Bond Enthalpies
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Topic Summary & Highlights
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Core Concept
Bond enthalpy (or bond dissociation energy) is the amount of energy required to break one mole of a particular type of bond in the gas phase. Average bond enthalpies are used to estimate the enthalpy change of a reaction by considering the energy required to break bonds in the reactants and the energy released when new bonds form in the products.
$\Delta H_{\text{reaction}} \approx \sum \Delta H_{\text{bonds broken}} - \sum \Delta H_{\text{bonds formed}}$
Practice Tips
Use Average Values Carefully: Remember, these values are averages and may differ slightly depending on molecular environment.
Focus on Bonds Changed: Only consider bonds that are broken in the reactants and formed in the products.
Account for All Bonds: Double-check that all bonds in each molecule are accounted for, especially in large molecules.
Sign Conventions: Energy required to break bonds is positive, while energy released from forming bonds is negative.
Steps for Calculating Reaction Enthalpy Using Bond Enthalpies
Write the Balanced Equation:
Write the balanced chemical equation for the reaction to identify all bonds that will be broken and formed.
List Bonds Broken and Formed:
Identify and list all the bonds in the reactants that will be broken.
Identify and list all the bonds in the products that will be formed.
Find Average Bond Enthalpies:
Use a bond enthalpy table to find the average bond enthalpies for each type of bond involved.
Calculate the Total Energy for Bonds Broken and Formed:
Multiply each bond enthalpy by the number of that bond type in the molecule, and sum for all bonds broken and all bonds formed.
Calculate $\Delta H_{\text{reaction}}$:
Use the formula: $\Delta H_{\text{reaction}} \approx \sum \Delta H_{\text{bonds broken}} - \sum \Delta H_{\text{bonds formed}}$
This gives an estimate of the enthalpy change for the reaction.
Key Concepts
Bond Enthalpy (Bond Dissociation Energy):
The amount of energy needed to break 1 mole of bonds in a gaseous substance.
Represented in kJ/mol.
Always positive because energy is required to break bonds.
Average Bond Enthalpy:
Bond enthalpies can vary depending on the molecular environment, so the average bond enthalpy is used as an approximation.
For example, the C–H bond enthalpy in methane ($\text{CH}_4$) is slightly different from that in ethane ($\text{C}_2\text{H}_6$), so an average value is used for general calculations.
Bond Breaking and Bond Formation:
Bond Breaking: Endothermic process (ΔH > 0), energy is absorbed.
Bond Formation: Exothermic process (ΔH < 0), energy is released.
Using Bond Enthalpies to Estimate $\Delta H_{\text{reaction}}$:
The enthalpy change of a reaction can be estimated by subtracting the total bond energies of the bonds formed from the total bond energies of the bonds broken.
$\Delta H_{\text{reaction}} \approx \sum \Delta H_{\text{bonds broken}} - \sum \Delta H_{\text{bonds formed}}$
Example Problem
Calculate the enthalpy change ($\Delta H_{\text{rxn}}$) for the following reaction using the average bond energies:
$\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g)$
Average Bond Energies (kJ/mol):
C–H: 412
O=O: 498
C=O (in $\text{CO}_2$): 799
O–H: 463
Step-by-Step Solution:
Write the formula for enthalpy of reaction using bond energies:
ΔHrxn = Bonds Broken − Bonds Formed
Identify bonds broken (reactants):
In $\text{CH}_4$: 4 C–H bonds
Energy: 4⋅412=1648 kJIn 2 $\text{O}_2$: 2 O=O bonds
Energy: 2⋅498=996 kJ
Total Energy for Bonds Broken:
1648+996=2644 kJ
Identify bonds formed (products):
In $\text{CO}_2$: 2 C=O bonds
Energy: 2⋅799=1598 kJIn 2 $\text{H}_2\text{O}2$: 4 O–H bonds
Energy: 4⋅463=1852 kJ
Total Energy for Bonds Formed:
1598+1852=3450 kJ
Calculate ΔHrxn:
Using the formula:
ΔHrxn=Bonds Broken−Bonds Formed
Substitute values:
ΔHrxn=2644−3450
Solve:
ΔHrxn=−806 kJ