Why is energy required to break chemical bonds

Bond-breaking is an endothermic process. Energy is released when new bonds form. Bond-making is an exothermic process. Whether a reaction is endothermic or exothermic depends on the difference between the energy needed to break bonds and the energy released when new bonds form. When you break the bonds, you get energy. As Derek Muller from Veritasium notes, this idea of energy stored in the chemical bonds is very wrong.

To get a better understanding of energy in chemical bonds, let's consider a simplified model. When one hydrogen atom interacts with another hydrogen to form molecular hydrogen H 2 , many things are going on. Still, one of the fundamental interactions is due to the electrostatic force between protons and electrons.

Yes, there are quantum mechanical effects too—but let me stick with a simple model. In this model, I have two hydrogen atoms that experience some type of electric force attracting them. When they get really close, there is another force repelling the two atoms. To keep things calm, I add a drag force. Here is what it looks like when these two atoms interact. View Iframe URL.

If you want to think of this system in terms of energy, it might be useful to look at a sketch of the potential energy for these two hydrogen atoms. It would sort of look like this just a sketch. We can imagine the hydrogen atoms are like a ball rolling on a hill shaped like the potential curve. You can see that it would increase in speed as it goes down the hill, then slow down and move back as it went up the "hill". The strength of a bond between two atoms increases as the number of electron pairs in the bond increases.

Generally, as the bond strength increases, the bond length decreases. Thus, we find that triple bonds are stronger and shorter than double bonds between the same two atoms; likewise, double bonds are stronger and shorter than single bonds between the same two atoms. Average bond energies for some common bonds appear in Table 9.

When one atom bonds to various atoms in a group, the bond strength typically decreases as we move down the group. The bond energy is the difference between the energy minimum which occurs at the bond distance and the energy of the two separated atoms. This is the quantity of energy released when the bond is formed. Conversely, the same amount of energy is required to break the bond.

For the H 2 molecule shown in Figure 5. This may seem like a small number. However, as we will learn in more detail later, bond energies are often discussed on a per-mole basis. For example, it requires 7. A comparison of some bond lengths and energies is shown in Figure 5.

We can find many of these bonds in a variety of molecules, and this table provides average values. For example, breaking the first C—H bond in CH 4 requires As seen in Table 9. We can use bond energies to calculate approximate enthalpy changes for reactions where enthalpies of formation are not available. Calculations of this type will also tell us whether a reaction is exothermic or endothermic.

This can be expressed mathematically in the following way:. The bond energy is obtained from a table like Table 9. Thus, in calculating enthalpies in this manner, it is important that we consider the bonding in all reactants and products.

Because D values are typically averages for one type of bond in many different molecules, this calculation provides a rough estimate, not an exact value, for the enthalpy of reaction.

Because the bonds in the products are stronger than those in the reactants, the reaction releases more energy than it consumes:. This excess energy is released as heat, so the reaction is exothermic. Twice that value is — We can express this as follows:. Using the bond energy values in Table 9. Note that there is a fairly significant gap between the values calculated using the two different methods.

This occurs because D values are the average of different bond strengths; therefore, they often give only rough agreement with other data. Using the bond energies in Table 9.