 So let's look at a couple of examples. We'll compare hydrogen fluoride and hydrogen iodide. So let's first look at the electronegativities of the atoms involved. And we can then calculate the difference in electronegativity for each bond. So for hydrogen fluoride, which is made of a hydrogen and a fluorine atom, the difference in electronegativity will be 1.9. And for the hydrogen iodide, it's going to be 0.4. So you can see that both bonds are polar because there's a difference in electronegativity between the two atoms that make them up. But that the HF bond is more polar than the HI bond. If we now look at the molecules as they're rendered in the Molecules 360 database, it gives us a picture that indicates where the electron density is highest. As we saw on the last slide, you can see that for HF, there's a strong concentration of electron density around the fluorine. And if the electrons are spending more time at that end, it means that that end of the molecule has a partial negative charge. Not a full charge like an ion, resulting from the complete transfer of an electron, but a significant and measurable partial charge nonetheless. Hydrogen iodide, on the other hand, looks green, meaning that the electron density is fairly evenly distributed. You can see there's a slight yellowish hue to the larger iodine atom and a slight bluish tint to the hydrogen end, indicating that the bond is not completely non-polar, but this is a very weak polarity. Now these pretty diagrams are all very well, but we need a quick way of indicating a bond dipole on paper if we don't have rainbow pens. So we do it like this. If I draw out the bond like this, and I've got it in that direction to match the picture, then we draw an arrow along the length of the bond and the pointy bit of the arrow goes towards the more electronegative end and then we put a little sort of cross on the other end like a plus to indicate the more positive end. So in other words, the arrow is pointing in the direction that the electron density is shifting. We can also use a small delta plus and delta minus to indicate the partial charges. The deltas in particular meant to indicate that we're not talking about ions here, it's not a complete charge. And if we do the same thing for hydrogen iodide, we draw the arrow to indicate the bond dipole and we draw in our partial charges. So you can see that polar bonds are not all alike. Some are more polar than others.