Intermolecular Forces (2) Permanent Dipole Bonding

So this is Chemical Bonding (7) and we are not finished yet by any means.

Let’s discuss today permanent dipole bonding. 

You can call van der Waals forces temporary dipoles because these dipoles are forever changing due to the excited movement of electrons

Electrons shifting around atoms leaves charge imbalances that set up these temporary dipoles even in neutral atoms like those of the Noble gases.

Question now is what kind of molecular conditions are going to have to exist for these kinds of dipoles to be permanent?

Now if you are a fan like me of that massive chemist Linus Pauling, him of the advocate of Vitamin C as a cancer cure (he did live to be around 90+ by the way) then you will know a little about that phenomenon called electronegativity.

His great book called– no marks for guessing this: “The Chemical Bond” (I have a first edition 4th printing!!) is worth selling your shirt for if you haven’t yet obtained a copy.

They are available here

Here’s Linus Pauling’s table of electronegativity values.

Electronegativity values increase with decrease in atomic radius across the periodic table left to right.

You can also see from the above Periodic Table that electronegativity values decrease with increasing atomic radius down a group.

The Halogens Group 7 shows this effect most markedly.  

The highest electronegativity value is 4 for Fluorine, the lowest 0.8 for Francium.

Metals tend to have lower values than non-metals.

Why have we turned to this concept?

Well, if you put together two atoms (not literally you understand) of differing electronegativity then one is going to pull the bonding electrons more than the other atom and, hey presto!! set up a permanent dipole!!

How will this be the case?

It all depends on the meaning of electronegativity.

Pauling defined electronegativity as the power of an atom within a covalent bond to attract the bonding electrons to itself. 

So the greater the atom’s electronegativity the greater its ability to attract electrons to itself.

The greater the difference in electronegativity between atoms of a given bond the more polar that bond is.

You can see in the diagram below how this principle works out for the fluorides of Period 2.

In fact, as an approximation, if the difference is >1.7 then the bond is probably ionic and if < 1.7 the bond is polar covalent but if the values are the same then the bond is pure equal sharing of electrons i.e. covalent. 

Check for yourself in these examples:

a) Sodium at 1.0 chlorine at 3.0 the NaCl bond is ionic which we know to be the case

b) Chlorine at 3.0 hydrogen at 2.1 the H—Cl bond is polar covalent hydrogen chloride is, we know, composed of polar molecules.

c) Chlorine at 3.0 means molecular chlorine itself is non–polar covalent since both atoms in Cl₂ have the same electronegativity. 

Here are the diagrams of these substances and their bonds

The result of these differences in electronegativity is that molecules like water or hydrogen halides contain polar covalent bonds.

E.g.

We also need to remember that the dipole of the molecule is a result of the combined dipoles of each polar bond 

That's why carbon dioxide (CO2 )or tetrachloromethane (CCl4 ) though containing polar bonds are not polar molecules.  

The effect of each polar bond cancels out the effect of each other polar bond.

See carbon dioxide in the diagram below

The final result is that these molecules attract each other through these partial positive and negative charges.

So let's compare the boiling points of haloalkanes with comparable molecular mass alkanes:

What we find is this: the haloalkane has the higher boiling point because of the stronger attraction between polar covalent bonds that create polar intermolecular attractions. 

You should be able to draw out a diagram of two haloalkane molecules using these two polar bonds to attract each other via their weak electrostatic attractions.

Here is an example for you to use as a template for other polar molecules

Note the δ+ and δ– to indicate the position of the significant polar bond and the evident polarity of the molecule.

Note too how these molecules can arrange to provide the position of maximum attraction.