Basics of Thermodynamics

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Typical thermodynamic system, showing input from a heat source (boiler) on the left and output to a heat sink (condenser) on the right. Work is extracted, in this case by a series of pistons.

In physics and chemistry, thermodynamics (from the Greek θέρμη therme, meaning "heat"[1] and δύναμις, dynamis, meaning "power") is the study of energy conversion between heat and mechanical work, and subsequently the macroscopic variables such as temperature, volume and pressure. Its progenitor, based on statistical predictions of the collective motion of particles from their microscopic behavior, is the field of statistical thermodynamics (or statistical mechanics), a branch of statistical physics. Historically, thermodynamics developed out of a need to increase the efficiency of early steam engines.


The Four Laws

The present article is focused on classical thermodynamics, which is focused on systems in thermodynamic equilibrium. It is wise to distinguish classical thermodynamics from non-equilibrium thermodynamics, which is concerned with systems that are not in thermodynamic equilibrium.

In thermodynamics, there are four laws that do not depend on the details of the systems under study or how they interact. Hence these laws are very generally valid, can be applied to systems about which one knows nothing other than the balance of energy and matter transfer. Examples of such systems include Einstein's prediction, around the turn of the 20th century, of spontaneous emission, and ongoing research into the thermodynamics of black holes.

These four laws are:

  • Zeroth law of thermodynamics, about thermal equilibrium:

This law was considered so obvious it was added as a virtual afterthought, hence the designation Zeroth, rather than Fourth. In short, if the temperature of material A is equal to the temperature of material B, and B is equal to the temperature of material C. then A and C must also be equal.

  • First law of thermodynamics, about the conservation of energy:
The change in the internal energy of a closed thermodynamic system is equal to the sum of the amount of heat energy supplied to or removed from the system and the work done on or by the system. So, we can say (1) "Energy is neither created nor destroyed" and (2) "There is no free lunch."[2]
  • Second law of thermodynamics, about entropy:
The total entropy of any isolated thermodynamic system always increases over time, approaching a maximum value or we can say, "In an isolated system, the entropy never decreases". Another way to phrase this: Heat cannot spontaneously flow from a colder location to a hotter area - work is required to achieve this.
  • Third law of thermodynamics, about the absolute zero of temperature:
As a system asymptotically approaches absolute zero of temperature all processes virtually cease and the entropy of the system asymptotically approaches a minimum value; also stated as: "the entropy of all systems and of all states of a system is smallest at absolute zero" or equivalently "it is impossible to reach the absolute zero of temperature by any finite number of processes". Absolute zero, at which all activity would stop if it were possible to happen, is −273.15 °C (degrees Celsius), or −459.67 °F (degrees Fahrenheit) or 0 K (kelvins, formerly sometimes degrees absolute).

System models

A diagram of a thermodynamics system.

An important concept in thermodynamics is the “system”. Everything in the universe except the system is known as surroundings. A system is the region of the universe under study. A system is separated from the remainder of the universe by a boundary which may be imaginary or not, but which by convention delimits a finite volume. The possible exchanges of work, heat, or matter between the system and the surroundings take place across this boundary. Boundaries are of four types: fixed, moveable, real, and imaginary.

Basically, the “boundary” is simply an imaginary dotted line drawn around a volume of something when there is going to be a change in the internal energy of that something. Anything that passes across the boundary that effects a change in the internal energy of the something needs to be accounted for in the energy balance equation. That something can be the volumetric region surrounding a single atom resonating energy, such as Max Planck defined in 1900; it can be a body of steam or air in a steam engine, such as Sadi Carnot defined in 1824; it can be the body of a tropical cyclone, such as Kerry Emanuel theorized in 1986 in the field of atmospheric thermodynamics; it could also be just one nuclide (i.e. a system of quarks) as some are theorizing presently in quantum thermodynamics.

For an engine, a fixed boundary means the piston is locked at its position; as such, a constant volume process occurs. In that same engine, a moveable boundary allows the piston to move in and out. For closed systems, boundaries are real while for open system boundaries are often imaginary. There are five dominant classes of systems:

1. Isolated Systems – matter and energy may not cross the boundary
2. Adiabatic Systems – heat must not cross the boundary
3. Diathermic Systems - heat may cross boundary
4. Closed Systems – matter may not cross the boundary or fixed mass
5. Open Systems – heat, work, and matter may cross the boundary (often called a control volume in this case)

As time passes in an isolated system, internal differences in the system tend to even out and pressures and temperatures tend to equalize, as do density differences. A system in which all equalizing processes have gone practically to completion, is considered to be in a state of thermodynamic equilibrium.

In thermodynamic equilibrium, a system's properties are, by definition, unchanging in time. Systems in equilibrium are much simpler and easier to understand than systems which are not in equilibrium. Often, when analysing a thermodynamic process, it can be assumed that each intermediate state in the process is at equilibrium. This will also considerably simplify the situation. Thermodynamic processes which develop so slowly as to allow each intermediate step to be an equilibrium state are said to be reversible processes.

States & processes

When a system is at equilibrium under a given set of conditions, it is said to be in a definite state. The thermodynamic state of the system can be described by a number of intensive variables and extensive variables. The properties of the system can be described by an equation of state which specifies the relationship between these variables. State may be thought of as the instantaneous quantitative description of a system with a set number of variables held constant.

A thermodynamic process may be defined as the energetic evolution of a thermodynamic system proceeding from an initial state to a final state. Typically, each thermodynamic process is distinguished from other processes, in energetic character, according to what parameters, as temperature, pressure, or volume, etc., are held fixed. Furthermore, it is useful to group these processes into pairs, in which each variable held constant is one member of a conjugate pair. The seven most common thermodynamic processes are shown below:

1. An isobaric process occurs at constant pressure.
2. An isochoric process, or isometric/isovolumetric process, occurs at constant volume.
3. An isothermal process occurs at a constant temperature.
4. An adiabatic process occurs without loss or gain of energy by heat.
5. An isentropic process (reversible adiabatic process) occurs at a constant entropy.
6. An isenthalpic process occurs at a constant enthalpy.
7. A steady state process occurs without a change in the internal energy of a system.

Intensive and extensive properties

In the physical sciences, an intensive property (also called a bulk property), is a physical property of a system that does not depend on the system size or the amount of material in the system: it is scale invariant. By contrast, an extensive property of a system is directly proportional to the system size or the amount of material in the system (see examples below). Some intensive properties, such as viscosity, are empirical macroscopic quantities and are not relevant to extremely small systems.

For example, density is an intensive quantity (it does not depend on the quantity), while mass and volume are extensive quantities. Note that the ratio of two extensive quantities that scale in the same way is scale-invariant, and hence an intensive quantity.

Thermodynamic variables

Thermodynamic variables describe the momentary condition of a thermodynamic system. Regardless of the path by which a system goes from one state to another — i.e., the sequence of intermediate states — the total change in any state variable will be the same. This means that the incremental changes in such variables are exact differentials. Examples of state variables include:

  • Density (ρ)
  • Energy (E)
  • Helmholtz free energy (A)
  • Gibbs free energy (G)
  • Enthalpy (H)
  • Internal energy (U)
  • Mass (m)
  • Exergy
  • Pressure (p)
  • Entropy (S)
  • Temperature (T)
  • Volume (V)
  • The amounts of each of the chemical components { ni }, expressed as numbers of moles or molecules

Equilibrium state

Systems found in nature are often dynamic and complex, but in many cases their states are amenable to description based on proximity to ideal conditions. One such ideal condition is that of a stable equilibrium state. Based on many observations, thermodynamics postulates that all systems having no effect on the external environment will change in such a way as to approach unique stable equilibrium states.


  • Modell, Michael; Robert C. Reid (1974). Thermodynamics and Its Applications. Englewood Cliffs, NJ: Prentice-Hall.
  • Cengel, Yunus A.; Boles, Michael; (2001). Thermodynamics: An Engineering Approach. McGraw-Hill Science.

External Links

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