Chapter 9 Covalent Bonding: Orbitals

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1 Chapter 9 Covalent Bonding: Orbitals Atomic Orbital (Valence Bond) Approach Advantages of Lewis Dot structures 1.) Predict geometries 2.) Predict polarities of molecules Disadvantages 1.) No information about energies of electrons 2.) No information about orbitals used in bonding. Valence bond approach is helpful for these. (Basic idea in valence bond) A covalent bond is formed from a pair of electrons with opposite spins in overlapped atomic orbitals. Bond formed from half-filled valence orbitals. Atom 1 Molecule 2 1s Atom 9F Molecule F 9F Filled orbitals do not form bonds e, Ne Problems with Valence Bond Approach: Be (no bonds?) B 1s 2s 2p (1 bond?) C ( 2 bonds?) To explain bonding for atoms like these it is assumed electrons are promoted and hybrid orbitals are formed. Be ( bonds form hybrid) B ( bonds form hybrid) 1s 1s 1s 2s 2p C ( bonds form hybrid) 1

2 ybrid orbitals have new properties that are different from the orbitals used to form them. The geometries are the same as predicted from electron repulsion theory. # electron groups ybridization Geometry angle Two types of Covalent Bonds: 1.) Sigma ( ) bonds form from hybrid orbitals. They have the electron density with bond axis. 2.) Pi ( ) bonds - form from unhybridized p-orbitals They have the electron density but of bond axis. build it! C = C Single bonds have sigma bond(s) and pi bond(s) Double bonds have sigma bond(s) and pi bond(s) Triple bonds have sigma bond(s) and pi bond(s) build it! O = C = O 2

3 build it! Which s are in the same plane? OCC=C 2 2 C=C=C=C 2 build it! N 2 YBRIDIZATION IN MULTIPLE BONDS: The extra electron pairs in multiple bonds (1 extra pair in double, 2 extra pairs in triple) are NOT hybridized. TO DETERMINE AMOUNT OF YBRIDIZATION OF ATOM: ybridize enough orbitals to contain: - all unshared electron pairs - electron pairs to form single bonds - one and only one pair in multiple bonds 3

4 Molecular Orbitals Results of Valence Bond approach to bonding & Lewis Dot Structures Weakness: Inability to predict the correct magnetic properties, O 2 & B 2 Need for resonance to handle special problems Gives no direct information on bond energies Reason - Assumed electrons stayed in atomic orbitals of the individual atoms. Another approach - Linear combination of atomic orbitals to give molecular orbital L.C.A.O. = M.O. BASIC ASSUMPTION OF MOLECULAR ORBITAL APPROAC Orbitals are properties of the molecule not the atoms. Main Ideas of Molecular Orbital Model: 1) Same number of molecular orbitals as the # atomic orbitals that were combined 2) Molecular orbitals can hold two e- with opposite spins 3) square of molecular orbital function indicates e- probability 4) Important properties of orbitals: size, shape, and energy (See fig for shape and size) 5) Molecular orbital configurations can be written much like e- configurations for atoms Ex. ydrogen 1s A 1s * 1s 1s B The orbitals described above are both sigma ( ) molecular orbitals bonding molecular orbital - ( 1s ) antibonding molecular orbital - ( 1s *) ex. 2 : 1s 2 4

5 Example: Predict the molecular orbital configuration in 2 - using the diagram below 1sA 1s * 1s 1sB a) Is this ion stable (does it have lower energy that its separated parts)? b) ow do you expect this bond strength to compare to 2? Bond Order - the difference between the number of bonding e- and the number of antibonding e- divided by two Bond Order = (# bonding e- ) - (# antibonding e-) 2 Calculate the Bond order for 2 and 2 - Bond order is an indication of _ The larger the bond order the the bond Ex. Predict the bond order and stability of e 2 Bonding in omonuclear Diatomic Molecules: omonuclear diatomic molecule - Ex. Predict the Bond Order and stability of Li 2 and Be 2 2s A 2s B 2s A 2s B NOTE: In order to participate in molecular orbitals, atomic orbitals must overlap. Pi ( ) molecular orbitals - (see Fig for shapes of and bonding p-orbitals) ow would you expect the orbitals to compare in energy to the orbitals? 5

6 Expected MO E diagram: What is the molecular orbital configuration for a) F 2? b) B 2? Paramagnetic - Diamagnetic - The expected E level diagram needs to be modified slightly to account for the magnetic properties of B 2. This change results from p-s mixing, thus the 2p and 2p orbitals are reversed Because the importance of p-s mixing becomes less important across the period, the 2p and 2p orbitals revert to the order expected in absence of p-s mixing for O 2 and F 2. 6

7 Bonding in eteronuclear Diatomic Molecules - heteronuclear diatomic molecule - When two atoms are near each other in the periodic table, we can use the MO diagram for homonuclear molecules. Ex. NO (like N 2 ) CN - When the two atoms are different, a new diagram must be used. Ex. F Because 2p is lower in energy than the hydrogen 1s orbital, the electrons prefer to be closer to the fluorine atom. Combining the Localized Electron and Molecular Orbital Models - In molecules that require resonance, the bond is localized while the bonding is delocalized. Ex. NO 3 - and C 6 6 7

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