We've spoken about energy changes before. We've used terms like exothermic and endothermic to describe both chemical and physical changes. We're going to delve a little bit deeper into those terms in this unit. Before we do that, I would like to look at a couple more vocabulary words. First, this entire unit is about thermodynamics. The word dynamic means movement, and thermo is referring to heat. And so we're going to be talking about the movement of heat throughout this unit. We saw that already when we talked about specific heat. We talked about heat moving from hot metal to cold water, for example, in the lab. Today, I want to focus on the word enthalpy. Enthalpy is the energy stored in chemical bonds. It is abbreviated with a capital H because there is an H in the word enthalpy. I don't know. We're running out of letters. Now, enthalpy itself is an important term, but it's really hard to measure the enthalpy of a bond. What's more important to us as chemist is to study the change in enthalpy, how the enthalpy will change in a chemical reaction. We're going to be looking at delta H values of reactions. Dr. Trow defines the delta H of a reaction like this. Earlier in the year, we were drawing energy diagrams that looked like this. We said an exothermic reaction was when the energy of the reactants decreased as it turned into products. In other words, energy was released into the environment. A better term for that would be enthalpy. We would say in an exothermic reaction, enthalpy decreases. The amount of energy stored in the reactants is greater than the amount of energy that's finally stored in the products. That enthalpy got turned into heat and was released into the surroundings. An exothermic reaction is a reaction that releases energy. It has a negative change in enthalpy. Enthalpy is lost in the process. Likewise, an endothermic process would be one where energy is gained. And more correctly, we would say the enthalpy is gained, where heat is coming from the surroundings and being stored within the products. So an endothermic process is one where energy is absorbed from the surroundings and would have a positive delta H of the reaction. Let's look at some specific reactions and let's see if we can diagram what happens to their changes in enthalpy. The first we want to look at is the combustion of methane. When one mole of methane is combined with two moles of oxygen, you create a mole of carbon dioxide in two moles of water. We know that already. And we know from personal experience that this is an exothermic process. Conversely, if you want to make solid hydrogen iodide, you could take half a mole of hydrogen and half a mole of iodine to make one mole of hydrogen iodide. This process is endothermic. So if we were to look at the enthalpy diagrams for these, the first one for the methane would have a negative delta H. So the change in enthalpy would decrease. And it turns out that when you burn a mole of methane with two moles of oxygen, you end up getting a delta H of negative 890.3 kilojoules. Conversely, for the endothermic process, your delta H is increasing. You are absorbing heat. And in this case, in order to make one mole of hydrogen iodide solid, you must absorb 26.5 kilojoules of energy. So for the exothermic process, we could say that the delta H is negative 890.3 kilojoules. And for the endothermic process, we could say the delta H is positive 26.5 kilojoules. Let's bring this topic together with the work we are doing at the end of last trimester. Let's do a little bit of stoichiometry with these enthalpy terms. So let's start with the combustion of liquid ethanol. So ethanol is C2H6O. And if it's a combustion reaction, we're going to combine it with oxygen and make carbon dioxide and water. When we balance it, we get a statement like this. One mole of ethanol will combine with three moles of oxygen to make two moles of carbon dioxide and three moles of water. Now it's a good question reaction just like with the methane. So this is going to be an exothermic process. And it turns out that it releases 1,371 kilojoules of energy when we follow the balanced reaction above. So we can combine that to make this statement. One mole of ethanol will combine with three moles of oxygen to make two moles of carbon dioxide and three moles of water. And the delta H of reaction is negative 1,371 kilojoules. So my question for you is if we have 2.35 moles of ethanol burning with an excess of oxygen, how much energy is released? So again, this reaction says one mole of ethanol will react with three moles of oxygen to make two moles of carbon dioxide, three moles of water. And this will release 1,371 kilojoules. Remember that negative sign means it's releasing energy. So the reaction says one mole of ethanol if reacted completely will release 1,371 kilojoules. And this problem we have 2.35 moles of the ethanol. C2H6O. And we know that 1,371 kilojoules are released for every 1 C2H6O. So just like we did with our other stoichiometry problems, we can cancel things out. And we are left with 3,222. But looking at 3 sig figs, I'm going to write this at 3,220 kilojoules. Now you might be asking where is the negative sign? The negative sign is telling you that energy is released. And the question is asking how much energy is released. So if we were to put the negative sign in there, it would be a little bit redundant. We don't really talk about releasing negative energy.