MOLECULAR ORBITAL THEORY.pptx

Theories of bonding

1. Molecular Orbital Theory (MOT) Human body cells are made up of many thousands of compounds, most of which consist of atoms covalently bonded by sharing electron pairs. We already discuss electron-pair bonds to explain bonding in molecules in term of overlapping of atomic orbitals. Where more than one Lewis structure could be drawn for the same arrangement of atomic nuclei, as in SO 2 the concept of resonance. A major weakness of the VBT is that the theory does not predict the magnetic properties of substances. An example is O 2 which is paramagnetic. It means that O 2 molecules must have unpaired electrons. Diatomic oxygen has an even number of valence electrons (12) and the octet rule predicts that all of these electron should be paired. According to VBT, O 2 should be diamagnetic, but experimental show it is not. Molecular Orbitals The ambiguity b/w experimental and theory for O 2 can be resolved by using alternative model of covalent bonding , the molecular orbital (MO) approach. MOT treats bonding in term of orbitals that extend over an entire molecule. The MO approach involves three basic operations . Step 1: combine valence atomic orbitals to give a new set of molecular orbitals (MOs). The number of MOs formed are equal to the number of atomic orbitals combined . Step 2: Arrange the MOs in order of increasing energy, as will be illustrated later. Step 3: Distribute the valence electrons of the molecule among available MOs, filling the lowest-energy MO first and continuing to build up the electron configuration . (a) Each MO can hold a maximum of two electrons (Pauli exclusion principle). (b) Electrons go into lowest-energy MO available . (c) Hund’s rule is obeyed .

2. Molecular Orbital Theory Homonuclear molecules with MOs originating from s orbitals Homonuclear molecules with MOs originating from s and p orbitals Heteronuclear diatomic molecules In MO theory, valence electrons are distributed over the entire molecule. Molecular orbitals of diatomic molecules arise from adding together (superimposing) atomic orbitals. A linear combination of atomic orbitals (LCAO) creates molecular orbitals (bonding orbitals and antibonding orbitals) n molecular orbitals can be constructed by n atomic orbitals.

3. LCAO-MO (Linear Combination of Atomic Orbitals-Molecular Orbitals) Orbitals split into a bonding (lower) and anti-bonding (higher) orbitals. Electron fill from lowest energy up Types of bonds: s = no nodal plane separates nuclei eg. s + s, p z + p z , s + p P = a nodal plane separates nuclei eg. p y + p y , p x + p x s Designates a molecular orbital that is cylindrically symmetric about the bond axis (with no nodal plane along the bond axis) s 1s When waves interfere constructively , the amplitude increases where they overlap. Increased amplitude in the internuclear region translate to an enhanced probability density ( ψ 2 ) b/w the nuclei. As electrons are in the form of waves

4. Shape of antibonding molecular orbital with a nodal plane. Destructive interference results in antibonding

5. Linear combination of two 1s atomic orbitals form two molecular orbitals. Bonding Molecular orbital known as sigma 1s Antibonding Molecular orbital known as sigma star 1s Molecular diagram for Hydrogen and Helium

6. Hydrogen has one electron in 1s orbital, two hydrogen atoms will share one electron each and will form 1 MO and 1 Anti-MO According to Hund’s rule bonding MO will accommodate 2 electrons which is lowest in energy. Bond Order Bond order = no. of e - in BMO – no. of e - in Anti-BMO/2 B.O = 2 – 0/2 B.O = 1 So, H 2 molecule has single covalent bond.

7. Helium has 2 electrons in 1s orbital, two hydrogen atoms will share 2 electrons each and will form 1 MO and 1 Anti-MO According to Hund’s rule, 2 electrons go to BMO while 2 electrons go to Anti-BMO. Bond Order Bond order = no. of e - in BMO – no. of e - in Anti-BMO/2 B.O = 2 – 2/2 B.O = 0 So, He 2 molecule do not exist or unstable.

8. Let’s find He 2 + molecule exist or not? Here, we have 3 electrons as 1 electron has lost because of positive charge. By filling MOs, 2 electrons will go to BMO while 1 electron will go to Anti-BMO. Bond Order Bond order = no. of e - in BMO – no. of e - in Anti-BMO/2 B.O = 2 – 1/2 B.O = 0.5 There is half bond in He 2 + molecule. It means, this molecule exist.

9. MO diagram for Alkali metals, like H 2 , Li 2 , K 2 etc Bond order for alkali metals is 1. LCAO for p-orbital, resulting in MOs (BMO and Anti-BMO) denoted by pi ( p ) and pi star ( p * ). When p-atomic orbital overlap contructively and destructively.

10. Example of p z atomic orbital overlaping Overlap of p z AO to form MO Sigma bond Energy of MOs formed by overlap of 2s and 2p AOs increase from bottom to top. Order of energy applies to Li 2 through N 2 of the second period.

11. Pz AOs linearly combine and form two MOs BMO and Anti-BMO When constructively overlap form π 2px bonding molecular orbital When destructively overlap from π * 2px antibonding molecular orbital In Pi bond charge density is 0 at the center while charge density is maximum above and below while in sigma bond, the charge density is maximum at the center.

12. Example from another book, Px and Py form Pi bond. Predicted and observed properties of second-period elements

13. MO diagram for diatomic molecules like N 2 , B 2 , C 2 Nitrogen has 5 valence electrons, 2s 2 2p 3 . LCAO of N 2 molecules, 2(2s 2 ) will form BMO and anti-BMO which will accommodate 4 electrons. 2(2p 3 ) will form 3 BMO and 3 anti-BMO.

14. Order has changed for O 2 and F 2 , s 2p falls down because of lower energy. It does not effect the filling of orbitals.

15. Molecular orbital trends for the interaction of s- p z atomic orbitals. The trend changes for Oxygen and Fluorine because of maximum overlap of s-p orbitals.

16. Reason for Para magnetism of O 2 molecule. It has 2 unpaired electrons in antibonding MOs. In figure you can see the liquid O 2 is suspended b/w magnetic field instead of falling down.