Ch2102 - Vsepr Theory And Coordination Chemistry

CH2102 - VSEPR Theory and an Introduction to Coordination Chemistry VSEPR – Valence Shell Electron Pair Repulsion VSEPR offers a simple method for predicting the shape of molecular compounds. The combination of p atomic orbitals may result in σ or π molecular orbitals, similarly d atomic orbitals may form σ, π or δ molecular orbitals. These terms reflect the type of overlap between the atomic orbitals to produce the molecular orbitals. The basis of VSEPR theory is that the shape of the molecule is determined by the repulsion between σ valence shell electron pairs. The electron pairs will repel each other and thus move as far from each other as possible and the molecular geometry is determined by the number of electron pairs.

VSEPR applies; 1) Only to valence shell electrons 2) To σ-bonding electrons (not π or δ) 3) To full orbital electrons (i.e. it is not applicable to unpaired electrons)

Using VSEPR Theory; 1. 2. 3. 4. 5. 6. 7.

Count the number of valence electrons of the central atom Add donated electrons from other atoms Adjust for the charges on species Calculate the total number of electrons and divide by 2 to find the number of electron pairs Select the basic geometry Determine which electron pairs are bonding or lone pairs Modify the basic geometry to account for the electron pair interactions In order of increasing repulsion; bonding pair-bonding pair < bonding pair-lone pair < lone pair-lone pair

since lone pairs are localised on the central atom and act as a more concentrated source of charge than the shared bonding pairs which are less localised. It is important to note that when describing the shape of a molecule lone pairs are not ‘visible’, and so ammonia, NH3, is pyramidal and not tetrahedral despite its co-ordination geometry being based upon a tetrahedral arrangement. When determining the shape of a molecule, in the case of several lone pairs, it is important to ensure the repulsions are minimised. Since the coordination geometry is determined only using σ electrons it is necessary to disregard the central atom electrons involved in π or δ bonds. Since each π bond is a shared electron pair with one electron rising from each atom, subtract one electron from each π bond involving the central atom. The π bond will affect the geometry, but only slightly, by pushing the angle to its adjacent atoms slightly over the ideal amount. i.e. in propene the Me-C=CH2 bond is 124.8⁰, slightly higher than the ideal angle of 120⁰. As a d-orbital is essentially 2 π bonds and a σ bond, the overall effect is -1; again its effect on the geometry is only to push adjacent atoms slightly further than the ideal angle. For VSEPR any charge on the molecule is always assigned to the central atom, even if it would seem better placed elsewhere. As such a negative charge is considered an extra electron for the central atom, and a positive charge is shown as the subtraction of an electron from the central atom electron count. In the case of radicals there will be a non-integral number of electron pairs, in this case the 1

unpaired electron takes up its own orbital (i.e. 5 electrons, or 22 electron pairs would take up 3 orbitals). The orbital containing the unpaired electron exerts less repulsion than it would if full, and so the bond angles change accordingly. If an electron were to be added, filling the half-filled orbital, it would then act as expected and the angles would revert to those expected.

An Introduction to Co-ordination Chemistry Werner postulated in the early 1900’s 3 things; 1) Most elements exhibit two types of valence, a) primary valence (or oxidation state) and b) secondary valence (or co-ordination number) 2) Every element tends to satisfy both its primary and secondary valence. 3) The secondary valence is directed toward fixed positions in space (forming the basis of stereochemistry for metal complexes) We have since advanced this area of knowledge and as such there are several modern terms to define; 

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Coordination chemistry – the area concerned with structures, reactivity’s and physical properties of molecules formed by the combination of metal centres (Lewis acids) and electron donors (Lewis bases), Complex – a complex is a species formed by the association of two or more simpler species, each, normally, capable of independent existence Ligand – a ligand is any negative ion or polar(isable) neutral molecule bound to a metal atom. This includes all Lewis bases (nucleophiles and reducing agents). Square brackets are commonly used to denote the complex entity formed between a metal and its ligands. Oxidation number – this is the charge that the central atom in a coordination compound would have if all of the ligands were to be removed along with the electron pairs they donated. It is represented by a Roman numeral. Although not technically the same, the term oxidation number is often used interchangeable with the term oxidation state (the two are only usually different when the ligand atom is less electronegative than the central atom). Coordination number – commonly the number of donor atoms or ligands bound to the metal. However this definition can be misleading when more complex ligands are involved (such as the cyclopentadienyl ligand). It is therefore better to define the coordination number as the number of two electron bonding pairs between a metal and its ligands (the number of coordinate bonds). Coordination bond – a covalent bond in which both electrons are supplied by one of the two components (the ligand) of the bond.

The Coordinate Bond In the case of metal-ligand interactions we use the concept of coordinate (or dative) bonds as defined above. Some characteristics of a coordination bond; -coordination bonds have enthalpies of the same magnitude as those of other covalent bonds. -one atom (the ligand) donates both of the electrons involved in the bond. -the electron pair donor, or ligand, is a Lewis base. the electron pair acceptor (typically a metal atom or ion) is a Lewis acid. -the pair of electrons donated by the Lewis base is most often a lone pair. In a neutral complex the ligands are listed in alphabetical order, followed by the metal atom. The number of ligands is indicated by; bi, tri, tetra, penta, hexa, hepta, octa, nona, deca etc

In anionic complexes the metal ends in –ate, anionic ligands end in –ato Molecular Structure of Coordination Compounds In describing the structure of coordination compounds there are several key concepts, the coordination number, the coordination geometry, the classification of ligands, isomerism and electronic structure. 1) Coordination number – for metal complexes the most common coordination number is 4 or 6, though values from 1-14 are possible 2) Coordination geometry – also called the complex or coordination stereochemistry. This describes the special distribution of ligands in a complex. Each coordination number has an associated coordination geometry (or several) as follows CN=2 – this coordination number is common for the late transition metal complexes e.g. Cu, Ag and Au complexes. The coordination geometry in this case is linear. [H3N-Ag-NH3]+ CN=3 – this coordination number is rare, an example could be Pt(PPh3)3, the associated geometry is trigonal planar Ph3P

120o Pt – PPh3

Ph3P CN=4 – a very common coordination number with two main geometries Tetrahedral; very common, particularly with simple anionic ligands e.g. halides, all bond angles are 109.5o

Square planar; common particularly for late transition metals with d7, d9 and especially d8 configurations

CN=5 – this coordination number is rare for transition metal complexes, but important as many 4 coordinate complexes undergo ligand substitution reactions via 5-coordinate intermediates. Again there are two main geometries Trigonal bypyramidal

Square pyramidal

Square pyramidal and trigonal bypyradmidal are very close in energy and can easily interconvert

CN=6 – another very common coordination number for transition metal complexes, there are, again, two associated geometries Octahedral; in which the L-M-L angles are all 90o

Trigonal prismatic; much rarer

CN>6 – although possible the geometries of coordination numbers higher than 6 tend to be more and more distorted, they’re most commonly found when considering lanthanides and actinides. 3) Classification of ligands – a ligand must be able to donate at least one pair of electrons to a Lewis acid, they are classified according to the number of coordination bonds that are formed with a metal centre a) Monodentate ligands – from mono (one) and dentate (bite), also called unidentates. These donate one pair of electrons to a metal centre. Ligands may be neutral, negatively charged and very rarely positively charged. Neutral ligands;

Other examples include ethers, thioethers, arsines etc.

Anionic ligands;

b) Bidentate ligands – these ligands donate two pairs of electrons to a metal centre and are hence capable of forming two coordinate bonds. These are known as chelate rings, and are very stable structures, one bidentate ligand will form a more stable structure than two monodentate ligands, the enhanced stability is known as the chelate effect, chelating ligands are classified by the size of the chelate ring that they form. 4 member chelate rings;

5 member chelate rings

In many cases two/three molecules of this class of ligand can form two/three chelate rings at one metal centre, forming bis and tris chelates respectively.

6 member chelate rings

c) Tridentate ligands – these ligands donate three pairs of electrons to a metal centre

There is also another class of tridentate ligand (and tetradentate etc.) known as a macrocylic ligand.

d) Tetradentate ligands – these donate four pairs of electrons, and are thus capable of forming four coordinate bonds

e) Hexadentate ligands – these donate six electron pairs and therefore are capable of forming six coordinate bonds

Isomerism in Coordination Compounds There are several types of possible isomerism in coordination compounds; Alternative coordination polyhedral – a very rare form of isomerism, it arises when a coordination compounds can adopt two or more geometries (i.e. square planar and tetrahedral) Coordination sphere isomerism – a common form of isomerism, this occurs where two complexes of identical empirical formulae have differing sets of ligands attached to the central metal ion e.g. [Cr(H2O)6]3+Cl3- -> [Cr(H2O)5Cl]2+Cl2-.H2O -> etc. Violet Green one water molecule of crystallisation Geometric isomerism – again very common, these are isomers which involve a different special arrangement of ligands about a central atom

Ionisation isomerism – a rare form of isomerism, these isomers possess different combinations of ligands in the coordination sphere and as a counter ion, they give rise to different ions when they dissociate in solution and hence may have different conductivities [Co(NH3)5Br]SO4 – violet [Co(NH3)5 SO4]Br – red Linkage isomerism – common, ambidentate (those which have more than one type of donor atom) ligands may bond to a metal atom through different atoms

Optical isomers – common, if the mirror images of a molecule are non-superimposable then they form optical isomers, this is most commonly encountered in bis and tris chelates of octahedral complexes

Polymerisation isomerism – rare, monomers and polymers may have the same empirical formula e.g. [Pt(NH3)2Cl2] and [Pt(NH3)4][PtCl4] (monomer and polymer respectively)

Crystal Field Theory One of the most commonly used descriptions for bonding in transition metal complexes is derived from the crystal field theory. In the crystal field theory the metal complex is represented as a point positive charge surrounded by a set of point negative charges (representing the ligand electron pairs). As well as assuming the charges to occupy points, crystal field theory assumes that electrostatic interactions alone are responsible for the complex formation. Electrostatic repulsions between the point negative charges of the ligands and the valence d-orbitals of the metal atom are responsible for the energy splitting of the d-orbitals, making them no longer degenerate. Crystal field splitting for an octahedral ML6 complex; In the presence of an octahedral crystal field d-orbitals are split into a lower energy triply degenerate set (the t2g orbitals) and a higher energy doubly degenerate set (the eg) separated by an energy gap, ∆o, the crystal field splitting parameter varies with the identity of the ligand and metal, and the charge of the metal atom. The eg set of orbitals (the dz2 and dx2d-y2) point directly along the Cartesian axis while the t2g axis (the dxy, dxz, and the dyz) point in between the axis.

This leads to the following energy profile;

∆o is dependent on the metal, it’s oxidation state and the type of ligand involved.

Crystal field splitting for a tetrahedral ML4 complex; In order to understand the crystal field splitting in a tetrahedral ligand field it is easier to imagine that the four point negative charge (ligands) are occupying the four opposite corners of a cube. In this case the e (dz2, dx2-y2) orbitals are further from the ligands than the t2 (dxy, dxz, and the dyz) orbitals and, as such, are repelled less. ∆t is the crystal field splitting parameter in a tetrahedral field and it is the energy difference between these two sets. Assuming constant bond lengths for the same ligands the relationship between ∆t and ∆o is ∆t=

4 9

in a tetrahedral complex.

∆o due to the lower number of ligands

Counting Electrons in Transition Metal Complexes

Once the number of delectrons has been determined the spin-state for the ion can be determined. 0 1 2 3 8 9 For d , d , d , d , d , d 10 and d configurations there is only one spin state possibility, only one way in which to fill d-orbitals regardless of crystal field strength. 4

5

6

7

For d , d , d and d configurations two possibilities, high-spin and low-spin exist;

High spin; P>∆o (weak field case), here the pairing energy of electrons in the d-orbital is greater than the crystal field splitting value (∆o) – hence electrons fill d-orbitals according to the Aufbau principle. For most, but not all, first row transition metal complexes this situation applies. Low spin; P<∆o (strong field case), here the pairing energy of electrons in the d-orbital is less than the crystal field splitting value (∆o) – hence electrons fill t2g orbitals first and then the eg orbitals. Lowspin complexes are common for 2nd and 3rd row complexes and for 1st row complexes with high field ligands such as CN-) Once ligand field splitting has been accounted for the t2gnegm (octahedral coordination) or emt2n (tetrahedral coordination) configuration can be assigned (where m and n represent the number of electrons in each orbital set). To assign these configurations the number of d-electrons, the coordination geometry and the spin-state must be known.

Once the d-electron configuration has been assigned the spin-only magnetic moment (µs.o.) can be calculated (in units of Bohr magnetons); 1

µs.o. = 2 S S + 1 𝑤ℎ𝑒𝑟𝑒 S = total spin = 2 x the no. of unpaired electrons µs.o. is also sometimes referred to as the effective magnetic moment, µeff.