July 23, 2014

Homework Help: Chemistry

Posted by Gabriela on Friday, April 27, 2007 at 11:33pm.

I am really confused about polar and non polar molecules. I really don't understand how to know if something is polar or not. Or about bond polarity vs. molecular polarity (net dipole moment).

Also, if it has a lone pair when you draw a lewis structure does this automatically make it polar right?!?

I am also confused about pi and sigma bonds. I know that a single bond is a pi bond and that a double bond has one pi and one sigma and that a triple bond has two pi and one sigma.

Polar and non-polar organic molecules.
Look at the electronegativity. The difference in electronegativity between two atoms will tell you if that particular bond is polar. For example, C and H form bonds in organic compounds and the bond is slightly polar. The EN for H is 2.1 and that for C is 2.5. Not very polar but slightly. The bond between O and C is more polar since the EN of O is 3.5 and the difference is 1.0 instead of 0.4. Therefore, those individual bonds would have dipoles associated with them. And that's about the story for individual bonds. What about when these bonds are in molecules? If the molecule is completely symmetrical, then all of the individual dipoles cancel and there is no net dipole for the molecule as a whole. If the molecule is not symmetrical, the individual dipoles don't cancel and the molecule has a net dipole moment. For example. CH4 is a tetrahedral molecule with bond angles of 109o28'. Therefore, although EACH C-H bond is slightly polar, the net dipole moment of the CH4 molecule is zero. It may not be so easy to see spatially, especially with a tetrahedral molecule such as CH4. So let's try an easier one. CO2. That is a linear molecule. O=C=O.
So there is a dipole between the left O and C and another between the right O and C, BUT the left CO dipole is pulling one way, the right CO dipole is pulling the other way, and the net effect is that they cancel each other out and the net dipole moment for CO2 is zero. That's how we know H2O is not a linear molecule. H2O could be drawn as
H-O-H and (and it was drawn as a linear molecule back in the OLD OLD days) IF HOH was a linear molecule, it would have no net dipole moment although there are relatively strong dipoles between O and each individual H. But we can measure the dipole moment for H2O and since it has one, we KNOW it can't be linear. That's why we always draw it as an angular molecule, something like this.
For all these reasons, we look at CH3OH and we can see that the left end is H-C and the right end is O-H; therefore, we make an educated guess that CH3OH is a polar molecule, like water. The longer the chain, the more like a hydrocarbon, and the less polarity as a solvent. So CH3CH2CH3OH is polar but not as polar as CH3OH. And both of the alcohols are more polar than something like pentane. Pentane is CH3CH2CH2CH2CH3. One end looks just like the other so pentane is not a polar solvent.

Second. In most cases, a lone pair of electrons will make the molecule have a net dipole moment. NH4^+, for example is symmetrical and has no net dipole moment, but NH3 has that lone pair hanging on (from the H^+ we pulled off), so those individual N-H bond dipoles are not canceled out because of symmetry. NH3 has a net dipole moment, NH4^+ does not.

Third. You are a little confused about the sigma and pi bonds. A single bond is a sigma bond, not a pi bond. You are correct that a double bond consists of a sigma and a pi bond. And a triple bond consists of a sigma and two pi bonds. You don't tell me what you are confused about so I can't explain more; however, knowing that a single bond is a sigma bond may be enough to clear up that part of your confusion.
I hope this helps.


On the subject of lone pairs of electrons, note that both O atoms in the O=C=O molecule have lone pairs of electrons but the molecule still does not have a net dipole moment.

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