Video: Deducing the Polar Molecule from a Set

Which of the following molecules is polar? [A] NH₃ [B] CO₂ [C] F₂ [D] CH₄ [E] CCl₄.


Video Transcript

Which of the following molecules is polar? A) NH₃, B) CO₂, C) F₂, D) CH₄, or E) CCl₄.

A molecule is polar if it has an overall dipole. For instance, a molecule of water has a permanent dipole because it has delta negative and delta positive portions and the individual bond dipoles don’t cancel out. Polar bonds arise when the two atoms on either side of covalent bond have different electronegativities. The electronegativity of an element is a measure of the strength of attraction of atoms of that element to shared electrons in covalent bonds. The most common scale for electronegativity is the Pauling scale, which gives elements a value between zero and four. Values closer to four indicate a stronger force of attraction between atoms and shared electrons.

Now, let’s go back to look at water. Oxygen has a higher electronegativity than hydrogen does. So in the hydrogen–oxygen bonds of water, more of the electron density is drawn towards the oxygen than the hydrogen. This makes the oxygen slightly more negative and the hydrogen slightly more positive. It’s this phenomenon and the geometry of a water molecule that gives rise to the permanent dipole. We can start analyzing our five candidate molecules looking at their geometry, their bond polarity, and then we can figure out if they’re polar or not.

NH₃ is the symbol for ammonia. We start constructing our Lewis structure by putting the atom that can form the most bonds, that’s the nitrogen, in the middle. And we place our hydrogens around it. Nitrogen is in group 15 of the periodic table. It has five valence electrons. Hydrogen is in group one and has one valence electron. We have three hydrogens and one nitrogen, giving us eight electrons to play with, in total, for our Lewis diagram. A single covalent bond between the nitrogen and each of the hydrogens satisfies the outer shell of hydrogen. Hydrogen’s valence shell is the first electron shell, with a maximum occupancy of two electrons. But nitrogen’s valence shell is the second electron shell, which has a maximum occupancy of eight.

We’ve used six electrons in our single covalent bonds. So we have two electrons left. So we can fill the valence shell of the nitrogen atom by adding a lone pair. This uses up our eight electrons. So we have three bonding pairs of electrons and one lone pair around the nitrogen atom. This means that the ammonia molecule is trigonal pyramidal. Now, we can look at the polarity of the bonds to see whether ammonia is polar.

Nitrogen has a greater electronegativity than hydrogen does. So more of the electron density rests on the nitrogen than the hydrogen. And the nitrogen atom is delta negative and the hydrogens are delta positive. These individual bond dipoles and the effect of the lone pair mean that ammonia has a strong permanent dipole. We can clearly divide the molecule in half between delta negative areas and delta positive areas. Meaning that we have a permanent dipole. So almost straightaway, we have our answer. Of the five molecules, the one that’s polar is ammonia, NH₃. But let’s have a look at the others just to be safe.

CO₂ is the symbol for carbon dioxide. Here, we’re putting the atom which can form the most bonds in the middle and the oxygens on either side. The carbon atom contributes four valence electrons to the Lewis structure, while each of the oxygens contribute six. This gives us 16 electrons to play with. We use four electrons in single bonds to bind together the carbon with the oxygens. We can use up the other 12 by adding in lone pairs to the oxygens. But this leaves the carbon four electrons short of a full octet. Instead, we use a lone pair from each oxygen atom and share it with a carbon atom, producing double bonds between the carbon and each oxygen.

This gives us a total of four bonding pairs to the carbon atom, shared equally between two substituents. Without extra lone pairs on the carbon, this enforces a linear structure on the carbon dioxide molecule. The two oxygen atoms have greater electronegativities than carbon does. So carbon dioxide has two polar bonds, a delta positive carbon and two delta negative oxygens. But the two individual dipoles for the bonds cancel each other out. This means that carbon dioxide is nonpolar.

Our next candidate is molecular fluorine, F₂. Regardless of how these two fluorine atoms are bonded together, there is no way that they could have a permanent dipole. The fluorine atoms have identical electronegativities. Therefore, electron density is not going to be one sided in either direction, so we don’t get a dipole. So fluorine is most definitely out of the running.

CH₄ is the symbol for methane. We placed the atom that can form the most bonds in the middle and the hydrogens around that. The carbon atom contributes four valence electrons. The hydrogens contribute one each, giving us a total of eight valence electrons to play with. That’s sufficient to cover the four bonds we need to connect the hydrogens to the carbon. This means that at our central carbon, we have four bonding pairs of electrons and four bonding atoms. This enforces a tetrahedral geometry on the methane molecule. Carbon is more electronegative than hydrogen. So the carbon ends up delta negative while the hydrogens end up delta positive. But all the dipoles cancel each other out.

You can show this by rotating the methane molecule. You’ll always be able to rotate one hydrogen into the position of the others without changing the shape or configuration of the molecule. So all the hydrogen atoms are equivalent. All the hydrogen carbon atom bonds are equivalent. And because it’s tetrahedral geometry, there is no net dipole.

This leaves us with carbon tetrachloride which has a very similar story to methane. Carbon goes in the middle of the Lewis diagram as the one that can form most bonds. And chlorines go around the periphery. The carbon contributes four valence electrons. Each of the chlorines, being in group 17, contribute seven valence electrons, giving us a total of 32 electrons. We use up eight electrons forming single bonds between the carbon and each chlorine, which satisfies carbon’s octet. And the remaining 24 electrons can be used up adding lone pairs to each of the chlorine atoms for filling their octets.

This enforces the same tetrahedral geometry we saw in methane. Chlorine is more electronegative than carbon. So the polarity of the bonds is reversed as compared to methane. Meaning carbon is delta positive and chlorines are delta negative. However, we still end up with pretty much the same result with a nonpolar molecule because of the tetrahedral geometry and the equivalence of the chlorines. Therefore, of our five candidates, the only one that displays a permanent dipole, which therefore can be described as polar, is NH₃.

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