I got a 1st in my first year studying chemistry at the University of Birmingham using these revision notes that I have uploaded. They include detail on orbital hybridisation, isolobality, Molecular Orbital (MO) Theory, the LCAO approach, bond order, types of non-bonding and bonding atomic orbital o...
Valence Bond and Molecular Orbital Theories
25 September 2017 15:29
Covalent, Metallic and Ionic Bonding Electronegativity
• Cl2 and O2 are homonuclear diatomic molecules with no electric dipoles, hence the • Electronegativity is an elemental property that can be used to predict
electrons that form the bond(s) are shared equally between the two atoms. This is the relative amount of covalent or ionic bonding in a compound.
covalent bonding. • There is no way to measure electronegativity so Pauling scale is used.
• Elements such as Na or Fe or alloys such as brass are non-molecular solids where one • The difference in X values of elements in the Pauling scale is related to
expects no, or very little, electron transfer between atoms. These are held together by the percentage ionic character of the bonds in the compound. For atom
metallic bonding and this can be considered a type of covalent bonding as there is uniform pair XY:
electron sharing. • Cl2: Therefore the bond is 0% ionic
• LiF and CsCl are non-molecular solids and the bonding can largely be accounted for using • LiF: Therefore the bond is 85% ionic.
electrostatic principles (e.g. Li , Cs , F and Cl ). These compounds transfer electrons to
+ + - - There is always a degree of covalent character.
form anions and cations so are ionic.
Lewis Theory Valence Bond (VB) Theory
• Originated 1916 • Developed by Linus Pauling and superseded Lewis theory.
• Cross dot diagrams e.g. where there is one electron from each atom shared, and the • The first quantum mechanical theory of covalent bonding i.e. it is
electrons are localised. based on atomic orbitals on atoms interacting to form bonds.
○ A problem with the theory is that in reality the electrons delocalise in the bond. Also the • Although Valence Bond Theory has now been superseded by
theory doesn't predict geometry e.g. why CH4 is tetrahedral rather than square planar. Molecular Orbital (MO) theory, it is useful to discuss as it introduces
• Species which in terms of Lewis structures demand the presence of more than an octet of some important concepts (hybridisation or mixing of orbitals, electron
electrons around at least one atom, are called hypervalent. promotion) and some of the terminology that is still used today.
Orbital approach to bonding examples • Promotion is the excitation of an • Predicts the shapes of molecules.
• H2O - central atom is O: [He]2s22p4 electron to an orbital of higher Orbital Hybridisation
energy in the course of bond • Pauling used orbital hybridisation to predict shapes.
formation. • If the central atom uses s- and p-orbitals:
• Electron promotion on energetic • s+p creates 2 sp hybrids - orient themselves to make a linear atom
2 3
• NH3 - central atom is N: [He]2s 2p grounds is favourable as the energy • s+2p creates 3 sp2 hybrids - trigonal bonding around central atom
of formation of 2 extra C-H bonds • s+3p creates 4 sp3 hybrids - tetrahedral bonding
that it would otherwise not be able ○ Each sp3 orbital has the same composition so all four bonds are identical
to form in the ground state is greater apart from their orientation in space
2 2
• CH4 - central atom is C: [He]2s 2p than the promotion energy needed • The large lobes of the hybrid orbitals orient themselves to be as far away from
for carbon to be in the excited state. each other as possible, forming each shape.
• Promotion and formation of four • Hybrid orbitals have pronounced directional character, in the sense that it has
bonds is a characteristic feature of enhanced amplitude in the internuclear region.
• Carbon will not stay in the ground state carbon because the promotional • This directional character arises from the constructive interference between the
to form 4 C-H bonds, goes in an excited energy is quite small to move an s-orbital and positive lobes of the p-orbitals
state where there are four unpaired electron from doubly occupied s to a • As a result of the enhanced amplitude in the internuclear region, the bond
electrons in separate orbitals. vacant p, reducing the electron strength is greater than for an s- or p- orbital alone.
repulsion it experiences in the • The increased strength helps to repay promotion energy.
Orbital Hybridisation Examples
ground state.
Isolobality
• Structurally similar molecular fragments are described as isolobal. The name originates
from the concept that the lobes of the hybrid orbitals have the same symmetry and overlap
of lobes is the key to the geometry of how two fragments combine to form a bond.
• These are all isolobal fragments: ○ Lone pair on Br
○ H 1s orbital As far as X is
concerned they
3
look the same, so ○ Methylene and oxygen are isolobal
○ sp lobe from a CH3 fragment the geometry of
the reaction should
be the same. Molecular Orbital (MO) Theory
• Current approach to describe bonding in molecules and predict
+
trends such as bond length and dissociation energy differences,
Formation of H : magnetism and chemical reactivity.
• ψ(H2+) is the wavefunction that will • MO theory does not predict geometry, in fact the geometry must be
mathematically describe the electron in the known in advance.
molecule • The basic premise is that just as atoms have atomic orbitals (s, p, d, f,
• 4πr2ψ2(H2+) defines the probability of finding etc), some with electrons and some without, molecules have
the electron at some position r in the molecule. molecular orbitals. These form over more than one atom and are
Determining molecular wavefunctions: combinations of atomic orbitals.
• Solve Schrodinger's equation. This can be done • Completely delocalised approach to bonding, electrons go into
exactly for one electron systems like H2+. molecular orbitals and these are filled from the lowest energy state
• When we move to H2 or something more progressively upwards.
complicated, use the Linear Combination of • A quantum mechanical approach - as an electron in an atomic orbital
Inorganic Chemistry I Page 1
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