Bond Dissociation Energy

In this tutorial I want to talk about bond dissociation energy and how we use it in organic chemistry.

So, what exactly is bond dissociation energy? By definition, it’s the amount of energy it takes to break a chemical bond. And here’s something extremely important I want to emphasize: it always takes energy to break a chemical bond. In other words, bond breaking is always an endothermic process, without exception.

When we talk about bond dissociation energy, we assume we’re dealing with homolytic cleavage, which only applies to molecules with covalent bonds. If we were talking about ionic compounds, we’d instead deal with lattice energy, which is very different.

Breaking Bonds vs Making Bonds

For example, take a chlorine molecule. If I apply heat or light, it splits into two chlorine atoms, each carrying one of the electrons they used to share. If you’ve already studied radical reactions, this should look familiar. If not, don’t worry—it’s not essential here. From an energy perspective, this bond breaking requires 239 kJ/mol. Notice the number is positive, because breaking bonds requires energy input.

I want to repeat this one more time because it’s such a common misconception: breaking a bond requires energy. If bond breaking released energy, there would be no reason for bonds to exist in the first place.

Now, what about bond making? For example, if two hydrogen atoms come together and share their electrons, forming an H–H bond, that process releases 435 kJ/mol of energy. Bond making is exothermic.

Using the Bond Dissociation Energy to Estimate Change in Enthalpy

So, how is this useful? Well, bond dissociation energy values are tabulated, so we can easily look them up. Since these values are readily available, we can use them to estimate enthalpy changes in reactions and predict whether those reactions are exothermic or endothermic.

Let’s go through an example. Suppose we want to estimate the enthalpy change for this reaction. First, I would draw the full Lewis structures for the starting materials and products to see which bonds are broken and which are formed.

On the reactants side, I’m breaking a carbon–hydrogen bond and a bromine–bromine bond. On the products side, I’m making a carbon–bromine bond and a hydrogen–bromine bond.

Looking up the values, the bonds I break require 613 kJ/mol of energy input. The bonds I form release 671 kJ/mol. Notice how here I’m showing the bond-making values as negative. That’s intentional. Normally, bond dissociation energy values are given as positive because they describe bond breaking, but when we’re making bonds, we should use negative numbers to represent energy release.

In general chemistry, you might have been taught to keep all values positive and then subtract products from reactants. But in practice, many students get confused about what to subtract from what. Using positive numbers for breaking and negative numbers for making keeps it straightforward. Bond breaking = positive. Bond making = negative.

Adding everything together gives us −58 kJ/mol. This reaction is exothermic overall.

Another Example

Let’s do one more example. Suppose I react an alkene with hydrogen bromide. I start by drawing the Lewis structures. On the reactant side, I’m breaking the π bond of the alkene and the H–Br bond. On the product side, I’m making a carbon–hydrogen bond and a carbon–bromine bond.

Looking up the bond dissociation energy values, it takes 611 kJ/mol to break the bonds. On the product side, bond formation releases 732 kJ/mol. Putting it all together, the enthalpy change is −121 kJ/mol. This reaction, like the previous one, is exothermic.

Limitations of the Bond Dissociation Energy Calculations

Now, are there cases where we cannot use bond dissociation energy values? Yes. Bond dissociation energy is only reliable for neutral molecules. If we’re dealing with ions, regular bond dissociation energy tables won’t help. There are some tables that estimate values for charged species, but you’re unlikely to encounter those in class or on exams.

Another limitation is that the same type of bond can have slightly different energies depending on where it is in the molecule. This is especially true for carbon–hydrogen bonds. For instance, C–H bonds in primary (1°), secondary (2°), and tertiary (3°) carbons all have slightly different bond dissociation energy values. The difference between a primary and tertiary C–H bond can be almost 20 kJ/mol. That may not sound huge, but it can significantly affect reaction mechanisms, regioselectivity, or reaction rates.

Finally, remember that the values you see in tables are usually gas-phase values. Solvent and phase effects can shift them slightly, but in an introductory organic chemistry course, we usually ignore those effects and just use the tabulated numbers.

And that’s really all you need to know about bond dissociation energy. Now, go ahead and practice some problems to make sure you’re comfortable with the calculations.