Posts Tagged ‘Nature’

Newer Entries »

The “hydrogen bond”; its early history.

Saturday, December 31st, 2016

My holiday reading has been Derek Lowe’s excellent Chemistry Book setting out 250 milestones in chemistry, organised by year. An entry for 1920 entitled hydrogen bonding seemed worth exploring in more detail here.

As with many historical concepts, it can often take a few years to coalesce into something we would readily recognise today, and hydrogen bonds are no exception. Wikipedia is another source of the history and it cites a 1912 article as the origin of the term in relation to the solvation of amines[1] but also notes that the better known setting of water occurs later in 1920.[2] Here I try to capture the essence of the concept with a few diagrams taken from these two articles.

 Firstly “The state of amines in aqueous solution“[1] which is mostly concerned with the measurement of ionization constants of primary, secondary and tertiary amines. It boils down to the below:

and the connection to ionization is laid out as:

Since in 1912, Lewis’ electron pair theory of the covalent bond had not yet emerged, the authors use the terms “strong union” and “weak union”, and of course it is the “weak union” that we now know of as the hydrogen bond. Some other comments about this seminal diagram:

  1. The article contains the very explicit and modern term stereochemical, which is used in a manner that suggests it was already common.‡ But there is only a hint at most that the nitrogen atoms might be tetrahedral, or that the “weak union” between (what we now think of as the lone pair on) the nitrogen and the hydrogen of the water is directional.
  2. The second weak union between the tetramethyl ammonium (which we now describe as a cation) and the hydroxide (now described as an anion; both terms are however implied by the description strong electrolyte) is probably not what we would now call a hydrogen bond, more an intimate ion-pair.

The second article in 1920 on water itself[2] is post-Lewis, but perhaps applied in a manner which we would not entirely agree with nowadays. Thus dinitrogen, N≡N is shown as below with just a single connecting bond.

Then we get the interaction between ammonia and water,† analogous to the example shown above:

and for water itself:â™ 

which in each case shows the central hydrogen having what we now call a valence shell of four electrons,♣ and hence more equivalent to the “strong unions” above. This shows that in 1920 chemists were rapidly adopting Lewis’ representations, but not always entirely successfully.

On balance, I think the 1912 article sets out the modern concept of a hydrogen bond representing a weak union to a hydrogen rather better than the Latimer and Rodebush attempt (at least diagrammatically).

‡Stereochemical notation is discussed in this post, and it dates from the 1930s.

†The modern take is explored here, in which the equilibrium set up between a “weak union” between ammonia and water (the weak electrolyte) and an isomeric “strong union” which represents ionization into an ammonium hydroxide ion-pair (the strong electrolyte) is favoured for the former by ΔG ~6 kcal/mol.

â™ The equilibrium between a “weak union” of two water molecules and the fully ionized strong union of hydronium hydroxide favours the former by ΔG ~23 kcal/mol.

♣ This 1920 representation does imply symmetry for the hydrogen, being ~equally disposed between the two oxygens. We now know that such symmetric hydrogen bonding is not unusual (see this post for how to fine-tune a hydrogen bond into this situation) but rather than requiring four electrons as implied in the diagram above, it is now better described as a three-centre-two-electron bond instead.

References

  1. T.S. Moore, and T.F. Winmill, "CLXXVII.—The state of amines in aqueous solution", J. Chem. Soc., Trans., vol. 101, pp. 1635-1676, 1912. https://doi.org/10.1039/ct9120101635
  2. W.M. Latimer, and W.H. Rodebush, "POLARITY AND IONIZATION FROM THE STANDPOINT OF THE LEWIS THEORY OF VALENCE.", Journal of the American Chemical Society, vol. 42, pp. 1419-1433, 1920. https://doi.org/10.1021/ja01452a015

Tags:, , , , , , , , , , ,
Posted in Historical | 2 Comments »

The "hydrogen bond"; its early history.

Saturday, December 31st, 2016

My holiday reading has been Derek Lowe’s excellent Chemistry Book setting out 250 milestones in chemistry, organised by year. An entry for 1920 entitled hydrogen bonding seemed worth exploring in more detail here.

As with many historical concepts, it can often take a few years to coalesce into something we would readily recognise today, and hydrogen bonds are no exception. Wikipedia is another source of the history and it cites a 1912 article as the origin of the term in relation to the solvation of amines[1] but also notes that the better known setting of water occurs later in 1920.[2] Here I try to capture the essence of the concept with a few diagrams taken from these two articles.

 Firstly “The state of amines in aqueous solution“[1] which is mostly concerned with the measurement of ionization constants of primary, secondary and tertiary amines. It boils down to the below:

and the connection to ionization is laid out as:

Since in 1912, Lewis’ electron pair theory of the covalent bond had not yet emerged, the authors use the terms “strong union” and “weak union”, and of course it is the “weak union” that we now know of as the hydrogen bond. Some other comments about this seminal diagram:

  1. The article contains the very explicit and modern term stereochemical, which is used in a manner that suggests it was already common.‡ But there is only a hint at most that the nitrogen atoms might be tetrahedral, or that the “weak union” between (what we now think of as the lone pair on) the nitrogen and the hydrogen of the water is directional.
  2. The second weak union between the tetramethyl ammonium (which we now describe as a cation) and the hydroxide (now described as an anion; both terms are however implied by the description strong electrolyte) is probably not what we would now call a hydrogen bond, more an intimate ion-pair.

The second article in 1920 on water itself[2] is post-Lewis, but perhaps applied in a manner which we would not entirely agree with nowadays. Thus dinitrogen, N≡N is shown as below with just a single connecting bond.

Then we get the interaction between ammonia and water,† analogous to the example shown above:

and for water itself:â™ 

which in each case shows the central hydrogen having what we now call a valence shell of four electrons,♣ and hence more equivalent to the “strong unions” above. This shows that in 1920 chemists were rapidly adopting Lewis’ representations, but not always entirely successfully.

On balance, I think the 1912 article sets out the modern concept of a hydrogen bond representing a weak union to a hydrogen rather better than the Latimer and Rodebush attempt (at least diagrammatically).

‡Stereochemical notation is discussed in this post, and it dates from the 1930s.

†The modern take is explored here, in which the equilibrium set up between a “weak union” between ammonia and water (the weak electrolyte) and an isomeric “strong union” which represents ionization into an ammonium hydroxide ion-pair (the strong electrolyte) is favoured for the former by ΔG ~6 kcal/mol.

â™ The equilibrium between a “weak union” of two water molecules and the fully ionized strong union of hydronium hydroxide favours the former by ΔG ~23 kcal/mol.

♣ This 1920 representation does imply symmetry for the hydrogen, being ~equally disposed between the two oxygens. We now know that such symmetric hydrogen bonding is not unusual (see this post for how to fine-tune a hydrogen bond into this situation) but rather than requiring four electrons as implied in the diagram above, it is now better described as a three-centre-two-electron bond instead.

References

  1. T.S. Moore, and T.F. Winmill, "CLXXVII.—The state of amines in aqueous solution", J. Chem. Soc., Trans., vol. 101, pp. 1635-1676, 1912. https://doi.org/10.1039/ct9120101635
  2. W.M. Latimer, and W.H. Rodebush, "POLARITY AND IONIZATION FROM THE STANDPOINT OF THE LEWIS THEORY OF VALENCE.", Journal of the American Chemical Society, vol. 42, pp. 1419-1433, 1920. https://doi.org/10.1021/ja01452a015

Tags:, , , , , , , , , , ,
Posted in Historical | 2 Comments »

The dipole moments of highly polar molecules: glycine zwitterion.

Saturday, December 24th, 2016

The previous posts produced discussion about the dipole moments of highly polar molecules. Here to produce some reference points for further discussion I look at the dipole moment of glycine, the classic zwitterion (an internal ion-pair).

Dielectric relaxation studies of glycine–water mixtures yield values that range from 15.7D[1] to 11.9D[2] although these have to be derived using various approximations and assumptions for up to 4 independent Debye processes. Before proceeding to calculations, I looked at the properties of ionized amino acids in the solid state, using the following search query for the Cambridge structure database (CSD). 

The distance measures hydrogen bonds to the carboxylate oxygens and the torsion their orientation. The O…H hydrogen bond distances vary between 1.7-1.85Ã…, which are short. The orientation of the hydrogen bond can be to the in-plane oxygen “σ-lone pair” (torsion 0 or 180°) and also an out-of-plane ~Ï€ form (torsion ~60-90°).

In aqueous solution, it is normally assumed that glycine sustains five such strong H-bonds (three to the H3N+ group and two[3] to the carboxylate anion), forming a polarised “salt bridge” across the ion-pair. Two model types were subjected to calculation using ωB97XD/Def2-TZVPP/SCRF=water. Aqueous glycine without any added explicit water molecules yields a dipole moment of 12.9D (DOI: 10.14469/hpc/2000), which is within the range noted above.‡

The solvated form is shown below, in one specific conformation of the three studied (ωB97XD/Def2-TZVPP/SCRF=water). The calculated O…H hydrogen bond lengths fall into the range revealed from crystal structures. The calculated dipole moments range from 12.6 (DOI: 10.14469/hpc/2007), 15.3 (DOI: 10.14469/hpc/2006) and 14.9D (DOI: 10.14469/hpc/2005), which is a modest increase over the model with no explicit water molecules. The actual dipole is of course a Boltzmann average over these and other as yet unexplored conformations, as well as other values for the number of water molecules.

Given the difficulties in interpreting the dipole moment of a complex Debye system such as hydrated glycine, the agreement between the limited range of solvated models and the measured values seems reasonable, and provides at least some measure of “calibration” for the polar molecules commented on previously.

‡Optimized with the solvent field on. If a vacuum model is used, the proton transfers from the N to the O.

References

  1. M.W. Aaron, and E.H. Grant, "Dielectric relaxation of glycine in water", Transactions of the Faraday Society, vol. 59, pp. 85, 1963. https://doi.org/10.1039/tf9635900085
  2. T. Sato, R. Buchner, Å. Fernandez, A. Chiba, and W. Kunz, "Dielectric relaxation spectroscopy of aqueous amino acid solutions: dynamics and interactions in aqueous glycine", Journal of Molecular Liquids, vol. 117, pp. 93-98, 2005. https://doi.org/10.1016/j.molliq.2004.08.001
  3. T. Shikata, "Dielectric Relaxation Behavior of Glycine Betaine in Aqueous Solution", The Journal of Physical Chemistry A, vol. 106, pp. 7664-7670, 2002. https://doi.org/10.1021/jp020957j

Tags:, , , , , , , , , , , , ,
Posted in crystal_structure_mining, Interesting chemistry | 1 Comment »

Molecules of the year? Pnictogen chains and 16 coordinate Cs.

Monday, December 19th, 2016

I am completing my survey of the vote for molecule of the year candidates, which this year seems focused on chemical records of one type or another.

The first article[1] reports striving towards creating a molecule covering a complete column of the period table. In this case, group 7, containing N, P, As, Sb, Bi and Mc. Only the first four of these were incorporated, although the prospects of extending this to five seem good (and to six extremely unlikely).  The structure of this pnictogen chain is referenced here: DOI: 10.5517/CCDC.CSD.CC1LHPJ9 and I have demurred from a calculation.

The second article[2] relates to what might be called hypercoordination, and the achievement of what is felt is a maximum value of 16 to a single metal. I thought I might approach this one by searching the Cambridge structure database (CSD) by specifying any metal with a coordination number 16 as the search query. However, I was foiled in this query because the search software (Conquest) allows a maximum value of only 15! So instead I list the total number of hits retrieved for coordination numbers of 10-15: 25224, 4753, 8856, 2492, 839, 348 respectively.  

These totals have to be taken with some caution; the coordination number of what may often be very weak interactions may be often determined by human chemical perception rather than hard and fast rules. Nevertheless, the assignment of 348 molecules to having a coordination number of 15 is still a remarkably high number. If I can persuade CCDC to allow searches with 16, who knows what other candidates might emerge to rival this one, DOI: CCDC.CSD.CC1KFCQ2

The final candidate[3] is the only one where no measured coordinates are reported, with the title “Preparation of an ion with the highest calculated proton affinity: ortho-diethynylbenzene dianion”. There high level theoretical and computational modelling is reported to which I cannot add anything useful.

The common theme emerging of my review is that most of the candidates have crystal structures to which I have been able to occasionally add some computed quantum mechanical properties to try to tease out some other aspects of their character. It is also nice to be able to cite a persistent identifier (DOI) that leads directly to the 3D coordinates for the structures. My first ever post to this blog in 2008 addressed one solution on how such immediacy might be achieved and it is nice to see this now as a mainstream aspect of chemical publishing.

References

  1. A. Hinz, A. Schulz, and A. Villinger, "Synthesis of a Molecule with Four Different Adjacent Pnictogens", Chemistry – A European Journal, vol. 22, pp. 12266-12269, 2016. https://doi.org/10.1002/chem.201601916
  2. D. Pollak, R. Goddard, and K. Pörschke, "Cs[H<sub>2</sub>NB<sub>2</sub>(C<sub>6</sub>F<sub>5</sub>)<sub>6</sub>] Featuring an Unequivocal 16-Coordinate Cation", Journal of the American Chemical Society, vol. 138, pp. 9444-9451, 2016. https://doi.org/10.1021/jacs.6b02590
  3. B.L.J. Poad, N.D. Reed, C.S. Hansen, A.J. Trevitt, S.J. Blanksby, E.G. Mackay, M.S. Sherburn, B. Chan, and L. Radom, "Preparation of an ion with the highest calculated proton affinity: ortho-diethynylbenzene dianion", Chemical Science, vol. 7, pp. 6245-6250, 2016. https://doi.org/10.1039/c6sc01726f

Tags:, , , , , , , , ,
Posted in crystal_structure_mining, Interesting chemistry | 2 Comments »

Long C=C bonds.

Thursday, December 1st, 2016

Following on from a search for long C-C bonds, here is the same repeated for C=C double bonds.

sq

The query restricts the search to each carbon having just two non-metallic substituents. To avoid conjugation with these, they each are 4-coordinated; the carbons themselves are three-coordinated. Further constraints are the usual no disorder, no errors and R < 0.1 and the C=C distance > 1.4Å (the standard value is ~1.32-1.34Å). The search query is deposited as DOI: 10.14469/hpc/1959[1]

c_c

The apparent longest example is LIRVEN, DOI: 10.5517/CC4R2MK[2] with a value of 1.589Å, longer than most C-C single bonds! Closer inspection reveals the presence of lithium cations, and so the molecule bearing the C=C bond must sustain two negative charges. So this apparent C=C bond is in fact anionic, with one electron going into each of the π* orbitals, thus lengthening the CC bond.‡ Not a true example of a neutral C=C bond[3] but it now becomes interesting for what its spin state might be. Is it a biradical or a triplet for example? One to be investigated further I fancy! Another example of this type is QUKCEE[4]

10-5517cc4r2mk-lirven

This next FAZWIM has a C=C length of 1.546Å. It is an old structure (1986), and comes without attached hydrogen atoms. Although drawn with no hydrogens on the central C=C bond, the length suggests this molecule is simply mis-assigned.†fazwin-diag fazwim

The final example I will highlight is pretty ordinary looking and published in 2016 as a private communication; ALOVOO, DOI: 10.5517/CCDC.CSD.CC1LJSWS[5] with a C=C length of 1.443Å. Again no obvious reason for the bond to be longer than normal.‡†

10-5517ccdc-csd-cc1ljsws-alovoo

In hunting for such unusual deviations from the norm, the most obvious explanation is normally some anomaly in the crystallographic analysis. Although the CSD (crystal structure database) is a very heavily curated resource, it seems unlikely that each deposition would be carefully inspected for its chemistry, and this must be our task here. But such anomalies can themselves point to interesting or unusual chemistry, which in  turn can be subjected to quantum computation to see if either the unusual value can be replicated or other reasons identified.  In this case, this exercise can been conducted by a human, but one can easily envisage the entire process being automated on a far larger scale.  The future?

‡ In fact the stoichiometry shows each “double bond” is actually a di-anion, with two electrons entering each of the the Ï€* orbitals.

†A calculation on the singlet state for the structure as drawn (ωB97XD/Def2-TZVPP, DOI: 10.14469/hpc/1960) gives a bond length of 1.342Å, i.e. that expected for a double bond. The triplet state is similar in energy, but with a much longer central bond length of 1.476Å, DOI: 10.14469/hpc/1962 but the geometry at the carbons is planar and not bent as shown above. The quintet state is 1.45Å and is again planar, doi 10.14469/hpc/1963. So calculations on FAZWIM strongly suggest the structure as shown is an error.

‡†The computed value is 1.324Å, perfectly normal. DOI: 10.14469/hpc/1966[6]

References

  1. H. Rzepa, "Long C=C bonds", 2016. https://doi.org/10.14469/hpc/1959
  2. Matsuo, T.., Watanabe, H.., Ichinohe, M.., and Sekiguchi, A.., "CCDC 141348: Experimental Crystal Structure Determination", 2000. https://doi.org/10.5517/cc4r2mk
  3. T. Matsuo, H. Watanabe, M. Ichinohe, and A. Sekiguchi, "Reduction of the 1,4,5,8-tetrasila-1,4,5,8-tetrahydroanthracene derivative with lithium metal. Isolation and characterization of the tetralithium salt of a tetraanion, and observation of an Si–H⋯Li+ interaction", Inorganic Chemistry Communications, vol. 2, pp. 510-512, 1999. https://doi.org/10.1016/s1387-7003(99)00136-7
  4. T. Matsuo, H. Watanabe, and A. Sekiguchi, "A Novel Tetralithium Salt of a Tetraanion and a Dilithium Salt of a Dianion, Formed by the Reduction of the Tetrasilylethylene Moiety. Synthesis, Characterization, and Observation of an Si-H···Li+ Interaction", Bulletin of the Chemical Society of Japan, vol. 73, pp. 1461-1467, 2000. https://doi.org/10.1246/bcsj.73.1461
  5. M.E. Light, S. Bain, and J. Kilburn, "CCDC 1475906: Experimental Crystal Structure Determination", 2016. https://doi.org/10.5517/ccdc.csd.cc1ljsws
  6. H. Rzepa, "ALOVOO", 2016. https://doi.org/10.14469/hpc/1966

Tags:, , , , , , , , , , , ,
Posted in crystal_structure_mining, Interesting chemistry | 2 Comments »

Molecule orbitals as indicators of reactivity: bromoallene.

Thursday, September 1st, 2016

Bromoallene is a pretty simple molecule, with two non-equivalent double bonds. How might it react with an electrophile, say dimethyldioxirane (DMDO) to form an epoxide?[1] Here I explore the difference between two different and very simple approaches to predicting its reactivity. bromoallene

Both approaches rely on the properties of the reactant and use two types of molecule orbitals derived from its electronic wavefunction. The first of these is very well-known as the molecular orbital (MO), which has the property that it tends to delocalise over all the contributing atoms (the “molecule”). MOs are often used in this context; the highest energy occupied MO is thought of as being associated with the most nucleophilic (electron donating) regions of the molecule and so such a HOMO would be expected to predict the region of nucleophilic attack. The second is known as the natural bond orbital (NBO), which is evaluated in a manner that tends to localise it on bonds (the functional groups or reaction centres) and atom centres. What do these respective orbitals reveal for bromoallene? 

The MOs
HOMO, -0.3380 HOMO-1, -0.3692 au
Click for 3D

Click for 3D
Click for 3D

Click for 3D
The NBOs
HONBO, -0.3769 HONBO-2, -0.3898
Click for 3D

Click for 3D
Click for 3D

Click for 3D

The table above shows the energies (in Hartrees) of the four relevant orbitals. The less negative (less stable) the orbital, the more nucleophilic it is. The (heavily) delocalized HOMO is located on the C=C bond bond carrying the C-Br bond, Δ1,2 alkene, but it also has a large component on the Br. The more stable HOMO-1 is located on the C=C bond located away from the Br, the Δ2,3 alkene and also with a (different type of) component on the Br.

In contrast, the HONBO is located on the Δ2,3 alkene and it is the HONBO-2 that is on the Î”1,2 alkene. Both these orbitals have very little “leakage” onto other atoms, they are almost completely localised.

Well, now we have a problem since these two analyses lead to diametrically opposing predictions! So what does experiment say? A recent article[1] addresses this issue by isolating the initially formed epoxide from reaction with DMDO and characterising it using crystallography. But here comes the catch; such isolation only proved possible if the allene was also substituted with large sterically bulky groups such as t-butyl or adamantyl. And the isolated product was the Î”1,2 epoxide. So does that mean that the MO method was correct and the NBO method wrong? Well, not necessarily. Those large groups play an additional role via steric effects. To factor in such effects one has to look at the transition state model for the reaction rather than depending purely on the reactant properties. And the steric effects in this case appear to win out over the electronic ones.[1]

The Klopman[2]-Salem[3] equation (shown in very simplified, and original, form below for just the covalent term) casts some light on what is going on. This term is a double summation over occupied/unoccupied (donor-acceptor) orbital interactions, involving the coefficients of the orbitals (the overlap integrals in effect) in the numerator and the energy difference between the occupied/unoccupied orbital pair as denominator.

KS1

Performing such a double summation is rarely attempted; instead the equation is reduced to just one single term involving the donor of highest energy and the acceptor of lowest energy, ensuring the energy difference is a minimum and hence the term itself is (potentially) the largest in the summation. There is still the issue of the orbital coefficients, and here we get to the crux of the difference between the use of MOs and NBOs. You can see by inspection that the two π-MOs for bromoallene have different coefficients on the two atoms of interest, the two carbons of the double bond. One really has to evaluate the size of this term in the summation by using quantitative values for the respective coefficients and to very probably include the further terms in the summation for any other orbitals which also have significantly non-zero coefficients on these two atoms. But with the NBOs, the localisation procedure used to derive them has reduced the coefficients to just the carbon atoms and effectively no other atoms; all the other terms in the double summation in effect do drop out entirely. So with NBOs, the only number that matters is the energy difference between the occupied/empty orbitals (the denominator). But since the acceptor (the electrophile, DMDO in this case) is the same for both regiochemistries, things reduce even further to just comparing the donor energies for the two alternatives (Table above). The higher/less stable of these will have the greater contribution in the Klopman-Salem equation.

This little molecule teaches the important lesson that electronic and steric effects both play a role in directing reactions, and in this system they may well oppose each other. Simple interpretations based on reactant orbitals may give only a partial and even potentially misleading answer.

References

  1. D. Christopher Braddock, A. Mahtey, H.S. Rzepa, and A.J.P. White, "Stable bromoallene oxides", Chemical Communications, vol. 52, pp. 11219-11222, 2016. https://doi.org/10.1039/c6cc06395k
  2. G. Klopman, "Chemical reactivity and the concept of charge- and frontier-controlled reactions", Journal of the American Chemical Society, vol. 90, pp. 223-234, 1968. https://doi.org/10.1021/ja01004a002
  3. L. Salem, "Intermolecular orbital theory of the interaction between conjugated systems. I. General theory", Journal of the American Chemical Society, vol. 90, pp. 543-552, 1968. https://doi.org/10.1021/ja01005a001

Tags:, , , , , , , , , , , , , ,
Posted in reaction mechanism | No Comments »

Newer Entries »