In this example, we use bond energies to approximate the change in enthalpy for a reaction. Our problem reads, "Calculate the change in enthalpy for the reaction below as written using the bond energies provided. Use the best Lewis structure for each molecule. In OSF4, the oxygen and all four fluorine atoms are bonded to the central sulfur atom." We can approximate the change in enthalpy for a reaction as the enthalpy change to break all the bonds of the reactants and the enthalpy change to form all the bonds of the products. Breaking bonds always requires energy, so the enthalpy change is positive; forming bonds releases energy, so the enthalpy change is negative. So our equation is the change in enthalpy for a reaction is the summation of the enthalpy changes for bonds broken, subtract the summation for enthalpy changes of bonds formed. If we're breaking all the bonds of the reactants and forming all the bonds of the products, we'll need the best Lewis structure for each substance in the reaction. Let's start by breaking all the bonds of the reactants. In our balanced chemical equation, we have 2 SF4 molecules, and each SF4 has 4 sulfur-fluorine single bonds, so we write 2 SF4 times 4 sulfur-fluorine single bonds. We also have one oxygen gas molecule as a reactant which has an oxygen-oxygen double bond, so we add 1 oxygen molecule times 1 oxygen-oxygen double bond. That's all of our reactants. Next, we need to make the bonds of our products. We have 2 OSF4 molecules, and each OSF4 has 4 sulfur-fluorine single bonds and 1 sulfur-oxygen double bond. Notice that the "2" out front is applied to both 4 sulfur-fluorine single bonds and 1 sulfur-oxygen double bond, so we used parentheses correctly. If we look at our Lewis structures closely, we'll see that the 4 sulfur-fluorine single bonds do not change from reactants to products. So in our equation, they show up as being added from the reactants and subtracted from the products. These can be removed from the equation since they're not changing. The overall effect is that in our reaction, we are breaking 1 oxygen-oxygen double bond from the oxygen molecule and four main two sulfur-oxygen double bonds. When we're using bond energies to calculate enthalpy change, we can always break all the bonds of the reactants and form all the bonds of the products, and we should get the correct change in enthalpy. Or if we are very careful, we can identify just what bonds are changing in the reaction, and we should get the same results if we correctly take into account coefficients in the balanced chemical equation. Now we can plug in our bond energies from the table. We have 1 mole of oxygen-oxygen double bonds, and the bond energy is 495 kilojoules per mole. We have 2 moles of sulfur-oxygen double bonds, and the bond energy is 523 kilojoules per mole. Completing the calculation, we should get -551 kilojoules for the change in enthalpy of the reaction. This is an approximation since the bond energies in the table are only an average; the actual bond energies for a given bond depend on the context of the molecule.