of neutral species elsewhere.83 Results from further applications of the calculated solvation free energies by the IB-CCM are also quite encouraging: e.g., (1) the intriguing inversion of the pKa orders between some common acids (e.g., benzoic acid and thiophenol) in molecular solvent DMSO and in IL [Bmim][NTf2] was well reproduced by computation; and (2) the weaker acidity of the acids in an environment full of the charged particles (i.e., ILs) than in DMSO can also be well explained based on computation. Moreover, the theoretically estimated solvation free energy of proton in [Bmim][NTf2] (255.8 kcal/mol, a predictive average value from our work79) has more recently been verified by the experiment of Matsubara et al. (255.3 ± 2.2 kcal/mol in [Emim][NTf2].84 The amazing agreement between theory and experiment ( 32), which is corroborated by a large discrepancy found for the reported autoprotolysis constant, pKauto = 33.3−44.0, from different groups. For details, see: (a) Schwesinger, R.; Schlemper, H. Angew. Chem., Int. Ed. Engl. 1987, 26, 1167−1169. (b) Kolthoff, I. M.; Chantooni, M. K., Jr J. Phys. Chem. 1968, 72, 2270−2272. (53) Matthews, W. S.; Bares, J. E.; Bartmess, J. E.; Bordwell, F. G.; Cornforth, F. J.; Drucker, G. E.; Margolin, Z.; McCallum, R. J.; McCollum, G. J.; Vanier, N. R. J. Am. Chem. Soc. 1975, 97, 7006−7014. (54) (a) Bordwell, F. G.; McCallum, R. J.; Olmstead, W. N. J. Org. Chem. 1984, 49, 1424−1427. (b) Kolthoff, I. M.; Chantooni, M. K.; Bhowmik, S. J. Am. Chem. Soc. 1968, 90, 23−28. 8622


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