Question
Use the table.
Bond H-H Bond energy 432
CI-CI. Bond energy 239
H-CI. Bond energy 427
Explain how you could calculate the total change in bond energy for the reaction H2 + CI 2 > 2HCI. How would you know if the reaction was endothermic or exothermic?
Bond H-H Bond energy 432
CI-CI. Bond energy 239
H-CI. Bond energy 427
Explain how you could calculate the total change in bond energy for the reaction H2 + CI 2 > 2HCI. How would you know if the reaction was endothermic or exothermic?
Answers
GPT-4o mini
To calculate the total change in bond energy for the reaction \( \text{H}_2 + \text{Cl}_2 \rightarrow 2 \text{HCl} \), you can follow these steps:
1. **Identify the bonds broken and formed:**
- In the reactants, you have one \( \text{H}_2 \) molecule and one \( \text{Cl}_2 \) molecule. Breaking these bonds will require energy.
- You need to break 1 \( \text{H-H} \) bond (bond energy = 432 kJ) and 1 \( \text{Cl-Cl} \) bond (bond energy = 239 kJ).
- In the products, you will form two \( \text{H-Cl} \) bonds. Since there are two \( \text{HCl} \) molecules produced, you will have two \( \text{H-Cl} \) bonds (bond energy = 427 kJ each).
2. **Calculate the total energy for bonds broken:**
- Energy required to break bonds:
\[
\text{Energy for bonds broken} = \text{Bond energy of } \text{H-H} + \text{Bond energy of } \text{Cl-Cl}
\]
\[
= 432 \, \text{kJ} + 239 \, \text{kJ} = 671 \, \text{kJ}
\]
3. **Calculate the total energy for bonds formed:**
- Energy released from bonds formed:
\[
\text{Energy for bonds formed} = 2 \times \text{Bond energy of } \text{H-Cl}
\]
\[
= 2 \times 427 \, \text{kJ} = 854 \, \text{kJ}
\]
4. **Calculate the total change in bond energy:**
- The total change in bond energy (ΔH) for the reaction is given by:
\[
\Delta H = \text{Total energy for bonds broken} - \text{Total energy for bonds formed}
\]
\[
= 671 \, \text{kJ} - 854 \, \text{kJ} = -183 \, \text{kJ}
\]
5. **Determine if the reaction is endothermic or exothermic:**
- If ΔH is negative (as it is here: -183 kJ), the reaction is exothermic, meaning that it releases energy to the surroundings.
- If ΔH were positive, the reaction would be endothermic, meaning it absorbs energy from the surroundings.
In summary, by calculating the total energies for bonds broken and formed, you determined that the reaction releases energy, indicating it is exothermic.
1. **Identify the bonds broken and formed:**
- In the reactants, you have one \( \text{H}_2 \) molecule and one \( \text{Cl}_2 \) molecule. Breaking these bonds will require energy.
- You need to break 1 \( \text{H-H} \) bond (bond energy = 432 kJ) and 1 \( \text{Cl-Cl} \) bond (bond energy = 239 kJ).
- In the products, you will form two \( \text{H-Cl} \) bonds. Since there are two \( \text{HCl} \) molecules produced, you will have two \( \text{H-Cl} \) bonds (bond energy = 427 kJ each).
2. **Calculate the total energy for bonds broken:**
- Energy required to break bonds:
\[
\text{Energy for bonds broken} = \text{Bond energy of } \text{H-H} + \text{Bond energy of } \text{Cl-Cl}
\]
\[
= 432 \, \text{kJ} + 239 \, \text{kJ} = 671 \, \text{kJ}
\]
3. **Calculate the total energy for bonds formed:**
- Energy released from bonds formed:
\[
\text{Energy for bonds formed} = 2 \times \text{Bond energy of } \text{H-Cl}
\]
\[
= 2 \times 427 \, \text{kJ} = 854 \, \text{kJ}
\]
4. **Calculate the total change in bond energy:**
- The total change in bond energy (ΔH) for the reaction is given by:
\[
\Delta H = \text{Total energy for bonds broken} - \text{Total energy for bonds formed}
\]
\[
= 671 \, \text{kJ} - 854 \, \text{kJ} = -183 \, \text{kJ}
\]
5. **Determine if the reaction is endothermic or exothermic:**
- If ΔH is negative (as it is here: -183 kJ), the reaction is exothermic, meaning that it releases energy to the surroundings.
- If ΔH were positive, the reaction would be endothermic, meaning it absorbs energy from the surroundings.
In summary, by calculating the total energies for bonds broken and formed, you determined that the reaction releases energy, indicating it is exothermic.