Structural Biochemistry/Chemical Bonding/ Electronic Structure Theories


Valence bond theory – this is used to describe hybrid orbitals and electron pairs. It is an extension of the electron dot and bybrid orbital representations. Crystal field theory – this is used to describe the split in metal d-orbitals, which approximates the energy levels from the ultraviolet and visible spectra, but it does not describe bonding Ligand field theory – this is an expansion of the crystal field theory that can be used to describe bonding between the metal ion and the ligands by focusing on the orbital interactions. Angular overlap method – this is used to estimate the orbital energies in a molecular orbital calculation.


Valence Bond TheoryEdit

Valence bond theory was originally proposed by Pauling as a way of hybridization between atomic orbitals. This was one of the major theory used to describing the bonding of coordination compounds, but it's rarely used today. The valence bond theory is filled by the Aufbau principle, which states that electrons are most stable when they filled each orbital with one electron before further filling it. This leads to the Madelung rule, which states that orbitals filled by n+l are filled first before orbitals filled by n+l of higher energy orbitals.

Period 4   Period 5   Period 6   Period 7
Element Z Electron Configuration   Element Z Electron Configuration   Element Z Electron Configuration   Element Z Electron Configuration
        Lanthanum 57 [Xe] 6s2 5d1   Actinium 89 [Rn] 7s2 6d1
        Cerium 58 [Xe] 6s2 4f1 5d1   Thorium 90 [Rn] 7s2 6d2
        Praseodymium 59 [Xe] 6s2 4f3   Protactinium 91 [Rn] 7s2 5f2 6d1
        Neodymium 60 [Xe] 6s2 4f4   Uranium 92 [Rn] 7s2 5f3 6d1
        Promethium 61 [Xe] 6s2 4f5   Neptunium 93 [Rn] 7s2 5f4 6d1
        Samarium 62 [Xe] 6s2 4f6   Plutonium 94 [Rn] 7s2 5f6
        Europium 63 [Xe] 6s2 4f7   Americium 95 [Rn] 7s2 5f7
        Gadolinium 64 [Xe] 6s2 4f7 5d1   Curium 96 [Rn] 7s2 5f7 6d1
        Terbium 65 [Xe] 6s2 4f9   Berkelium 97 [Rn] 7s2 5f9
Scandium 21 [Ar] 4s2 3d1   Yttrium 39 [Kr] 5s2 4d1   Lutetium 71 [Xe] 6s2 4f14 5d1   Lawrencium 103 [Rn] 7s2 5f14 7p1
Titanium 22 [Ar] 4s2 3d2   Zirconium 40 [Kr] 5s2 4d2   Hafnium 72 [Xe] 6s2 4f14 5d2   Rutherfordium 104 [Rn] 7s2 5f14 6d2
Vanadium 23 [Ar] 4s2 3d3   Niobium 41 [Kr] 5s1 4d4   Tantalum 73 [Xe] 6s2 4f14 5d3    
Chromium 24 [Ar] 4s1 3d5   Molybdenum 42 [Kr] 5s1 4d5   Tungsten 74 [Xe] 6s2 4f14 5d4    
Manganese 25 [Ar] 4s2 3d5   Technetium 43 [Kr] 5s2 4d5   Rhenium 75 [Xe] 6s2 4f14 5d5    
Iron 26 [Ar] 4s2 3d6   Ruthenium 44 [Kr] 5s1 4d7   Osmium 76 [Xe] 6s2 4f14 5d6    
Cobalt 27 [Ar] 4s2 3d7   Rhodium 45 [Kr] 5s1 4d8   Iridium 77 [Xe] 6s2 4f14 5d7    
Nickel 28 [Ar] 4s2 3d8 or
[Ar] 4s1 3d9 (disputed)[1]
  Palladium 46 [Kr] 4d10   Platinum 78 [Xe] 6s1 4f14 5d9    
Copper 29 [Ar] 4s1 3d10   Silver 47 [Kr] 5s1 4d10   Gold 79 [Xe] 6s1 4f14 5d10    
Zinc 30 [Ar] 4s2 3d10   Cadmium 48 [Kr] 5s2 4d10   Mercury 80 [Xe] 6s2 4f14 5d10    

Crystal Field TheoryEdit

Crystal field theory was originally developed to describe the structure of metal ions in crystals. The energies of the d orbitals are split by electrostatic field. This was developed in the 1930s, which ignored the covalent bonding since ionic crystals didn't describe it. When the ligands come close to the metal, a big destabilization occurs. The dx^2-y^2 and dz^2 have eg symmetry whereas the dxy, dxz, dyz have t2g symmetry, which implies that they're triply degenerate. eg accounts for 40% of the energy difference while t2g accounts for 60% of the energy. While eg's energy rises, t2g's energy falls to accounts for the difference in rising so that their energies cancel out. This energy difference is called crystal field stabilization energy. This also leads to the idea of high spin and low spin. High spin occurs when Δo > the pairing energy with a weak ligand field and low spin occurs when Δo < pairing energy with a strong ligand field.

Factors that influence Δo: 1) charge : the greater the charge on the central ion, the greater the pull from the ion, which results in an increase in oxidation state.

2) identity of the metal : the greater #d orbitals will have a greater Δo. For example, the 5d orbital can interact more efficiently than 3d, which results in a greater Δo.

3) identity of the ligand : stronger Lewis bases have greater Δo

Ligand Field TheoryEdit

Ligand field theory is used to describe ligand-metal orbital interactions. The metal's d-orbitals match the irreducible representations Eg and T2g. The metal's s and p orbitals have the symmetry A1g and T1u. This leads to a total of 3 bonding orbitals.

Coordination NumbersEdit

Factors involved in determining the overall shape of a coordination compound

1) the number of bonds – since most bonds are exothermic, the creation of bonds will lead to a greater stabilization of the structure

2) VSEPR considerations

3) Occupancy of d-orbitals – the number of d electrons affect the geometry of the coordination compound

4) Steric interference – occurs when large ligands surround the central metal

5) Crystal packing effects – the regular shape is distorted when packed into a crystal. This would hinder the understanding of the original structure since it'd be unclear as if the distortion came from the crystallization process or naturally.

Coordination number 4 – this is usually obtained from square planer and tetrahedral structures with four ligands. Some examples are: CrO4 2-, Ni(CO)4, [Cu{py)4]+. Coordination number 5 – this is usually obtained from trigonal bipyramid, square pyramid, and pentagonal planes. Some examples are: Fe(CO)5 and PF5. Coordination number 6 – this is the most common coordination number since it is usually obtained from the structure of an octahedral. Some examples are: [Co(en)3]3+ and [Co(NO2)6]3-. Coordination number 7 – this is usually obtained from pentagonal bipyramid, capped trigonal prism, and capped octahedron. Some examples of molecules in this structure are: [NiF7]2- and [NbF7]2-. Coordination number 8 – this is usually only obtained in simple ionic lattice structures like CsCl.


Miessler, Gary. Inorganic Chemistry. 4th Edition.

  1. Scerri, Eric R. (2007). The periodic table: its story and its significance. Oxford University Press. pp. 239–240. ISBN 0-19-530573-6.