Combustion. So combustion is the reaction here that you see at the top, is an example of combustion. It's the simplest alkane, so methane, one carbon, four hydrogen. So one carbon is meth, and then obviously they're single bonded, I guess, technically. But there's only one carbon, so it's not making really any bonds.
So the ending is A-N-E, methane. And then the combustion is the part where we add carbon dioxide and then we produce water and, I mean, we add oxygen, sorry, and produce water and carbon dioxide. That's the equation that you see here at the top. We're going to talk about just the reactants side for a second. So if you go down where we see lower energy reactants, we have methane plus our two oxygen molecules.
And in order to turn them into carbon dioxide and water later, we have to break those molecules apart. into their individual atoms, that takes a lot of energy. So we're going to add energy to the reactant side, to the left side of the equation, in order to have the individual carbon, the individual hydrogens, and individual oxygens with no bonds. They need a lot of energy to not be bonded together.
They have to be able to move fast enough, basically, that the bonds aren't going to stick them together. We call that the activated complex. Every reaction goes through this where we start with reactants, they all break apart, and then we form products. So when they're all broken apart, we call that the activated complex. We'll see that on a graph at the end.
Okay, so once we're at the activated complex, everybody's in their atom forms rather than in compounds. We start the second half of the reaction, which is forming the compounds. So that's what you see on this slide.
So we're starting here with the higher energy activated complex, all of the elements in their individual form, and then we're forming our products. So now they're forming different bonds. We have the carbon dioxide and then the two waters. So they all broke apart into their individual atoms, then they rearranged and formed new bonds. When they did that, when they form bonds, they can release that energy.
It's almost like... They take in a bunch of energy, almost like you're breathing in and you're breaking up all those bonds. And then you can release your air and exhale or you release your energy and we create new molecules.
So for this part, we have energy on the right side of the equation, which is the product side. It's being released. So in every reaction, you have energy being added and energy leaving the system. But since the energy was released from the... products, and that means it's on the product side, that means it's going to exist at a lower energy level because we released that energy.
So our products are at a lower energy level, and we consider that exothermic. We compare how much energy was absorbed and how much energy was released to determine if something is endo or exothermic. So if a combustion, the type of reaction of combustion is an exothermic reaction because it releases more energy than it absorbs.
So you can see here in this exothermic. graph. The reactants are starting at a particular energy. They are going to absorb energy till we get to the activated complex where again, everything is broken up. And then we release energy down to the products.
Since the difference in energy between our reactants and our products, the equation for finding that is right here on this bullet point. It says delta H is the difference between the energy of the reactants and the energy of the products. So we do energy of products minus energy of reactants. So say the products were, I don't know, at five and the reactants were at 10. If we say five minus 10, that gives us a negative five. Energy cannot be a negative number.
If you have a negative energy, it just means energy is being released. If you have a positive energy, that means energy is being absorbed. So our overall, our net energy for an exothermic reaction is going to be negative because we're going to release more energy than we absorbed.
You can see down here in this other little picture, we have an endothermic reaction. Our reactants are starting at a low level of energy. We're adding a bunch of energy and then we're releasing some. So you're always going to add and release energy.
But we added way more than we released. So if we did products minus reactants, now let's say products are at 10 and reactants are at 5. 10 minus 5 is positive 5. That means our net, how much we added overall or subtracted overall, we added. more energy than we released.
So that's an endothermic reaction. Exothermic reactions tend to feel hot. So when you hear the word combustion, you tend to think like an explosion or fire or something like that, which is a good association. It's not necessarily.
true. Usually it doesn't contain fire, but it does definitely produce heat. And this graph tells you why, because we're releasing more energy.
We're releasing energy in the form of heat. So you're going to feel it as heat. Endothermic reactions, you feel as cold.
They absorb energy, so you feel them as cold. A good example of that is those ice packs that you can break. When you break that ice pack, you are activating a chemical reaction that is endothermic. So it's absorbing a bunch of heat. So it's taking heat out of the bag, which is why you start to feel the bag as cold because it's taking away the heat from that bag because it's using that heat for the reaction.
Okay, so that's combustion, the reaction, the difference between endothermic and exothermic, and then enthalpy. The delta H, the term for that is enthalpy. But again, you calculate that by the energy of the products minus the energy of the reactants. That's also called Hess's law.
We'll get to that later in the year maybe. But yeah, we call that Hess's law that calculation. For our purposes, you'll be given it. I won't make you calculate it. But that delta H is called enthalpy.
Sometimes we call that heat of reaction as well.