Based on the provided reference, you can change the internal energy of a system, and according to the reference, this change is equal to the change in the system's kinetic energy when heat is supplied at constant volume.
The Relationship Explained
Internal energy ($U$) represents the total energy contained within a thermodynamic system, including the kinetic energy of its molecules (due to translation, rotation, and vibration) and the potential energy associated with intermolecular forces. Changing a system's internal energy is typically done through processes involving heat transfer and work.
The provided reference highlights a specific scenario governed by the First Law of Thermodynamics, which states:
$\Delta U = Q - W$
Where:
- $\Delta U$ is the change in internal energy.
- $Q$ is the heat added to the system.
- $W$ is the work done by the system.
Constant Volume Conditions
The reference specifies a condition where the volume of the container remains unchanged. In thermodynamics, work done by a system against an external pressure is related to the change in volume ($W = P\Delta V$). If the volume doesn't change ($\Delta V = 0$), then the work done by the system is zero ($W = 0$).
Under this constant volume condition, the First Law of Thermodynamics simplifies to:
$\Delta U = Q$
This means that any heat added to the system directly increases its internal energy, and any heat removed decreases its internal energy.
Internal Energy and Kinetic Energy Link
The reference then makes a key statement connecting internal energy and kinetic energy:
"Therefore , the change in internal energy is equal to the change in Kinetic energy of the system."
While the phrasing "Kinetic energy of the system" could potentially refer to the bulk motion of the entire system, in the context of internal energy and heat transfer within a stationary, constant-volume container, it most likely refers to the internal kinetic energy of the particles within the system. This internal kinetic energy is directly related to the system's temperature.
How Heat Changes Internal Energy and Kinetic Energy
Putting it all together, based on the reference:
- Start with a constant volume system: The container holding the substance is rigid, preventing it from expanding or contracting.
- Supply heat ($Q > 0$): Energy is transferred into the system, for example, by placing the container on a hot plate.
- Work done is zero ($W = 0$): Because the volume doesn't change.
- Internal energy increases ($\Delta U = Q$): According to the First Law of Thermodynamics under constant volume.
- Internal kinetic energy increases: As heat is added, the particles (atoms or molecules) absorb this energy, causing them to move faster, rotate more, or vibrate more intensely. This increases their average microscopic kinetic energy, which is a primary component of internal energy.
- Change in internal energy equals change in kinetic energy (according to the reference): The reference explicitly states that the increase in internal energy caused by the heat is equal to the increase in the "Kinetic energy of the system" (interpreted here as the internal kinetic energy of the particles).
Therefore, according to the provided reference, adding heat to a system held at constant volume is the method described for changing its internal energy, and this change corresponds directly to a change in the internal kinetic energy of the system's particles.
Summary
| Process | Conditions | First Law Simplified | Internal Energy Change | Kinetic Energy Change (per Reference) |
|---|---|---|---|---|
| Heat Transfer | Constant Volume | $\Delta U = Q$ | $\Delta U = Q$ | Equal to $\Delta U$ |
In essence, supplying heat to a fixed volume increases the system's internal energy, primarily by increasing the kinetic energy of its constituent particles.
