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The bond energy is a measure of the amount of energy needed to break apart one mole of covalently bonded gases. The SI units used to describe bond energy is kiloJoules per mole of bonds (kJ/mol).
Atoms bond together to form compounds because in doing so they attain lower energies than they possess as individual atoms. A quantity of energy, equal to the difference between the energies of the bonded atoms and the energies of the separated atoms, is released, usually as heat. That is, the bonded atoms have a lower energy than the individual atoms do. When atoms combine to make a compound, energy is always given off, and the compound has a lower overall energy.
When a chemical reaction occurs, molecular bonds are broken and other bonds are formed to make different molecules. For example, the bonds of two water molecules are broken to form hydrogen and oxygen.
\[ 2H_2O \rightarrow 2H_2 + O_2\]
Bonds do not break and form spontaneously-an energy change is required. The energy input required to break a bond is known as bond energy. While the concept may seem simple, bond energy serves a very important purpose in describing the structure and characteristics of a molecule. It can be used to determine which Lewis Dot Structure is most suitable when there are multiple Lewis Dot Structures.
When a bond is strong, there is a higher bond energy because it takes more energy to break a strong bond. This correlates with Bond Order and Bond Length. When the Bond Order is higher, bond length is shorter, and the shorter the bond length means a greater the Bond Energy because of increased electric attraction. In general, the shorter the bond length, the greater the bond energy.
Think about it this way: it is easy to snap a pencil, but if you keep snapping the pencil it gets harder each time since the length of the pencil decreases. A higher bond energy (or a higher bond order or shorter bond length) means that a bond is less likely to break apart. In other words, it is more stable than a molecule with a lower bond energy. With Lewis Structures then, the structure with the higher bond energy is more likely to occur.
Although each molecule has its own characteristic bond energy, some generalizations are possible. For example, although the exact value of a C–H bond energy depends on the particular molecule, all C–H bonds have a bond energy of roughly the same value because they are all C–H bonds. It takes roughly 100 kcal of energy to break 1 mol of C–H bonds, so we speak of the bond energy of a C–H bond as being about 100 kcal/mol. A C–C bond has an approximate bond energy of 80 kcal/mol, while a C=C has a bond energy of about 145 kcal/mol. We can calculate a more general bond energy by finding the average of the bond energies of a specific bond in different molecules to get the average bond energy.
Average Bond Energies (kJ/mol)
|Single Bonds||Multiple Bonds|
|C = C|| |
|C ≡ C|| |
|O = O|| |
| ||C = O*|| |
|C ≡ O|| |
| ||N—O|| |
|N = O|| |
|N = N|| |
|N ≡ N|| |
|C ≡ N|| |
| ||C = N|| |
| ||Si—H|| |
| || |
| ||Cl—Cl|| |
| || |
| ||Cl—Br|| |
| || |
| ||Br—Br|| |
| || |
*C == O(CO2) = 799
When more bond energies of the bond in different molecules that are taken into consideration, the average will be more accurate. Keep in mind that:
When a chemical reaction occurs, the atoms in the reactants rearrange their chemical bonds to make products. The new arrangement of bonds does not have the same total energy as the bonds in the reactants. Therefore, when chemical reactions occur, there will always be an accompanying energy change.
In some reactions, the energy of the products is lower than the energy of the reactants. Thus, in the course of the reaction, the substances lose energy to the surrounding environment. Such reactions are exothermic and can be represented by an energy-level diagram like the one shown below. In most cases, the energy is given off as heat (although a few reactions give off energy as light).
Figure: Exothermic Reactions. For an exothermic chemical reaction, energy is given off as reactants are converted to products.
In chemical reactions where the products have a higher energy than the reactants, the reactants must absorb energy from their environment to react. These reactions are endothermic and can be represented by an energy-level diagram like the one shown below.
Figure: Endothermic Reactions. For an endothermic chemical reaction, energy is absorbed as reactants are converted to products.
Exothermic and endothermic reactions can be thought of as having energy as either a product of the reaction or a reactant. Exothermic reactions give off energy, so energy is a product. Endothermic reactions require energy, so energy is a reactant.
Is each chemical reaction exothermic or endothermic?
What is the enthalpy change for this reaction and is it endothermic or exothermic?
\[H_2(g)+I_2(g) \rightarrow 2HI(g)\]
First look at the equation and determine what bonds exist.
Because we're dealing with net change, we only need to look at 1 mol of H-H, I-I, and H-I bond. Then examine the bond breakage which is located in the reactant side:
Then we look at the bond formation which is on the product side:
Hess's Law relates to this equation as it depicts how the energy of the overall reaction is equal to the sum of the individual steps involving energy change.
The breakage and formation of bonds is similar to a relationship: you can either get married or divorced and it is more favorable to be married.
which is an output (released) energy = 872.8 kJ/mol + 498.7 kJ/mol = 1371.5 kJ/mol.
Total energy difference is 1840 kJ/mol – 1371.5 kJ/mol = 469 kJ/mol, which indicates that the reaction is endothermic and that 469 kJ of heat is needed to be supplied to carry out this reaction.
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