 Recall that when two atoms bond, their relative electronegativities will determine whether the bond is polar or not. For instance, the NN triple bond in a nitrogen molecule is non-polar because the two atoms have identical electronegativities, and this means the electron density is equally shared between them. However, in a water molecule the bonds are between different atoms, oxygen and hydrogen. In each bond the oxygen is the more electronegative atom, so that end of the bond has a greater electron density and therefore a partial negative charge. We draw the bond dipole arrow like this to indicate where the greater electron density is. When you look at a molecule as a whole the polarities of all the bonds combine to give the overall polarity of the molecule. So in water these two bond dipoles, which are at an angle to each other, combine to give a molecular dipole that looks like this. In contrast, carbon dioxide has two polar bonds, but because the dipoles are pointing in opposite directions they cancel each other out, meaning that CO2 is a non-polar molecule. It is these molecular dipoles, or lack of them, that cause intermolecular forces, that is the attractions between molecules. And there are three kinds of intermolecular forces that you know about. Dipole-dipole forces, van der Waalsf or dispersion forces, and hydrogen bonding. Let's look at dipole-dipole attractions. These occur in molecules with an overall molecular dipole, a slightly positive end and a slightly negative end. In this picture here from the UC Davis-Kemwicky these little diagrams are meant to represent HCl molecules. When the molecules encounter each other the positive end of one molecule can attract the negative end of another. This attraction is called a dipole-dipole force or attraction. Note that similar ends will repel each other since they have the same charge. When you put a lot of these molecules together, as long as the sum of the attractive forces is greater than the repulsive ones, they will hang together as a liquid or a solid. Next let's look at van der Waalsf or dispersion forces. All molecules have electron clouds, and these clouds are constantly fluctuating. These fluctuations can cause an uneven distribution of electron density, and hence give the molecule a transient or instantaneous dipole. This occurs even in atoms or molecules that have no permanent dipole, such as a helium atom or a hydrogen molecule, as is shown here. An instantaneous dipole on one particle can then cause an induced dipole on a neighboring one. For instance, the slight positive charge on the helium atom here attracts a little bit of electron density on the left of the neighboring atom, causing uneven electron density and hence a dipole on it. The constant appearance and disappearance of these instantaneous and induced dipoles leads to tiny attractions between molecules. And across many molecules, these attractions can hold them together as a liquid or a solid. All molecules and atoms show van der Waalsf forces because all molecules have electron clouds. Looked at individually, van der Waalsf forces are much weaker than dipole-dipole attractions. However, larger atoms or molecules with bigger electron clouds and more surface area across which these fluctuations can occur will have significant van der Waalsf forces. So the trend to remember is that as the surface area of an atom or molecule increases, the greater the van der Waalsf forces. Lastly, we have hydrogen bonding. This is an even stronger form of dipole-dipole forces. It occurs in molecules that have HF, HO, or HN bonds. These are the most polar covalent bonds that you can make. One of those, HF doesn't exist in organic molecules because HF is a complete molecule. Neither H nor F want to form any more bonds. So organic molecules that contain OH and NH bonds are the ones that will show hydrogen bonding. So hydrogen bonds are rather like dipole-dipole attractions in that they arise from a permanent dipole in the molecule with positive and negative ends. But they are significantly stronger than dipole-dipole attractions and also more directional. That is, they occur between specific atoms rather than over the molecule as a whole. For instance, when we look at water, it's not organic but it's a good example of hydrogen bonding. The hydrogen bonding occurs specifically between the oxygen of one water molecule and a hydrogen atom in the next. Now just as a final note, it's important to remember that the hydrogen bond is not the same as a covalent bond. It's a transient attraction between atoms in neighboring molecules. Some people get mixed up and think that the OH-covalent bond in water, for instance, is a hydrogen bond. It's not. It's the attraction between the oxygen of one water molecule and the hydrogen of a neighboring water molecule.