LỊCH SỮ HÌNH THÀNH ĐIỀU HÒA KHÔNG KHÍ

Refrigeration Fundamentals Throughout History: Methods Used to
Obtain Colder Temperatures, and Principles Governing Them

A literature Seminar
Presented by
Jesse N. Lawrence

1:00 pm, Tuesday, February 25, 2003
Department of Chemistry
The University of Alabama


There has always been a need or at least a desire to cool some environments below
ambient temperatures. This paper will explore some of the methods used to produce
refrigeration, and to give some history on the subject. The properties of many of the refrigerants
used in evaporative systems will be discussed, such as what characteristics qualify a liquid as a
good candidate for evaporative refrigeration, why the CFCs were successful as refrigerants, and
what new compounds are in place to phase them out. There will also be some discussion on the
cryogenic field, the history of scientist’s attempts to liquefy the components of air and other
gases, as well as some methods used to approach near zero absolute temperature.
Records dating back to around 2000 B.C indicate that people knew of the preserving
effects colder temperatures had on food. Alexander the Great served his soldiers snow-cooled
drinks around 300 BC, and as far back as 755 AD Khalif Madhi provided refrigerated transport
across the desert to Mecca using snow as refrigerant.1 Some of the earliest methods of producing
cold made use of naturally occurring ice and freezing mixtures such as salt and snow.2 The fact
that sodium nitrate lowers the temperature of water upon dissolving was known in the 14th
century A.D. Naturally occurring ice was either shipped from colder climates or collected in the
winter and stored in cold houses, the earliest mention of them dating to about 1000 B.C. in some
ancient Chinese poems called the Shi Ching.3 These houses were made of various insulating
materials, such as straw dirt, and even manure. In the 18th century, this ice was typically only
available to the rich and powerful. In 1806 a man by the name of Frederick Tudor began an ice

business by cutting it from the Hudson River and other nearby bodies of water, and selling it at a
price that made it available to a greater number of people.4 He eventually shipped it to various
localities throughout the world, his first venture being a shipment of 130 tons of ice to the port of
St. Pierre, Martinique. Ice was unknown there, and there were no facilities to store it. The
attempt would have been disastrous, but Tudor worked with the proprietor of a local eating house,
and they concocted and sold ice creams, which had previously been unknown in the West Indies.
As it was he lost $3500, but construction of an icehouse at St. Pierre and the use of sawdust while
shipping made the business profitable.3 He also had active shipments to the southern states until
they were halted by the civil war. Others joined the ice trade, and many entrepreneurs began to
ship ice from other places. In 1854 it was reported that 156,000 tons of ice were shipped from
Boston. Ice houses in the U.S. typically made use of sawdust as insulation, having walls 1 meter
thick. This ice trade continued even after the development of artificial ice, generally fueled by
claims of natural ice’s superior qualities to a man-made product. The trade finally gave way in
1930.
The main method used to provide refrigeration relies on the evaporation of some liquid or
other, and the bulk of this paper will be devoted to its principles. As of 1755, the fact that ether
chilled the skin was already known. At that time, a Professor of Chemistry, William Cullen,
demonstrated the formation of ice in water that was in thermal contact with a body of ether. By
reducing the pressure above the ether, he caused it to boil, and lower in temperature to the point
that it formed ice.
This was the first half of the refrigeration cycle, and it still remained to find a way to recirculate the ether, for to simply refill the liquid ether vessel when the previous had boiled away
would not be cost effective. Information on methods to liquefy vapors through compression was
gathered in the last half of the 18th century. In 1780, two men by the names of J.F.Clouet and
G.Monge liquefied sulfur dioxide. Ammonia was liquefied in 1787 by van Marum and van
Troostwijk. The idea of putting condensation and evaporation techniques together to create a
cyclic system seems to have been first suggested by Oliver Evans of Philadelphia, but the first
cyclic refrigeration machine was made by Jacob Perkins. Its description can be found in a patent
specification of 18344. There were earlier patents given for refrigeration machines, the first
appearing in 1790, but Perkins’ seems to be the first to have been built and put to any use.3 It was
intended to be used with any volatile fluid as refrigerant, especially ether. It consisted of four

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main components: an evaporator, compressor, condenser, and expansion valve. The compressor
pumps vapor into the condenser, which is essentially a heat exchanger. The vapor heats upon
compression. This heated, compressed vapor is cooled by outside air or water, causing it to
condense. The liquid then flows through the expansion valve, which is essentially a portion of the
piping where flow is restricted by some means or other to create a pressure differential across it.
At the lower pressure side of the valve, the liquid is then in a superheated state, and some of it
boils to cool the rest to the boiling point of the lower pressure, which then flows into the
evaporator. The evaporator is also a heat exchanger, and heat from the substance meant to be
cooled causes the refrigerant to evaporate. This evaporated vapor is then pulled into the
compressor, and the cycle begins again. Obviously if the ambient temperature at the condenser is
greater than the temperature of the compressed vapor, or if the desired temperature of the
substance which needs to be cooled is lower than the boiling temperature of the liquid in the
evaporator, the system will be useless. Perkins’ invention generated as much excitement as
Crystal Pepsi. It made no appearance in the literature of the time, and was only casually
referred to by Bramwell 50 years later.
The person most responsible for putting the refrigeration machine into use was James
Harrison of Scotland. He received little technical training save for some chemistry classes he
took while he studied to be a printer at college. He noticed the chilling effect of ether when using
it to wash type, and around 1850 he invented a hand operated machine that would produce ice. In
1856 and 57 he applied for British patents, and built better machines in England. These machines
were shipped to various places for applications such as the making of ice and the crystallization
of paraffin wax from shale oil. They were regularly manufactured until the advent of ammonia
and carbon dioxide systems, but remained popular in places such as India, where the water used
to cool the condenser was generally at a higher temperature than other localities. Ether’s normal
boiling point is 34.5°C, thus when used to make ice, there must be a pressure less than
atmospheric on the low pressure side of the system. This can be dangerous because any leaks will

bring in oxygen, thus creating a potentially explosive environment. Dimethyl ether, with a
normal boiling point of –23.6°C was introduced by Carles Tellier in 1864. Sulfur dioxide, with a
boiling point of -10°C, was introduced in 1874. Although dimethyl ether was never in general
use, sulfur dioxide was used extensively for about 60 years. Carl Von Linde was the first to
introduce ammonia as a refrigerant in the 1870’s. It was advantageous in that having an
atmospheric bp of -33.3°C, it could provide much colder temperatures than previously available,
but it also exhibited pressures of ten atmospheres or more in the condenser, which required
sturdier construction. Carbon dioxide systems were developed for use in 1886, and because of its
nontoxic nature, it was used extensively on ships until 1955, when it was then replaced by the
CFCs.4
Another system which relies upon the evaporation of a liquid is the absorption system. It
essentially operates in the same manner as a compression system, but instead of going through a
compressor the refrigerant is absorbed in some substance or other, and then released to a
condenser by heating. Thus the absorption chamber, also called the generator, takes the place of
the compressor. The principles of such a system were demonstrated by Sir John Leslie in 1810.
He placed two vessels in a bell jar, one of water and one of sulfuric acid, then evacuated the
vessel somewhat with a pump. Over a period of time, ice was seen to form on the top of the water
vessel. The pump was not absolutely necessary in this case; it simply sped along the process.
Water vapor is readily absorbed by sulfuric acid, and by so doing, the sulfuric vessel continued to
‘pump’ vapor from the water vessel, lowering its temperature until it froze. Early systems
developed around 1878 using this principle weren’t cyclic; they were designed to have the
sulfuric acid removed from the system in order to re-concentrate it by boiling off the absorbed
water. In this system water acted as the refrigerant. A better absorption system was developed in
1859, by Ferdinand Carré employing ammonia as the refrigerant and water as the absorbent. In

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order to make such a system cyclic, a system of valves and containers are built wherein a vessel
containing a strong solution of the refrigerant, in this case ammonia, is heated to produce

pressurized ammonia vapor which goes to a condenser to become liquid. The rest of the system is
the same as the compression system, the vapor being re-absorbed in a weak ammonia solution.
The system of valves and vessels insure that the vapor formed on heating can only go in a
‘forward’ direction. Other systems were developed using an aqueous solution of lithium Bromide
as absorbent and water as refrigerant.4
So what characteristics make a ‘good’ evaporative refrigerant? Since the system operates
on the energy required to evaporate the liquid, it is good to have a high enthalpy of evaporation.
The refrigerant must also exhibit boiling points at pressures that are appropriate to the materials
and seals used to construct the system. It is also more economic if the operating pressure in the
condenser be well below the critical pressure of the gas, for the closer to the critical pressure one
approaches, the more power that is required to compress it. Viscosity can also be a factor.
Typically a lower viscosity will allow smaller pipes, valves, and compressor passages.5 It is also
desirable that it be non-toxic, or at least not terribly so. Of the refrigerants mentioned, sulfur
dioxide is the most toxic. An exposure of 0.5-1% of air for five minutes can be fatal. Ammonia
is next, requiring 30 minutes to produce death at the same concentration. Carbon dioxide,
methane, ethane and propane are all of very low toxicity, as are the CFC’s. It is also best to use a
refrigerant that is not flammable. Any hydrocarbon such as methane suffers from this
requirement, but they otherwise exhibit all of the properties necessary to perform well as a
refrigerant. These find use in the petroleum industry as they are readily available, and there are
plenty of other things to set on fire anyway. Ammonia is somewhat flammable, and explosions
have occurred. If enough hydrogen atoms in an alkane are replaced with halogens, the refrigerant
will be non-flammable. Methyl chloride, for instance, is difficult to ignite, but it has also caused
explosions. The detection of leaks can be aided by a strong smell. Despite this fact, ammonia is
sometimes not used for this very reason, the predominating belief being that a strong smell from a
relatively small and harmless leak may induce panic. Sulfur dioxide is the only other refrigerant
that has a strong odor. Carbon dioxide can be tasted at certain concentrations, and some people
claim they can smell the CFCs, but the number is few. If desired, chemicals such as acrolein
have been added to indicate a leak by smell. It is also necessary that the refrigerant not react with
the materials used in constructing the refrigerating system. When using ammonia, copper must
be avoided unless water is strictly kept from the system. Methyl chloride is reactive with

aluminum. The CFC’s are generally inert, but the presence of water in the system can form
hydrochloric acid when either methyl chloride or R-12 is used. Also, the refrigerant should be
stable at the system’s operating conditions. All the refrigerants mentioned are rather stable,
especially carbon dioxide, sulfur dioxide, and the alkanes. Ammonia decomposes slightly under
certain circumstances. It is also necessary that the refrigerant behave appropriately with the
lubricating oil in the system. The behavior of the lubricating oil with the refrigerant must be
considered when designing the system. For instance, one must be certain the crankcase oil does
not lose too much viscosity if dissolved in the refrigerant, or carried away from the crankcase
entirely. The basic nature of ammonia tends to saponify some lubricating oils used in
refrigerating systems. Cost is also a factor. Despite its toxic nature and tendency to form sulfuric
acid in the presence of water, sulfur dioxide was used extensively because it was inexpensive.4
Obviously, it must also never be solid at any of the conditions of operation, or it will not cycle,
and must phase change after leaving the expansion valve. The freezing point of ammonia is -78°
C, far enough away from its normal bp of –33.3°C. Carbon dioxide can be troublesome in that if
too much pressure is lost through leakage, it will form a solid that can block the expansion valve.3
All of the earlier evaporative refrigerants had one or more problems associated with
them. Despite its status as the first refrigerant fluid to be used commercially, ether was not
popular for long due to its flammability. Some early refrigerants that stayed in use for a longer
time, some of which are still in use today, were methyl chloride, sulfur dioxide, carbon dioxide,

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ammonia, and some hydrocarbons. Other early refrigerants that never gained much use were
acetone and alcohol, also highly flammable.6 Ammonia has many characteristics that make it
useful as a refrigerant. Despite its drawbacks mentioned previously, with a boiling point of
–33.3° C it exhibits good working pressure, typically not exceeding 200 psi, and has a large latent
heat of evaporation. It is still widely used. Carbon dioxide is a good refrigerant in that it is not
toxic or reactive to metals. However, the pressures necessary to use it are quite high; generally in
the area of 1200 psi or more, and this calls for stronger equipment and better seals. Its latent heat

of vaporization is not as good as that of ammonia. Because of the higher pressures it requires, a
smaller volume of vapor is formed, and thus a smaller capacity compressor can be used with it.
Sulfur dioxide operates at pressures much lower than ammonia, typically about 60 psi, and thus a
very large volume of vapor is formed, requiring larger compressors. Its latent heat of
vaporization is a bit greater than that of CO2. Its lower operating pressure is also disadvantageous
because the low pressure side may be below atmospheric pressure, which will allow in air and
water vapor at any leak points. Water can then freeze and plug up the line, or combine with any
SO3 present to form sulfuric acid, which will then corrode the pipes.2,3 Methyl Chloride, CH3Cl,
was introduced in the U.S. in the 1920’s, and was popular for a time as a replacement for the
CFCs during WW II, during which they were not as available. It had good operating qualities,
but was somewhat flammable and toxic. Nevertheless, it had a place even in residential
applications at least until the late 1950’s. Water has been used at times as a refrigerant, but is
limited in its cooling ability by its freezing point of 0° C.3
In light of some of the difficulties exhibited by the earlier refrigerant fluids, it was
necessary to find fluids that had better qualities. This was done by Midgely, Henne, and McNary
of the General Motors Corporation.6 In 1928 Thomas Midgeley was given an assignment to find
a refrigerant that had a low toxicity, low flammability, good stability, and an atmospheric boiling
point between –40 and 32° F. It took him and his associates three days. They synthesized the first
fluorocarbon refrigerant, using guinea pigs to demonstrate its low toxicity. They synthesized all
15 combinations of one carbon with various combinations of chlorine, fluorine, and hydrogen.
They tested their properties, and finally chose dichlorodifluoromethane as having the most
desirable characteristics. The new refrigerant was announced at the American Chemical Society’s
meeting of April, 1930, at which all other sections of the society adjourned their meetings to
attend Midgely’s presentation. At the end of his talk, Midgley demonstrated some of the good
properties of R-12 by inhaling a lungful of the gas, then using it to extinguish a candle.6 In
general the properties of the HCFCs will tend toward increased flammability when the number of
hydrogens are increased, and increased toxicicity when the number of chlorine atoms are
increased.7
There is a numerical nomenclature system that names all evaporative refrigerants as
follows:

1. For the chlorofluorocarbons, the last digit on the right is the number of fluorine atoms
contained in the molecule.
2. The next to last digit is the number of hydrogen atoms in the molecule plus one.
3. The first digit, if the molecule contains more than 1 carbon atom, is the number of
carbon atoms minus 1. If the molecule only has one carbon atom, the digit is
generally omitted, though some texts may include a zero. Thus, R-12 has 2 fluorine
atoms, 0 hydrogen atoms, and 1 carbon atom. Any open valence spots are filled by
chlorine atoms, giving R-12 a molecular formula of CF2Cl2.
4. If some or all of the possible chlorine atoms are replaced with bromine atoms, an
extra digit preceded by a B is added to the end, representing the number of bromine
atoms, thus R-13 is CClF3, and R-13B1 is CBrF3.6
5. Isomers are designated by a lower case letter following the digits. The most
symmetric isomer is named as above, while the next most symmetric isomer will
have an ‘a’ following the digits. If a still less symmetric isomer is possible, the digits

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will be followed by a ‘b’, and so on. Thus R-134 is CHF2CHF2, whereas R-134a is
CH2FCF3
6. Cyclic compounds are represented by an upper case ‘C’ placed before the digits.4
7. Azeotropic mixtures of refrigerants are assigned to the 500 series according to the
order in which they became commercially available. (R-500, R-501, etc.)
8. Zeotropic mixtures are assigned to the 400 series.
9. Hydrocarbons are assigned according to rules 1-4, except for butane and isobutane,
which are in the 600 series, along with other organic refrigerants.
10. Inorganic refrigerants are assigned to the 700 series, using the molecular weight in
the numbering. Thus ammonia is R-717.6
11. Unsaturated compounds are named according to rules 1-6, but the resulting number
has an extra 1 placed in front of it. Thus R-1113 is FClC=CF2.8

The initial R- (for refrigerant) that comes before the number portion of the name may be replaced
by Freon, CFC, HCFC, and some others.
Refrigerant mixtures may be used for different reasons. For instance, a mixture of R-12
and R-114 is used in systems where the ambient temperature may be excessively high, such as
refineries. This calls for a stable refrigerant that does not exhibit very high pressure at ambient
temperatures. Both of these refrigerants are quite stable, and the presence of R-114 (bp 3.77°C)
will lower the higher pressures associated with R-12(bp -29.79°C). Also, the R-12 creates a
greater cooling capacity than R-114 exhibits alone. R-502 was developed to take the place of R22 in systems that had extraordinarily long return lines to the compressor. Under these conditions,
the vapor can gain more heat than normal on the return, and will then cause unacceptably high
temperatures upon compression, resulting in system failure. The switch to R-502 reduced the
temperature at the compressor by as much as 100°F. It had three characteristics that enabled this.
First, its vapor had a larger heat capacity than that of R-22’s. Its heat of vaporization is also less
than that of R-22, causing a need to use a higher vapor flow rate to provide the same amount of
cooling, which reduces the time the vapor has to pick up additional heat on its return, as well as
the amount of heat given off upon condensation. Lastly, its vapor also exhibits a lower rise in
temperature due to compression than R-22. That is to say, its Joule-Thompson coefficient is not
as large. R-502 is simply R-22 with R-115 added to it, and one could use R-115 by itself, but this
would result in a significant loss in cooling capacity.6 Non-azeotropic blends will exhibit a
change in boiling temperature as the composition of the liquid changes. This phenomon is known
as glide in the refrigeration industry. It can enhance a system’s performance. If a countercurrent
heat exchanger is used in the evaporator, the first refrigerant the medium to be cooled will contact
will be the warmer than the last bit of refrigerant it contacts. As the medium to be cooled travels
along the exchanger, it will be colder, and will experience even colder refrigerant, as the
refrigerant entering the heat exchanger will have more of its colder boiling substituent present. In
this manner the medium receives maximum refrigeration from the refrigerant. Counter flow heat
exchangers are difficult to build, however, and demand a higher cost.7
The CFCs, of course, have also exhibited some undesirable properties in their interactions
when in the upper atmosphere. Any molecules having chlorine or bromine undergo chemistry
that depletes the ozone present, allowing greater amounts of harmful UV radiation to reach the
surface of the planet. The process begins by having a chlorine or bromine radical separated from

its parent molecule under the influence of UV radiation. This radical can then combine with
ozone, yielding oxygen and XO, X=Br or Cl. Any free oxygen atoms that would have created
ozone may instead come upon the XO molecule, forming oxygen and the starting radical all over
again. The most practical solution to remedy this has been the proposals to lessen the use of
CFC’s and eventually halt it altogether. Some other ideas have been proposed, one being by a
professor at UCLA. He suggested providing extra electrons for the radicals by floating a large
metal sheet in the stratosphere with balloons. The photoelectric effect would provide electrons
for the radicals, rendering them inert. His suggestions have not been tried. The most popular

5


replacements so far for the CFCs have been hydrofluorocarbons. Temporary replacements have
been molecules with a lower number of chlorine and bromine atoms. Thus, R-134a can be a
permanent replacement for R-12, as it has no bromine or chlorine. Other distant alternatives may
lie in the family of fluoroiodocarbons.7
In 1834 Jean Peltier noticed that if current was passed through a junction of dissimilar
metals, the junction would either become cooler or hotter,depending on the direction of the
current across the junction. The thermoelectric effect is best between alloys, most pure
substances providing weak results.4
Another novel method to provide colder temperatures appeared in France in1931.
George Ranque observed the separation of a stream of compressed air into two streams, one hot
and the other cold. The device he used has no moving parts, and has come to be known as the
Ranque tube, vortex tube, or werbelrhor. It works by introducing a stream of compressed air
perpendicularly and tangentially into a pipe, where it then forms a vortex. One end of the pipe
has a plug in it with a hole drilled just in the center of the plug so that only air in the center of the
pipe can escape. A cone is inserted point first into the other end of the pipe so only air flowing on
the outer portion of the pipe can escape at this end. Depending on the amount in which the cone
is inserted, a stream of air significantly warmer than that introduced by the compressed pipe will
issue from this end and a colder stream will issue from the other end. If the cone is completely

inserted to stop any flow at that end, the temperature at the other will be that of the compressed
air. As the cone is removed, the flow rate reduces, as does the temperature.4 While no
quantitative theory yet exists to explain this effect, the dimensions have been optimized to
provide a cold air stream 100° F colder than the compressed air stream, and they are still in use
today.3 They are convenient in places such as machine shops where compressed air is readily
available.
Still another method which can be used to provide cooling is the expansion of a gas. It
was known as early as the 18th century that a reduction in pressure would provide a drop in the
temperature of air. Richard Trevithik was the first to propose using this phenomenon in
refrigeration applications in 1828. The first system to use this principle was made by John Gorrie
in Florida in 1844. Many improvements were made upon the design, the process being rather
simple. First, air is drawn from a room or vessel that one desires to cool. It enters a compressor,
and increases in pressure and temperature. This hot compressed air is then cooled by outside air
or water, as the case may be, after which it is allowed to expand again, preferably against a piston
or turbine which provides work to the compressor, upon which it cools further, then re-enters the
vessel to be cooled. A technique similar to this is used to cool jet aircraft.
For the physicist, cryogenic may mean temperatures close to zero K. For our purposes, it
will refer to temperatures below about 120 K, which is a bit above the boiling point of natural
gas. Much of the earliest work at cryogenic temperatures was simply an attempt to liquefy and/or
separate various gases. The knowledge that had been developed to reach colder temperatures
with vapor compression systems enabled scientists to cool compressed gases that had formerly
been uncondensable at ambient temperatures. As of 1854 Michael Faraday had liquefied many of
the gases known at that time. He did so by compressing them and cooling to 163 K with solid
carbon dioxide and ether. All attempts to liquefy gases such as hydrogen, nitrogen, oxygen, and
some others had failed. In Vienna, Johannes Natterer showed that oxygen, nitrogen, and
hydrogen were still gaseous at 3600 atmospheres and 195 K. It was considered for a time that
these gases were uncondensable until Thomas Andrews discovered the significance of the critical
temperature of carbon dioxide during his work from 1861-1869. Efforts were begun again to
liquefy these gases, and in December 1877 Louis Cailletet in France and Raoul Pictet in Geneva
formed liquid oxygen at almost the same time. Cailletet compressed oxygen to 300 atmospheres

and cooled it to 244 K with sulfur dioxide. Upon expansion, a mist of liquid appeared. Neither
man made enough to see in a container. This was done in 1883 by Sigmund von Wroblewski and
Karek Olszewski in Poland. They compressed the oxygen and cooled it to 137 K with liquid

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ethylene boiling under vacuum. In 1885 Wroblewski liquified air, finding that liqiud nitrogen
and oxygen are completely miscible. By the early 1890’s liquid air, oxygen and nitrogen were
produced in sufficient quantities to experiment with, and there began to be a commercial demand
for these liquids, especially oxygen. The methods used to liquefy these gases employed cascade
refrigeration systems where one refrigerant was used to cool the condenser of a lower boiling
refrigerant, and so on, to reach the temperatures necessary to condense these gases. At most these
methods produced about 14 liters per hour. Various scientists worked independently to produce
greater amounts of these liquids, among them being Carl Linde in Germany, and William
Hampson in England. They all developed similar methods in 1895. All were based on the
reduction of temperature experienced by a gas when throttled from a high pressure to a lower one.
This was known as the Joule-Thompson effect. The reduction seen in air is small, only about _ K
per atmosphere. Linde and Hampson applied for patents within a few days of each other, but
Linde was the first to produce enough liquid air to sell. His system essentially worked by first
compressing the air and cooling it by use of a low temperature refrigerant. The gas traveled
through an expansion valve, causing it to lower more in temperature, and entered a holding vessel
for any liquid that formed on expansion. The expanded gas then ran back along the pressurized
gas just before the expansion valve, providing additional cooling, and was then sent back to the
pump to be re-pressurized, and the cycle would begin again.
By 1898 all of the known gases had been successfully liquefied with the exception of
hydrogen. Pictet had claimed to have produced a mist of hydrogen, but this was not generally
believed. The struggle to reach lower temperatures continued, and help in this regard was made
available in 1892 when James Dewar invented the vacuum flask. Calculations of the critical state
of hydrogen estimated its critical temperature at 30K. The coldest temperature then possible was

achieved by boiling liquid oxygen at reduced temperatures. It was also known that Hydrogen
could not be liquefied by the Linde-Hampson process, for its behavior was the opposite of other
gases, heating upon expansion instead of cooling. Olszewski found that when cooled to a
temperature below 190 K, compressed hydrogen would then reverse its behavior, and would cool
upon expansion. This is known as the inversion temperature for hydrogen. Dewar succeeded in
liquefying hydrogen in 1898 by cooling the compressed gas to below 190 K with liquefied air,
then condensing it with the Linde-Hampson process.
New gases were being found in the atmosphere, some being easier to liquefy than others.
Argon was discovered in 1894, and had a boiling point between that of oxygen and nitrogen.
Helium was extracted from the mineral cleveite in 1895, and the rest of the noble gases were
found by separation from liquid air. Krypton and xenon could be condensed by liquid air, and
neon by liquid hydrogen, but all attempts to liquefy helium failed. It has a critical temperature of
5.2 K. The freezing point of hydrogen is 13.8 K, and it could thus not be used to liquefy helium
by boiling it under reduced pressure. The inversion temperature of Helium is 40 K, however, and
it was eventually liquefied through the Linde-Hampson process by Kamerlingh Onnes in 1908.
Evaporation of helium under reduced pressure produced a temperature of 0.83 K in
1922, and this seemed to be the limit unless a new method for reducing temperature was found.
Other methods were employed to reach still colder temperatures, such as isothermal
magnetisation of a material, gadolinium sulphate for instance, followed by demagnetization. This
resulted in a temperature of 0.25 K in 1933 and 0.01 K later. Other demagnetisation methods
resulted in temperatures on the order 0f 10-3 K. Even colder temperatures have been reached by
the use of lasers on gas molecules, using the momentum of the photons to slow the molecule’s
translational speed Zero K is unnattainable as a result of the third law of thermodynamics.4
What causes most gases to cool upon expansion, and why do the others, namely hydrogen
and helium, heat upon expansion, until they are cooled to their inversion temperature? What is the
significance of the inversion temperature? In 1843 James Joule investigated the internal energy
change associated with the expansion of a gas by immersing two chambers attached by a valve in
a water bath with adiabatic walls. One chamber contained a pressurized gas, and the other was

7

evacuated. A thermometer was inserted in the water to note a temperature change. Since the gas
expands into a vacuum, no work, w, is done by the gas as it expands. The walls are adiabatic, so
the heat, q, is also zero. Hence the change in internal energy _U = q + w = 0. This experiment
would determine the change in temperature with respect to volume at constant U, _T/_V.
Typically, the change in temperature with change in volume is very small. That coupled with the
very small heat capacity of the gas compared to the large heat capacity of the water gave Joule a
_T of zero for every reading. He concluded that _T/_V was zero regardless of the change in
volume, or more specifically, that (∂T/∂V)U was zero. The value (∂T/∂V)U is known as the Joule
coefficient, commonly given the symbol µJ. In 1924 Keyes and Sears used an improved setup
and repeated the experiment, this time noting a small but noticeable temperature change. In 1853
Joule and William Thompson performed the Joule-Thompson experiment, which was similar to
his previous experiment. In this experiment, a gas was pushed from a higher pressure cylinder to
a lower pressure cylinder through a rigid porous plug. At the end of each cylinder was a piston
providing the different pressures. A thermometer was inserted in each cylinder, and the system
had adiabatic walls. The gas would be brought to some initial Pi and Ti, and would then be
allowed to flow through the plug to the lower pressure Pf, and the final temperature Tf would be
read. As the walls are adiabatic, the heat exchanged, q, is zero. There are two points where work
is performed on the system. At the high pressure piston the work is performed on the system, and
if all of the gas is throttled through, this amount can be calculated as PhVh Where Vh is the
volume of the gas when all of it is on the high pressure side. The system itself does work on the
low pressure piston, which can be calculated as PlVl. Thus the total work w on the system can be
calculated as PhVh - PlVl. For this experiment _U is equal to w instead of zero. From the identity
U + PV = H, it can be derived that Hh = Hl. This is a constant enthalpy process. Taking various
measurements of _T/_P by starting with the same initial temperature and pressure and varying the
final pressure will yield an isenthalpic curve on a pressure-temperature diagram, the slope of a
tangential line to which will yield (∂T/∂P)H, the Joule-Thompson coefficient, known as µJ-T.
Values for gases can range from 3 to –0.1 K/atm, depending on the gas and its initial temperature
and pressure. This value is critical to the liquefaction of gases in the Linde-Hampson process. If

the value is negative, the gas heats upon expansion, and liquefaction is not possible. The reason
why some gases cool upon expansion, while others will heat depends on the intermolecular forces
present. In any gas, there are attractive and repulsive forces present between the molecules or
atoms. Keep in mind that when a system goes to a more stable state, it gives off energy; the
process is exothermic, whereas if a system goes to a less stable state, energy must come from
somewhere for the process to occur. In the process of gas expansion, this energy is in the form of
heat. When cooling is observed upon expansion, the attractive forces predominate. A larger
average separation of molecules leads to a less stable state, and energy in the form of heat must
be taken from the surroundings. When heating is observed upon expansion, the repulsive forces
predominate, and a greater separation of molecules results in a more stable state, releasing energy
in the form of heat. What about the fact that a gas that heats on expansion will then cool on
expansion if its temperature is brought low enough? When the temperature is lowered, the
average kinetic energy of the gas molecules is also lowered. Collisions between molecules are
not as forceful, and where the collisions at a higher temperature were high energy enough to
make repulsion the predominant force, at lower temperature the attractive forces become more
significant, and expansion then yields a cooling effect. Not only the initial temperature, but the
initial pressure chosen will have a bearing on whether or not the Joule or Joule-Thompson
coefficient will be positive or negative. If you start at a pressure on the isenthalpic curve where
heating is observed instead of cooling, you will first observe heating until you reach the inversion
point, Where (∂T/∂P)H = 0. After this point on the isenthalpic expansion will yield a cooling
effect instead of heating.9 There are many inversion pressures and temperatures possible,
depending on the initial conditions of the gas, and a plot of a family of isenthalpic curves will

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yield a locus of inversion points that will have a maximum inversion temperature at P = 0. Above
this temperature it is not possible to observe cooling upon expansion regardless of what pressure
is used. The following table lists some maximum inversion temperatures:
Helium

40
Hydrogen
195
Oxygen
49.7
Nitrogen
621
Argon
723
Carbon Dioxide
15034
So what is the significance of the inversion point? It can be derived from the ideal gas law that
∂U/∂V = ∂H/∂P = 0 for an ideal gas. It can also be shown that ∂U/∂V = -CvµJ and
∂H/∂P = -CPµJ-T. Cv and CP are the heat capacities at constant volume and pressure, which will
always be positive. Thus when ∂U/∂V = ∂H/∂P = 0, µJ = µJ-T = 0, which is the inversion point.
At this point the attractive and repulsive forces are balanced, in a word, and the gas behaves very
much like an ideal gas, even though P is not close to zero.
How does higher pressure affect the temperature change upon expansion? In the same way that
higher temperatures correspond to higher energy collisions, making the repulsive forces more
significant and thus increasing U of the system, higher pressures can also squeeze the molecules
together to a point where the repulsive forces become significant, also increasing U of the system.
Thus upon expansion this higher U goes into heating the system. Of course these same attractive
and repulsive forces that change the temperature as the molecules in a gas separate upon
expansion also determine the cooling effect observed when a liquid boils, The only difference
being the fact that since the molecules have strong enough attractive forces to form a liquid, it is
guaranteed that they are the predominate force, and a cooling effect will always be observed upon
boiling of a liquid, that is, when the liquid molecules undergo an expansion.9
References.
1. Koelet, P. C. Industrial Refrigeration Principles, Design and Application.
Marcell dekker, Inc., New York, 1992.

2. Daniels, G. W. Refrigeration in the Chemical Industry. D. Van Nostrand Company,
New York, 1926.
3. Jordan, R. C. and Priester, G. B. Refrigeration and Air Conditioning. Prentice-Hall
Inc., Englewood Cliffs, NJ. 1956.
4. Gosney, W. B. Principles of Refrigeration. Cambridge University Press, New York,
1982.
5. King, G. R. Modern Refrigeration Practice. McGraw-Hill Book Company, New
York, 1971.
6. Downing, R. C. Fluorocarbon Refrigerants Handbook. Prentice-Hall Inc., Englewood
Cliffs, NJ. 1988.
7. Wylie, D. and Davenport, J. W. New Refrigerants for Air conditioning and
Refrigeration Systems. The Fairmont Press, Inc., Lilburn, GA. 1996.
8. Meacock, H. M. Refrigeration Processes. Pergamon Press, New York, 1979.
9. Levine, I. N. Physical Chemistry, 4th Ed. Mcgraw-Hill Inc., New York, 1995.

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