If a system has 4.50 × 102 kcal of work done to it, and releases 5.00 x 10² kJ of heat into its surroundings, what is the change in internal energy (AE or AU) of the system? | ΔΕ = I Enter numeric value kJ
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.
![### Thermodynamics Problem
**Problem Statement:**
If a system has \(4.50 \times 10^2\) kcal of work done to it, and releases \(5.00 \times 10^2\) kJ of heat into its surroundings, what is the change in internal energy (\(\Delta E\) or \(\Delta U\)) of the system?
**Solution:**
To solve this problem, we can use the First Law of Thermodynamics, which is expressed as:
\[
\Delta E = Q - W
\]
Where:
- \(\Delta E\) is the change in internal energy.
- \(Q\) is the heat absorbed by the system.
- \(W\) is the work done on the system.
**Given:**
- Work done on the system, \(W = 4.50 \times 10^2\) kcal.
- Heat released by the system, \(Q = -5.00 \times 10^2\) kJ.
**Conversion:**
1 kcal = 4.184 kJ
Convert the work done to kJ:
\[ W = 4.50 \times 10^2 \, \text{kcal} \times 4.184 \, \text{kJ/kcal} = 1882.8 \, \text{kJ} \]
**Calculation:**
Substitute the values into the equation:
\[
\Delta E = (-5.00 \times 10^2 \, \text{kJ}) - 1882.8 \, \text{kJ}
\]
Calculate \(\Delta E\):
\[
\Delta E = -500 - 1882.8 = -2382.8 \, \text{kJ}
\]
**Conclusion:**
The change in internal energy of the system is \(-2382.8\) kJ.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F76884ef0-0050-4a4d-a83f-5c6d84fc3fcd%2Fd6265d82-a12c-4fbc-9044-a4b20423ea5e%2Fdtiq8x7_processed.jpeg&w=3840&q=75)
Trending now
This is a popular solution!
Step by step
Solved in 2 steps with 2 images
