I have been trying to figure out these challenge problems in my Chemistry 30 class for the past week but have found no luck. here goes:

 

molecules of NH3 are polar but molecules of BF3 aren't.

 

SbCl3 has a measurable dipole, while SbCl5 doesn't.

 

Explain both problems by means of atomic and molecular structures and also regarding principles of bonding.

Think about the arrangement of bonding and non-bonding electron pairs around the central atoms. Are there any non-bonding pairs left over? is this responsible for the dipole?

I have been trying to figure out these challenge problems in my Chemistry 30 class for the past week but have found no luck. here goes:

 

molecules of NH3 are polar but molecules of BF3 aren't.

 

SbCl3 has a measurable dipole' date=' while SbCl5 doesn't.

 

Explain both problems by means of atomic and molecular structures and also regarding principles of bonding.[/quote']

NH3 has three bonded electrons and one lone pair (N has 5 electrons in its outer shell). All of these 4 entities (approximately) want to go as far as possible from each other due to electrostatic forces. This results in a molecule, with a central N-atom inside a tetrahedron. Three H atoms are at three vertices of the tetrahedron and the fourth vertex of the tetrahedron is occupied by the lone pair. So, there is a negative charge-bias at one vertex of the tetrahedron and a slight positive bias at the other three vertexes.

 

Now, take BF3. Boron has three electrons in its outer shell. All three are involved in a bond. Only three entities now need to be as far as possible from each other. This is achieved by occupying the vertices of a triangle with the fluorine atoms and the boron at the centre of the triangle. Now we have a planar molecule and all three fluorine atoms having the same (small) negative charge bias.

 

Now you should be capable of understanding why SbCl3 is measurably polar and SbCl5 is not (hint: Sb has 5 bonding electrons in its outer shell).