Methane-producing bacteria convert liquid acetic acid (CH3CO₂H) into CO₂(g) and CH₂(g). AS° ΔΗ, Ο (kJ/mol) (J/mol K) CH3CO₂H (1) -484.5 159.8 CO₂(g) -393.5 213.8 CH₂(g) -74.8 186.2 Calculate ΔΗ. = kJ Calculate AG rxn = rxn (round to 3 sig figs) Is this process endothermic or exothermic under standard conditions? Is the reaction spontaneous under standard conditions? kJ (round to 3 sig figs)
Thermochemistry
Thermochemistry can be considered as a branch of thermodynamics that deals with the connections between warmth, work, and various types of energy, formed because of different synthetic and actual cycles. Thermochemistry describes the energy changes that occur as a result of reactions or chemical changes in a substance.
Exergonic Reaction
The term exergonic is derived from the Greek word in which ‘ergon’ means work and exergonic means ‘work outside’. Exergonic reactions releases work energy. Exergonic reactions are different from exothermic reactions, the one that releases only heat energy during the course of the reaction. So, exothermic reaction is one type of exergonic reaction. Exergonic reaction releases work energy in different forms like heat, light or sound. For example, a glow stick releases light making that an exergonic reaction and not an exothermic reaction since no heat is released. Even endothermic reactions at very high temperature are exergonic.
![## Conversion of Liquid Acetic Acid to CO₂ and CH₄ by Methane-Producing Bacteria
This educational page covers the thermodynamic calculations involved in converting liquid acetic acid (CH₃CO₂H) into carbon dioxide (CO₂) and methane (CH₄).
### Thermodynamic Data Table
Below is a table of standard enthalpy (ΔHf°) and standard entropy (ΔS°) values for the reactants and products:
| Compound | ΔHf° (kJ/mol) | ΔS° (J/mol∙K) |
|----------------|---------------|---------------|
| CH₃CO₂H (l) | -484.5 | 159.8 |
| CO₂ (g) | -393.5 | 213.8 |
| CH₄ (g) | -74.8 | 186.2 |
### Calculations
#### Enthalpy Change (ΔH°rxn)
To calculate the enthalpy change of the reaction (ΔH°rxn), use the equation:
\[ \Delta H^\circ_\text{rxn} = \Sigma \Delta H^\circ_\text{products} - \Sigma \Delta H^\circ_\text{reactants} \]
\[
ΔH^\circ_\text{rxn} = [ (ΔH^\circ_{f,\ CO₂}) + (ΔH^\circ_{f,\ CH₄}) ] - [ ΔH^\circ_{f,\ CH₃CO₂H} ]
\]
\[
ΔH^\circ_\text{rxn} = [ (-393.5 \ \text{kJ/mol}) + (-74.8 \ \text{kJ/mol}) ] - [ -484.5 \ \text{kJ/mol} ]
\]
\[
ΔH^\circ_\text{rxn} = -468.3 \ \text{kJ/mol} + 484.5 \ \text{kJ/mol} = 16.2 \ \text{kJ/mol}
\]
#### Free Energy Change (ΔG°rxn)
To calculate the free energy change of the reaction (ΔG°rxn), we use the Gibbs free energy equation:
\[ \Delta G^\circ_\text{rxn} = \Delta H^\circ_\text{rxn} - T \Delta S^\circ_\text](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F29f12fd2-649f-46fc-9623-49951444b106%2Ff5638ea7-ae16-4811-94f2-9426509012ab%2Fgryy3vm_processed.jpeg&w=3840&q=75)
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