When my high school chemistry class gets to polar molecules, we run into trouble because most of the students have never heard of the word “vector” and, even though we’ve covered VSEPR theory, they don’t have the spatial reasoning abilities to see how polar bonds could cancel out or contribute to a dipole by looking at drawings on the board. After doing my traditional explanation of how polar bonds can create polar molecules using Lewis diagrams and arrows, we move to the lab portion of the classroom where we have more space. I ask a couple of volunteers to model a simple molecule like HCl using a long (1-2 m) piece of tubing.
The rules for the molecules are as follows:
- The tubing represents the bonding electrons.
- Students represent the atoms.
- 3D shapes of molecules are determined by VSEPR theory.
- Students tug on the tubing based on the electronegativity of the atoms they represent.
- Lone pairs are not depicted.
Since the tubing is stretchy and can be pulled by the students, we have the ability to depict polar bonds and thus, polar molecules. With help from the spectators, we figure out the effect the electronegativity of chlorine and hydrogen have on how hard each person tugs on the tubing. I note that hydrogen can get only so close to chlorine due to repulsion of their
electron clouds, so the hydrogen person should not be moved by the tugging chlorine, just leaning toward the chlorine person. Then I ask the audience if this is a polar molecule, and they usually all get that it is, indeed, polar. To make sure everyone thoroughly understands, I give an explanation, sometimes referring to a Lewis structure drawn on a nearby white board. It’s usually helpful to ask the volunteers to turn themselves into a diatomic molecule such as H2 or F2 to show that no matter how hard the atoms tug on the electrons in this tug-of-war match, the key is how hard they pull relative to their competition.
Next, we model CO2, a more complex situation. I ask for three volunteers and give them four pieces of tubing, noting that two pieces of tubing represent a double bond. As they get themselves into position, often with the help of a Lewis structure, I review VSEPR theory by reminding everyone that electrons position themselves so that they are as far apart as possible.
The audience critiques the model until all three people are positioned in a straight line a couple meters long. Next, we factor in the electronegativity of each atom so the oxygen people start pulling on the central carbon person (the oxygens might need to be reminded that they are essentially twins with equal electronegativity/strength). I ask the carbon person if she is being pulled one way or another. After some thought, she usually concludes “no”. I explain to the class that this is a situation in which polar bonds cancel each other out and refer to a Lewis structure of CO2 showing two dipole arrows. Said in simpler terms, since the carbon person is not being forced to lean in one direction or another, this molecule is nonpolar.
Moving on to an even more complex situation, I ask for four volunteers, give them three pieces of tubing and ask them to model BF3, a trigonal planar molecule. It’s a little tricky to get all the bond angles to be 120° and for everyone to pull equally, but with a stretch of the imagination, everyone, especially the central boron person, can tell that boron’s electrons are not being pulled in any one direction and that this molecule is nonpolar. This is easier to accept if I give this same group another piece of tubing and tell them to model CH2O, another trigonal planar molecule. As the volunteers consider their electronegativity and adjust their tugging strength, everyone, including the audience, can see that the central carbon is pulled in one direction, toward the oxygen, making this variation of a trigonal planar molecule polar.
Now that the class understands how to model molecules and has warmed up with simple linear or planar particles, we move on to even more challenging molecules: those with tetrahedral and pyramidal shapes. Since we have already discussed VSEPR theory and the impact lone pairs and bonding electrons have on molecular geometry, we can focus on the effect that shape has on molecule polarity.
First, I ask a group of students to model NH3, a trigonal pyramidal molecule. At first, they usually make a planar molecule, so I remind them that the 3D shape is really important in determining polarity. After a little thinking, often times with suggestions from the spectators, the hydrogen students squat down and the nitrogen student stands tall, often holding the tubing above his head.
For the final challenge, I ask a group of volunteers to model CF4. It takes a bit of creativity to arrange themselves in a tetrahedron, but once the molecule is properly formed and the fluorine people start pulling on the tubing, the carbon person realizes that there is no clear direction in which he is being pulled. (It’s unlikely that the carbon will feel no net tugging direction since it’s really hard to get the angles and pulling strength perfect in a tetrahedron).
Once everyone understands why this is a nonpolar molecule, we go through related molecules by substituting hydrogens for fluorines, asking the carbon to report on the net tugging direction that he feels.
To get the most mileage out of this activity, I don’t randomly assign people to portray the atoms. Instead, I pick the people who have a very weak understanding of polarity (or chemistry in general) and ask those people to be the central atom so that they can benefit the most from the first-hand, kinesthetic experience of this activity. By repeatedly asking for a new set of volunteers, nearly all the students are able to experience
polarity. It’s helpful to model particles whose Lewis structures we have already drawn and are displayed somewhere in the room. I also have some ball-and-stick models handy to remind the students what these shapes look like and help position themselves accurately.
Once we’ve finished this activity, the students definitely have an improved comprehension of polarity and I find it is a memorable activity that I can refer to later in the year when someone asks me a question about dipoles. Not only does the activity help my students understand polarity, but it’s also very effective at deepening their comprehension of electronegativity, VSEPR theory and molecular geometry, plus, it improves their ability to
look at a 2D Lewis structure and visualize what these particles look like in 3D.