# without using electronegativity, which one of these molecules is polar? [closed]

If we are given 5 molecules without their electronegativity

BrF5

PBr5

CCl4

XeF2

XeF4

How to determine that one of them is polar or not ?

I know by using the difference in electronegativity we can determine the polarity But again, electronegativity is not given ?? /:

## closed as off-topic by G M, ron, Martin - マーチン♦, Greg E., Jannis AndreskaJul 13 '14 at 17:03

This question appears to be off-topic. The users who voted to close gave this specific reason:

If this question can be reworded to fit the rules in the help center, please edit the question.

• Hi Maher, Use the dissenter answer as a tip, but please show your effort and what you think according to our homework policy! Thanks!:-) – G M Jul 12 '14 at 16:16
• I think he has shown some effort - he says "How to determine that one of them is polar or not ? I know by using the difference in electronegativity we can determine the polarity But again, electronegativity is not given ?? /:" This shows that he understands the basic concept of polarity but is stuck when applying it to molecules. – thomij Jul 13 '14 at 18:12
• @thomij Maher has been encouraged to elaborate a bit beyond just a sentence in his attempts. As such, I'm leaving this closed for the time being. – jonsca Jul 13 '14 at 21:45
• @jonsca Maybe we should include in the guidelines a statement about the required length of the elaboration. Otherwise the standard is not very clear. – thomij Jul 13 '14 at 22:05
• @thomij I think that's a good idea. I'm addressing it on a case-by-case basis at this point, and going more by the spirit of the law than the letter of the law. – jonsca Jul 14 '14 at 0:12

When we say something is polar, it has a very specific meaning. It means that there is a spatial separation of charges, or a dipole moment. In the simplest case, this is given by:

$$\mu = Qr$$

Where $\mu$ is the dipole moment, $Q$ is the charge on each pole (equal and opposite) and $r$ is the distance between the charges.

Anything with charges can have a dipole moment, and therefore be polar. In chemistry, we typically worry about whether or not two things are polar:

1. Bonds
2. Molecules

# Bond Polarity

Chemists are concerned with the polarity of bonds because polarity affects the character of the bond - the more polar the bond is, the more it behaves like an ionic bond. The less polar, the more it behaves like a covalent bond. This has implications for everything from naming rules to reactivity, but all you need to know for this problem is:

All bonds between two atoms are polar, unless the atoms are identical.

(This is including whatever else the atoms may be bonded to). The only question after that is "how polar is the bond," and that is where tables of electronegativity come in. Using an arbitrarily defined scale, we can measure the relative polarity of bonds by comparing the difference in electronegativity between the two elements.

# Molecule Polarity

For molecules, polarity isn't quite so simple. We know that all bonds between dissimilar atoms are polar, but in a molecule, sometimes the dipole moments add up to form a net dipole moment of zero. This is a little bit difficult to explain if you don't have a mathematics and physics background that includes vectors and summations of forces/moments.

If you do, then when I say something like:

The net dipole moment for a molecule is equal to the sum of dipole moments over each bond.

It will make sense to you. If you don't, then you have to get a little more creative. I tell my students to imagine that there are ropes connecting each outer molecule to the central molecule, along each bond. Then imagine pulling each rope towards the side that is more electronegative (has a higher electron density). The more polar the bond, the harder you pull on the rope. Then imagine whether the molecule moves. If the forces balance, it stays put, so the net dipole moment is zero and it is not polar. If it does move, there was a net dipole moment, so it is polar.

This works pretty well - as long as you can visualize the molecular geometry. That's the hard part. To know how the bonds are oriented in space, you have to have a strong grasp of Lewis structures and VSEPR theory. Assuming you do, you can look at the structure of each one and decide if it is polar or not - whether or not you know the individual atom electronegativity. This is because you know that all bonds between dissimilar elements are polar, and in these particular examples, it doesn't matter which direction the dipole moment vectors are pointing (out or in).

Let's look at each one, using wikipedia's geometric pictures.

$\ce{BrF5}$

As you can see, there is a lone pair at the bottom (this is a square pyramidal geometry). The net dipole moment will be pointing "up", which makes this a polar molecule.

$\ce{PBr5}$

This one has a trigonal bipyramidal geometry, with each bond symmetrically opposed to each of the others. If you add up the vectors, they result in a net dipole moment of zero. Therefore, it is non-polar.

EDIT: As ron points out in the comments below, this is not really a binary molecular compound, and in reality forms an ionic crystal structure with $\ce{PBr4+}$ and $\ce{Br-}$. I don't think the author of the question intended you to worry about that, but if they did, the answer is still "non polar," although we wouldn't really call it a "molecule" any more.

$\ce{CCl4}$

Carbon tetrachloride has a tetrahedral geometry, and all the dipole moment vectors cancel. Therefore, it is non-polar.

$\ce{XeF2}$

This one is linear - no net dipole moment, so it is non-polar.

$\ce{XeF4}$

I could only find the ball-and-stick model for this one. It is square planar - there is a lone pair of electrons (not shown) on each "face" of the square on Xe. There is no net dipole moment, so it is also non-polar.

Out of all of these, the only polar molecule is $\ce{BrF5}$ (the first one).

To review the steps:

1. Draw the Lewis structure
2. Figure out the geometry (using VSEPR theory)
3. Visualize or draw the geometry
4. Find the net dipole moment (you don't have to actually do calculations if you can visualize it)
5. If the net dipole moment is zero, it is non-polar. Otherwise, it is polar.
• Isn't $\ce{PBr_5}$ ionic, with a tetrahedral $\ce{PBr_4}$? – ron Jul 12 '14 at 17:26
• Yes, but for the purposes of the homework, I doubt they were looking for students to know that. It's hard to find examples of trigonal bypyramidal molecules without using the same ones over and over, so this one ends up being in lots of assignments. At any rate, you would still call it non-polar (polar and ionic are different) and if you looked at the tetrahedral PBr4, that would also be non-polar. Usually the point of these types of questions is to test whether students can draw Lewis structures, know the geometries, and visualize them in space to decide if its polar. – thomij Jul 12 '14 at 17:36

Use molecular geometries. Some geometries are asymmetrical, so they necessarily are polar. Other molecules are symmetrical, so even though they contain elements of great differences in electronegativities, these molecules are non-polar.

Consider carbon dioxide - which has the elements carbon and oxygen in its composition. The EN of carbon is 2.5; the EN of oxygen is 3.5. Clearly, there is a big difference in electronegativity, which suggests that the oxygens are isolating electron density away from the carbon.

As we can see in the above picture, the two permanent dipoles cancel each other out since they are equal in magnitude and opposite in direction. So carbon dioxide might have elements of great difference in EN, but is still non-polar.

Now apply this knowledge to your problem. To determine molecular geometries, understand that both electrons and attached atoms influence geometry. Try looking up and applying the AXE method. I will leave my answer at here for now, since you have not shown much work in tackling this problem.