A generic solid, X, has a molar mass of 59.1 g/mol. In a constant-pressure calorimeter, 15.8 g of X is dissolved in 311 g of water at 23.00 °C. X(s) ->> X(aq) The temperature of the resulting solution rises to 26.50 °C. Assume the solution has the same specific heat as water, 4.184 J/(g-°C), and that there is negligible heat loss to the surroundings. How much heat was absorbed by the solution? q= What is the enthalpy of the reaction? AHrxn= kJ kJ/mol
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.
![**Calorimetry and Enthalpy Experiment**
**Objective:**
To determine the heat absorbed by a solution and the enthalpy of a reaction involving a generic solid.
**Materials and Data:**
- A generic solid, X, with a molar mass of 59.1 g/mol.
- Mass of X: 15.8 g.
- Water volume: 311 g.
- Initial water temperature: 23.00 °C.
- Final solution temperature: 26.50 °C.
- Specific heat capacity of water: 4.184 J/(g·°C).
**Reaction:**
\[ \text{X(s)} \rightarrow \text{X(aq)} \]
**Assumptions:**
- The solution’s specific heat is the same as water.
- Negligible heat loss to surroundings.
**Calculations:**
1. **Heat Absorbed by Solution:**
Calculate the heat absorbed using the formula:
\[
q = m \cdot c \cdot \Delta T
\]
where:
- \( m \) = mass of the solution (15.8 g + 311 g),
- \( c \) = specific heat capacity (4.184 J/g°C),
- \( \Delta T \) = change in temperature (26.50 °C - 23.00 °C).
\[
q = \text{_____} \text{ kJ}
\]
2. **Enthalpy of Reaction (\( \Delta H_{\text{rxn}} \)):**
Determine the enthalpy change per mole of X:
\[
\Delta H_{\text{rxn}} = \frac{q}{\text{moles of } X}
\]
\[
\Delta H_{\text{rxn}} = \text{_____} \text{ kJ/mol}
\]
**Discussion:**
The experiment provides insight into thermodynamic principles by measuring how a reaction’s energy exchange affects temperature. Understanding these concepts is crucial for applications in chemistry and engineering.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fbfab9a7f-9ea4-456a-bb57-050ea1ff4d68%2F4982062e-3d44-40b1-9b14-7e7fe77b6451%2Fl2xa98_processed.jpeg&w=3840&q=75)
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