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Rules for Predicting Molecular Geometry 1. Sketch the Lewis structure of the molecule or ion 2. Count the electron pairs and arrange them in the way that minimizes electron-pair repulsion. 3. Determine the position of the atoms from the way the electron pairs are shared.
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1. Sketch the Lewis structure of the molecule or ion
2. Count the electron pairs and arrange them in the way that minimizes electron-pair repulsion.
3. Determine the position of the atoms from the way the electron pairs are shared.
4. Determine the name of the molecular structure from the position of the atoms.
5. Double or triple bonds are counted as one bonding pair when predicting geometry.
A dipole arises when two electrical charges of equal magnitude but opposite sign are separated by distance.
The sum of these vectors will give us the dipole for the molecule
equilibrium bond distance
The point at which the potentialenergy is a minimum is called the equilibrium bond distance
The combination of an s orbital and a p orbital produces 2 new orbitals called sp orbitals.
These new orbitals are called hybrid orbitals
The process is called hybridization
What this means is that both the s and one p orbital are involved in bonding to the connecting atoms
Everything we have talked about so far has only dealt with what we call sigma bonds
Sigma bond (s) A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.
When a molecule has two or more resonance structures, the pi electrons can be delocalized over all the atoms that have pi bond overlap.
Benzene is an excellent example. For benzene the p orbitals all overlap leading to a very delocalized electron system
In general delocalized p bonding is present in all molecules where we can draw resonance structures with the multiple bonds located in different places.
These two new orbitals have different energies.
The one that is lower in energy is called the bonding orbital,
The one higher in energy is called an antibonding orbital.
Molecular Orbitals (MO’s) from Atomic Orbitals (AO’s)
1. # of Molecular Orbitals = # of Atomic Orbitals
2. The number of electrons occupying the Molecular orbitals is equal to the sum of the valence electrons on the constituent atoms.
3. When filling MO’s the Pauli Exclusion Principle Applies (2 electrons per Molecular Orbital)
4. For degenerate MO’s, Hund's rule applies.
5. AO’s of similar energy combine more readily than ones of different energy
6. The more overlap between AOs the lower the energy of the bonding orbital they create and the higher the energy of the antibonding orbital.
1) 1 sigma bond through overlap of orbitals along the internuclear axis.
2) 2 pi bonds through overlap of orbitals above and below (or to the sides) of the internuclear axis.
The s2s and s2p molecular orbitals interact with each other so as to lower the energy of the s2s MO and raise the energy of the s2p MO.