Final answer:
According to the first law of thermodynamics, the change in internal energy of a system (ΔU) is the net heat added to the system (Q) minus the work done by the system (W). The provided examples illustrate how these energy changes are calculated. The relationship between heat and work in a thermodynamic process depends on the specific conditions and states of the system. So, the answer is: a) True b) False.
Step-by-step explanation:
In physics, especially in the context of thermodynamics, energy can be transferred into or out of a system in two forms: as work or as heat. The change in internal energy (ΔU) of a system is determined by the net heat added to the system (Q) and the work done by the system (W), according to the first law of thermodynamics, which is ΔU = Q - W. For part (a), if a system does 49 J of work (which means it is work done by the system on the surroundings, and hence we consider it as -49 J) and its internal energy decreases by 73 J, we can say that the net heat must have been removed from the system. Similarly, for part (b), if a gas releases 38 J of heat (which is -38 J since it is leaving the system) and has 109 J of work done on it (which increases the system's internal energy), then overall, despite the work done on the gas increasing the internal energy, the energy flowing out due to heat leads to a net decrease in internal energy. To address the example given with calculations: If there is a heat transfer of 40.00 J to a system (positive since it's added to the system) and the system does 10.00 J of work (negative since the system is doing work), then the change in internal energy after these interactions would be ΔU = Q - W = 40 J - (-10 J) = 50 J. Later, if there is a heat transfer of 25.00 J out of the system (-25 J since it's leaving the system) and 4.00 J of work is done on the system (positive), the additional change in internal energy would be ΔU = -25 J - 4 J = -29 J. The net change in internal energy would then be 50 J - 29 J = 21 J.