A system does 506 kJ of work and loses 271 kJ of heat to the surroundings. What is the change in internal energy, Δ?ΔE, of the system? Note that internal energy is symbolized as Δ?ΔU in some sources.
A system does 506 kJ of work and loses 271 kJ of heat to the surroundings.
What is the change in internal energy, Δ?ΔE, of the system? Note that internal energy is symbolized as Δ?ΔU in some sources.
Please find your solution below :
According to First Law of Thermodynamics, heat can neither we created nor be destroyed, it can only be converted from one form to the another. Therefore, energy of universe remains same and can be exchanged between system and surroundings. This law give the relation between changes in energy states due to work and heat transfer.
ΔU = q + W
where,
ΔU is change in internal energy of the system
q is algebraic sum of heat transfer between system and surroundings
W is total interaction of work of the system with its surroundings
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