Proton Held Between Two Lewis Bases
Hydrogen bonding occurs when a proton Lewis acid, H+, is held between two Lewis bases.
The hydrogen bond is a weak type of complexation deemed responsible for the high boiling points of water, alcohols, carboxylic acids etc. and the high solubility of (low molecular weight) alcohols, carboxylic acids and sugars in water.
There is a long and detailed discussion on the Wikipedia hydrogen bond page, however, while this page gives lots of description and illustration it avoids an explanation of the true nature of hydrogen bonding.
While hydrogen bonding does occur when a proton Lewis acid, H+, is held between two Lewis bases there a severe problem. The simple ‘Lewis base–proton–Lewis base’ model cannot exist on simple LACO MO grounds as there are too many electrons!
Two electrons are provided by each Lewis base, giving a total of four electrons. The presence of four electrons must result in the formation of (at least) two molecular orbitals due the Pauli exclusion principle. The first MO will be bonding but the second MO will be antibonding. As the higher energy antibonding MO will dominate, the net result will be antibonding so no bond will form.
Lewis base-H+ bonds are generally highly polar structures and it is easy to consider the hydrogen bond to result from dipole-dipole electrostatic attraction. This is certainly the model and used by many text books. "The lone pair of electrons on one water's δ– oxygen attracts the δ+ hydrogen on an adjacent water molecule:
There is a clues to the true nature of the hydrogen bond, and it comes from reaction chemistry: All Brønsted acid proton transfer reactions pass through a hydrogen bonded intermediate transition state.
Ammonia reacts with hydrogen chloride to produce ammonium chloride, NH4Cl. However, if the reaction is performed at -269°C, 4K, the NH3/HCl hydrogen bonded complex can be trapped – by matrix isolation – and studied by infrared vibrational spectroscopy. This experimental system shows there to be a structure in which the hydrogen atom rapidly moves, vibrates, between the chloride and amine Lewis base centres:
When the proton vibrates in the hydrogen bond it moves from a state in which it is bonded to one Lewis base to a state in which it is bonded to the other, there are two energy minima:
The time averaged effect, the superposition, is for the two Lewis bases to be attracted to each other through the hydrogen atom.
A Hydrogen Bonding Analogy From Any Soap Opera
The girl cannot decide between two boys, who hate each other, yet the boys find themselves strangely drawn together because of the girl... they are attracted to each other through the girl...
And so it is with hydrogen bonding: the two Lewis base centres should repel, but are drawn together through the proton Lewis acid...
Because, with the exception of the electron e– and the photon hν, the proton, H+, is smallest and lightest of all chemical entities. In the hydrogen bond the proton quantum tunnels between the two Lewis bases. The proton buzzes between the two Lewis base centres, associating with both, so drawing them together.
Water, oxygen hydride, is a liquid at room temperature yet all of the other main group hydrides close to water in the periodic table are gases at 25°C and 1.0 atm pressure:
While NH3 and HF do exhibit hydrogen bonding and elevated boiling points, it seems that H2O is ideally suited [the sweet spot] to exhibit strong hydrogen bonding.
Water hydrogen bonds with ammonia, and either molecule can behave as the H+ donor or acceptor.
More complicated molecules can have different types of hydrogen bonding function:
Hydrogen bonding is seen with all molecules possessing -OH functions, including alcohols, carboxylic acids and sugars such as glucose. Carboxylic acids such as acetic acid exist as gas phase dimers:
β-Diketones partially exist the hydrogen bonded cyclic-enol form:
Hydrogen bonding is of immense importance in molecular biology as it constitutes the glue which holds together the twin strands of the DNA double helix and is responsible for secondary, alpha–helix & beta-sheet, and tertiary protein structure.
Side view of an α-helix of alanine residues in atomic detail. Two hydrogen bonds to the same peptide group are highlighted in magenta:
Hydrogen bonding can be studied by substituting D+ for H+, but the congeneric series concept is not really very useful.
Three-Centre Two-Electron Bridge Bonding Ligands
Lewis Acid–Lewis Base–Lewis Acid Complex
A 3-centre-2-electron bond, 3c–2e, is an electron deficient chemical bond where three atoms share two electrons, Wikipedia.
A bridging ligand, Wikipedia, is a ligand that connects two or more atoms, usually metal ions. The ligand may be atomic or polyatomic. Virtually all complex organic compounds can serve as bridging ligands, so the term is usually restricted to small ligands such as hydride, halide or pseudohalides or to ligands linking two metals.
Bridge bonding occurs when a Lewis base is held between a pair of vacant p or d orbital Lewis Acids. The system has two electrons which lead to the formation of a single bonding MO.
Borane, BH3, does not exist at room temperature because it dimerises to diborane. Hydride bridging bonds are found in diborane, B2H6, where the two central hydrogen atoms are simultaneously bonded to both boron atoms in 3c-2e bonds, Wikipedia:
Likewise, the compound commonly called "trimethylaluminium, Al(CH3)3," is actually the dimer with the formula Al2(CH3)6:
Halogen anion bridging bonds as in palladium[II] chloride, PdCl2:
In the ruthenium complex, (η6-C6H6)2Ru2Cl2(μ-Cl)2, two chloride ligands are terminal and two are μ2 bridging. The η in the beginning of the formula denotes the hapticity of the benzene ligands.
Virtually all ligands are known to bridge, with the exception of amines and ammonia. Particularly common inorganic bridging ligands, from Wikipedia, are:
- Hydride, H–
- Halides, Cl–, Br– & I–
- Hydroxide, OH–
- Oxide, O2–
- Sulfide, S2–
- Hydrogen sulfide, SH–
- Carbon monoxide, CO
Many of these reactive reagents behave as if their structures are the simple molecular lobe LUMO Lewis acids: BH3 and Al(CH3)3.
The congeneric series concept is not really very useful here.
van der Waals Attraction
There are three sub-classes of van der Waals attraction:
Induced-dipole/Induced-dipole Attraction (London Dispersion Forces)
It is tempting to consider these forces to be of different strengths, but it is the distance range that is more important. Dipole/dipole attraction is relatively long range in action while the London spontaneous-dipole/Induced-dipole attraction requires contact between the van der Waals surfaces: the molecules need to touch.
Molecules with permanent dipole moments, polar molecules, such as iodine chloride, ICl, exhibit dipole-dipole attraction. The iodine end of iodine chloride is δ+ and the chlorine end is δ–. Molecules interact with each other so that the dipoles line up end-to-end:
All molecules with a permanent dipole exhibit permanent-dipole/permanent-dipole attraction. At temperatures below the material’s melting point, the structure will show long range order and crystallinity.
Molecular dipoles (polar molecules) are able to induce weak dipoles in adjacent non-polar species. The effect gives rise to a weaker attraction than dipole-dipole attraction:
London Dispersion Force (LDF) Attraction:
The very fact that it is possible to liquefy helium – and indeed all molecular materials – demonstrates that there must be some type of inter-molecular attraction taking place between the helium atoms. (Helium is a molecular material, where the helium molecule consists of just one atom.)
The attraction is known as the London dispersion force and is deemed to arise from short time scale fluctuations in the electronic structure of species which results in the formation of instantaneous dipoles.
The instantaneous-induced-dipole/induced-dipole London dispersion forces (LDF) are surprisingly strong but they only act at very short range. It is as if the surface of even neutral, non-polar molecules like methane are 'sticky'.
Soccer Balls Covered in Velcro
Imagine a room filled with 50 soccer balls or so covered in Velcro and half a dozen four year old children.
The kids will kick the balls about, and the balls will fly around. But as the children become tired the balls will slow down and stick together.
So it is with a molecular gas. As the temperature is lowered the molecules will stick to each other via London dispersion forces – instantaneous-induced-dipole/Induced-dipole attractions – to give a condensed phase.
All molecules exhibit London dispersion forces and the strength increases with the size/surface area of the molecule. This logic can be used to explains the increasing boiling and sublimation temperatures of the halogens.
Going down the periodic table the atoms become larger, so the diatomic molecules become larger and their surface area becomes larger. Thus, the van der Waals forces increase and so do the boiling points:
F2 < Cl2 < Br2 < I2
Likewise, longer chain alkanes have higher boiling points than shorter chain alkanes. Branching, which decreases surface area, reduces boiling point.
Which is stronger: dipole/dipole attraction or the London force?
Consider the molecular halogen bromine, Br2, and the interhalogen iodine chloride, ICl.
Both have a mass of close to 160, both are are 70 electron systems, but Br2 is non-polar and ICl is polar. Yet they have rather similar boiling points of 59 ° and 97° respectively:
This implies that the dipole/dipole attraction makes only a minor contribution to the net attraction, and the most of the molecular stickiness is due to the the London dispersion force.
Gecko Toes, Setae and van der Waals Forces:
"The toes of the gecko have developed a special adaptation that allows them to adhere to most surfaces. Recent studies of the spatula tipped setae on gecko footpads demonstrate that the attractive forces that hold geckos to surfaces are van der Waals interactions between the finely divided setae and the surfaces themselves. Every square millimetre of a gecko's footpad contains about 14,000 hair-like setae." Wikipedia:
- Hydrogen bonding substances are soluble in hydrogen bonding solvents like water or methanol.
- Polar substances are soluble in polar solvents like chloroform, CHCl3, or toluene, C6H5CH3.
- Non-polar substances are soluble in non-polar solvents like cyclohexane, C6H12.
The congeneric concept is not useful here.
Molecular Shape Recognition Complex
"An inclusion compound is a complexing which one chemical compound, the host, forms a cavity in which molecules of a second guest compound are located. The definition of inclusion compounds is very broad, extending to channels formed between molecules in a crystal lattice in which guest molecules can fit. If the spaces in the host lattice are enclosed on all sides so that the guest species is ‘trapped’ as in a cage, the compound is known as a clathrate. In molecular encapsulation a guest molecule is actually trapped inside another molecule." Wikipedia
Solid urea can accommodate octane or 1-bromoctane, but not 2-methyl heptane or 2-bromooctane, as guest molecules in a hexagonal host lattice structure which contains long guest channels about 500pm in diameter.
Zeolites are aluminosilicate minerals (natural and synthetic) which have open structures able to accommodate a wide range of guest molecules. The microporous molecular structure of a zeolite ZSM-5:
Molecular recognition plays an important role in biological systems and is observed in between receptor-ligand, antigen-antibody, DNA-protein, sugar-lectin, RNA-ribosome, etc.
An example of molecular recognition is the antibiotic vancomycin that selectively binds with the peptides with terminal D-alanyl-D-alanine in bacterial cells through five hydrogen bonds. The vancomycin is lethal to the bacteria since once it has bound to these particular peptides they are unable to be used to construct cell wall.