6.8 Describing a Reaction: Bond Dissociation Energies

Bond Dissociation Energies and Reaction Energetics

Bond dissociation energies (BDEs) let you predict whether a reaction will release or absorb energy. By comparing the strength of bonds broken versus bonds formed, you can estimate the overall enthalpy change of a reaction. This is one of the most practical tools in organic chemistry for evaluating reaction thermodynamics.

Calculation of Bond Dissociation Energy

Bond dissociation energy (BDE) is the energy required to homolytically break one specific covalent bond in a molecule, producing two radicals. It's measured in kJ/mol or kcal/mol, and every bond type has a characteristic value. For example, a typical C–H bond has a BDE of about 413 kJ/mol (105 kcal/mol).

To estimate the overall enthalpy change ($\Delta H$) of a reaction using BDEs:

1. Identify all bonds broken in the reactants and look up their BDE values.
2. Identify all bonds formed in the products and look up their BDE values.
3. Apply the formula:

$\Delta H = \sum \text{BDEs of bonds broken} - \sum \text{BDEs of bonds formed}$

A negative $\Delta H$ means the reaction is exothermic; a positive $\Delta H$ means it's endothermic.

If you're just summing energy input, the math is straightforward. Breaking two C–H bonds requires $2 \times 413 = 826$ kJ/mol. But the key insight for reactions is that you always compare energy in (breaking) versus energy out (forming).

Bond Strengths and Reaction Thermodynamics

The relationship between bond strength and reaction energy comes down to one idea: stronger bonds release more energy when they form, and require more energy to break.

• Exothermic reactions ($\Delta H < 0$): The bonds formed in the products are collectively stronger than the bonds broken in the reactants. More energy is released during bond formation than was consumed during bond breaking.
• Endothermic reactions ($\Delta H > 0$): The bonds broken in the reactants are collectively stronger than the bonds formed in the products. The reaction absorbs more energy than it releases.

Think of it this way: if you break weak bonds and form strong ones, you come out ahead energetically. If you break strong bonds and form weaker ones, you have to put in more energy than you get back.

High-Energy Compounds in Biochemistry

The term "high-energy compound" can be misleading. ATP doesn't store energy in one particular "high-energy bond" that snaps and releases a burst of power. Instead, the energy released during ATP hydrolysis comes from the fact that the products (ADP and inorganic phosphate) form bonds that are collectively more stable than those in ATP. Electrostatic repulsion between the negatively charged phosphate groups in ATP also contributes: once the bond is broken, the products are lower in energy because that repulsion is relieved.

$ATP + H_2O \rightarrow ADP + P_i \quad \Delta G \approx -30.5 \text{ kJ/mol}$

Some patterns in biomolecule reactivity relate to bond strengths:

• Carbohydrates like glucose are readily metabolized because their bonds can be broken and reformed into more stable products ($CO_2$ and $H_2O$) with a net release of energy.
• Lipids (fats) have many strong C–C and C–H bonds. Per gram, they release more energy than carbohydrates during metabolism because they contain more of these bonds and are more reduced (less oxidized), meaning more oxidation can occur.
• Proteins have strong peptide bonds (C–N), making them relatively stable. The body preferentially uses carbohydrates and fats for energy rather than breaking down proteins.

Thermochemistry and Reaction Kinetics

BDEs are a thermodynamic quantity: they tell you about the overall energy change of a reaction, not how fast it happens. That distinction matters.

Thermochemistry deals with heat changes. BDEs let you estimate $\Delta H$ for a reaction, which tells you whether the reaction is energetically favorable in terms of enthalpy.

Kinetics deals with reaction rate, which depends on activation energy ($E_a$), the minimum energy barrier the reactants must overcome to reach the transition state. A reaction can be highly exothermic but still very slow if $E_a$ is large.

Bond strength does influence activation energy in a general sense: reactions that require breaking very strong bonds tend to have higher activation energies. But $\Delta H$ and $E_a$ are separate quantities. A favorable $\Delta H$ does not guarantee a low activation energy, and vice versa.