Glossary
Adiabatic process
A thermodynamic process where no heat is transferred into or out of the system (Q=0), so any change in internal energy is solely due to work.
Example:
Rapidly compressing air with a bicycle pump is an adiabatic process; the air gets hot because the work done on it increases its internal energy without time for heat to escape.
Boltzmann constant (k_B)
A physical constant relating the average kinetic energy of particles in a gas to the gas's absolute temperature.
Example:
The Boltzmann constant is used when relating the energy of individual particles to the macroscopic temperature of a system.
Closed System
A thermodynamic system that can exchange energy (heat and work) but not matter with its surroundings.
Example:
A sealed balloon filled with air is a closed system; it can expand or contract (doing work) and change temperature (exchanging heat), but no air enters or leaves.
First Law of Thermodynamics
A fundamental principle stating that energy is conserved; it describes how a system's internal energy changes due to heat transfer and work done on or by the system.
Example:
When you inflate a bicycle tire, the air inside gets warmer because the work done on the gas increases its internal energy according to the First Law of Thermodynamics.
Heat (Q)
Energy transferred between systems or a system and its surroundings due to a temperature difference.
Example:
When you place an ice cube in a warm drink, heat flows from the drink to the ice, causing the ice to melt.
Ideal Gases
A theoretical gas composed of many randomly moving point particles that do not interact with each other except through elastic collisions, simplifying the calculation of internal energy.
Example:
For an ideal gas in a sealed container, its internal energy depends only on its temperature, not its volume or pressure.
Ideal gas constant (R)
A physical constant that relates the energy scale to the temperature scale, used in the ideal gas law.
Example:
The ideal gas constant (R) is a universal value that helps connect pressure, volume, temperature, and the amount of gas.
Internal Energy (U)
The total energy contained within a thermodynamic system, comprising the kinetic and potential energies of its constituent particles.
Example:
Heating a pot of water increases the average kinetic energy of its molecules, thereby increasing the water's internal energy.
Isobaric process
A thermodynamic process that occurs at constant pressure, where work done is simply -PΔV.
Example:
Boiling water in an open pot is an isobaric process because the pressure remains constant at atmospheric pressure while the volume changes as steam forms.
Isolated System
A thermodynamic system that cannot exchange either energy (heat or work) or matter with its surroundings.
Example:
A perfectly insulated thermos containing hot coffee approximates an isolated system, as it minimizes heat transfer to the outside.
Isothermal process
A thermodynamic process that occurs at a constant temperature, requiring heat exchange with the surroundings to maintain this temperature.
Example:
A gas expanding slowly in contact with a large heat reservoir undergoes an isothermal process, where any work done is compensated by heat transfer to keep the temperature steady.
Isotherms
Lines on a PV diagram that represent processes occurring at a constant temperature.
Example:
On a PV diagram, an isotherm for an ideal gas will appear as a hyperbola, showing that pressure is inversely proportional to volume at constant temperature.
Isovolumetric (Isochoric) process
A thermodynamic process that occurs at constant volume, meaning no work is done by or on the system (W=0).
Example:
Heating a gas in a rigid, sealed container is an isovolumetric process because the volume cannot change, so all energy added goes directly to internal energy.
Kinetic energy (of particles)
The energy associated with the motion of the atoms and molecules within a system.
Example:
As a gas heats up, its particles move faster, increasing their average kinetic energy.
Number of moles (n)
A unit of measurement for the amount of substance, representing Avogadro's number of particles.
Example:
To calculate the volume of a gas at standard temperature and pressure, you often need to know its number of moles.
Number of particles (N)
The total count of individual atoms or molecules within a given system.
Example:
The pressure exerted by a gas in a container is directly proportional to the number of particles colliding with the walls.
PV Diagrams
Graphs that plot pressure (P) versus volume (V) for a thermodynamic process, useful for visualizing changes and calculating work.
Example:
Engineers use PV diagrams to analyze the efficiency of heat engines by tracing the cycle of a working fluid.
Potential energy (of particles)
The energy stored due to the interactions or positions of the atoms and molecules within a system.
Example:
In a liquid, the attractive forces between molecules contribute to their potential energy, which is less significant in an ideal gas.
Temperature (T)
A measure of the average kinetic energy of the particles within a system, typically expressed in Kelvin for thermodynamic calculations.
Example:
When you touch a hot stove, the high temperature indicates that the stove's particles have a high average kinetic energy.
Work (W)
Energy transferred between a system and its surroundings by means other than heat, often involving a force acting over a distance.
Example:
A piston compressing a gas performs work on the gas, increasing its internal energy.
Work done by a system
Work performed by the system on its surroundings, typically resulting in a decrease in the system's internal energy if no heat is added.
Example:
When a gas expands and pushes a piston, it is performing work done by a system, which is conventionally negative in the First Law equation.