Enthalpy - Online Tutor, Practice Problems & Exam Prep

Bond dissociation energies are crucial for calculating the enthalpy of reactions, indicating whether a reaction is exothermic or endothermic. Exothermic reactions release energy (negative ΔH) when bonds form, while endothermic reactions absorb energy (positive ΔH) when bonds break. As atoms increase in size, bond strength generally decreases due to greater distance. Understanding these concepts helps determine reaction spontaneity, which also involves temperature and entropy (ΔS). Analyze reactions by identifying broken and formed bonds, using provided bond energy values to calculate ΔH.

How to calculate enthalpy using bond dissociation energies.

All right guys. So let's go ahead and talk about one of the most important topics in thermodynamics, which is using bond dissociation energies to calculate the enthalpy of a reaction. Alright, so bond dissociation energies are values that are going to describe the strength of chemical bonds. Alright? Now, in chemistry courses that are way beyond the level of this course, like organic chemistry, we can actually use math to figure out exactly why two atoms have the exact strength of bond that they do. Okay? But in this course, you're never going to be asked that. You're never going to be asked to explain why the bond between H and H is 436 kilojoules per mole. Okay? That's not going to happen. These are going to be essentially experimental values that you're going to be given that you're going to use as a reference to just figure out one question. Is this going to be an exothermic reaction, or is this going to be an endothermic reaction?

Now, keep in mind, just because we know the enthalpy, doesn't mean that we know the spontaneity. Remember that the enthalpy is just one component of the spontaneity. We're also going to need to know the temperature and the ΔS. But for right now, all we're worrying about is the ΔH. We're just asking if this is exothermic or endothermic. Alright. So let's remind ourselves of what exothermic means. It just means that I'm going to get a negative energy at the end and what that means is that I'm making bonds. Remember that it's favorable to make bonds. You release energy as you save energy between those orbitals. And then, if I get a positive value, that means I'm breaking bonds because remember it takes energy. I have to put energy into the system to make these atoms fly apart again.

So here I've given you guys my own chart of bond dissociation energies. All these values should be exactly the same as whatever chart you're going to be given. The only thing I did is I just didn't want to take up a full page with every single value, so I just gave you guys some of the most important ones here. As you can see, they're basically broken up into bonds of hydrogen. So I'm just going to put hydrogen bonds, even though don't get that confused with hydrogen bonding. That's just something different. Okay. Then I put here methyls. Okay. So things bonded to methyl. And then finally, basically like things bonded to each other. I don't know how to say that better. But just like methyl to methyl, oxygen to oxygen, stuff like that. So you guys can see just the general trends.

What you can see is that there's a general trend that as my atoms get bigger, the bonds are going to get weaker. Okay? So you can see how it's not a perfect trend. And like I said, you have to use very complicated math to defend it. But as a general rule, what we can say is that as these atoms that are attached to each other get bigger, the bonds are a little bit weaker and that has to do with the distance between them. These bigger atoms are going to have a farther distance than the smaller ones. So one of the weakest bonds here is going to be II. Why? Because those are two humongous atoms that are trying to make a bond to each other. They're going to be much further apart than hydrogen to hydrogen will be. That can partially explain the huge difference in energy between hydrogen to hydrogen and iodine to iodine. All right?

So as I just told you, if we make a bond, that's going to be a negative value. So that means that I would assign a negative value to whatever bonds I'm making. So if I make an HOH, that means I'll put a negative value on that. If I'm breaking a bond, then I'm going to assign a positive value to that. Okay? So what I want you guys to do for the next two problems is I want you guys to analyze the reaction that's taking place. You don't need to know the mechanism. You don't need to know what's going on. What you do need to know is which bonds are being broken, which bonds are being made. Alright? And then assign values to those using the chart that I gave you above. At the end, after you sum all of that together, you're going to know the enthalpy of the reaction, and you're going to be able to tell me is this exothermic or endothermic. So go ahead and try it for yourself.

Calculating Enthalpies

All right, guys. So like I said, you don't need to know the mechanism of this reaction. What we do need to know is what bonds are being broken and what bonds are being made. All right? So on my reagents side, reagents, what I have is the bonds that are breaking. Cause what I can see is that I used to have a CH bond here, but now that's turning into CBr. So that means that at some point, I must have broken that bond. So that means that this is going to be the breaking side. If this is the breaking side, what sign of enthalpy should this have? Should it be positive or negative? It should be positive, right? Because it's taking energy to make these bonds come apart. So, all of the values that I put over here are going to be positive. All right? Then all the values I put over here for the making side are going to be negative because these are the ones that are going to release energy into the system. So now all I have to do is identify these bonds. Okay? One of them I know already is going to be CH3H2. I know that bond is getting broken because, like I said, it used to be there and now it isn't. Another one that has to get broken is Br₂ because Br₂ is turning into HBr. So this is going to be Br₂. That's another one that's going to require energy to break. Now let's look at my negative side, my exothermic side. So basically, the bonds that I'm making are equal to CH3Br1 because that's a new bond that didn't exist before. And another one is HBr. That's a bond that didn't exist before. So now all I have to do is pull down the values from my chart and whichever ones are on my reagent side are going to be or my breaking side are going to be positive. Whichever ones are on my product side or my making side are going to be negative. So let's go ahead and look up. So what are the values that I'm bringing down? Well, the CH3H2 is going to be 436, so it's going to be positive 436. Br₂ is going to be it looks like where is that? It's right here. So that's going to be 192, positive 192. Then on the negative side or on the making side, CH3Br1 I'm going to just take myself out of the screen for a second so you guys can see. So on the negative side, CH3Br1 is going to be right here. That's going to be 923, so I'm going to give that a negative 923. And then HBr is going to be negative 368. Okay? So now I'm pulling it back down. We have all the values we need. Hi, I'm back. So now all we need to do is plug all these values into a calculator, and whatever our end answer is, that's going to be the end answer and I'm here. Sorry about that. Stop auto-focusing. Okay. So let's go ahead. I've got my calculator here. Let's go ahead and enter these values. I have 436 +192 minus 923 minus 368. And what that should give you at the end is that your enthalpy, your delta H, is equal to negative 33 kilojoules per mole. So that's going to answer a really big question for us. Is this exothermic or is this endothermic? And the answer is that this has to be exothermic because, at the end, the sum of my bond association energies is actually going to be a negative number, which means that I'm getting more stable at the end. I'm making more bonds, or I'm gaining more energy than I'm losing by breaking those bonds. So this would be an exothermic reaction. Do you know if this would be a spontaneous reaction? No. We have no clue yet. We still need the temperature and the Delta S to figure that out. But we know with certainty that this is an exothermic reaction. Okay? Now I want you guys to do the same thing for the second question. Once again, all on your own, try to figure it out, and then I'll give you guys the answer.

Predict the sign and magnitude of ∆H^o in kj/mol for the following reaction. Identify the reaction as either exothermic or endothermic.

• A. -234 Kj/mol
• B. -1118 Kj/mol
• C. +56 Kj/mol
• D. +348 Kj/mol

Here’s what students ask on this topic:

What is enthalpy and how is it related to bond dissociation energies?

Enthalpy (ΔH) is a measure of the total energy of a thermodynamic system, including internal energy and the energy required to make room for it by displacing its environment. In the context of chemical reactions, bond dissociation energies are used to calculate the enthalpy change. When bonds are broken, energy is absorbed (positive ΔH), and when bonds are formed, energy is released (negative ΔH). By summing the bond dissociation energies of all bonds broken and formed in a reaction, you can determine whether the reaction is exothermic (releases energy) or endothermic (absorbs energy).

How do you calculate the enthalpy change (ΔH) of a reaction using bond dissociation energies?

To calculate the enthalpy change (ΔH) of a reaction using bond dissociation energies, follow these steps: 1) Identify all the bonds broken in the reactants and sum their bond dissociation energies (positive values). 2) Identify all the bonds formed in the products and sum their bond dissociation energies (negative values). 3) Add the total energy of bonds broken to the total energy of bonds formed. The resulting value is the ΔH of the reaction. If ΔH is negative, the reaction is exothermic; if positive, it is endothermic.

What is the difference between exothermic and endothermic reactions in terms of enthalpy?

Exothermic reactions release energy, resulting in a negative enthalpy change (ΔH < 0). This occurs because the energy released from forming new bonds in the products is greater than the energy required to break the bonds in the reactants. Endothermic reactions absorb energy, leading to a positive enthalpy change (ΔH > 0). In these reactions, the energy required to break the bonds in the reactants is greater than the energy released from forming new bonds in the products.

Why do larger atoms generally form weaker bonds?

Larger atoms generally form weaker bonds because the distance between their nuclei and the shared electrons increases as atomic size increases. This greater distance reduces the electrostatic attraction between the nuclei and the bonding electrons, resulting in weaker bonds. For example, the bond dissociation energy of an I-I bond is lower than that of an H-H bond because iodine atoms are much larger than hydrogen atoms, leading to a weaker bond.

How does bond dissociation energy affect the spontaneity of a reaction?

Bond dissociation energy affects the enthalpy change (ΔH) of a reaction, which is one component of spontaneity. However, spontaneity also depends on entropy change (ΔS) and temperature (T). The Gibbs free energy equation, ΔG = ΔH - TΔS, combines these factors. A reaction is spontaneous if ΔG is negative. While a negative ΔH (exothermic reaction) can contribute to spontaneity, a positive ΔS (increase in disorder) and high temperature can also drive a reaction to be spontaneous, even if ΔH is positive (endothermic reaction).