Gas is one of the four fundamental states of matter (the others being
solid, liquid, and plasma). A pure gas may be made up of individual
atoms (e.g. a noble gas like neon), elemental molecules made from one
type of atom (e.g. oxygen), or compound molecules made from a variety
of atoms (e.g. carbon dioxide). A gas mixture would contain a variety
of pure gases much like the air. What distinguishes a gas from liquids
and solids is the vast separation of the individual gas particles.
This separation usually makes a colorless gas invisible to the human
observer. The interaction of gas particles in the presence of electric
and gravitational fields are considered negligible as indicated by the
constant velocity vectors in the image. One type of commonly known gas
The gaseous state of matter is found between the liquid and plasma
states, the latter of which provides the upper temperature boundary
for gases. Bounding the lower end of the temperature scale lie
degenerative quantum gases which are gaining increasing
attention. High-density atomic gases super cooled to incredibly low
temperatures are classified by their statistical behavior as either a
Bose gas or a Fermi gas. For a comprehensive listing of these exotic
states of matter see list of states of matter.
1 Elemental gases
3 Physical characteristics
4.3 Specific volume
5.1 Kinetic theory
5.2 Brownian motion
5.3 Intermolecular forces
6 Simplified models
6.1 Ideal and perfect gas models
6.2 Real gas
7 Historical synthesis
7.1 Boyle's law
7.2 Charles's law
7.3 Gay-Lussac's Law
7.4 Avogadro's law
7.5 Dalton's law
8.2 Reynolds number
8.5 Boundary layer
8.6 Maximum entropy principle
9 See also
12 Further reading
The only chemical elements that are stable diatomic homonuclear
molecules at STP are hydrogen (H2), nitrogen (N2), oxygen (O2), and
two halogens: fluorine (F2) and chlorine (Cl2). When grouped together
with the monatomic noble gases – helium (He), neon (Ne), argon (Ar),
krypton (Kr), xenon (Xe), and radon (Rn) – these gases are called
This article may require cleanup to meet's quality
standards. The specific problem is: "pronounced like ch in "loch"" --
and how is that ch pronounced (IPA?)? Sources give "[lɒk]" and
"[lɒx]" as IPA for loch. Is it [x]? Please help improve this article
if you can. (March 2018) (Learn how and when to remove this template
The word gas was first used by the early 17th-century Flemish chemist
J.B. van Helmont. He identified carbon dioxide, the first known gas
other than air. Van Helmont's word appears to have been simply a
phonetic transcription of the
Ancient Greek word χάος Chaos –
the g in Dutch being pronounced like ch in "loch" – in which case
Van Helmont was simply following the established alchemical usage
first attested in the works of Paracelsus. According to Paracelsus's
terminology, chaos meant something like "ultra-rarefied water".
An alternative story is that Van Helmont's word is corrupted from
gahst (or geist), signifying a ghost or spirit. This was because
certain gases suggested a supernatural origin, such as from their
ability to cause death, extinguish flames, and to occur in "mines,
bottom of wells, churchyards and other lonely places".
Drifting smoke particles provide clues to the movement of the
Because most gases are difficult to observe directly, they are
described through the use of four physical properties or macroscopic
characteristics: pressure, volume, number of particles (chemists group
them by moles) and temperature. These four characteristics were
repeatedly observed by scientists such as Robert Boyle, Jacques
Charles, John Dalton,
Joseph Gay-Lussac and
Amedeo Avogadro for a
variety of gases in various settings. Their detailed studies
ultimately led to a mathematical relationship among these properties
expressed by the ideal gas law (see simplified models section below).
Gas particles are widely separated from one another, and consequently,
have weaker intermolecular bonds than liquids or solids. These
intermolecular forces result from electrostatic interactions between
gas particles. Like-charged areas of different gas particles repel,
while oppositely charged regions of different gas particles attract
one another; gases that contain permanently charged ions are known as
plasmas. Gaseous compounds with polar covalent bonds contain permanent
charge imbalances and so experience relatively strong intermolecular
forces, although the molecule while the compound's net charge remains
neutral. Transient, randomly induced charges exist across non-polar
covalent bonds of molecules and electrostatic interactions caused by
them are referred to as Van der Waals forces. The interaction of these
intermolecular forces varies within a substance which determines many
of the physical properties unique to each gas. A comparison of
boiling points for compounds formed by ionic and covalent bonds leads
us to this conclusion. The drifting smoke particles in the image
provides some insight into low-pressure gas behavior.
Compared to the other states of matter, gases have low density and
Pressure and temperature influence the particles within a
certain volume. This variation in particle separation and speed is
referred to as compressibility. This particle separation and size
influences optical properties of gases as can be found in the
following list of refractive indices. Finally, gas particles spread
apart or diffuse in order to homogeneously distribute themselves
throughout any container.
Shuttle imagery of re-entry phase
When observing a gas, it is typical to specify a frame of reference or
length scale. A larger length scale corresponds to a macroscopic or
global point of view of the gas. This region (referred to as a volume)
must be sufficient in size to contain a large sampling of gas
particles. The resulting statistical analysis of this sample size
produces the "average" behavior (i.e. velocity, temperature or
pressure) of all the gas particles within the region. In contrast, a
smaller length scale corresponds to a microscopic or particle point of
Macroscopically, the gas characteristics measured are either in terms
of the gas particles themselves (velocity, pressure, or temperature)
or their surroundings (volume). For example,
Robert Boyle studied
pneumatic chemistry for a small portion of his career. One of his
experiments related the macroscopic properties of pressure and volume
of a gas. His experiment used a J-tube manometer which looks like a
test tube in the shape of the letter J. Boyle trapped an inert gas in
the closed end of the test tube with a column of mercury, thereby
making the number of particles and the temperature constant. He
observed that when the pressure was increased in the gas, by adding
more mercury to the column, the trapped gas' volume decreased (this is
known as an inverse relationship). Furthermore, when Boyle multiplied
the pressure and volume of each observation, the product was constant.
This relationship held for every gas that Boyle observed leading to
the law, (PV=k), named to honor his work in this field.
There are many mathematical tools available for analyzing gas
properties. As gases are subjected to extreme conditions, these tools
become more complex, from the
Euler equations for inviscid flow to the
Navier–Stokes equations that fully account for viscous effects.
These equations are adapted to the conditions of the gas system in
question. Boyle's lab equipment allowed the use of algebra to obtain
his analytical results. His results were possible because he was
studying gases in relatively low pressure situations where they
behaved in an "ideal" manner. These ideal relationships apply to
safety calculations for a variety of flight conditions on the
materials in use. The high technology equipment in use today was
designed to help us safely explore the more exotic operating
environments where the gases no longer behave in an "ideal" manner.
This advanced math, including statistics and multivariable calculus,
makes possible the solution to such complex dynamic situations as
space vehicle reentry. An example is the analysis of the space shuttle
reentry pictured to ensure the material properties under this loading
condition are appropriate. In this flight regime, the gas is no longer
Main article: Pressure
The symbol used to represent pressure in equations is "p" or "P" with
SI units of pascals.
When describing a container of gas, the term pressure (or absolute
pressure) refers to the average force per unit area that the gas
exerts on the surface of the container. Within this volume, it is
sometimes easier to visualize the gas particles moving in straight
lines until they collide with the container (see diagram at top of the
article). The force imparted by a gas particle into the container
during this collision is the change in momentum of the particle.
During a collision only the normal component of velocity changes. A
particle traveling parallel to the wall does not change its momentum.
Therefore, the average force on a surface must be the average change
in linear momentum from all of these gas particle collisions.
Pressure is the sum of all the normal components of force exerted by
the particles impacting the walls of the container divided by the
surface area of the wall.
Air balloon shrinks after submersion in liquid nitrogen
The symbol used to represent temperature in equations is T with SI
units of kelvins.
The speed of a gas particle is proportional to its absolute
temperature. The volume of the balloon in the video shrinks when the
trapped gas particles slow down with the addition of extremely cold
nitrogen. The temperature of any physical system is related to the
motions of the particles (molecules and atoms) which make up the [gas]
system. In statistical mechanics, temperature is the measure of
the average kinetic energy stored in a particle. The methods of
storing this energy are dictated by the degrees of freedom of the
particle itself (energy modes).
Kinetic energy added (endothermic
process) to gas particles by way of collisions produces linear,
rotational, and vibrational motion. In contrast, a molecule in a solid
can only increase its vibrational modes with the addition of heat as
the lattice crystal structure prevents both linear and rotational
motions. These heated gas molecules have a greater speed range which
constantly varies due to constant collisions with other particles. The
speed range can be described by the Maxwell–Boltzmann distribution.
Use of this distribution implies ideal gases near thermodynamic
equilibrium for the system of particles being considered.
Main article: Specific volume
The symbol used to represent specific volume in equations is "v" with
SI units of cubic meters per kilogram.
The symbol used to represent volume in equations is "V" with SI units
of cubic meters.
When performing a thermodynamic analysis, it is typical to speak of
intensive and extensive properties. Properties which depend on the
amount of gas (either by mass or volume) are called extensive
properties, while properties that do not depend on the amount of gas
are called intensive properties.
Specific volume is an example of an
intensive property because it is the ratio of volume occupied by a
unit of mass of a gas that is identical throughout a system at
equilibrium. 1000 atoms a gas occupy the same space as any other
1000 atoms for any given temperature and pressure. This concept is
easier to visualize for solids such as iron which are incompressible
compared to gases. Since a gas fills any container in which it is
placed, volume is an extensive property.
Main article: Density
The symbol used to represent density in equations is ρ (rho) with SI
units of kilograms per cubic meter. This term is the reciprocal of
Since gas molecules can move freely within a container, their mass is
normally characterized by density.
Density is the amount of mass per
unit volume of a substance, or the inverse of specific volume. For
gases, the density can vary over a wide range because the particles
are free to move closer together when constrained by pressure or
volume. This variation of density is referred to as compressibility.
Like pressure and temperature, density is a state variable of a gas
and the change in density during any process is governed by the laws
of thermodynamics. For a static gas, the density is the same
throughout the entire container.
Density is therefore a scalar
quantity. It can be shown by kinetic theory that the density is
inversely proportional to the size of the container in which a fixed
mass of gas is confined. In this case of a fixed mass, the density
decreases as the volume increases.
If one could observe a gas under a powerful microscope, one would see
a collection of particles (molecules, atoms, ions, electrons, etc.)
without any definite shape or volume that are in more or less random
motion. These neutral gas particles only change direction when they
collide with another particle or with the sides of the container. In
an ideal gas, these collisions are perfectly elastic. This particle or
microscopic view of a gas is described by the Kinetic-molecular
theory. The assumptions behind this theory can be found in the
postulates section of Kinetic Theory.
Main article: Kinetic theory of gases
Kinetic theory provides insight into the macroscopic properties of
gases by considering their molecular composition and motion. Starting
with the definitions of momentum and kinetic energy, one can use
the conservation of momentum and geometric relationships of a cube to
relate macroscopic system properties of temperature and pressure to
the microscopic property of kinetic energy per molecule. The theory
provides averaged values for these two properties.
The theory also explains how the gas system responds to change. For
example, as a gas is heated from absolute zero, when it is (in theory)
perfectly still, its internal energy (temperature) is increased. As a
gas is heated, the particles speed up and its temperature rises. This
results in greater numbers of collisions with the container per unit
time due to the higher particle speeds associated with elevated
temperatures. The pressure increases in proportion to the number of
collisions per unit time.
Random motion of gas particles results in diffusion.
Main article: Brownian motion
Brownian motion is the mathematical model used to describe the random
movement of particles suspended in a fluid. The gas particle
animation, using pink and green particles, illustrates how this
behavior results in the spreading out of gases (entropy). These events
are also described by particle theory.
Since it is at the limit of (or beyond) current technology to observe
individual gas particles (atoms or molecules), only theoretical
calculations give suggestions about how they move, but their motion is
Brownian motion because
Brownian motion involves a
smooth drag due to the frictional force of many gas molecules,
punctuated by violent collisions of an individual (or several) gas
molecule(s) with the particle. The particle (generally consisting of
millions or billions of atoms) thus moves in a jagged course, yet not
so jagged as would be expected if an individual gas molecule were
When gases are compressed, intermolecular forces like those shown here
start to play a more active role.
Main articles: van der Waals force and Intermolecular force
As discussed earlier, momentary attractions (or repulsions) between
particles have an effect on gas dynamics. In physical chemistry, the
name given to these intermolecular forces is van der Waals force.
These forces play a key role in determining physical properties of a
gas such as viscosity and flow rate (see physical characteristics
section). Ignoring these forces in certain conditions allows a real
gas to be treated like an ideal gas. This assumption allows the use of
ideal gas laws which greatly simplifies calculations.
Proper use of these gas relationships requires the kinetic-molecular
theory (KMT). When gas particles experience intermolecular forces they
gradually influence one another as the spacing between them is reduced
(the hydrogen bond model illustrates one example). In the absence of
any charge, at some point when the spacing between gas particles is
greatly reduced they can no longer avoid collisions between themselves
at normal gas temperatures. Another case for increased collisions
among gas particles would include a fixed volume of gas, which upon
heating would contain very fast particles. This means that these ideal
equations provide reasonable results except for extremely high
pressure (compressible) or high temperature (ionized) conditions. All
of these excepted conditions allow energy transfer to take place
within the gas system. The absence of these internal transfers is what
is referred to as ideal conditions in which the energy exchange occurs
only at the boundaries of the system. Real gases experience some of
these collisions and intermolecular forces. When these collisions are
statistically negligible (incompressible), results from these ideal
equations are still meaningful. If the gas particles are compressed
into close proximity they behave more like a liquid (see fluid
Main article: Equation of state
An equation of state (for gases) is a mathematical model used to
roughly describe or predict the state properties of a gas. At present,
there is no single equation of state that accurately predicts the
properties of all gases under all conditions. Therefore, a number of
much more accurate equations of state have been developed for gases in
specific temperature and pressure ranges. The "gas models" that are
most widely discussed are "perfect gas", "ideal gas" and "real gas".
Each of these models has its own set of assumptions to facilitate the
analysis of a given thermodynamic system. Each successive model
expands the temperature range of coverage to which it applies.
Ideal and perfect gas models
Main article: Perfect gas
The equation of state for an ideal or perfect gas is the ideal gas law
where P is the pressure, V is the volume, n is amount of gas (in mol
units), R is the universal gas constant, 8.314 J/(mol K),
and T is the temperature. Written this way, it is sometimes called the
"chemist's version", since it emphasizes the number of molecules n. It
can also be written as
displaystyle P=rho R_ s T,
displaystyle R_ s
is the specific gas constant for a particular gas, in units
J/(kg K), and ρ = m/V is density. This notation is the "gas
dynamicist's" version, which is more practical in modeling of gas
flows involving acceleration without chemical reactions.
The ideal gas law does not make an assumption about the specific heat
of a gas. In the most general case, the specific heat is a function of
both temperature and pressure. If the pressure-dependence is neglected
(and possibly the temperature-dependence as well) in a particular
application, sometimes the gas is said to be a perfect gas, although
the exact assumptions may vary depending on the author and/or field of
For an ideal gas, the ideal gas law applies without restrictions on
the specific heat. An ideal gas is a simplified "real gas" with the
assumption that the compressibility factor Z is set to 1 meaning that
this pneumatic ratio remains constant. A compressibility factor of one
also requires the four state variables to follow the ideal gas law.
This approximation is more suitable for applications in engineering
although simpler models can be used to produce a "ball-park" range as
to where the real solution should lie. An example where the "ideal gas
approximation" would be suitable would be inside a combustion chamber
of a jet engine. It may also be useful to keep the elementary
reactions and chemical dissociations for calculating emissions.
21 April 1990 eruption of Mount Redoubt, Alaska, illustrating real
gases not in thermodynamic equilibrium.
Main article: Real gas
Each one of the assumptions listed below adds to the complexity of the
problem's solution. As the density of a gas increases with rising
pressure, the intermolecular forces play a more substantial role in
gas behavior which results in the ideal gas law no longer providing
"reasonable" results. At the upper end of the engine temperature
ranges (e.g. combustor sections – 1300 K), the complex fuel
particles absorb internal energy by means of rotations and vibrations
that cause their specific heats to vary from those of diatomic
molecules and noble gases. At more than double that temperature,
electronic excitation and dissociation of the gas particles begins to
occur causing the pressure to adjust to a greater number of particles
(transition from gas to plasma). Finally, all of the thermodynamic
processes were presumed to describe uniform gases whose velocities
varied according to a fixed distribution. Using a non-equilibrium
situation implies the flow field must be characterized in some manner
to enable a solution. One of the first attempts to expand the
boundaries of the ideal gas law was to include coverage for different
thermodynamic processes by adjusting the equation to read pVn =
constant and then varying the n through different values such as the
specific heat ratio, γ.
Real gas effects include those adjustments made to account for a
greater range of gas behavior:
Compressibility effects (Z allowed to vary from 1.0)
Variable heat capacity (specific heats vary with temperature)
Van der Waals forces (related to compressibility, can substitute other
equations of state)
Non-equilibrium thermodynamic effects
Issues with molecular dissociation and elementary reactions with
For most applications, such a detailed analysis is excessive. Examples
where real gas effects would have a significant impact would be on the
Space Shuttle re-entry where extremely high temperatures and pressures
were present or the gases produced during geological events as in the
image of the 1990 eruption of Mount Redoubt.
Main article: Boyle's law
Boyle's Law was perhaps the first expression of an equation of state.
Robert Boyle performed a series of experiments employing a
J-shaped glass tube, which was sealed on one end. Mercury was added to
the tube, trapping a fixed quantity of air in the short, sealed end of
the tube. Then the volume of gas was carefully measured as additional
mercury was added to the tube. The pressure of the gas could be
determined by the difference between the mercury level in the short
end of the tube and that in the long, open end. The image of Boyle's
Equipment shows some of the exotic tools used by Boyle during his
study of gases.
Through these experiments, Boyle noted that the pressure exerted by a
gas held at a constant temperature varies inversely with the volume of
the gas. For example, if the volume is halved, the pressure is
doubled; and if the volume is doubled, the pressure is halved. Given
the inverse relationship between pressure and volume, the product of
pressure (P) and volume (V) is a constant (k) for a given mass of
confined gas as long as the temperature is constant. Stated as a
formula, thus is:
Because the before and after volumes and pressures of the fixed amount
of gas, where the before and after temperatures are the same both
equal the constant k, they can be related by the equation:
displaystyle qquad P_ 1 V_ 1 =P_ 2 V_ 2 .
Main article: Charles's law
In 1787, the French physicist and balloon pioneer, Jacques Charles,
found that oxygen, nitrogen, hydrogen, carbon dioxide, and air expand
to the same extent over the same 80 kelvin interval. He noted that,
for an ideal gas at constant pressure, the volume is directly
proportional to its temperature:
displaystyle frac V_ 1 T_ 1 = frac V_ 2 T_ 2
Main article: Gay-Lussac's Law
Joseph Louis Gay-Lussac
Joseph Louis Gay-Lussac published results of similar, though
more extensive experiments. Gay-Lussac credited Charles' earlier
work by naming the law in his honor. Gay-Lussac himself is credited
with the law describing pressure, which he found in 1809. It states
that the pressure exerted on a container's sides by an ideal gas is
proportional to its temperature.
displaystyle frac P_ 1 T_ 1 = frac P_ 2 T_ 2 ,
Main article: Avogadro's law
Amedeo Avogadro verified that equal volumes of pure gases
contain the same number of particles. His theory was not generally
accepted until 1858 when another Italian chemist Stanislao Cannizzaro
was able to explain non-ideal exceptions. For his work with gases a
century prior, the number that bears his name Avogadro's constant
represents the number of atoms found in 12 grams of elemental
carbon-12 (6.022×1023 mol−1). This specific number of gas
particles, at standard temperature and pressure (ideal gas law)
occupies 22.40 liters, which is referred to as the molar volume.
Avogadro's law states that the volume occupied by an ideal gas is
proportional to the number of moles (or molecules) present in the
container. This gives rise to the molar volume of a gas, which at STP
is 22.4 dm3 (or litres). The relation is given by
displaystyle frac V_ 1 n_ 1 = frac V_ 2 n_ 2 ,
where n is equal to the number of moles of gas (the number of
molecules divided by Avogadro's Number).
Main article: Dalton's law
John Dalton published the Law of Partial Pressures from his
work with ideal gas law relationship: The pressure of a mixture of non
reactive gases is equal to the sum of the pressures of all of the
constituent gases alone. Mathematically, this can be represented for n
Pressuretotal = Pressure1 + Pressure2 + ... + Pressuren
The image of Dalton's journal depicts symbology he used as shorthand
to record the path he followed. Among his key journal observations
upon mixing unreactive "elastic fluids" (gases) were the
Unlike liquids, heavier gases did not drift to the bottom upon mixing.
Gas particle identity played no role in determining final pressure
(they behaved as if their size was negligible).
Compressibility factors for air.
Thermodynamicists use this factor (Z) to alter the ideal gas equation
to account for compressibility effects of real gases. This factor
represents the ratio of actual to ideal specific volumes. It is
sometimes referred to as a "fudge-factor" or correction to expand the
useful range of the ideal gas law for design purposes. Usually this Z
value is very close to unity. The compressibility factor image
illustrates how Z varies over a range of very cold temperatures.
Main article: Reynolds number
In fluid mechanics, the
Reynolds number is the ratio of inertial
forces (vsρ) to viscous forces (μ/L). It is one of the most
important dimensionless numbers in fluid dynamics and is used, usually
along with other dimensionless numbers, to provide a criterion for
determining dynamic similitude. As such, the
Reynolds number provides
the link between modeling results (design) and the full-scale actual
conditions. It can also be used to characterize the flow.
Satellite view of weather pattern in vicinity of Robinson Crusoe
Islands on 15 September 1999, shows a turbulent cloud pattern called a
Kármán vortex street
Main article: Viscosity
Viscosity, a physical property, is a measure of how well adjacent
molecules stick to one another. A solid can withstand a shearing force
due to the strength of these sticky intermolecular forces. A fluid
will continuously deform when subjected to a similar load. While a gas
has a lower value of viscosity than a liquid, it is still an
observable property. If gases had no viscosity, then they would not
stick to the surface of a wing and form a boundary layer. A study of
the delta wing in the Schlieren image reveals that the gas particles
stick to one another (see
Boundary layer section).
Delta wing in wind tunnel. The shadows form as the indices of
refraction change within the gas as it compresses on the leading edge
of this wing.
Main article: Turbulence
In fluid dynamics, turbulence or turbulent flow is a flow regime
characterized by chaotic, stochastic property changes. This includes
low momentum diffusion, high momentum convection, and rapid variation
of pressure and velocity in space and time. The satellite view of
weather around Robinson Crusoe Islands illustrates one example.
Main article: Boundary layer
Particles will, in effect, "stick" to the surface of an object moving
through it. This layer of particles is called the boundary layer. At
the surface of the object, it is essentially static due to the
friction of the surface. The object, with its boundary layer is
effectively the new shape of the object that the rest of the molecules
"see" as the object approaches. This boundary layer can separate from
the surface, essentially creating a new surface and completely
changing the flow path. The classical example of this is a stalling
airfoil. The delta wing image clearly shows the boundary layer
thickening as the gas flows from right to left along the leading edge.
Maximum entropy principle
Main article: Principle of maximum entropy
As the total number of degrees of freedom approaches infinity, the
system will be found in the macrostate that corresponds to the highest
multiplicity. In order to illustrate this principle, observe the skin
temperature of a frozen metal bar. Using a thermal image of the skin
temperature, note the temperature distribution on the surface. This
initial observation of temperature represents a "microstate". At some
future time, a second observation of the skin temperature produces a
second microstate. By continuing this observation process, it is
possible to produce a series of microstates that illustrate the
thermal history of the bar's surface. Characterization of this
historical series of microstates is possible by choosing the
macrostate that successfully classifies them all into a single
When energy transfer ceases from a system, this condition is referred
to as thermodynamic equilibrium. Usually this condition implies the
system and surroundings are at the same temperature so that heat no
longer transfers between them. It also implies that external forces
are balanced (volume does not change), and all chemical reactions
within the system are complete. The timeline varies for these events
depending on the system in question. A container of ice allowed to
melt at room temperature takes hours, while in semiconductors the heat
transfer that occurs in the device transition from an on to off state
could be on the order of a few nanoseconds.
List of gases
^ This early 20th century discussion infers what is regarded as the
plasma state. See page 137 of American Chemical Society, Faraday
Society, Chemical Society (Great Britain) The Journal of Physical
Chemistry, Volume 11 Cornell (1907).
^ The work by T. Zelevinski provides another link to latest research
about Strontium in this new field of study. See Tanya Zelevinsky
(2009). "84Sr—just right for forming a Bose-Einstein condensate".
Physics. 2: 94. Bibcode:2009PhyOJ...2...94Z.
^ for links material on the
Bose–Einstein condensate see Quantum Gas
Microscope Offers Glimpse Of Quirky Ultracold Atoms. ScienceDaily. 4
^ J. B. van Helmont, Ortus medicinae. … (Amsterdam, (Netherlands):
Louis Elzevir, 1652 (first edition: 1648)). The word "gas" first
appears on page 58, where he mentions: "…
Gas (meum scil. inventum)
…" (… gas (namely, my discovery) …). On page 59, he states: "…
in nominis egestate, halitum illum,
Gas vocavi, non longe a Chao …"
(… in need of a name, I called this vapor "gas", not far from
^ Ley, Willy (June 1966). "The Re-Designed Solar System". For Your
Information. Galaxy Science Fiction. pp. 94–106.
^ Harper, Douglas. "gas". Online Etymology Dictionary.
^ Draper, John William (1861). A textbook on chemistry. New York:
Harper and Sons. p. 178.
^ The authors make the connection between molecular forces of metals
and their corresponding physical properties. By extension, this
concept would apply to gases as well, though not universally. Cornell
(1907) pp. 164–5.
^ One noticeable exception to this physical property connection is
conductivity which varies depending on the state of matter (ionic
compounds in water) as described by
Michael Faraday in 1833 when he
noted that ice does not conduct a current. See page 45 of John
Tyndall's Faraday as a Discoverer (1868).
^ John S. Hutchinson (2008). Concept Development Studies in Chemistry.
^ Anderson, p.501
^ J. Clerk Maxwell (1904). Theory of Heat. Mineola: Dover
Publications. pp. 319–20. ISBN 0-486-41735-2.
^ See pages 137–8 of Society, Cornell (1907).
^ Kenneth Wark (1977). Thermodynamics (3 ed.). McGraw-Hill.
p. 12. ISBN 0-07-068280-1.
^ For assumptions of
Kinetic Theory see McPherson, pp.60–61
^ Anderson, pp. 289–291
^ John, p.205
^ John, pp. 247–56
^ McPherson, pp.52–55
^ McPherson, pp.55–60
^ John P. Millington (1906). John Dalton. pp. 72, 77–78.
Anderson, John D. (1984). Fundamentals of Aerodynamics. McGraw-Hill
Higher Education. ISBN 0-07-001656-9.
John, James (1984).
Gas Dynamics. Allyn and Bacon.
McPherson, William; Henderson, William (1917). An Elementary study of
Wikimedia Commons has media related to Gases.
Philip Hill and Carl Peterson. Mechanics and Thermodynamics of
Propulsion: Second Edition Addison-Wesley, 1992.
National Aeronautics and Space Administration (NASA). Animated Gas
Lab. Accessed February 2008.
Georgia State University. HyperPhysics. Accessed February 2008.
Antony Lewis WordWeb. Accessed February 2008.
Northwestern Michigan College The Gaseous State. Accessed February
States of matter (list)
Gas / Vapor
Quantum spin liquid
Enthalpy of fusion
Enthalpy of sublimation
Enthalpy of vaporization
Latent internal energy
Equation of state
Macroscopic quantum phenomena
Order and disorder (physics)