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Introduction to the Lewis Acid/Base Interaction Matrix

A Deeper Exploration into the Lewis Acid/Base Interaction Matrix, one cell at a time:

Type 1 Lewis Acid/Base Complexation Chemistry

Type 1 complexes, typified by hydrogen H₂, are covalently bonded. The bonding is frontier molecular orbital (FMO) controlled. Hydrogen, H₂, has a 1σ² MO structure, ie they have two electrons in their 1σ molecular orbital.

Find out more about the bonding diatomic species elsewhere in this webbook, here.

Protons complex with hydride ions to form molecular hydrogen, H₂, a uniquely simple and much studied diatomic molecule.

The H⁺  +  H^–   →  H₂  reaction is not reversible: H₂ does not act as a proton donor (although at high temperature, when exposed to high energy UV radiation or when absorbed onto a metallic surface, H₂ can homolytically dissociate: radical cleavage).

As protons and hydride ions do not exist as independent species, they require "delivery" by donor complexes, ie reagents. Protons, H⁺ ions, are supplied by Brønsted acids and hydride ions by hydride donor complexes.

For example, hydrogen chloride an H⁺ donor reacts with sodium hydride an H^– donor to give diatomic hydrogen and sodium chloride:

   HCl + NaH    →  H₂ + NaCl

[H₃]⁺, the product of H⁺ and H₂, is the simplest possible triatomic molecular ion – it has only two electrons – and is of considerable theoretical interest.

The [H₃]⁺ molecular ion occupies an important position in theoretical models of interstellar chemistry as the [H₃]⁺ forms in hydrogen-rich interstellar gas clouds. The [H₃]⁺ ion can protonate carbon, oxygen and other atoms, thereby initiating the interstellar synthesis of nearly 100 molecules including: hydroxyl radicals (HO•), carbon monoxide, ethanol, linear polyacetylenes and cyclopropenylidene.

(In this author's opinion the [H₃]⁺ ion should be called the 'hydronium ion', and [OH₃]⁺ should be the 'oxonium ion'.)

Type 2 Lewis Acid/Base Complexation Chemistry

Protons do not complex efficiently with complex anion Lewis bases as the proton must disrupt the anion’s high symmetry HOMO when a 1:1 complex forms. Hence, the resulting complex is a very powerful proton donor.

Complexes are ionic/highly polar covalent.

Superacids are neutral.

Proton plus complex anion complexes must be prepared in ‘exotic’ solvents such as liquid SO₂ and where the complex anion has fluoride ion ligands.

George Olah's 'magic acid' is prepared by mixing antimony pentafluoride and fluorosulfonic acid. Hydrocarbon wax will dissolve in a magic acid solution.

Such complexes are the strongest Brønsted acids known. Super [Brønsted] acids, 'superacids', have pKa values in the region -15 to -25, ie they are up to 20 orders of magnitude more [Brønsted] acidic than sulfuric acid.

Superacids are able to fully protonate all Lewis base organic functional groups (as opposed to protonating a low equilibrium concentration). Alkenes, benzene, carbonyls and nitro functional groups are all protonated by superacids. For example:

Read more on the Wikipedia superacid page.

Type 3 Lewis Acid/Base Complexation Chemistry

The point-charge of the proton, H⁺, is able to efficiently penetrate and complex with the directional lobe shaped sp³/sp²/sp orbitals of all Lobe-HOMO Lewis bases. The resulting complexes are either covalent or polar-covalent.

The vast majority of Brønsted acids are Type 3 Proton/Lobe-HOMO Lewis acid/base complexes, including all of the common mineral acids H₂SO₄, HCl, HNO₃, H₃PO₄ etc.

Complexes show regular Brønsted acid behavior across congeneric series, ie across and down the periodic table:

Brønsted acidity increases with increasing conjugation and resulting π-stabilisation.

Read more elsewhere in this webbook, here.

Type 4 Lewis Acid/Base Complexation Chemistry

When a proton complexes with a hydrocarbon π-system (as opposed to complexing with a single heteroatom atomic centre) the conjugated π-system is reduced in length by one p-orbital to form a Hückel distinct MO system: 2p alkenes are protonated to give give 1p carbenium ions, 6p benzene is protonated gives the corresponding 5p carbenium ion, etc.

Bonding in the complexes is covalent.

When a proton complexes with a hydrocarbon π-system the conjugated π-system is reduced in length by one p-orbital to form a Hückel distinct MO system: 2p alkenes give 1p carbenium ions, 6p benzene gives the corresponding 5p carbenium ion, etc.

If there is a heteroatom Lewis base centre present, it is likely that protonation will occur at that centre. For example, 2p carbonyl functions protonate on oxygen and 6p (aromatic) pyridine protonates on the nitrogen lone pair which is not part of the conjugated π-system:

While the action of a strong acid on benzene is null, benzene is deuterated with deutero sulfuric acid, D₂SO₄. In time the benzene will become fully (or per) deuterated:

Type 5 Lewis Acid/Base Complexation Chemistry

Group I & II alkali and alkaline earth hydrides exist as ionic lattice solids. However, molecular 1:1 (LiH) and 1:2 (MgH₂) complexes are formed in the vapor phase, although with difficulty as the compounds are liable to decompose back to elemental form.

LiH is a much theoretically studied diatomic species. Studies show that bonding involves more than just s-LUMO/s-HOMO overlap. In valence bond (VB) terminology, the lithium’s 2s-LUMO mixes with (ie hybridizes with) a higher energy empty 2p orbital (ie the LUMO + 1 MO) to generate a directional sp hybrid bond.

MO calculations show that the bonding in LiH involves 66% Li 2s/H1s overlap and 34% Li 2p/H 1s overlap.

The charge on a Type 5 complex is always neutral.

Saline hydride complexes either act as strong proton abstracting Bases or as donors of nucleophilic hydride ion. Organic chemists employ saline hydride complexes as proton abstracting Bases, with H₂ being the conjugate Brønsted acid.

Sodium hydride readily abstracts a proton from dimethyl sulfoxide (DMSO), pKa 35, to form sodium dimsyl:

Inorganic chemists are more likely to use saline hydride reagents as a source of nucleophilic (and reducing) hydride ion, for example in the synthesis of sodium borohydride:

LiH             NaH              KH               RbH              CsH

BeH₂          MgH₂            CaH₂            SrH₂              BaH₂

Type 6 Lewis Acid/Base Complexation Chemistry

In polar solvents species dissociate to give independently solvated ions. While in solution, the complex anion part of a Type 6 complex is susceptible to ligand transfer or substitution. The s-LUMO ion part of the complex usually acts as a spectator ion, however, choice of ion influences solubility.

Reagents such as lithium aluminium hydride and sodium borohydride act as hydride ion donor complexes. Type 6 hydride complexes are weaker proton abstracters than the very basic Type 5 complexes and they are correspondingly more suitable for use by organic chemists as sources of reducing nucleophilic hydride ion:

Many important ligand replacement congeneric series are known, for example:

Type 7 Lewis Acid/Base Complexation Chemistry

The s-LUMO Lewis acids like to be multiply complexed by Lobe-HOMO Lewis base ligands to achieve maximum spherical symmetry about the cation.

Complexes form as ionic crystal lattice (cesium fluoride) or they may generate polar-covalent bonds (methyl lithium).

With a few important exceptions – for example, amines and the saline hydrides – most of the proton abstracting (Brønsted base) reagents are Type 7 Lewis acid/base complexes:

NaOH     KNH₂     LiCH₃     CH₃COONa    NaHCO₃   Na₂CO₃     etc.

Other Type 7 complexes, in which the Lobe-HOMO Lewis Base is the conjugate base of a strong Brønsted acid, are non-basic ionic salts:

KCl     NaI     LiBr   etc.

In polar solvent solution (water, THF, DMF or DMSO), Type 7 complexes may show differing degrees of ion pairing:

Polar solvation requires the solvent to have lone pair Lobe-HOMO centres which compete with the Type 7 Complexes’ Lobe-HOMO Lewis base centre.

Methyl lithium, for example, is usually considered to be a 1:1 polarised covalent compound, however, in diethyl ether or THF solvents the lithium cation is co-complexed with solvent molecules.

In alkane solvents alkyl lithium reagents aggregate into hexamers and higher structures, Wikipedia.

There are many important congeneric series: alkyl carbanion reagents become more basic as the alkyl groups are replaced by hydrogen ligands, and as the metal counter ion becomes congenerically heavier.

Regular changes in physical properties occur as pairs of congeneric Lewis acid/base series interact to form congeneric series, planars and volumes of complexes:

Type 8 Lewis Acid/Base Complexation Chemistry

s-LUMO Lewis acids complex poorly with π-HOMO Lewis bases:

• The s-LUMO Lewis acid is "looking for" ligands, the π-HOMO Lewis base provides a large delocalised anionic π-system.

Thus, the bonding interaction is charge controlled and ionic even though the participating species are "looking for" FMO controlled bonding. The problem is that the FMOs are of dissimilar geometry.

Species may only be stable in [non-aqueous] solution where the s-LUMO Lewis acid is complexed by the solvent. Typical solvents are diethyl ether or THF.

s-LUMO Lewis acids are employed as non-electrophilic spectator counter ions to π-HOMO systems with net negative charge: ie the allyl, benzyl and cyclopentadienyl anions. The effect of such complexation is to make the π-HOMO anion to appear ‘naked’ for spectroscopic and reactivity studies.

Bonding in Type 8 complexes is strongly influenced by environment: polar solvents compete with the π-system ligand at complexing the hard s-LUMO Lewis acid cation. s-LUMO/π-anion chemistry is usually carried out in diethyl ether, THF or glyme (polyether) solvents where the ethereal Lewis bases are the true s-LUMO complexing agents.

Most type 8 complexes are strong proton abstracting bases because the conjugate Brønsted acids of the π-HOMO systems are weakly [Brønsted] acidic π-hydrocarbons. The conjugate Brønsted acid of benzyl lithium, for example, is toluene, pKa of 41.

The nature of the s-LUMO cation/π-system bond changes congenerically with cation: a Li⁺/π^– complex is less ionic (therefore more covalent and therefore more strongly ion-paired in solution) than the equivalent Cs⁺/π^– complex.

Type 9 Lewis Acid/Base Complexation Chemistry

If the onium ion has a proton ‘ligands’ the onium ion will act as a Brønsted acid and donate a proton to the hydride ion to form H₂:

H₃O⁺      +     H^–      →    H₂     +    H₂O:

If the onium ion has alkyl ligands (such as a trialkyl oxonium ion or a tetraalkyl ammonium ion) a nucleophilic substitution will occur in which electropositive nucleophilic hydride ion reduces one of the onium ion ligands to the corresponding alkane:

Type 10 Lewis Acid/Base Complexation Chemistry

Few series of much interest.

Type 11 Lewis Acid/Base Complexation Chemistry

Onium ions exhibit three types chemistry with Lobe-HOMO Lewis bases:

• Ionic salt formation:
  
  [NH₄]⁺      +     Cl ^–      →    NH₄Cl

• Proton transfer from the onium ion to lobe-HOMO Lewis base:
  
  [NH₄]⁺      +     OH^–      →    NH₄OH

• Alkyl carbenium ion transfer from the onium ion to the lobe-HOMO Lewis base:
  
  [(CH₃)₃O]⁺   + :NR₃   → (CH₃)₂O:  +  [CH₃–NR₃]⁺

There are many congeneric series and two dimensional congeneric planars, for example the alkyl ammonium halide planar:

Type 12 Lewis Acid/Base Complexation Chemistry

Charge controlled ionic bonding, but... a proton transfer reaction is likely to occur rather than the formation of a complex.

Onium ions generally do not often form complexes with π-HOMO Lewis bases because H⁺ or R⁺ transfer (from the onium ion to the π-HOMO Lewis base) is likely to occur.

The acetate ion can be dual classified as both as a Lobe-HOMO and π-HOMO Lewis base (the acetate ion is isoelectronic with the allyl anion). Thus, ammonium acetate can be considered to be a Type 12 complex:
     [NH₄]⁺ [CH₃COO]^–
(but it is probably better to consider it as a Type 11 complex.)

Type 13 Lewis Acid/Base Complexation Chemistry

s-HOMO Lewis bases are reducing agents, so complexation with hydride or hydrogen is also classed as reduction.

However, with respect to carbenium ions and organic chemistry, as C-H bonds are polarised C^δ––H^δ+, ie as weak Brønsted acids, Type 13 complexes reverse their C-H bond polarisation:

C⁺ +   H^–    →   C–H    →   C^δ––H^δ+

and so become Type 3 complexes.

Carbenium ion (ie where the Lewis acid = R₃C⁺ + hydride ion) complexation reactions are rare as both species interact strongly with their counter ions.

• The true X + Y → X-Y complexation reaction would have to be performed under high vacuum vapor phase conditions with the reactive species generated by an experimentally exotic process such as laser induced dissociation.
• Alkyl halides, such as 3-bromohexane, are readily reduced by H₂ plus a catalyst, or by hydride ion donor reagents, as are carbonyl functions.
• Proton abstraction may occur if there is an acidic proton present and a strongly basic hydride donor reagent is used.
• A few C-H functions can act donors of hydride ion: aldehydes in the Cannizzaro reaction and isopropoxide ions during Meerwein-Ponndorf-Varley reduction.

Type 14 Lewis Acid/Base Complexation Chemistry

Type 14 complexes include electrophilic haloenium ion or carbenium ion reagents with non-electrophilic counter ions.

Lobe-LUMO Lewis acids have a strong symbiotic affinity for halogen anion ligands, close in type to those they already possess. Thus:

•   AlCl₃ has a high affinity for Cl^– to give [AlCl₄]^–

•   FeBr₃ has a high affinity for Br^– to give [FeBr₄]^–

Type 14 complexes are not prepared from the cation plus anion, but from the dihalogen, organic halide or acyl halide plus halophilic Lewis acid:

Cl₂        +      AlCl₃      →        Cl⁺   [AlCl₄]^–

R–Cl     +      AlCl₃      →        R⁺    [AlCl₄]^–

The resulting Type 14 complexes act as sources of "naked" electrophilic Cl⁺, Br⁺ and R₃C⁺, etc. that are able to undergo electrophilic aromatic substitution with benzene and other aromatics:

Type 15 Lewis Acid/Base Complexation Chemistry

The normal strong covalent and polarised covalent s-bonds of main group chemistry are generally Lobe-LUMO/Lobe-HOMO complexes:

C-C    C-N    C-O    C-F    C-Br     C-Nu    C-Nfg    Cl-Cl   O-O   S-Cl     P-Br    ^–B-N⁺   etc.

In solution Lobe-LUMO/Lobe-HOMO complexes may be covalently bonded, polar-covalently bonded or strongly ion-paired, weakly ion-paired or solvent separated.

Prochiral tertiary carbenium ions – carbenium ions with three different ligands and a solvent separated Lobe-HOMO Lewis Base counter ion – react with nucleophiles to give racemic mixtures of enantiomers:

Chiral sp³ carbon/nucleofuge complexes react with nucleophiles via a concerted S_n2 mechanistic process in which the chiral centre inverts ‘like an umbrella’, a 'Walden inversion'.

Strong ion-pairing makes concerted second order substitution more favorable than step-wise first order reactivity.

Acyl chlorides and other carbonyl/Nfg complexes also undergo concerted nucleophilic substitution, although the mechanism is rather different:

There are many congeneric series and planars, for example:

Type 16 Lewis Acid/Base Complexation Chemistry

Classical Organic Chemistry

Many vacant p orbital Lobe-LUMO Lewis acids, particularly carbenium ions and acylium ions, are aggressive electrophiles, E⁺, able to react with π-HOMO organics, such as alkenes, via electrophilic addition or electrophilic addition-followed-by-elimination. For example:

An electrophile, E⁺, may react with benzene and other aromatics to give (Friedel-Crafts or nitration) electrophilic aromatic substitution products:

Few congeneric series or planars of interest.

Type 17 Lewis Acid/Base Complexation Chemistry

π-System Lewis acids are reduced by s-HOMO Lewis bases:

Hydride ion donor reagents, such as lithium aluminium hydride or sodium borohydride, both Type 6 complexes:

H₂ plus transition metal catalyst:

During such reductions the hydrogen is likely to be in the form of a transition metal/hydride Type 21 complex.

The saline hydride reagents, NaH etc., Type 5 complexes, are generally too Brønsted basic to be used as organic reducing agents.

Type 18 Lewis Acid/Base Complexation Chemistry

Charge controlled ionic complexes which form solvent separated ions in polar solvents so making the interesting π-cations appear ‘naked’ for spectroscopic study.

Species can only be truly naked under high vacuum vapour phase conditions or in mathematical in silico computer models.

Exotic cationic π-systems, such as allyl & pentatrienyl cations:

cyclopropenyl, cyclobutdienyl, tropylium & cyclooctatetrenely cations:

are reactive electrophilic entities that require very non-nucleophilic anionic counter ions, and tetrafluoroborate, [BF₄]^–, and hexafluorantimonate, [SbF₆]^–, type ions are ideal.

Type 19 Lewis Acid/Base Complexation Chemistry

A number of interacting congeneric series are known, for example, phenyl methane carbenium ion/halogen anion complexes:

Type 20 Lewis Acid/Base Complexation Chemistry

If the π-LUMO Lewis acid and π-HOMO Lewis base species have suitable:

• Shape (geometry)
• Charge
• and Phase matched FMOs

an orbital phase-symmetry controlled multi-centre π/π complexation reaction can take place with various the participating atoms rehybridizing to give a largely σ-bonded product. The transition state is deemed to proceed via a pericyclic intermediate, where "peri-" is a prefix meaning around or surrounding.

The classic pericyclic interaction is the Diels-Alder cycloaddition between a diene and a dienophile:

The orbital phase symmetry arguments required for cycloaddition to take place are interchangeable. Cycloaddition can either involve the HOMO of the diene + the LUMO of the dieneophile or the LUMO of the diene + the HOMO of the dieneophile:

Pericyclic reactions are concerted: they take place in a single step. As a consequence, concerted processes provide allow for great Stereochemical control, and pericyclic processes are amongst the most useful of all synthetic methodologies available to the synthetic organic chemist.

Pericyclic chemistry is discussed in detail elsewhere in this webbook.

Electron-poor δ⁺ π-systems can interact with electron-rich δ^– π-systems. The initial attraction between is electrostatic (ionic) in nature to form a π/πcharge transfer complex.

FMO Controlled Pericyclic Interactions

Diels-Alder cycloaddition can be considered as multicentre π-LUMO plus π-HOMO Lewis acid/base Complexation chemistry. However, it is sometimes difficult to decide which species is acting as the Lewis acid and which is the Lewis base.

In normal electron demand Diels-Alder cycloaddition chemistry the diene is electron rich, implying Lewis base character, and the alkene, the dieneophile, is electron deficient implying that the species is the Lewis acid:

In reverse electron demand cycloaddition, the diene acts as the electron deficient Lewis acid and the dieneophile is the electron rich Lewis base:

There are four general classes of pericyclic reaction process:

• Cycloaddition
• Electrocyclic Reactions
• Sigmatropic Rearrangements
• Group Transfer Reactions

Pericyclic chemistry is discussed in detail elsewhere in this webbook.

Charge Transfer Complexes

π/π-Interactions can also lead to the formation of a charge-transfer complex, for example between the electron poor 1,3,5-trinitrobenzene and electron rich benzidine.

π/π-Interactions can also lead to the formation or conductive organic metals such as stacked TTF/TCNQ materials:

Type 21 Lewis Acid/Base Complexation Chemistry

• Metallic or interstitial hydrides are formed by many d-block and f-block elements when heated with hydrogen under pressure. The hydrides tend to be non-stoichiometric and they may be of variable composition.
• There is a hydride gap where elements do not form hydrides. This roughly maps to the siderophile elements.
• Intermediate hydrides have properties between metallic and covalent.

Metallic properties and non-stoichiometric compositions are common. The bonded hydrogen may be mobile. In various models the hydrogen may be present as H⁺ with the lost electrons going to the metal’s d-orbitals. In other models the hydrogen is assumed to have acquired electrons from the metallic conduction band and so be present as H^–.

Note that even though a metal may not form a hydride, it can still form a hydride complex ion, for example rhenium:

Re⁷⁺    +   9H^–      →     [ReH₉]^–

Hydride and hydrogen donors. Transition metal hydrides, MH_n, can sometimes be formed by ‘dissolving’ hydrogen gas into a bulk metal to from a metal/hydride phase. This phase is the active ‘hydrogenating agent’ employed during catalytic hydrogenation. Palladium metal is able to absorb 900 times its own volume of hydrogen gas.

Transition metal/hydride coordination complexes can be formed by nucleophilic displacement of a halogen (or other nucleofugal ligand) by hydride ions supplied by LiAlH₄ or NaBH₄:

Metallic hydrides materials are usually dark powders or brittle solids. The bonding mechanism is important because heavy metal hydrides are being actively considered as hydrogen storage materials for hydrogen powered automobiles.

Type 22 Lewis Acid/Base Complexation Chemistry

There are few common heavy metal/complex anion complexes. That said, Pearson states in his early HSAB publications that transition metal ions of high oxidation state are harder than those of low oxidation state.

Thus, we would only expect transition ionic metal/complex anion complexes to form with the harder higher oxidation state transition metal ions.

This is what is found: the hard/hard copper(II) tetrafluoroborate complex, Cu[II] (BF₄)₂, is known, but the mixed soft/hard copper(I) tetrafluoroborate, Cu[I] BF₄, is not.

However, mixed soft/hard complex silver tetrafluoroborate, AgBF₄, is known and is a useful chemical reagent.

Type 23 Lewis Acid/Base Complexation Chemistry

The nature of bonding in heavy metal complexes, particularly transition metal complexes, is a vast subject with a long history and is of great technological importance: indeed the term ‘complex’ was first used to describe coordinated transition metal ions.

The two most important bonding models are the electrostatic ligand-field theory and molecular orbital (MO) theory.

Ligand-field theory assumes that the metal cation + anionic ligand interaction is primarily ionic. However, the more demanding MO model is computationally more accessible.

The geometry of heavy metal complexes can usually be predicted by the valence shell electron pair repulsion VSEPR method, however, the effect of d and f orbitals must also be considered: [FeCl₄]^– is tetrahedral but the d⁸ complex [PdCl4]^– is square planar.

Transition metal complex ions may exhibit Jahn-Tellar distortion: Wikipedia, Robert J. Lancashire page & ScienceWorld.

Lobe-HOMO Lewis bases form a multitude of complexes with heavy metal cations, indeed a great deal of classical inorganic coordination chemistry involves Type 23 complexation. Complexes may be linear, [AgCl₂]^–, tetrahedral, Ni(CO)₄ or octahedral:

Heavy metals show multiple oxidation states due to loss of different numbers of d and f orbital electrons. Redox considerations can be as important as HOMO/LUMO interactions when predicting reactivity.

Heavy metal complexes are usually coloured due to electronic transitions involving d or f orbitals and spectroscopic study can be used to probe the nature of the bonding in the complex. Sometimes small changes in ligands give rise to large spectral changes, and sometimes not.

As a rule, d-d electronic transitions give rise to pale colors whereas charge-transfer transitions result in complexes with dark colours.

Transition metals are employed in many biological systems where a protein molecule has a heavy metal ion at the active site, including: Hemoglobin, zinc finger protein & cytochrome:

Structure of heme b 

Cytochrome c with heme c.

Industrial process catalysts may be heterogeneous (solid catalyst, liquid or gaseous reaction mixture) or homogeneous with respect to the reaction mixture. Wilkinson's catalyst, chlorotris(triphenylphosphine)rhodium(I), facilitates the homogeneous hydrogenation of alkenes:

The phosphine ligands can be modified so as to give a chiral versions of Wilkinson's catalyst that are able to catalyse asymmetric hydrogenation. This approach has been developed into the Monsanto method for the production of L-DOPA, from Wikipedia:

Note that the hydrogenation catalyst is in this reaction: [Rh(R,R')-DiPAMP, COD]⁺ [BF4]^–, has a non-nucleophilic tetrafluoroborate (complex anion Lewis base) counter ion.

Type 24 Lewis Acid/Base Complexation Chemistry

The bonding in π-organometallics must be considered using LCAO MO theory because the normal rules of valency break down. For example, is the valency of the Fe(II) ion in the metallocene in ferrocene, Fe(C₅H₅)₂? Is the iron 2 valent or 10 valent?

The hapto nomenclature, η^x, where the hapacity (Wikipedia) gives the number of conjugated p orbitals which ligate to a metal, rather than the number of electrons.

There is a very extensive π-heavy metal chemistry known: all heavy metals are known to form π-organometallic compounds.

• Organotitanium chemistry
• Organochromium chemistry
• Organomanganese chemistry
• Organoiron chemistry
• Organocobalt chemistry
• Organonickel chemistry
• Organocopper chemistry
• Organozinc chemistry
• Organopalladium chemistry
• Organosilver chemistry
• Organocadmium chemistry
• Organotin chemistry
• Organoiridium chemistry
• Organoplatinum chemistry
• Organogold chemistry
• Organolead chemistry

For example, chromium metallocenes:

Crabtree's catalyst:

Hydrogen Bonding

Proton Held Between Two Lewis Bases: Lewis Base/Proton/Lewis Base Complex

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:

The time averaged effect, the superposition, is for the two Lewis bases to be attracted to each other through the hydrogen atom.

But why?

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 to exhibit hydrogen bonding.

Another strange manifestation of hydrogen bonding is that water has its maximum density at 4°C. Thus, water ice is less dense than liquid water and it floats. Most solids have greater density than the liquid. Crystal structure of hexagonal ice. Gray dashed lines indicate hydrogen bonds (Wikipedia):

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 (from Wikipedia):

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:

Diagram of a section of β-pleated sheet with H-bonding between protein strands:

Chemical structure of DNA. Hydrogen bonds between A=T and between C≡G are shown as dotted lines:

Three-Center Two-Electron Bridge Bonding Ligands

A 3-center-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:

The compound commonly called "trimethylaluminium, Al(CH₃)₃," is actually the dimer with the formula Al₂(CH₃)₆, Wikipedia:

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, Wikipedia:

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, O^2–
• Sulfide, S^2–
• 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.

van der Waals Complexation

There are several types of van der Waals attraction:

• Permanent-dipole/Permanent-dipole
• Permanent-dipole/Induced-dipole
• Instantaneous-induced-dipole/Induced-dipole (the London force)

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.

Permanent-Dipole/Permanent-Dipole Attraction:

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. 

Permanent-Dipole/Induced-Dipole Attraction :

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, CH₄, are 'sticky'.

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 or London forces?

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 millimeter of a gecko's footpad contains about 14,000 hair-like setae." Wikipedia:

As well as giving an explanation of the forces involved in the melting, sublimation and boiling of molecular materials, van der Waals forces can help explain dissolution of a solute in a solvent.

As a first approximation: "Like Dissolves Like":

• 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.

Guest-Host Complexation

Guest-Host Complexation

"Complexes composed of two or more molecules or ions held together in unique structural relationships by hydrogen bonding, ion pairing or by van der Waals force, but not by full covalent bonds." Wikipedia, p-xylylenediammonium bound within a cucurbituril:

Molecular Recognition

"The term molecular recognition refers to the specific interaction between two or more molecules through non covalent bonding such as hydrogen bonding, metal coordination, hydrophobic forces, van der Waals forces, pi-pi interactions, electrostatic and/or electromagnetic effects. The host and guest involved in molecular recognition exhibit molecular complementarity." Wikipedia

Inclusion Compounds

"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, (NH₂)₂C=O, 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:

Biological Systems

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.

Another, out of countless possible examples:

Zinc finger proteins coordinate one or more zinc ions to help stabilize their folds. The diagram represents the Cys2His2 zinc finger motif, consisting of an α-helix and an antiparallel β-sheet. The zinc ion (green) is coordinated by two histidine residues and two cysteine residues.

Zinc finger proteins typically bind DNA, RNA, proteins or small molecules. The diagram represents the protein Zif268 (blue) containing three zinc fingers in complex with DNA (orange). The coordinating amino acid residues and zinc ions (green) are highlighted.

All biochemistry and molecular biology, totally outside the scope of this webbook!