Archive for the 'Kass' Category

Is the cyclopropenyl anion antiaromatic?

The concept of antiaromaticity is an outgrowth of the well-entrenched notion or aromaticity. While 4n+2 π-electron systems are aromatic, 4n π-electron systems should be antiaromatic. That should mean that antiaromatic systems are unstable. The cyclopropenyl anion 1a has 4 π-electrons and should be antiaromatic. Kass has provided computational results that strongly indicate it is not antiaromatic!1

Let’s first look at the 3-cyclopropenyl cation 1c. Kass has computed (at both G3 and W1) the hydride affinity of 1c-4c. The hydride affinities of the latter three compounds plotted against the C=C-C+ angle is linear. The hydride affinity of 1c however falls way below the line, indicative of 1c being very stable – it is aromatic having just 2 π-electrons.

A similar plot of the deprotonation enthalpies leading to 1a-4d vs. C=C-C- angle is linear including all four compounds. If 1a where antiaromatic, one would anticipate that the deprotonation energy to form 1a would be much greater than expected simply from the effect of the smaller angle. Kass suggests that this indicates that 1a is not antiaromatic, but just a regular run-of-the-mill (very) reactive anion.

A hint at what’s going on is provided by the geometry of the lowest energy structure of 1a, shown in Figure 1. The molecule is non-planar, having Cs symmetry. A truly antiaromatic structure should be planar, really of D3h symmetry. The distortion from this symmetry reduces the antiaromatic character, in the same way that cyclobutadiene is not a perfect square and that cyclooctatraene is tub-shaped and not planar. So perhaps it is more fair to say that 1a has a distorted structure to avoid antiaromaticity, and that the idealized D3h structure, does not exist because of its antiaromatic character.

Figure 1. G3 optimized geometry of 1a.


(1) Kass, S. R. "Cyclopropenyl Anion: An Energetically Nonaromatic Ion," J. Org. Chem. 2013, 78, 7370-7372, DOI: 10.1021/jo401350m.


1a: InChI=1S/C3H3/c1-2-3-1/h1-3H/q-1

1c: InChI=1S/C3H3/c1-2-3-1/h1-3H/q+1

Aromaticity &Kass Steven Bachrach 20 Jan 2014 2 Comments

Hydrogen-bonded-assisted acidity

Can a hydrogen bonding network affect acidity? Kass has examined the polyol 1 whose conjugate base 1cb can potentially be stabilized by a large hydrogen bonding network.1 Kass had previously found a significant acidy enhancement in comparing t-butanol (ΔG(deprotonation) = 369.2 kcal mol-1) with that of 2G(deprotonation) = 334.4 kcal mol-1).2

Table 1 lists the computed and experimental free energies of deprotonation of 1. The experimental values are computed at M06-2x/maug-cc-pVT(+d)Z. The structure of 1cb is drawn in Figure 1.

Table 1. Computed and experimental free energies (kcal mol-1 of deprotonation of some alcohols.













The difference in the acidity of t-butanol and 2, some 30 kcal mol-1, reflects the stability afforded by three intramolecular hydrogen bonds to the oxyanion. In going from 2cb to 1cb, each of the hydroxyl groups that donate to the oxyanion act as the acceptor of a hydrogen bond from the more removed hydroxyl groups. There is in effect a first and second layer of hydrogen bond network in 1cb. These secondary hydrogen bonds lead to further stabilization of the anion, as reflected in the diminished DPE of 1 over 2: 320.2 vs. 335.0 kcal mol-1. Note that this secondary layer does not stabilize the anion to the same degree as the primary layer, but nonetheless its effect is large and quite striking.


Figure 1. M06-2x/maug-cc-pVT(+d)Z optimized structure of 1cb.

Even in solution these more remote hydrogen bonds can stabilize the anion. So, using the CPCM approach and modeling DMSO, 2 is predicted have a pKa that is 15 units below that of t-butanol, and 1 is predicted to be 3 pKa units more acidic than 2. Experiments verify this prediction with the pKas of 16.1 for 2 and 11.4 for 1.


(1) Shokri, A.; Abedin, A.; Fattahi, A.; Kass, S. R. "Effect of Hydrogen Bonds on pKa Values: Importance of Networking," J. Am. Chem. Soc. 2012, 134, 10646-10650, DOI: 10.1021/ja3037349

(2) Tian, Z.; Fattahi, A.; Lis, L.; Kass, S. R. "Single-Centered Hydrogen-Bonded Enhanced Acidity (SHEA) Acids: A New Class of Brønsted Acids," J. Am. Chem. Soc. 2009, 131, 16984-16988, DOI: 10.1021/ja9075106


1: InChI=1S/C13H28O7/c14-4-1-10(17)7-13(20,8-11(18)2-5-15)9-12(19)3-6-16/h10-12,14-20H,1-9H2

1cb: InChI=1S/C13H27O7/c14-4-1-10(17)7-13(20,8-11(18)2-5-15)9-12(19)3-6-16/h10-12,14-19H,1-9H2/q-1

2: InChI=1S/C7H16O4/c8-4-1-7(11,2-5-9)3-6-10/h8-11H,1-6H2

2cb: InChI=1S/C7H15O4/c8-4-1-7(11,2-5-9)3-6-10/h8-10H,1-6H2/q-1

Acidity &Hydrogen bond &Kass Steven Bachrach 28 Aug 2012 1 Comment

Acidity of lithium acetylide

I have had a long-standing interest in organolithium compounds, dating back to my graduate student days. Thus, I was excited to read Kass and Radom’s latest work on the computational and experimental evaluation of the acidity of lithium acetylide LiCCH.1

The gas phase experimental acidity is accomplished by preparing the conjugate base of lithium acetylide through a procedure of collision-induced dissociation with loss of CO2, as in Scheme 1. By reacting this anion with a variety of different acids, they were able to bracket the acidity and determine that ΔHacid is 391.0 ± 1.3 kcal mol-1. This is about 13 kcal mol-1 less acidic than acetylene itself. The reduction is acidity understandable in terms of the C-Li being essentially ionic, and thereby loss of the proton builds up negative charge on a carbon adjacent to a carbon that already has a great deal of negative charge.

Scheme 1

Computations support this enthalpy for deprotonation. The G3, G4 and W1 values for the enthalpy deprotonation of lithium acetylide are 389.1, 388.9, and 390.4 kcal mol-1, respectively. It should also be noted that the conjugate base of lithium acetylide posses a non-classical bridging geometry 1, which is well-known for organolithium species.2


(1) Meyer, M. M.; Chan, B.; Radom, L.; Kass, S., R., "Gas-Phase Synthesis and Reactivity of Lithium Acetylide Ion, LiCC-," Angew. Chem. Int. Ed., 2010, 49, 5161-5164, DOI: 10.1002/anie.201001485

(2) Streitwieser, A.; Bachrach, S. M.; Dorigo, A.; Schleyer, P. v. R. In Lithium Chemistry: A Theoretical and Experimental Overview; Sapse, A.-M., Schleyer, P. v. R., Eds.; Wiley-Interscience: New York, 1995, p 1-45.

Acidity &Kass Steven Bachrach 18 Oct 2010 No Comments

Acidity of remote protons

The α-proton of ketones and aldehydes are acidic, thanks to delocalization of the resulting anion. However, α-protons at a bridgehead position are much less acidic – the resulting anion is not delocalized as the enolate would be an anti-Bredt alkene. So, what about more remote protons from the carbonyl – would they exhibit enhanced acidity due to inductive or field effects?

Kass has examined the deprotonation of 2-adamantone 1 via experiment and computation.1 The relative energies of the five different anions are listed in Table 1. Previous H/D exchange experiments indicate that the relative reactivity is βax > βeq > α, and this is well reproduced by computations.2

Table 1. Relative energies (kcal mol-1) of the enolates of 1.






















Kass’ bracketing experiments indicate the enthalpy for deptrotonation of 2-adamantone is 394.7 ± 1.4 kcal mol-1. This is in nice accord with the computational results for loss of the βax proton: 393.8 (M06-2x/aug-cc-pVDZ) and 396.8 kcla mol-1 (G3). One interesting computational result is a competive cyclic structure 2, whose stability is similar to that to the βax ion at M06-2x and is the optimized structure produced at MP2/6-31G(d) when searching for the βeq enolate.

So, to answer our question, protons remote from a carbonyl are more acidic than alkane
analogues, but much less acidic than typical α-protons of ketones.


(1) Meyer, M. M.; Kass, S. R., "Enolates in 3-D: An Experimental and Computational Study of Deprotonated 2-Adamantanone," J. Org. Chem., 2010, 75, 4274-4279, DOI: 10.1021/jo100953y

(2) Stothers, J. B.; Tan, C. T., "Adamantanone: stereochemistry of its homoenolization as shown by 2H nuclear magnetic resonance," J. Chem. Soc., Chem. Commun., 1974, 738-739, DOI: 10.1039/C39740000738


1: InChI=1/C10H14O/c11-10-8-2-6-1-7(4-8)5-9(10)3-6/h6-9H,1-5H2

2: InChI=1/C10H13O/c11-10-7-2-5-1-6(4-7)9(10)8(10)3-5/h5-9H,1-4H2/q-1

Acidity &Kass Steven Bachrach 17 Aug 2010 1 Comment

Protonation of 4-aminobenzoic acid

Molecular structures can differ depending on phase, particularly between the gas and solution phase. Kass has looked at the protonation of 4-aminobenzoic acid. In water, the amino is its most basic site, but what is it in the gas phase? The computed relative energies of the protonation sites are listed in Table 1. If one corrects the B3LYP values for their errors in predicting the proton affinity of aniline and benzoic acid, the carbonyl oxygen is predicted to be the most basic site by 5.0 kcal mol-1, in nice accord with the G3 prediction of 4.1 kcal mol-1. Clearly, the structure depends on the medium.

Table 1. Computed relative proton affinities (kcal mol-1) of 4-aminobenzoic acid.

C=O 0.0 0.0
NH2 7.9 4.1
OH 12.2 9.8

Electrospray of 4-aminobenzoic acid from 3:1 methanol/water and 1:1 acetonitrile/water solutions gave different CID spectra. H/D exchange confirmed that electrospray from the emthanol/water solution gave the oxygen protonated species while that from the acetonitrile/water solution gave the ammonium species.


(1) Tian, Z.; Kass, S. R., “Gas-Phase versus Liquid-Phase Structures by Electrospray Ionization Mass Spectrometry,” Angew. Chem. Int. Ed., 2009, 48, 1321-1323, DOI: 10.1002/anie.200805392.


4-aminobenzoic acid: InChI=1/C7H7NO2/c8-6-3-1-5(2-4-6)7(9)10/h1-4H,8H2,(H,9,10)/f/h9H

Acidity &amino acids &Kass &Solvation Steven Bachrach 30 Mar 2009 No Comments

Which is the Most Acidic Proton of Tyrosine?

Following on their prediction that the thiol of cysteine1 is more acidic than the carboxylic acid group (see this post), Kass has examined the acidity of tyrosine 1.2 Which is more acidic: the hydroxyl (leading to the phenoxide 2) or the carboxyl (leading to the carboxylate 3) proton?




Kass optimized the structures of tyrosine and its two possible conjugate bases at B3LYP/aug-cc-pVDZ, shown in Figure 1, and also computed their energies at G3B3. 2 is predicted to be 0.2 kcal mol-1 lower in energy than 3 at B3LYP and slightly more stable at G3B3 (0.5 kcal mol-1). However, both computational methods underestimate the acidity of acetic acid more than that of phenol. When the deprotonation energies are corrected for this error, the phenolic proton is predicted to be 0.4 kcal mol-1 more acidic than the carboxylate proton at B3LYP and 0.9 kcal mol-1 more acidic at G3B3.




Figure 1. B3LYP/aug-cc-pVDZ optimized structures of tyrosine 1 and its two conjugate bases 2 and 3.2

Gas phase experiments indicate that deprotonation of tyrosine leads to a 70:30 mixture of the phenoxide to carboxylate anions. The computations are in nice agreement with this experiment. (A Boltzmann weighting of the computed lowest energy conformers makes only a small difference to the distribution relative to using simply the single lowest energy conformer.) This demonstrates once again the important role of solvent, since only the carboxylate anion is seen in aqueous solution.


(1) Tian, Z.; Pawlow, A.; Poutsma, J. C.; Kass, S. R., "Are Carboxyl Groups the Most Acidic Sites in Amino Acids? Gas-Phase Acidity, H/D Exchange Experiments, and Computations on Cysteine and Its Conjugate Base," J. Am. Chem. Soc., 2007, 129, 5403-5407, DOI: 10.1021/ja0666194.

(2) Tian, Z.; Wang, X.-B.; Wang, L.-S.; Kass, S. R., "Are Carboxyl Groups the Most Acidic Sites in Amino Acids? Gas-Phase Acidities, Photoelectron Spectra, and Computations on Tyrosine, p-Hydroxybenzoic Acid, and Their Conjugate Bases," J. Am. Chem. Soc., 2009, 131, 1174-1181, DOI: 10.1021/ja807982k.


1: InChI=1/C9H11NO3/c10-8(9(12)13)5-6-1-3-7(11)4-2-6/h1-4,8,11H,5,10H2,(H,12,13)/t8-/m0/s1/f/h12H

2: InChI=1/C9H11NO3/c10-8(9(12)13)5-6-1-3-7(11)4-2-6/h1-4,8,11H,5,10H2,(H,12,13)/p-1/t8-/m0/s1/fC9H10NO3/q-1

3: InChI=1/C9H11NO3/c10-8(9(12)13)5-6-1-3-7(11)4-2-6/h1-4,8,11H,5,10H2,(H,12,13)/p-1/t8-/m0/s1/fC9H10NO3/h11h,12H/q-1

Acidity &amino acids &Kass Steven Bachrach 04 Mar 2009 2 Comments

Which is the Most Acidic Proton of Cysteine?

Kass has once again uncovered a simple system that challenges our notions of basic chemical concepts. It is a well accepted notion that the most acidic proton of all of the amino acids is the carboxylic acid one. However, acidities are strongly influenced by the solvent, and the absence of solvent in the gas phase can dramatically alter things.

Kass and co-workers examined the gas-phase acidity of cysteine with computational and
experimental techniques.1 The lowest energy conformer of cysteine is 1a, characterized by having three intramolecular hydrogen bonds (Figure 1). The next lowest conformer, 1b, has only two intramolecular hydrogen bonds and is 1.5 kcal mol-1 higher in energy at G3B3.



Figure 1. B3LYP/aug-cc-pVDZ optimized structures of cysteine 1.1

They optimized a number of different configurations of the conjugate base of cysteine: two conformers from the loss of the carboxylate proton (2a and 2b), two conformers from the loss of the thiol proton (2c and 2d), and one conformer from the loss of the thiol proton of the zwitterion (2e). These structures are shown in Figure 2 along with their relative energies. All of these structures possess two intramolecular hydrogen bonds.








Figure 2. B3LYP/aug-cc-pVDZ optimized structures of the conjugate base of cysteine 2. Relative energies (kcal mol-1) in parenthesis computed at G3B3.1

The gas phase acidity of carboxylic acids is greater than thiols; the deprotonation energy of propanoic acid (CH3CH2CO2H) is 347.7 kcal mol-1 at G3B3 (347.2 expt.2), about 6 kcal mol-1 less than that of ethanethiol (CH2CH2SH: 355.0 at G3B3 and 354.2 expt.2). However, the computations indicate that 2c is the lowest energy structure of deprotonated cysteine, and 2c comes about by loss of the thiol proton! Te lowest energy cysteine conjugate base from loss of the carboxylate proton is 1a, which is 3.1 kcal mol-1 higher in energy. Apparently, the hydrogen bonding network in 2c is quite favorable, able to make up for the inherent favorability of a carboxylate over a thiolate anion.

The G3B3 computed deprotonation energy of cysteine is 333.3 kcal mol-1 (for removal of the thiol proton). Kass determined the deprotonation energy of cysteine using a kinetic and a thermodynamic method. The kinetic method gives a value of 332.9 ± 3.3 kcal mol-1­, while the thermodynamic method gives 334.4 ± 3.3 kcal mol-1­. These are in fine agreement with the computed value.

This study ably demonstrates the dramatic role that solvent can play in determining molecular properties. Kass titled the article “Are carboxyl groups the most acidic sites in amino acids?” and answers with “no” – in the gas phase the thiol group is more acidic. He ends the article with an indication that the alcohol of tyrosine may be competitive in acidity with its carboxylic group, too.


(1) Tian, Z.; Pawlow, A.; Poutsma, J. C.; Kass, S. R., "Are Carboxyl Groups the Most Acidic Sites in Amino Acids? Gas-Phase Acidity, H/D Exchange Experiments, and Computations on Cysteine and Its Conjugate Base," J. Am. Chem. Soc., 2007, 129, 5403-5407, DOI: 10.1021/ja0666194.

(2) NIST, NIST Chemistry WebBook, 2005,


1: InChI=1/C3H7NO2S/c4-2(1-7)3(5)6/h2,7H,1,4H2,(H,5,6)

Acidity &amino acids &G3 &Kass Steven Bachrach 16 Aug 2007 3 Comments

Bond Dissociation Enthalpies of Hydrocarbons

In Chapter 2.2, we suggest that the experimental deprotonation energy (DPE) of cyclohexane is in doubt. G2MP2 predicts the DPE of cyclohexane is 414.5 kcal mol-1, a figure significantly higher than the experimental1 value of 404 kcal mol-1. Given that the deviation between the G2MP2 computed DPE and experiment is about 2 kcal mol-1, we suggest that cyclohexane should be re-examined.

In a recent JACS article,2 Kass calls into question the experimental bond dissociation energies (BDE) of the small cycloalkanes. With his experimental determination of the BDE of both the vinyl and allylic positions of cyclobutene, Kass can compare experimental and computed BDEs for a range of hydrocarbon environments, as listed in Table 1. The two composite methods G3 and W1 provide excellent BDE values for the small alkanes, one acyclic alkene, and the small cyclic alkenes. These composite methods appear to accurately predict BDEs of hydrocarbons.

However, the small cyclic alkanes are dramatic outliers. The well-accepted experimental BDEs of cyclopropane, cyclobutane, and cyclohexane are 3-5 kcal mol-1 lower than those predicted by the composite methods. Given the strong performance of the computational methods, and the difficulties associated with experimental determinations of BDEs, Kass suggests that the BDEs of these cycloalkanes are in error. Further experiments are deserved.

Table 1. Computed and experimental BDEs (kcal mol-1) of some simple hydrocarbons.

  G3a W1a Expt.
Methane 104.2 104.3 105.0±0.1b
Ethane 101.2 101.2 100.5±0.3b
(CH3)CH2 98.9 98.4 98.1±0.7b
CH3CH2CH2CH3 98.8 98.8 98.3±0.5b
Z-2-butene (allyl) 86.0 87.0 85.6±1.5d
Cyclopropene (vinyl) 109.6 109.8 106.7±3.7e
Cyclopropene (allyl) 100.4 100.4 90.6±4.0f
Cyclobutene (vinyl) 111.9 112.4 112.5±2.5a
Cyclobutene (allyl) 90.6 91.7 91.2±2.3a
Cyclopentene (allyl) 84.2 85.0 82.3±1.1g
Cyclohexene (allyl) 83.9   85±1h
Cyclopropane 109.2 109.0 106.3±0.3b
Cyclobutane 100.5 99.9 96.8±1.0b
Cyclopentane 96.4 96.9 95.6±1.0b
Cyclohexane 100.0   99.5±1.2b

aRef. 2. bRef. 3. cRef. 4. dRef. 5. eRef. 6. fRef. 7. gRef. 8. hRef. 9


(1) NIST webbook, 2005,

(2) Tian, Z.; Fattahi, A.; Lis, L.; Kass, S. R., "Cycloalkane and Cycloalkene C-H Bond Dissociation Energies," J. Am. Chem. Soc. 2006, 128, 17087-17092, DOI: 10.1021/ja065348u

(3) Yao, Y.-R. Handbook of Bond Dissociation Energies in Organic Compounds; CRC Press: Boca Raton, FL, 2003.

(4) CRC Handbook of Chemistry and Physics; 85th ed.; Lide, D. R., Ed.; CRC Press: Boca Raton, FL, 2004.

(5) Tsang, W. In Energetics of Stable Molecules and Reactive Intermediates, NATO Science Series C; Minas da Piedade, M. E., Ed.; Kluwer Academic Publishers: Dordrecht, Netherlands, 1999; Vol. 535.

(6) Fattahi, A.; McCarthy, R. E.; Ahmad, M. R.; Kass, S. R., "Why Does Cyclopropene Have the Acidity of an Acetylene but the Bond Energy of Methane?," J. Am. Chem. Soc. 2003, 125, 11746-11750, DOI: 10.1021/ja035725s.

(7) DeFrees, D. J.; McIver, R. T., Jr.; Hehre, W. J., "Heats of Formation of Gaseous Free Radicals via Ion Cyclotron Double Resonance Spectroscopy," J. Am. Chem. Soc. 1980, 102, 3334-3338, DOI: 10.1021/ja00530a005.

(8) Furuyama, S.; Golden, D. M.; Benson, S. W., "Kinetic Study of the Gas-Phase Reaction c-C5H8+I2 c-C5H6+2HI. Heat of Formation and the Stabilization Energy of the Cyclopentenyl Radical," Int. J. Chem. Kinet. 1970, 2, 93-99.

(9) Alfassi, Z. B.; Feldman, L., "The Kinetics of Radiation-Induced Hydrogen Abstraction
by Trichloromethyl Radicals in the Liquid Phase: Cyclohexene," Int. J. Chem. Kinet. 1981, 13, 771-783.

Bond Dissociation Energy &G3 &Kass Steven Bachrach 11 Jul 2007 No Comments