Polar Bonds and Polar Molecules

When you see a Lewis structure like this: hopefully you have come to appreciate what it represents: It represents a molecule of NF3. And that molecule is comprised of one atom of nitrogen that is bonded to three individual atoms of fluorine. And these bonds N F are made up of electron pairs that are being shared between the atoms.

In addition to the three F atoms, the N atom in the center also possesses a pair of electrons that is not being shared with any other atoms. This pair is known as a “nonbonding electron pair.” N F

It’s also important to appreciate what this Lewis structure does not represent. It does not represent the shape of the molecule. In other words, just because the Lewis structure is shaped like a T… it does not mean the molecule is necessarily T-shaped. N F N F

Instead, the molecule will assume the shape that allows these four electron regions… to be as far apart from each other as possible. (Recall that electrons are always negatively charged and therefore tend to repel one another.) This pushes the atoms into a trigonal pyramidal arrangement. F N N F

But let’s take a closer look now at how these electrons are being “shared.” Due to their rather small size and their large effective nuclear charge, the fluorine atoms have a much stronger attraction for the electrons than does the nitrogen atom. This means the electrons in these bonds are not being shared equally between the two atoms, but instead are being hogged by the fluorine atoms. N F F N

Electronegativity is a value that indicates how much an atom tends to “hog” a pair of bonding electrons. So whenever two atoms are bonded together, you can simply compare their two electronegativity values. If one is significantly greater than the other (by 0.5 or more), then the bond between them is considered polar. If they are roughly the same (with a difference of 0.4 or less), then the bond is considered nonpolar. Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5

We can see that F (4.0) has a significantly higher electronegativity than N (3.0) [4.0 – 3.0 = 1.0, which is definitely greater than the 0.5 cut-off.] This means that F is hogging the electrons toward its end of the bond. That gives each F a partial negative charge: And it gives the N a partial positive charge: And it means that the NF3 molecule contains polar bonds. Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 F N d+ d- d- d-

The NF3 molecule contains polar bonds, but is the entire molecule polar? Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 To answer this we simply have to look at how the partial charges are distributed throughout the molecule. If they are distributed symmetrically, they will all cancel out and the molecule will be considered nonpolar. F N d+ d- d- d- But if the distribution is lop-sided, then they will not cancel out and the molecule as a whole will be considered polar.

So we can better see what’s going on, let’s simplify our drawing of the molecule by getting rid of the spheres and replacing the bonding electron pairs with lines: Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 F N F N d+ + Since there only is one positively charged atom (the N), the center of positive charge would be right in the center of the N: d- d- - d- There are three atoms with negative charges (the F’s), and their collective center would be right here below the N:

N F + - Now lets look at a different molecule: CH4. This shows that the charge distribution is lop-sided and therefore the molecule as a whole is polar, with the nitrogen end on top being partial positive, and the fluorine end on the bottom being partial negative. Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 F N + - Now lets look at a different molecule: CH4.

H C C H CH4 would have a Lewis structure like this. And with four bonded atoms repelling each other, the shape would be perfectly tetrahedral: Would the C-H bonds be polar? C (2.5) and H (2.1) have a difference of only 0.4. That’s not quite enough to make the bonds polar. Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5

H C H C And with no polar bonds, the molecule cannot be polar. So CH4 contains no polar bonds, and as a whole the molecule is nonpolar. But what about BF3? Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 H C

B F B F BF3 would have a Lewis structure like this. And with three bonded atoms (and no nonbonding electron pairs), the shape would be trigonal planar: Would the B-F bonds be polar? B (2.0) and F (4.0) have a difference of a whopping 2.0. That’s certainly a big enough difference to make the bonds polar. Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 B F B F

And since the F has the higher electronegativity, the F’s are the ones that are hogging the electrons, so they would get the partial negs: And the B would get the partial pos: So the molecule definitely contains polar bonds. But is the molecule as a whole polar? Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 B F d- d- d+ d-

Since, once again, there is only one atom with a positive charge (the B), the center of positive charge would be right in the center of the B. And again there are three atoms with a negative charge (the F’s), and their center would be… Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5 right in the exact same place as the center of positive charge. This means that the charge distribution is completely symmetrical B F d- d- - d+ + and that the molecule as a whole is nonpolar. d-

Now try problems 1-5 on the polarity tutorial worksheet Now try problems 1-5 on the polarity tutorial worksheet. First determine whether or not the molecule contains polar bonds; then decide whether or not the molecule as a whole is ploar. Remember: for a molecule to be polar it must contain polar bonds and have a lopsided charge distribution Electro-negativity values F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 Se 2.4 Te 2.1 H 2.1 P 2.1 As 2.0 B 2.0 Si 1.8 Be 1.5