Internal energy in thermodynamics refers to the total energy contained within a system at a microscopic level. It represents the energy stored by the particles that make up a substance, encompassing their motions and interactions. This concept focuses on the energy inside the system itself, rather than the energy of the system as a whole relative to its surroundings.
What Makes Up Internal Energy
Internal energy comprises two primary forms of energy at the molecular scale: microscopic kinetic energy and microscopic potential energy. These components account for the movements of individual particles and the forces between them.
Microscopic kinetic energy is associated with the motion of atoms and molecules within the system. This includes translational motion, where particles move from one point to another, rotational motion, where molecules spin around an axis, and vibrational motion, where atoms within a molecule oscillate back and forth. Higher temperatures correspond to greater average kinetic energy of these particles.
Microscopic potential energy, on the other hand, is stored in the forces that bind particles together. This includes the energy associated with intermolecular forces and the energy stored in chemical bonds between atoms. Changes in the arrangement or phase of matter, such as melting ice into water or boiling water into steam, involve changes in this potential energy component. For example, heat added during a phase change primarily increases the potential energy between molecules rather than their kinetic energy.
How Internal Energy Increases or Decreases
A system’s internal energy can change through two main mechanisms: the transfer of heat and the performance of work.
When heat is transferred to a system, its internal energy typically increases. This added heat causes the particles within the system to move more vigorously, increasing their average kinetic energy and, consequently, the system’s temperature. Conversely, removing heat from a system leads to a decrease in its internal energy.
Work done on a system also increases its internal energy. For instance, compressing a gas involves work being done on the gas, which adds energy to its molecules and raises its internal energy. Conversely, when a system does work on its surroundings, such as an expanding gas pushing a piston, its internal energy decreases as energy is expended. The principle of energy conservation dictates that any change in a system’s internal energy is a direct result of the balance between heat transferred and work performed.
Internal Energy Compared to Other Energies
Internal energy is distinct from other forms of energy a system might possess, particularly macroscopic forms.
Macroscopic kinetic energy describes the energy of a system due to its overall motion, like a moving car or a thrown ball. This differs from internal energy, which relates to the random, disordered motion of particles within the system, not the organized movement of the entire system. A glass of water sitting still on a table has no macroscopic kinetic energy, but its molecules are in constant, rapid motion contributing to its internal energy.
Similarly, internal energy differs from macroscopic potential energy, such as gravitational potential energy. Gravitational potential energy depends on a system’s position in an external force field, like an object’s height above the ground. In contrast, the potential energy component of internal energy arises from the forces and interactions between the particles within the system, not external forces acting on the system as a whole. While changes in internal energy are the focus in thermodynamics, its absolute value, much like gravitational potential energy, is not typically measured, but rather the changes are important.
Everyday Examples of Internal Energy
Everyday phenomena demonstrate how internal energy changes in response to heat and work.
Consider heating water in a kettle. As heat is added, the water molecules gain kinetic energy, causing the temperature to rise. If heating continues to the boiling point, the added energy then increases the potential energy between water molecules, allowing them to break free from liquid bonds and become steam, a phase change that occurs without a temperature increase.
Rubbing your hands together on a cold day is a common example. The work done by friction between your hands is converted into internal energy, which increases the kinetic energy of the molecules in your skin, making your hands feel warm.
Compressing a gas, such as with a bicycle pump, causes it to become warm. The work done by the pump on the air molecules increases their internal energy, which manifests as an increase in temperature.
Melting ice also illustrates internal energy changes. When ice melts, heat is absorbed, but the temperature of the ice-water mixture remains constant. This absorbed heat energy primarily goes into increasing the potential energy between water molecules, allowing them to transition from a rigid solid structure to a more fluid liquid state.