Energy Effects

First things first: there's a Law of Conservation of Energy, just like how there's a Law of Conservation of Mass (which states that mass cannot be lost or gained, only changed to another form, which I alluded to in my post about chemical equations). The Law of Conservation of Energy essentially states that energy cannot be gained or lost, it can only be converted to another form or transferred elsewhere. For example, a chemical reaction (a.k.a. the "system") might absorb energy from the surroundings, or release it to the surroundings. In fact, each of these types of reactions has a special name- exothermic reactions release energy from the system to the surroundings (I think that "ex" or "exo" means "outer," like in "exoskeleton" and "export") while endothermic reactions absorb energy from the surroundings into the system.

Now there's this other term called "enthalpy," but I'm not entirely sure what it means because, apparently, the true meaning of this term is beyond the Chemistry 2AB course. So far it seems that enthalpy is sort of like potential energy. It's represented by the letter H.

The next thing I need to tell you is how bond breaking and bond making, the basis of pretty much all chemical reactions, relate to exothermic and endothermic reactions. Well, you see, most bonds need at least some energy to break. Energy must be absorbed into the system to break the bonds. The next step in the reaction is the formation of the bonds. Stronger bonds release more energy into the surroundings than weaker bonds. If the overall net effect is the absorption of energy into the system, then the reaction is endothermic. If the overall net effect is the release of energy into the surroundings, then the reaction is exothermic.

By the way, because the formation of stronger bonds release more energy into the surroundings, breaking stronger bonds requires more energy from the surroundings. The inverse is true for weaker bonds: forming weaker bonds releases less energy, while breaking weaker bonds doesn't require as much energy from the surroundings.

All of this can be summed up in some nice little enthalpy diagrams. Here's an enthalpy diagram for an endothermic reaction:

The delta H is the net change of enthalpy in the system.If it's positive, then the reaction is endothermic. If it's negative, the reaction is exothermic. When including enthalpy in a chemical equation, you can either add a "delta H = x kJ" at the end of the equation, or you can add it in the actual equation. If it's endothermic, you need to add the enthalpy change in the reactants side, as energy is being absorbed into the system; if it's exothermic, you need to add it in the products side as energy is being released. Just think about the nature of the reaction to help you remember which side to put the enthalpy change on.

The Ea, or activation energy, is the amount of energy required to break the bonds. The transition state or activated complex is where the original bonds are being broken and new bonds are being formed. In the above diagram, it's easier to see that the products have a higher enthalpy (potential energy?) than the reactants.

And that's pretty much it about energy diagrams, unless you can find anything else that I need to explain. Next up on the agenda is reaction rates! Yayyyyyyyyyyy.