So this tutorial is going to be about bond polarity and dipole moments. So let's look at this question. Show the bond polarity and dipole moments in the following bonds and molecules. So let's start with the carbon-oxygen bond.
So how can we show the bond polarity and dipole moment? And also, how can we determine if that bond is going to be polar or nonpolar? So what we have right now is a covalent bond. Carbon and oxygen are both non-metals.
Whenever you have two non-metals combined, typically a covalent bond is formed. And to determine if it's polar or non-polar, we need to look at the electronegativity values. If the E-N difference is less than 0.5, then what you have is a non-polar bond. If the electronegativity difference is greater than 0.5, then you have a polar covalent bond.
Carbon has an electronegativity value of 2.5, and oxygen is 3.5. If you need access to these numbers, I suggest that you go to Google Images and type in electronegativity values from the periodic table or something like that. Now going back to this, we can see that the electronegativity difference in this molecule is 3.5 minus 2.5.
So that's 1.0. So therefore, the carbon-oxygen bond is a polar bond, or a polar covalent bond. Now that we know that oxygen is more electronegative than carbon, we can indicate the bond polarity.
Oxygen has a partial negative charge, and carbon has a partial positive charge. Now to draw the dipole moment, you need to draw an arrow that starts from the positive part of the molecule and points towards the negative part of the molecule. So this is the dipole moment, and this represents the bond polarity.
Now whenever you have a net dipole moment, the molecule as a whole is polar. If all the dipole moments cancel, then the molecule as a whole is nonpolar. But we'll talk about that later. Let's move on to our second example.
The oxygen-fluorine bond. So go ahead and indicate the bond polarity and the dipole moments for this bond, and determine if it's polar or nonpolar. So let's look at the EN values first.
Oxygen has an electronegativity value of 3.5, and fluorine is 4.0. So the EN difference is 0.5, which makes this a polar covalent bond. Both oxygen and fluorine are nonmetals.
And whenever you have two nonmetals combined, typically they form covalent bonds. But we have a polar covalent bond because... fluorine and oxygen, they don't share the electrons equally.
Fluorine is going to pull the electrons toward itself. And because it's more electronegative, it's going to acquire a partial negative charge. Whereas oxygen is going to be electron deficient, it's going to have a partial positive charge. So that's the bond polarity of the oxygen-fluorine bond.
Now the last thing that we have to draw is the dipole moment. And just point the arrow to the element that is most electronegative, in this case, fluorine. And so that's the dipole moment of the OF bond. Now let's move on to part C.
The sulfur-hydrogen bond. Sulfur is a nonmetal and hydrogen is a nonmetal. So therefore this is going to be a covalent bond.
Now is it going to be a polar covalent bond or a nonpolar covalent bond? By the way if we had a metal and a nonmetal typically it would represent an ionic bond. Hydrogen has an election negativity value of 2.1, and sulfur, that is about 2.5, and so the election negativity difference is 0.4, which means that the hydrogen-sulfur bond is relatively nonpolar.
So what we have is a nonpolar covalent bond. Sulfur is more electronegative than hydrogen, so sulfur is going to have the partial negative charge, and hydrogen is going to bear the partial positive charge. So to draw the dipole moment, we're going to point the arrow towards the more electronegative sulfur atom. And so that's it for that example. So now what about molecules?
So here's how you draw water. So that's a simple Lewis structure of the H2O molecule. So given this molecule with the appropriate geometry, indicate the bond polarity and the dipole moment. So first, let's focus on the oxygen-hydrogen bond.
Oxygen is 3.5 and hydrogen has an En value of 2.1. So oxygen bears the partial negative charge and hydrogen bears the partial positive charge. So this bond is a polar covalent bond.
The electronegativity difference is 3.5 minus 2.1, that's 1.4. So it's much greater than 0.5. And the dipole moment of this bond is going to point towards the oxygen atom. So the OH bond is polar. Now if we draw the dipole moments from the molecule, it's going to point towards oxygen.
Now to indicate the bond polarity, we know that hydrogen is going to have a partial positive charge, and there's two of them. So oxygen needs a partial negative charge that's twice the value. We've got to put a 2 in front of it so it can neutralize these two partial positive charges. So make sure you put that 2 in front.
And so that's how you can indicate the bond polarity and the dipole moment of a molecule. So now we need to ask the final question. Is the molecule polar or nonpolar? Because you can have a molecule with polar bonds, but overall, the molecule itself may be nonpolar.
And so you need to look at the arrows, the dipole moments. So we have one arrow going this way, and the other arrow going this way. For the molecule as a whole to be polar, you need to have a net dipole moment. Based on the way this molecule is drawn, notice that the X components of the arrows cancel.
So these two are opposite in direction, they cancel. But the y components do not cancel. They point in the same direction. So therefore, this molecule has a net dipole moment in a positive y direction. So because there's a net dipole moment and these two arrows do not completely cancel, we could say that the molecule as a whole is polar.
Now let's consider methane, CH4. Methane has a tetrahedral structure, so it's not easy to represent it on this screen because it's a three-dimensional object, not a two-dimensional object. Now the carbon-hydrogen bond is nonpolar.
The electronegativity difference is less than 0.5. In fact, it's 0.4. However, carbon still bears the negative charge, and hydrogen bears the partial positive charge. So to indicate the bond polarity of methane, we need to show the four partial positive charge of every hydrogen atom.
For carbon, it's going to have a negative charge, but we're going to multiply it by 4 due to these four charges. So that's the bond polarity of the molecule. Now the dipole moment is going to point towards the electronegative carbon atom.
so we have one going here another going there and if you draw this molecule its 3d structure and you look at the arrows correctly or if you really analyze it the way this structure is set up it's set up in such a way that all of the dipole moments completely cancel each other so that the net dipole moment is zero. So what we have is a molecule that contains nonpolar bonds and at the same time the molecule as a whole is nonpolar overall. Now let's look at our last example, carbon dioxide. So I'm going to give you the Lewis structure of this molecule. It looks like this.
So go ahead and work on that problem. The electronegativity difference between these two atoms is 1.0. So what we have is a polar covalent bond. Carbon is partially positive with respect to oxygen because oxygen is more electronegative than carbon. And so the dipole moment is going to point towards the electronegative oxygen atom.
So this is going to be partially negative, and carbon is partially positive times 2. So now we have a dipole moment that points towards the oxygen atom, and because there's two bonds, we're going to have two arrows. Now notice that these arrows are opposite in direction, which means that they completely cancel out. So the dipole moment for this entire molecule is zero.
So even though the carbon-oxygen bond is polar, the carbon dioxide molecule is nonpolar. So this is one example where you can have a molecule with polar bonds, but the whole molecule is nonpolar overall. Due to the fact that there is no net dipole moment.