when heat is added to a system and no external work is done by the system, how does the internal energy of the system change?
The correct answer and explanation is:
When heat is added to a system and no external work is done by the system, the internal energy of the system increases.
This is a consequence of the first law of thermodynamics, which states that the change in internal energy of a system is equal to the heat added to the system minus the work done by the system. Mathematically, this is expressed as: ΔU=Q−W\Delta U = Q – W
Where:
- ΔU\Delta U is the change in internal energy
- QQ is the heat added to the system
- WW is the work done by the system
In this case, since no external work is done by the system (W=0W = 0), the equation simplifies to: ΔU=Q\Delta U = Q
Therefore, the change in internal energy is directly equal to the heat added to the system. As heat is added, the internal energy increases because the added heat increases the kinetic energy of the particles in the system. This leads to an increase in temperature if the system is a gas or a liquid.
In systems like gases, this additional heat causes the molecules or atoms to move faster, which directly increases the temperature. In solids and liquids, added heat may lead to increased vibration or movement of the particles within the substance. The overall effect is an increase in the internal energy, reflected as a rise in temperature or, in the case of phase changes, a change in the state of the substance.
It is important to note that the nature of the system matters. For example, if the system undergoes a phase change (such as melting or boiling), the added heat may be used to overcome intermolecular forces, rather than increasing temperature, but still results in an increase in internal energy.