 So we've made connections between enthalpy and heat, in particular enthalpy is the same thing as heat when we do something at constant pressure. And we motivated that conversion of heat into enthalpy in part by saying that it's more convenient to use enthalpy as a state function rather than heat which is a path function. So now we begin to get the real benefit of using enthalpy as a state function because it will allow us to calculate the enthalpy change for processes that we can't actually physically perform or at least not perform easily. So let me show you what I mean by that with a few examples. As the first relatively simple example, let's suppose the chemical process we're interested in is the decomposition of iron III oxide, Fe2O3, decomposing that into pure solid iron and oxygen gas. So if I balance that reaction, two irons and three halves of an oxygen II molecule balance this Fe2O3. So let's call that reaction A. So reaction A that I'm interested in, maybe I want to know what is the delta H for that reaction. That's not the kind of thing that we can necessarily terribly easily measure in a calorimeter because it's a little difficult to force that reaction to happen. Iron II, iron III oxide, that's essentially rust. Spontaneously iron and oxygen will combine and form Fe2O3 but it's a little more difficult to make it go the other way around. But what that suggests is we can, if I just take and call this reaction B, if I flip this reaction around and say iron solid and oxygen gas forming Fe2O3, the reverse of the reaction I was originally interested in, I can make that reaction happen without much difficulty. If I take some iron, it'll rust very slowly in the presence of oxygen but if I shave the iron down into iron filing, so it's got a lot more surface area, perhaps give it a little spark to initiate this combustion process, I can cause iron II pretty quickly combust and form Fe2O3. So that's the type of thing I could do in a calorimeter. So I can measure delta H for that reaction relatively easily but since we know that enthalpy is a state function, the difference in enthalpy between these products in this reaction is exactly the negative of the enthalpy change between these products and these reactants. So what that means is delta H for reaction A is the negative of delta H for reaction B. So like I said, that's a relatively simple example. Perhaps we can't do the reaction we're interested in but we can do the reverse reaction, just throw a negative sign in front of the enthalpy and that makes us able to use a reaction that we can measure the properties of more easily. As a slightly more involved example, let's take, let's say now that what we're interested in is hydrogen peroxide. So let's say that what we're interested in doing is talking about the, let's talk about the formation of hydrogen peroxide, so there's a reaction that's fairly easy to balance. Hydrogen peroxide could be formed by combining H2 and O2 gas together. So let's say that's the reaction that we're interested in. That reaction looks like something we might be able to do in a calorimeter. Take hydrogen, combust it in the presence of oxygen and perhaps we would form hydrogen peroxide except in fact we wouldn't. What actually happens if you put hydrogen gas in the presence of oxygen gas and react them is it will combust, hydrogen will burn in oxygen, but what you form is the more stable water, let's say water liquid product. So balancing this reaction, hydrogen and only half an oxygen will form water liquid. It doesn't react to form hydrogen peroxide, it reacts to form water instead. So let's say this is reaction number one, not quite, so that's the reaction that nature will choose to have happen if we try to combust hydrogen in a calorimeter in order to measure the enthalpy change for this process. But we know some other things about hydrogen peroxide. We know that if we take hydrogen peroxide it's not terribly difficult to get hydrogen peroxide to decompose. As you know if you've ever poured hydrogen peroxide on a cut it bubbles and it's a very good oxidizing agent. It will reduce relatively easily. So this reaction is one that it's not crazy to consider causing to happen. Hydrogen peroxide decomposing to form water and giving off some oxygen gas. So between reaction number one and reaction number two we have all the pieces we need to understand this reaction. If I draw those out maybe graphically, let's say if I start with a mole of oxygen and a mole of hydrogen, the things that I want to combine to form H2O2, I can combine the mole of hydrogen and half of a mole of oxygen to form a mole of H2O and I've got left over a half of a mole of an O2. Half of a mole of O2 didn't react. So that's what we call reaction number one. But water and a half a mole of O2, those are exactly the products in this reaction number two. So as products hydrogen and peroxide decomposing into water and one half an O2, that's reaction number two. So again remembering that enthalpy is a state function, the delta H for this process converting H2O2 directly into H2O2, it doesn't matter whether we take the direct path or whether we take the indirect path of first converting it to H2 and a half an oxygen and then converting it further to H2O2, which is the reverse of reaction two, it doesn't matter which of these two paths I take, they're both going to have the same enthalpy. So the delta H of this reaction that I'm interested in is the delta H of step one and not delta H for step two but the reverse of that, so a negative delta H of step two. So very often if there's a chemical reaction that we can't literally perform in the laboratory, we can still find out something about the thermodynamics of that reaction, the thermochemistry by breaking it down into reactions that we can either measure the properties of or themselves can be broken down in other ways to measure the properties of. So the key observation here is something called Hess's law. Any time we do something like this or something like this, what we've used is Hess's law which says that we can treat chemical reactions essentially algebraically. If I turn a reaction around backwards, I just flip the sign of the delta H or indeed the delta U if I'm using energies instead of enthalpies. If I add two reactions together, then I add their enthalpies together. If I flip one around and then add them together, I can treat them algebraically, add them together but one of them having a negative sign. So I can combine reactions algebraically and their enthalpy changes for those reactions combine algebraically in the same way. So we could go through many other examples. If I double a reaction, I double its enthalpy. If I add three different reactions one after another, I just add the three enthalpies together. So that's essentially what Hess's law says is you combine the enthalpies algebraically the same way that you treat the chemical reactions if you can form a reaction that you're interested in by combining simpler chemical reactions. So it turns out this is especially useful when we use Hess's law to combine formation reactions and that's what we'll consider in the next lecture.