Bond Enthalpy — Revision Notes

⚡ 30-Second Revision

Definition: — Average energy to break 1 mole of bonds in gaseous state.
Sign Convention: — Bond breaking is endothermic (energy absorbed, +ve). Bond formation is exothermic (energy released, -ve).
Formula for $\Delta H_{rxn}$: — $\Delta H_{rxn} = \sum E_{\text{bonds broken}} - \sum E_{\text{bonds formed}}$
Factors affecting $E_{\text{bond}}$: — Bond order (triple > double > single), atomic size (smaller atoms = stronger bonds).
BDE vs. Average Bond Enthalpy: — BDE is specific, average bond enthalpy is generalized (for polyatomic molecules).

2-Minute Revision

Bond enthalpy, or bond energy, is a measure of the strength of a chemical bond, representing the average energy required to break one mole of a specific type of bond in the gaseous state. This process is always endothermic, meaning energy is absorbed, so bond enthalpy values are positive.

Conversely, bond formation is an exothermic process, releasing energy. For polyatomic molecules, we use average bond enthalpies because the energy to break identical bonds can vary depending on their molecular environment.

Key factors influencing bond enthalpy include bond order (triple bonds are strongest, then double, then single), and atomic size (smaller atoms generally form stronger bonds). The most crucial application for NEET is calculating the enthalpy change of a reaction ($\Delta H_{rxn}$) using the formula: $\Delta H_{rxn} = \sum (\text{bond enthalpies of bonds broken in reactants}) - \sum (\text{bond enthalpies of bonds formed in products})$.

Remember to correctly count all bonds and apply stoichiometric coefficients. A positive $\Delta H_{rxn}$ indicates an endothermic reaction, while a negative value indicates an exothermic reaction.

5-Minute Revision

Bond enthalpy is a fundamental thermodynamic concept that quantifies the strength of a chemical bond. It's defined as the average energy needed to break one mole of a particular bond type in the gaseous state. Since energy must be supplied to break bonds, bond enthalpy values are always positive (endothermic process). When bonds form, energy is released (exothermic process), and the energy change is negative of the bond enthalpy.

For diatomic molecules like H$_2$ or Cl$_2$, the bond enthalpy is equivalent to the bond dissociation energy (BDE). However, for polyatomic molecules (e.g., CH$_4$), the energy required to break successive identical bonds (like C-H) can differ. Therefore, we use 'average bond enthalpy' – an average value for that bond type across various molecules, which is typically provided in tables.

Several factors influence bond enthalpy:

1
Bond Order: — Higher bond order (e.g., C\(\equiv\)C > C=C > C-C) leads to stronger bonds and higher bond enthalpy.
2
Atomic Size: — Smaller atoms form stronger bonds due to greater orbital overlap and closer nuclei (e.g., H-F > H-Cl).
3
Electronegativity Difference: — A larger difference can increase bond strength due to partial ionic character.

The primary application for NEET is calculating the enthalpy change of a reaction ($\Delta H_{rxn}$) using bond enthalpies. The formula is:

$$\Delta H_{rxn} = \sum E_{\text{bonds broken}} - \sum E_{\text{bonds formed}}$$

Worked Example: Calculate $\Delta H$ for H$_2$(g) + I$_2$(g) \(\rightarrow\) 2HI(g). Given bond enthalpies: H-H = $436\,\text{kJ/mol}$, I-I = $151\,\text{kJ/mol}$, H-I = $299\,\text{kJ/mol}$.

Bonds broken (reactants):

1 H-H bond: $436\,\text{kJ/mol}$
1 I-I bond: $151\,\text{kJ/mol}$

Total energy for bonds broken = $436 + 151 = 587\,\text{kJ/mol}$

Bonds formed (products):

2 H-I bonds: $2 \times 299\,\text{kJ/mol} = 598\,\text{kJ/mol}$

Calculate $\Delta H_{rxn}$:

$\Delta H_{rxn} = 587\,\text{kJ/mol} - 598\,\text{kJ/mol} = -11\,\text{kJ/mol}$

This indicates an exothermic reaction. Always ensure you correctly identify and count all bonds, especially in polyatomic molecules and when stoichiometric coefficients are involved. Pay close attention to the sign convention.

Prelims Revision Notes

1
Definition: — Bond enthalpy ($E_{\text{bond}}$) is the average energy required to break one mole of a specific bond in the gaseous state. It's always positive (endothermic process).
2
Bond Formation: — The reverse of bond breaking; energy is released (exothermic process), so $\Delta H_{\text{formation}} = -E_{\text{bond}}$.
3
Bond Dissociation Energy (BDE): — Energy to break a specific bond in a particular molecule. For diatomic molecules, BDE = $E_{\text{bond}}$. For polyatomic molecules, BDEs for identical bonds can vary, hence average bond enthalpy is used.
4
**Factors Affecting Bond Strength (and $E_{\text{bond}}$):**

* Bond Order: Triple bonds > Double bonds > Single bonds (e.g., C\(\equiv\)C > C=C > C-C). * Atomic Size: Smaller atoms form stronger bonds (e.g., H-F > H-Cl). * Electronegativity Difference: Can increase bond strength due to partial ionic character. * Lone Pair Repulsions: Can weaken bonds (e.g., F-F is weaker than Cl-Cl).

1
**Calculation of Reaction Enthalpy ($\Delta H_{rxn}$):**

* Formula: $\Delta H_{rxn} = \sum (\text{Bond enthalpies of bonds broken in reactants}) - \sum (\text{Bond enthalpies of bonds formed in products})$ * Steps: 1. Draw Lewis structures or visualize bonds for all reactants and products.

2. Identify and count each type of bond broken in reactants (multiply by stoichiometric coefficients). 3. Identify and count each type of bond formed in products (multiply by stoichiometric coefficients).

4. Sum the bond enthalpies for bonds broken. 5. Sum the bond enthalpies for bonds formed. 6. Apply the formula. A negative $\Delta H_{rxn}$ means exothermic; a positive $\Delta H_{rxn}$ means endothermic.

1
Common Traps:

* Incorrectly counting bonds (e.g., 2 C=O bonds in CO$_2$, 2 O-H bonds in H$_2$O). * Forgetting to multiply by stoichiometric coefficients. * Reversing the sign of $\Delta H_{rxn}$. * Confusing BDE with average bond enthalpy.

Vyyuha Quick Recall

Break Endothermically, Form Exothermically. (Bonds Break Endothermically, Bonds Form Exothermically).

For the formula: Broken Minus Formed. ($\Delta H_{rxn} = \sum E_{\text{bonds broken}} - \sum E_{\text{bonds formed}}$)