Thermodynamics - the science that is concerned with energy, particularly ‘energy-in-transit’ in the forms of heat and work, and those properties of systems that are related to energy.

Energy – the ability to do work.  All energy is relative!  Energy-in-transit is not relative.

Three kinds of energy:

(1) potential - energy due to relative position,

(2) kinetic   - energy due to relative velocity,

(3) internal  - the sum of all potential and kinetic energies of constituent parts [atoms, molecules, etc.] of a system.

Two kinds of ‘energy-in-transit’:

(1) heat – energy transferred between system and surroundings because of a temperature difference, or gradient.

(2) work - energy transferred between system and surroundings because of a pressure difference, or gradient.



Thermodynamic System – just “the thing” that we are talking about!  Everything else is called the surroundings.  The sum of the system and the surroundings is the universe.

Three kinds of systems:

(1) closed system – a fixed quantity of material; energy can cross the system boundaries but mass can not.

(2) open system – a particular region of space; both mass and energy may cross the system boundaries.

(3) isolated system (not an important concept) – neither energy nor mass may cross the system boundaries.

In elementary thermodynamics all systems consist only of atoms and molecules where the net electric charge of the system is zero.  In addition, all electrical and magnetic and surface forces are generally neglected.



Thermodynamic Materials - Systems composed of atoms and molecules are called materials.

Two kinds of materials:

(1) pure materials - composed of only one molecular species, and

(2) mixtures - composed of two or more molecular species.

ideal mixtures - mixtures where the volume and enthalpy of the mixture are simply the sums of the volumes and enthalpies of the pure components at the temperature and pressure of the mixture.  Elementary thermodynamics deals only with ideal mixtures.  Advanced thermodynamics is concerned with non-ideal mixtures, in phase equilibrium and reaction equilibrium.

Four basic concepts of materials:

(1) Quantity

    (a) mass (or weight in a known gravitational field)

    (b) number of objects (one gram mole = 6.025 x 10^23 objects)

mean-molar-mass (molecular weight or atomic weight) is the mass of one mole of a particular collection of objects, and is the constant which allow conversion between these two measures of quantity.

(2) Composition of a mixture

    (a) fraction - quantity of a particular species per unit quantity of the mixture.

    (b) concentration - quantity of a particular species per unit volume of the mixture.

(3) Phase - a homogeneous quantity of material, characterized throughout by a single set of thermodynamic properties.

    (a) solids - materials which are capable of resisting shear stresses.

    (b) fluids - materials which exhibit continuous deformation under shear stress.

        (c) liquids - fluids which can conform to their containers without occupying them completely.

        (d) gases - fluids which conform to and completely occupy their containers.

        (e) vapors - gases at temperatures less than their critical temperature.

quality - ratio of quantity of vapor to the total quantity of material [vapor & liquid] or [vapor & solid] in a system.

(4) State - defined by the properties of a material.

    (a) subcooled liquid (or compressed liquid) - a liquid at a temperature below its saturation temperature or at a pressure above its saturation pressure.

    (b) superheated vapor - a vapor at a temperature above its saturation temperature or at a pressure below its saturation pressure.

    (c) saturated - if two or more phases exist within a system at equilibrium, the system is said to be saturated and all phases present are saturated.  In particular, if vapor and liquid phases are both present within a system, the vapor is said to be saturated vapor and the liquid is said to be saturated liquid.  Similarly, if two liquid phases exist within a system at equilibrium, both liquid phases are saturated.

saturation pressure (or vapor pressure) - the pressure at which a phase change will take place at a given temperature.

saturation temperature - the temperature at which a phase change will take place at a given pressure.

critical point - that state of a saturated system where the liquid and vapor phases become indistinguishable.  The properties of a material at its critical point are the same for both vapor and liquid phases.

equilibrium - the condition of a system in which no net change in the properties of the system occur with time.  A closed system is usually implied.

[steady state - no accumulations of matter or energy occur within the system.  An open system is implied.]


Thermodynamic Properties - any quantity that depends only on the state of a material and is independent of the process by which a material arrives at a given state.

Properties of a System - the average or homogeneous properties of a system at equilibrium.

Two kinds of properties:

(1) intensive - independent of the quantity of material [T, P, Cp and Cv], and all specific and molar properties.

(2) extensive - directly proportional to the quantity of material [V, S, U, H, etc.].

Pseudointensive properties - extensive properties expressed per unit quantity of material [v, s, u, h, etc.].

Two kinds of pseudointensive properties:

(1) specific properties - expressed on a unit mass basis, and

(2) molar properties - expressed on a unit mole basis.

Five basic thermodynamic properties:

(1) temperature [T] (thermal potential) - a measure of the relative hotness or coldness of a material.

(2) pressure [P] (mechanical potential) - the normal (perpendicular) component of force per unit area.

(3) volume [V] (mechanical displacement) - the quantity of space possessed by a material.

(4) entropy [S] (thermal displacement) - the quantity of disorder possessed by a material.

(5) internal energy [U] - the energy of a material which is due to the kinetic and potential energies of its constituent parts (atoms and molecules, usually).

Two secondary thermodynamic properties:

(1) enthalpy [H] - internal energy plus the pressure-volume product.

(2) heat capacity [Cp or Cv] (specific heat) - the amount of energy required to increase the temperature of one unit quantity of material by one degree, under specific conditions.

(a)    constant pressure Cp = dh/dT

(b)    constant volume   Cv = du/dT

Unlike gases, liquids and solids are nearly incompressible, and it is almost impossible to change their temperature while holding their volumes constant.  The specific heats of liquids and solids almost always imply their constant pressure heat capacity (usually on a unit mass basis), so that, in general, for liquids and solids we used Cp.

Gibbs Phase Rule:   F = 2 + Ns - Np

F - degrees of freedom of the system = the number of independent,

    intensive thermodynamic variables (properties or compositions) which

    must be specified to fix the intensive state of the system,

Ns - number of molecular species within the system, and

Np - the number of phases within the system.

The thermodynamic variables specified as degrees of freedom are normally temperature, pressure and compositions (mole fractions) of the phases.  Note that only [Ns - 1] compositions of each phase are independent.  To fix the extensive state of the system, an additional extensive variable must be specified (i.e. total moles of the system).


Thermodynamic Processes and Cycles

process - any succession of events.

chemical process - a chemical or physical operation, or series of operations, which transforms raw materials into products.

thermodynamic process - the path of succession of states through which the system passes in moving from an initial state to a final state.

polytropic process - a thermodynamic process for which [PVn] is constant.  These processes are usually associated only to systems for which the ideal gas assumption holds.

 Four special polytropic processes:

(1) isobaric  - - - - - - - constant pressure   [n = 0]

(2) isothermal  - - - - - - constant temperature[n = 1]

(3) isentropic  - - - - - - constant entropy    [n = gamma,(Cp/Cv)]

(4) isochoric (isometric) - constant volume     [n = infinity]

Two other important processes:

(1) adiabatic - no heat transfer.

(2) isenthalpic - constant enthalpy.  This is the same as isothermal for an ideal gas system.

reversible process – an idealized process in which the deviation from thermodynamic equilibrium is infinitesimal at any particular instant during the process.  All of the states through which a system passes during a reversible process may be considered to be equilibrium states.  This is an idealized situation that would require infinite time and/or equipment size to be realized.  The concept of a reversible process serves to set a maximum for the efficiency of a given process.  Note that an isentropic process is an adiabatic-reversible process, so that real isentropic processes are not possible.

thermodynamic cycle - a process for which the final and initial states are the same.

Four common ‘idealized’ thermodynamic cycles:

(1) Carnot cycle  - isothermal and isentropic compressions followed by

                    isothermal and isentropic expansions.

(2) Rankine cycle - isobaric   and isentropic compressions followed by

  isobaric   and isentropic expansions.

(3) Otto cycle    - isentropic and isochoric  compressions followed by

                    isentropic and isochoric  expansions.

(4) Diesel cycle  - isentropic compression followed by isobaric,

                    isentropic and isochoric expansions


Thermodynamic Data Presentation

Data, such as properties of pure materials, is generally acquired by experimentation and can be presented in three fundamentally different forms:

(1) Tables    [i.e. the steam tables]

(2) Graphs    [i.e. a T-s or P-h diagrams]

(3) Equations [i.e. the ideal gas equation]

Each of these forms of presentation has advantages and disadvantages.

(1) Tables are precise but discontinuous, so that interpolation is often required.  In addition, they can be bulky and can be difficult to use when implicit variables are specified.  They also can require large amounts of data storage when used with computer programs.

(2) Graphs are continuous in their explicit variables but suffer loss of precision when they are of convenient size.  In addition, they are discontinuous for implicit variables, so that imprecise visual interpolation is often required.  They also suffer in readability as the number of implicit variables displayed increases above three or four.  Although they can give an excellent overall "feel" for the data, they are virtually useless for computer purposes.

(3) Equations are in many ways the best form of presentation for data.  They allow mathematical manipulation, are easy to use with computer programs, and are as precise as the data used to generate their constants.  However, equations that accurately represent significantly large ranges of data can be very complex and usually employ a number of constant terms.  Complex equations are usually difficult to solve for their implicit variables and often require trial and error procedures in their use.  They are most suited for use in computer programs.


Thermodynamic Laws

A physical law is a simple statement of an observable physical phenomenon that has no underlying, more-basic reason for being except that the most accurate observations have always proved it to be true.

Laws of Thermodynamics

Zeroth:  Two bodies in thermal equilibrium with a third body are in thermal equilibrium with each other.  (This ?Law? simply states that ‘thermometers work’.)

First:   A Simple Statement: Heat and work are both forms of energy in transit, and energy is always conserved. or

         A Classical Statement: During any cyclic process on a closed system the cyclic integral of heat is always equal to the cyclic integral of work.

Second:  Simple Statement #1: Spontaneous fluxes always take place down their corresponding potential gradients. or

         Simple Statement #2: Heat and work are both forms of energy in transit, but they are not qualitatively equal forms of energy because work can always be converted entirely into heat, but heat can never be converted entirely into work. or

         Kelvin-Plank Statement:  It is impossible to construct a device which operates in a cycle and produces no effect other than the raising of a weight and the exchange of heat with a single reservoir. or

         Clausius Statement:  It is impossible to construct a device which operates in a cycle and produces no effect other than the transfer of heat from a cooler body to a hotter body.

Albert Einstein considered the Second Law of Thermodynamics to be the only real physical law.

Third:   The absolute entropy of a pure, crystalline material at a temperature of absolute zero is zero.  (This ?Law? is the second half of the definition of entropy.)

Other Laws of Importance in Thermodynamics

Conservation of Matter:  Matter can be neither created nor destroyed but only changed from one form to another.  Note that Albert Einstein showed that matter could be ‘destroyed’ by converting it into energy.

Joule's ?Law?:  The internal energy of an ideal gas is a function of temperature only.

Avagadro's ?Law?:  Equal volumes of different ideal gases at the same temperature and pressure contain the same number of molecules.