Table of Contents
IB Chemistry: Bond Enthalpies
Bond Enthalpies
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Define the term average bond enthalpy
Average Bond Enthalpy: Enthalpy changes when one mole of gaseous bonds is broken.
Note: I underlined “gaseous” because it’s absolutely crucial you add it in the definition, or you may lose marks.
The reverse is the amount of energy required to form 1 mole of a bond between two atoms.
There’s a reason we call this “average” bond enthalpy. For example, if you want to break a C-H bond, or a Li-F bond. The amount of energy required to break the bonds are not the same all the time, so therefore the value presented is an average of all calculated enthalpy values.
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Explain, in terms of average bond enthalpies, why some reactions are exothermic and others are endothermic
- Energy is required to break bonds (Endothermic)
- Energy is released when a bond is formed (Exothermic)
For the first one, think of it like this. When a couple are holding hands, it would take some energy to make them get off right?
Exothermic Reaction
In an exothermic reaction, the amount of energy required to break the bonds of the reactants is actually less than the total amount of energy released when the bonds form a product.
From this, we can conclude that the bonds in the reactants are weaker than the bonds in the products, and hence the products are more stable.
Endothermic Reaction
The opposite happens here. The amount of energy required to break the bonds is actually more than the total amount of energy released when the bonds form a product.
From this, we can conclude that the bonds in the reactants are stronger than the bonds in the products, and hence the reactants are more stable.
