 If I give you the formula for a covalent substance, CH4, how do you know what that molecule looks like? Which atoms are bonded to which? We need a way of working out the arrangement of atoms in a molecule, and a robust way of doing this is to use Lewis structures, which were invented by Gilbert Lewis in the early 20th century. Here are some examples of Lewis structures for some covalent molecules. You can see that they show how the atoms are attached to each other, where there are single, double and triple bonds, and where the non-bonding electrons are. The steps to drawing a Lewis structure are not complicated. The important thing is to remember to do them in order. I'm going to go through them and draw the structure for ammonia at the same time. So the first step, whenever you're drawing a Lewis structure, is to total up the valence electrons of all the atoms in your molecule. So we've got one nitrogen, which we know has five valence electrons because it's in group five, and we have three hydrogen atoms, which each have one valence electron, which equals a total of three electrons. Now if we add all of those up, we have a total of eight valence electrons from all of our atoms. The second step is to work out how many bonds each atom is able to form, and then to use that to draw a sort of skeleton structure for the molecule, where we join up the atoms with the bonds. So we know that because nitrogen is in group five, and it has a valence of minus three, it usually forms three bonds. And we know that hydrogen is in group one, and it usually forms one bond. So we need to draw a structure in which nitrogen forms three bonds, and each hydrogen only forms one. And the only way that we can do that is to put the nitrogen in the middle, and then use one bond each to join up the three hydrogens. The next step is to look at each of the atoms in the molecules so far and work out if they have a full outer shell. Once you've taken the shared electrons in the bonds into account. So we'll do the hydrogens first. Hydrogen is in the first period of the periodic table, which means that to have a full outer shell, it only needs two electrons. We also know that each of our covalent bonds has two electrons in it. So because each of the hydrogen atoms have formed one bond, they each have access to two electrons, and so they have a full outer shell. So the three hydrogens are fine. Now we'll look at the nitrogen. Nitrogen is in the second period of the periodic table. In order for it to have a full outer shell, it needs to have eight valence electrons. Now as it stands, it has formed three bonds. Each bond has two electrons in it. So it only has six electrons where it needs eight. So nitrogen does not yet have a full outer shell. If we find this, then we need to give it a couple of extra electrons. Now it can't form any more bonds. It's formed its three bonds, and not only that, we don't have any other atoms that it could form a bond to. So these extra electrons are not going to be bonding electrons. They're going to be non-bonding electrons, and we sometimes call them a lone pair. So a pair of electrons that are not taking part in a bond. We'll represent them by a little pair of dots. And then we recount the electrons that the nitrogen has access to. It's got the six electrons from the three bonds, and it has two electrons from the lone pair, which gives it eight. So it now has a full outer shell, and it's happy. The final step is to do a check. The point is that when you draw a Lewis structure, the number of electrons that you represent in the Lewis structure by bonds and lone pairs must equal the number of valence electrons that you started off with. So for ammonia, we determined at the beginning that there were eight valence electrons that we could play with. So when we've finished drawing our Lewis structure, we need to check that we have actually used exactly eight electrons, no fewer, no more. And we see when we look at our structure that we have three bonds. Each bond contains two electrons, so that's six. And we have a lone pair, that's another two, which makes eight. And that equals our original total. So we're all good.