I included a picture of equation 7.15 (Tro pg. 292) which is metioned below
a). What are the three major problems associated with the emissions of fossil fuel combustion?
- One renewable energy fuel source is hydrogen. Hydrogen fuel cells use the chemical energy of hydrogen to cleanly and efficiently produce electricity. However, hydrogen is usually “tied-up” in other molecules such as hydrocarbons and must be harvested. The two main ways this is done is by (1) “splitting” water to make hydrogen gas or (2) reforming natural gas (methane). However, process 2 means we are using fossil fuels (non-renewable energy sources) to create renewable energy.
Using the standard enthalpies of formation in the table below and equation 7.15 (Tro pg. 292), calculate ΔHrxn for reaction 2.
Substance
ΔHf°(kJ/mol)
CH4(g)
-74.6
H2O(g)
-241.8
H2O(l)
-285.8
CO2(g)
-393.5
CO2(aq)
-413.8
b. compare the reaction enthalpies for splitting water and for reforming natural gas.
Transcribed Image Text:### Understanding Enthalpy and Hess's Law
As we know from Hess's law, the enthalpy of reaction for the overall reaction is the sum of the enthalpies of reaction of the individual steps:
1. \( \text{CH}_4(g) + \text{O}_2(g) \rightarrow \text{C}(s, \text{graphite}) + 2\text{H}_2O(l) \) \quad \( \Delta H_1^\circ = +74.6 \, \text{kJ/mol} \)
2. (a) \( \text{C}(s, \text{graphite}) + \text{O}_2(g) \rightarrow \text{CO}_2(g) \) \quad \( \Delta H_2^\circ = -393.5 \, \text{kJ/mol} \)
2. (b) \( 2\text{H}_2O(l) + \text{O}_2(g) \rightarrow 2\text{H}_2O(g) \) \quad \( \Delta H_3^\circ = -483.6 \, \text{kJ/mol} \)
Final Equation:
\( \text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2O(g) \) \quad \( \Delta H_\text{rxn}^\circ = -802.5 \, \text{kJ/mol} \)
#### General Process for Calculating \( \Delta H_\text{rxn} \)
To calculate \( \Delta H_\text{rxn} \), subtract the enthalpies of formation of the reactants multiplied by their stoichiometric coefficients from the enthalpies of formation of the products multiplied by their stoichiometric coefficients.
Expressed in equation form:
\[
\Delta H_\text{rxn}^\circ = \sum n_p \Delta H_{f}^\circ (\text{products}) - \sum n_r \Delta H_{f}^\circ (\text{reactants})
\]
Eq. [7.15]
In this equation:
- \( n_p \) represents the stoichiometric coefficients of the products.
- \( n_r \) represents the stoichiometric coefficients of the reactants.
- \( \Delta H
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![### Understanding Enthalpy and Hess's Law
As we know from Hess's law, the enthalpy of reaction for the overall reaction is the sum of the enthalpies of reaction of the individual steps:
1. \( \text{CH}_4(g) + \text{O}_2(g) \rightarrow \text{C}(s, \text{graphite}) + 2\text{H}_2O(l) \) \quad \( \Delta H_1^\circ = +74.6 \, \text{kJ/mol} \)
2. (a) \( \text{C}(s, \text{graphite}) + \text{O}_2(g) \rightarrow \text{CO}_2(g) \) \quad \( \Delta H_2^\circ = -393.5 \, \text{kJ/mol} \)
2. (b) \( 2\text{H}_2O(l) + \text{O}_2(g) \rightarrow 2\text{H}_2O(g) \) \quad \( \Delta H_3^\circ = -483.6 \, \text{kJ/mol} \)
Final Equation:
\( \text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2O(g) \) \quad \( \Delta H_\text{rxn}^\circ = -802.5 \, \text{kJ/mol} \)
#### General Process for Calculating \( \Delta H_\text{rxn} \)
To calculate \( \Delta H_\text{rxn} \), subtract the enthalpies of formation of the reactants multiplied by their stoichiometric coefficients from the enthalpies of formation of the products multiplied by their stoichiometric coefficients.
Expressed in equation form:
\[
\Delta H_\text{rxn}^\circ = \sum n_p \Delta H_{f}^\circ (\text{products}) - \sum n_r \Delta H_{f}^\circ (\text{reactants})
\]
Eq. [7.15]
In this equation:
- \( n_p \) represents the stoichiometric coefficients of the products.
- \( n_r \) represents the stoichiometric coefficients of the reactants.
- \( \Delta H](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F1bf8ae83-2003-4837-b9fe-b672e6e8d8b8%2Fd1e67482-8e03-4813-b141-a2926c5d1012%2Fl7bcbvzq.jpeg&w=3840&q=75)
