Thermodynamics & Heat

Thermodynamics is the study of heat ("thermo"), which means "internal molecular energy" and its movement from object ("system") to object and heat effects (work) on bodies (systems) and their surroundings.   Thermodynamics studies the movement of heat and how it does work.   This study began in the late 18th century (see below) to determine the efficiency of the newly invented steam engine.   It has since been extended to assist in the understanding of many fields of science: chemical reactions, astronomy, physiology, materials, etc.

Heat is the internal molecular energy transferred from one body (system) to another because of a difference in temperature.   The SI standard of heat is the Joule.   4.186 J = 1 calorie, which is the amount of heat required to raise 1 gram of water 1 °C. from 14.5°C to 15.5°C.

Classical thermodynamics studies heat and work macroscopically and is summarized in the famous equation

Ideal Gas Law: P•V = n•R•T, where

P = pressure
V = volume
n = mass as measured in moles (1 mole = amount of substance containing the same number of atoms or molecules or ions as 12.00 grams of carbon-12)
R = universal gas constant = 8.315 Joules/mol·K
T = absolute (Kelvin) temperature

T is the measure of the amount of heat energy, i.e., internal molecular motion, in a body (system).   T is proportional to the average translational kinetic energy of molecules in a gas.

An important concept in thermodynamics is a volume called a “system”, which is an object or group of objects under consideration.   A system is a region in space.   It is separated from is surroundings (environment) by a boundary that defines the system.   The possible exchanges of work, heat, and matter between the system and the surroundings take place across the boundary.   There are 5 main types of systems:

Isolated System – mass (matter) and energy cannot cross the boundary; W = Q = ΔU = 0.
Adiabatic System – heat cannot cross the boundary, matter can; Q = 0 and ΔU = -U.
Diathermic System - heat can cross the boundary, mass cannot;
Closed System – matter cannot cross the boundary; energy can;
Open System – energy and mass can cross the boundary.

A thermodynamic process describes a change in the system.   The 6 most common thermodynamic processes are the following:

An isobaric process occurs at constant pressure;
An isochoric (isometric/isovolumetric) process occurs at constant volume; W = 0 and ΔU = Q.
An isothermal process occurs at constant temperature;
An isentropic process occurs at constant entropy;
An isenthalpic process occurs at a constant enthalpy;
An adiabatic process occurs without loss or gain of heat.

In an isolated system over time, internal pressures, temperatures and masses tend to equalize.   A system in which all equalizing processes have gone practically to completion is a state of thermodynamic equilibrium.   In thermodynamic equilibrium, a system's properties are, by definition, unchanging in time.   Thermodynamic processes which develop so slowly as to allow each intermediate step to be an equilibrium state are called reversible processes.

There are 4 general laws (postulates) of thermodynamic systems:

Zeroth law of thermodynamics states that if 2 systems are in thermal equilibrium with a third, they are also in thermal equilibrium with each other.   Since the internal energy in all these systems are the same, their temperatures will be the same.   This law is used to define temperature.

First law of thermodynamics states the conservation of energy.   The change in the internal energy of a closed (and open) thermodynamic system is equal to the sum of the amount of heat energy added to the system and the work done on the system.   Δ U = Q – W.   In an isolated system, W = Q = ΔU = 0.

Special version of Second law of thermodynamics states that heat flows naturally from a system (object) at higher temperture to a system system (object) of lower temperature.   Because some heat must always leave the system at some temperature, this law implies that no device can completely transform heat added (Q) to work (W).   This is the Kelvin-Planck version of the second law.

General version of Second law of thermodynamics states that entropy of an isolated system never decreases with time.   It increases or remains the same.   It can only remain the same in a reversible (ideal) process.   In an irreversible (real) process, entropy will always increase.   ΔS > 0.   Entropy can be considered a measure of a system's disorder.   Therefore, the second law can be restated as natural processes tend to move to a state of greater disorder.   This implies that for any natural process, some energy must always be so disordered (degraded) and so can do no work.

The First and Second Laws can be combined algebraically into the Combined Law of Thermodynamics:

U – T•ΔS + p•ΔV ≤ 0

Third law of thermodynamics states that absolute (Kelvin) zero temperature is a limit where all molecular motion ceases.   There can be no temperature below this limit.

HISTORY

The remarkable advances in combustion engines, such as the steam engine and hot air engine in the latter half of the 18th century proceeded without any knowledge between heat and work.   Joseph Black, M.D. (1728-99) was a professor of medicine at the U. of Glasgow at a time when James Watt worked there as an instrument maker.   Black was also a chemist who conducted experiments with heated objects in the 1760s to discover specific heat (heat capacity of objects).   He also discovered latent heat of vaporization (boiling) and fusion (freezing). ^{Cummins 28}

At the time of Dr. Black, the caloric theory of heat prevailed.   It was invented in 1783 by the famous chemist, Antoine Lavoisier.   Heat (caloric) was supposed to be a material substance, an elastic fluid, having particles that strongly repel each other.   Caloric occupied the space between atoms and molecules.   Caloric had little or no weight and was squeezed out by friction or flowed out by applying fire to the object.   (The language of this obsolete theory persists today when one speaks of heat "flowing" from one object to another and by a body "soaking up" heat.)   Caloric particles were attracted to matter, the attraction depending on the type of matter.   Caloric could be neither created nor destroyed.   Thus, there was no connection between heat and work. ^{Cummins 29}

The first person who questioned the caloric theory was Benjamin Thompson ("Count Rumford') (1753-1814) who produced famous cannon-boring experiments in 1798 as superintendent of the Munich arsenal.   He showed that no matter how long a dulled boring tool was worked against the sides of a cannon wall, the cannon would keep producing heat, even though hot metal chips were also removed.   Humphry Davy (1778-1829) performed friction experiments with ice and published them in 1799 cast more doubt on the caloric theory.   Generating heat by rubbing objects together, he concluded that heat was probably a vibration of particles. ^{Cummins 29-30}

The Frenchman, Nicholas Leonard Sadi Carnot (1796-1832) published his only book in 1824 that presented the modern theory of heat.   By studying steam engines, Carnot invented a method by which all heat-engines and refrigerating machines can be studied for their efficiency.   He studied intermediate stages of "closed cycle" systems where the working fluid returns to its original state and demonstrated the impossibility of a reversible cycle; that is, all cycles were irreversible.   Carnot said that the maximum efficiency of a heat-engine is determined by the temperature difference between the start and end of an expansion (power) stroke of the engine.   Carnot described the relationship between heat and work by what later became known as the First Law of Thermodynamics.   Although his book did not indicate his abandonment of the caloric theory, his later treatise, Reflexions did so.   In 1842, Joulius Robert Mayer (1814-78) and in 1843, James Prescott Joule (1818-89) independently published papers on the First Law.   Joule also showed how work is converted to heat by his famous paddle wheel experiments. After several refinements, he concluded with his mechanical equivalent of heat in 1850:   778.2 ft-lb = 1 Btu, which was close to its actual value of 772. ^{Cummins 30-34}

In 1851, Rudolph Clausius (1822-88) and William Thomson (Lord Kelvin, 184-1907) discovered the Second Law of Thermodynamics: Heat cannot be itself pass from a colder to a hotter object. ^{Cummins 36}

Thus, the upcoming inventors of gasoline engines would have access to a complete relationship between heat and work.

The Third Law of Thermodynamics was discovered by the German chemist, Walther Nernst (1864-1941), between 1906 and 1912.