Final answer:
Statements D ('The internal energy of a system is not a state function') and E ('Work being done by the system onto the surroundings is given a positive sign') are NOT true. Internal energy is a state function and work done by the system is considered negative in thermodynamic conventions.
Step-by-step explanation:
The statements to consider regarding a thermodynamic system are:
- A) The internal energy of a system is the sum of all energy in the system.
- B) Heat entering a system is given a positive sign.
- C) The internal energy of a system is equal to heat plus work.
- D) The internal energy of a system is not a state function.
- E) Work being done by the system onto the surroundings is given a positive sign.
Circling the statements that are NOT true, we must identify errors based on the first law of thermodynamics:
- Statement D is not true. The internal energy of a system is a state function, which means it's determined by the current state of the system, not the path the system took to reach that state.
- Statement E is not true. When the system does work onto the surroundings, the work is considered negative according to the convention most commonly used in physics.
The remaining statements, A, B, and C, are consistent with the principles of thermodynamics. Specifically, B) is correct because heat added to a system increases its internal energy and is thus considered positive. Statement C) is a simplified expression of the first law of thermodynamics, where the change in internal energy (ΔU) is equal to the heat added to the system (q) plus the work done on the system (w).