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Gas | Definition, State of Matter, Properties, Structure, & Facts | Britannica
Gas | Definition, State of Matter, Properties, Structure, & Facts | Britannica
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gas
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IntroductionStructureKinetic-molecular pictureNumerical magnitudesIntermolecular separation and average speedMean-free path and collision rateMolecular sizesSummary of numerical magnitudesFree-molecule gasContinuity of gaseous and liquid statesBehaviour and propertiesEquilibrium propertiesIdeal gas equation of stateInternal energyTransport propertiesViscosityHeat conductionDiffusionThermal diffusionKinetic theory of gasesIdeal gasPressureEffusionThermal transpirationViscosityThermal conductivityDiffusion and thermal diffusionBoltzmann equationDeviations from the ideal modelEquation of stateTransport properties
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Also known as: gaseous state
Written by
Edward A. Mason
Professor of Chemistry and Engineering, 1967–92; Newport Rogers Professor of Chemistry, 1983–92, Brown University, Providence, Rhode Island. Coauthor of Transport Properties of Ions in Gases and...
Edward A. Mason
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gas, one of the three fundamental states of matter, with distinctly different properties from the liquid and solid states. Structure The remarkable feature of gases is that they appear to have no structure at all. They have neither a definite size nor shape, whereas ordinary solids have both a definite size and a definite shape, and liquids have a definite size, or volume, even though they adapt their shape to that of the container in which they are placed. Gases will completely fill any closed container; their properties depend on the volume of a container but not on its shape. Kinetic-molecular picture Gases nevertheless do have a structure of sorts on a molecular scale. They consist of a vast number of molecules moving chaotically in all directions and colliding with one another and with the walls of their container. Beyond this, there is no structure—the molecules are distributed essentially randomly in space, traveling in arbitrary directions at speeds that are distributed randomly about an average determined by the gas temperature. The pressure exerted by a gas is the result of the innumerable impacts of the molecules on the container walls and appears steady to human senses because so many collisions occur each second on all sections of the walls. More subtle properties such as heat conductivity, viscosity (resistance to flow), and diffusion are attributed to the molecules themselves carrying the mechanical quantities of energy, momentum, and mass, respectively. These are called transport properties, and the rate of transport is dominated by the collisions between molecules, which force their trajectories into tortuous shapes. The molecular collisions are in turn controlled by the forces between the molecules and are described by the laws of mechanics. Thus, gases are treated as a large collection of tiny particles subject to the laws of physics. Their properties are attributed primarily to the motion of the molecules and can be explained by the kinetic theory of gases. It is not obvious that this should be the case, and for many years a static picture of gases was instead espoused, in which the pressure, for instance, was attributed to repulsive forces between essentially stationary particles pushing on the container walls. How the kinetic-molecular picture finally came to be universally accepted is a fascinating piece of scientific history and is discussed briefly below in the section Kinetic theory of gases. Any theory of gas behaviour based on this kinetic model must also be a statistical one because of the enormous numbers of particles involved. The kinetic theory of gases is now a classical part of statistical physics and is indeed a sort of miniature display case for many of the fundamental concepts and methods of science. Such important modern concepts as distribution functions, cross sections, microscopic reversibility, and time-reversal invariance have their historical roots in kinetic theory, as does the entire atomistic view of matter.
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Numerical magnitudes When considering various physical phenomena, it is helpful for one to have some idea of the numerical magnitudes involved. In particular, there are several characteristics whose values should be known, at least within an order of magnitude (a factor of 10), in order for one to obtain a clear idea of the nature of gaseous molecules. These features include the size, average speed, and intermolecular separation at ordinary temperatures and pressures. In addition, other important considerations are how many collisions a typical molecule makes in one second under these conditions and how far such a typical molecule travels before colliding with another molecule. It has been established that molecules have sizes on the order of a few angstrom units (1 Å = 10−8 centimetre [cm]) and that there are about 6 × 1023 molecules in one mole, which is defined as the amount of a substance whose mass in grams is equal to its molecular weight (e.g., 1 mole of water, H2O, is 18.0152 grams). With this knowledge, one could calculate at least some of the gas values. It is interesting to see how the answers could be estimated from simple observations and then to compare the results to the accepted values that are based on more precise measurements and theories. Intermolecular separation and average speed One of the easiest properties to work out is the average distance between molecules compared to their diameter; water will be used here for this purpose. Consider 1 gram of H2O at 100° C and atmospheric pressure, which are the normal boiling point conditions. The liquid occupies a volume of 1.04 cubic centimetres (cm3); once converted to steam it occupies a volume of 1.67 × 103 cm3. Thus, the average volume occupied by one molecule in the gas is larger than the corresponding volume occupied in the liquid by a factor of 1.67 × 103/1.04, or about 1,600. Since volume varies as the cube of distance, the ratio of the mean separation distance in the gas to that in the liquid is roughly equal to the cube root of 1,600, or about 12. If the molecules in the liquid are considered to be touching each other, the ratio of the intermolecular separation to the molecular diameter in ordinary gases is on the order of 10 under ordinary conditions. It should be noted that the actual separation and diameter cannot be determined in this way; only their ratio can be calculated.
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It is also relatively simple to estimate the average speed of gas molecules. Consider a sound wave in a gas, which is just the propagation of a small pressure disturbance. If pressure is attributed to molecular impacts on a test surface, then surely a pressure disturbance cannot travel faster than the molecules themselves. In other words, the average molecular speed in a gas should be somewhat greater than the speed of sound in the gas. The speed of sound in air at ordinary temperatures is about 330 metres per second (m/s), so the molecular speed will be estimated here to be somewhat greater, say, about 5 × 104 centimetres per second (cm/s). This value depends on the particular gas and the temperature, but it will be sufficient for the kind of estimates sought here. Mean-free path and collision rate The average molecular speed, along with an observed rate of the diffusion of gases, can be used to estimate the length and tortuosity of the path traveled by a typical molecule. If a bottle of ammonia is opened in a closed room, at least a few minutes pass before the ammonia can be detected at a distance of just one metre. (Ammonia, NH3, is a gas; the familiar bottle of “ammonia” typically seen is actually a solution of the gas in water.) Yet, if the ammonia molecules traveled directly to an observer at a speed somewhat faster than that of sound, the odour should be detectable in only a few milliseconds. The explanation for the discrepancy is that the ammonia molecules collide with many air molecules, and their paths are greatly distorted as a result. For a quantitative estimate of the diffusion time, a more controlled system must be considered, because even gentle stray air currents in a closed room greatly speed up the spreading of the ammonia. To eliminate the effect of such air currents, a closed tube—say, a glass tube one centimetre in diameter and one metre in length—can be used. A small amount of ammonia gas is released at one end, and both ends are then closed. In order to measure how long it takes for the ammonia to travel to the other end, a piece of moist red litmus paper might be used as a detector; it will turn blue when the ammonia reaches it. This process takes quite a long time—about several hours—because diffusion occurs at such a slow rate. In this case, the time will be taken to be approximately 3 hours, or roughly 104 seconds (s). During this time interval, a typical ammonia molecule actually travels a distance of (5 × 104 cm/s)(104 s) = 5 × 108 cm = 5,000 kilometres (km), roughly the distance across the United States. In other words, such a molecule travels a total distance of five million metres in order to progress a net distance of only one metre.
The solution to a basic statistical problem can be used to estimate the number of collisions such a typical diffusing molecule experienced (N) and the average distance traveled between collisions (l), called the mean free path. The product of N and l must equal the total distance traveled—i.e., Nl = 5 × 108 cm. This distance can be thought of as a chain 5,000 km long, made up of N links, each of length l. The statistical question then is as follows: If such a chain is randomly jumbled, how far apart will its ends be on the average? This end-to-end distance corresponds to the length of the diffusion tube (one metre). This is a venerable statistical problem that recurs in many applications. One of the more vivid ways of illustrating the concept is known as the “drunkard’s walk.” In this scenario a drunkard takes steps of length l but, because of inebriation, takes them in random directions. After N steps, how far will he be from his starting point? The answer is that his progress is proportional not to N but to N1/2. For example, if the drunkard takes four steps, each of length l, he will end up at a distance of 2l from his starting point. Gas molecules move in three dimensions, whereas the drunkard moves in two dimensions; however, the result is the same. Thus, the square root of N multiplied by the length of the mean free path equals the length of the diffusion tube: N1/2l = 102 cm. From the equations for Nl and N1/2l, it can readily be calculated that N = 2.5 × 1013 collisions and l = 2.0 × 10-5 cm. The mean time between collisions, τ, is found by dividing the time of the diffusion experiment by the number of collisions during that time: τ = (104)/(2.5 × 1013) = 4 × 10-10 seconds between collisions, corresponding to a collision frequency of 2.5 × 109 collisions per second. It is thus understandable that gases appear to be continuous fluids on ordinary scales of time and distance.
10.S: Gases (Summary) - Chemistry LibreTexts
10.S: Gases (Summary) - Chemistry LibreTexts
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10.1: Characteristics of Gases10.2: Pressure10.3: The Gas Laws10.4: The Ideal Gas Equation10.5: Further Applications of the Ideal-Gas Equations10.6: Gas Mixtures and Partial Pressures10.7: Kinetic-Molecular Theory10.8: Molecular Effusion and Diffusion10.9: Real Gases - Deviations from Ideal Behavior
10.1: Characteristics of Gases
Bulk matter can exist in three states: gas, liquid, and solid. Gases have the lowest density of the three, are highly compressible, and fill their containers completely. Elements that exist as gases at room temperature and pressure are clustered on the right side of the periodic table; they occur as either monatomic gases (the noble gases) or diatomic molecules (some halogens, N₂, O₂).
Gases expand spontaneously to fill containers in which they are held, equaling their volume. Consequently, they are highly compressible.
Gases form homogeneous mixtures with each other regardless of the identities or relative proportions of the component gases
The characteristic properties of gases arise because the individual molecules are relatively far apart, hence, acting largely as though they were alone.
10.2: Pressure
Pressure is defined as the force exerted per unit area; it can be measured using a barometer or manometer. Four quantities must be known for a complete physical description of a sample of a gas: temperature, volume, amount, and pressure. Pressure is force per unit area of surface; the SI unit for pressure is the pascal (Pa), defined as 1 newton per square meter (N/m²). The pressure exerted by an object is proportional to the force it exerts and inversely proportional to the area.
Pressure, P, is the Force, F, that acts on a given Area, A:
\[P = F / A \nonumber \]
Atmospheric Pressure and the Barometer
The force, F, exerted by any object is the product of its mass, m, times its acceleration, a: F = ma
SI unit for force is kg-m/s2 and is called the Newton (N)
SI unit of pressure is N/m2, called a Pascal
Standard atmospheric pressure: defined as 760 torr (760 mm Hg), or, in SI units, 101.325 kPa
Atmosphere: unit of pressure equal to 760 torr; 1 atm = 101.325 kPa
10.3: The Gas Laws
The volume of a gas is inversely proportional to its pressure and directly proportional to its temperature and the amount of gas. Boyle showed that the volume of a sample of a gas is inversely proportional to pressure (Boyle’s law), Charles and Gay-Lussac demonstrated that the volume of a gas is directly proportional to its temperature at constant pressure (Charles’s law), and Avogadro showed that the volume of a gas is directly proportional to the number of moles of gas (Avogadro’s law).
The Pressure-Volume Relationship: Boyle’s Law
Boyle’s law states that the volume of a fixed quantity of gas maintained at constant temperature is inversely proportional to the pressure. When two measurements are inversely proportional, one gets smaller as the other one gets larger.
\[PV = constant \nonumber \]
where \(P\) = pressure, \(V\) = volume
The Temperature-Volume Relationship: Charles’s Law
Charles’s law: states that the volume of a fixed amount of gas maintained at constant pressure is directly proportional to its absolute temperature. Thus, as the pressure gets higher, so does the temperature.
\[\dfrac{V}{T} = constant \nonumber \]
where \(V\) = volume, \(T\) = Temperature
The Quantity-Volume Relationship: Avogadro’s Law
Law of combining volumes: at a given pressure and temperature, the volumes of gases that react with one another are in the ratios of small whole numbers. (ie: 2H2 + O2 = 2H2O)
Avogadro’s hypothesis: equal volumes of gases at the same temperature and pressure contain equal number of molecules
Avogadro’s law: The volume of a gas maintained at constant temperature and pressure is directly proportional to the number of moles of the gas
\[V = constant \times n \nonumber \]
where V = volume, n = number of moles
10.4: The Ideal Gas Equation
The empirical relationships among the volume, the temperature, the pressure, and the amount of a gas can be combined into the ideal gas law, PV = nRT. The proportionality constant, R, is called the gas constant. The ideal gas law describes the behavior of an ideal gas, a hypothetical substance whose behavior can be explained quantitatively by the ideal gas law and the kinetic molecular theory of gases. Standard temperature and pressure (STP) is 0°C and 1 atm.
Ideal-gas equation
\[PV = nRT \nonumber \]
where P = pressure, V = volume, n = number of moles, R = gas constant, T = Temperature (always expressed on absolute-temperature scale, usually Kelvin)
The Gas constant (\(R\)) is the constant of proportionality in the ideal-gas equation. Some values of R are given below
Some values of R
Units
Numerical value
L-atm/mol-K
0.08206
Cal/mol-K
1.987
J/mol-K
8.314
M3-Pa/mol-K
8.314
L-torr/mol-K
62.36
Standard temperature and pressure (STP): 0°C and 1 atm. 1 mol of gas at STP has a volume of 22.41 L (molar volume)
10.5: Further Applications of the Ideal-Gas Equations
The relationship between the amounts of products and reactants in a chemical reaction can be expressed in units of moles or masses of pure substances, of volumes of solutions, or of volumes of gaseous substances. The ideal gas law can be used to calculate the volume of gaseous products or reactants as needed. In the laboratory, gases produced in a reaction are often collected by the displacement of water from filled vessels; the amount of gas can be calculated from the volume of water displaced.
Density of a gas (\*\rho\) = density, M = molar mass):
\[\rho = \dfrac{PM}{ RT} \nonumber \]
Molar mass of a gas:
\[M = \dfrac{ \rho RT }{ P} \nonumber \]
10.6: Gas Mixtures and Partial Pressures
The pressure exerted by each gas in a gas mixture is independent of the pressure exerted by all other gases present. Consequently, the total pressure exerted by a mixture of gases is the sum of the partial pressures of the components (Dalton’s law of partial pressures). The amount of gas in a mixture may be described by its partial pressure or its mole fraction. In a mixture, the partial pressure of each gas is the product of the total pressure and the mole fraction.
Partial pressure: the pressure exerted by a particular gas in a mixture
Dalton’s law of partial pressures: law stating that the total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone
The total pressure at constant temperature and volume is determined by the total number of moles of gas present, whether that total represents just one substance or a mixture
Partial Pressures and Mole Fractions
Mole fraction: the ratio of the number of one component of a mixture to the total moles of all components; abbreviated \(\chi\), with a subscript to identify the components.
The partial pressure of a gas in a mixture is its mole fraction times the total pressure
10.7: Kinetic-Molecular Theory
The behavior of ideal gases is explained by the kinetic molecular theory of gases. Molecular motion, which leads to collisions between molecules and the container walls, explains pressure, and the large intermolecular distances in gases explain their high compressibility. Although all gases have the same average kinetic energy at a given temperature, they do not all possess the same root mean square speed. The actual values of speed and kinetic energy are not the same for all gas particles.
Kinetic-molecular theory: set of assumptions about the nature of gases. These assumptions, when translated into mathematical form, yield the ideal-gas equation
Gases consist of large numbers of molecules that are in continuous, random motion
The volume of all the molecules of the gas is negligible compared to the total volume in which the gas is contained
Attractive and repulsive forces between gas molecules are negligible
Energy can be transferred between molecules during collisions, but the average kinetic energy of the molecules does not change with time, as long as the temperature of the gas remains constant
The average kinetic energy of the molecules is proportional to absolute temperature. At any given temperature, the molecules of all gases have the same average kinetic energy
The pressure of a gas is caused by collisions of molecules with the walls of the container. The magnitude of the pressure is determined by how often and how "hard" the molecules strike the walls.
If two different gases are at the same temperature, they have the same average kinetic energy. If the temperature of a gas is doubled, its kinetic energy also doubles. Hence, molecular motion increases with increasing temperature.
Root-mean-square (rms) speed: the square root of the squared speeds of the gas molecules in a gas sample. This quantity is the speed of a molecule possessing average kinetic energy.
The rms speed is important because the average kinetic energy of the gas molecules, \(ε\), is related directly to \(u^2\):
\[ ε = \dfrac{1}{2}mu^2 \nonumber \]
where \(m\) is the mass of the molecule
Because mass doesn’t change with temperature, the rms speed (and also the average speed) of molecules must increase as temperature increases
Applications to the Gas Laws
Effect of a volume increase at constant temperature: If the volume is increased, the molecules must move a longer distance between collisions. Consequently, there are fewer collisions per unit time with the container walls, and pressure decreases.
Effect of a temperature increase at constant volume: An increases in temperature means an increase in the average kinetic energy of the molecules. If there is no change in volume, there will be more collisions with the walls per unit time. Furthermore, the molecules strike harder, hence explaining how the observed pressure increases.
10.8: Molecular Effusion and Diffusion
Diffusion is the gradual mixing of gases to form a sample of uniform composition even in the absence of mechanical agitation. In contrast, effusion is the escape of a gas from a container through a tiny opening into an evacuated space. The rate of effusion of a gas is inversely proportional to the square root of its molar mass (Graham’s law), a relationship that closely approximates the rate of diffusion. As a result, light gases tend to diffuse and effuse much more rapidly than heavier gases.
A gas composed of light gas particles will have the same average kinetic energy as one composed of much heavier particles, provided that the two gases are at the same temperature. The mass, \(m\), of the particles in the lighter gas is smaller that that in the heavier gas. Consequently, the particles of the lighter gas must have a higher rms speed, \(u\), than the heavier one:
\[u =\sqrt{\dfrac{3RT}{ M}} \nonumber \]
Since M is in the denominator, the less massive the gas molecules, the higher the rms speed
Effusion: the escape of a gas through an orifice or hole. The rate of effusion depends on the molecular mass of the gas.
Diffusion: the spreading of one substance through a space or through another substance
Graham’s Law of Effusion
Graham’s law: law stating that the rate of effusion of a gas is inversely proportional to the square root of its molecular weight
\[\dfrac{r_1}{ r_2} = \sqrt{\dfrac{ M_2}{ M_1}} \nonumber \]
where \(r\) is the rate of effusion
The rate of effusion is also directly proportional to the rms speed of the molecules. This is because the only way for the molecule to escape is to "collide" with the opening. Hence, the faster the molecules are moving, the greater the likelihood that a molecule will hit the opening and effuse.
Diffusion and Mean Free Path
Diffusion, like effusion, is faster for light molecules than for heavy ones. The diffusion of gases is much slower than molecular speeds because of molecular collisions. Because of these collisions, the direction of motion of a gas molecule is constantly changing, making this a slow process.
Mean free path: average distance traveled by a molecule between collisions. The higher the density of a gas, the smaller the mean free path
10.9: Real Gases - Deviations from Ideal Behavior
No real gas exhibits ideal gas behavior, although many real gases approximate it over a range of conditions. Gases most closely approximate ideal gas behavior at high temperatures and low pressures. Deviations from ideal gas law behavior can be described by the van der Waals equation, which includes empirical constants to correct for the actual volume of the gaseous molecules and quantify the reduction in pressure due to intermolecular attractive forces.
The ideal gas equation may be rearranged as follows to understand deviations from ideal-gas behavior:
\[\dfrac{PV}{RT} = n \nonumber \]
For a mole of ideal gas (n = 1), the quantity PV / RT = 1 at all pressures. However, real gases do not behave in such a way. At high pressures, the deviation is very high, however it is less with lower pressures. In general, the deviations from idea behavior increase as temperature decreases, becoming significant near the temperature at which the gas is converted into a liquid.
Basic assumptions in the kinetic molecular theory suggest that molecules of an ideal gas occupy no space and have no attraction for each other. Real molecules, however, do have finite volumes, and they do attract one another.
Also, if the volume of the container in which the gas is contained is large, the molecules have plenty of free space, and do not take much of the container’s volume itself. However, as pressure increases, the gas molecules occupy a much larger fraction of the container’s volume.
In addition, the attractive forces between molecules come into play at short distances, as when molecules are crowded together at high pressures. Because of these attractive forces, the impact of a given molecule with the wall of the container is lessened.
Temperature determines how effective attractive forces between gas molecules are. As the gas is cooled, the average kinetic energy decreases, while intermolecular attractions remain constant.
The Van der Walls Equation
According to the ideal gas equation:
\[ \underbrace{P = \dfrac{nRT}{ V}}_{\text{ideal gas}} \nonumber \]
According to Van der Waals:
\[P = \dfrac{nRT}{ V – nb} – \dfrac{n^2a}{V^2} \nonumber \]
Correction for volume of molecules – Correction for molecular attractions
The Van der Waals constant b is a measure of the actual volume occupied by a mole of gas molecules; b has units of L/mol.
The Van der Walls constant a has units of L2-atm/mol2. The magnitude of a reflects how strongly the gas molecules attract each other
Van der Waals equation
\[\left[ P + \left(\dfrac{n^2a}{ V^2}\right) \right] (V – nb) = nRT \nonumber \]
The Van der Waals constants a and b are different for each gas. The values of these constants generally increase with an increase in mass of the molecule and with an increase in the complexity of its structures.
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1Elemental gases
2Etymology
3Physical characteristics
4Macroscopic view of gases
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4.1Pressure
4.2Temperature
4.3Specific volume
4.4Density
5Microscopic view of gases
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5.1Kinetic theory of gases
5.2Thermal motion and statistical mechanics
5.3Brownian motion
5.4Intermolecular forces - the primary difference between Real and Ideal gases
6Mathematical models
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6.1Ideal and perfect gas
6.2Real gas
6.3Permanent gas
7Historical research
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7.1Boyle's law
7.2Charles's law
7.3Gay-Lussac's law
7.4Avogadro's law
7.5Dalton's law
8Special topics
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8.1Compressibility
8.2Reynolds number
8.3Viscosity
8.4Turbulence
8.5Boundary layer
8.6Maximum entropy principle
8.7Thermodynamic equilibrium
9See also
10Notes
11References
12Further reading
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Gas
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From Wikipedia, the free encyclopedia
This is the latest accepted revision, reviewed on 4 March 2024.
State of matter
This article is about the state of matter. For liquified petroleum gas used as an automotive fuel, see autogas. For gasoline ("gas"), see gasoline. For the uses of gases, and other meanings, see Gas (disambiguation).
Drifting smoke particles indicate the movement of the surrounding gas.Gas is one of the four fundamental states of matter. The others are solid, liquid, and plasma.[1]
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, such as air, contains a variety of pure gases. What distinguishes gases 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 gaseous state of matter occurs between the liquid and plasma states,[2] the latter of which provides the upper-temperature boundary for gases. Bounding the lower end of the temperature scale lie degenerative quantum gases[3] which are gaining increasing attention.[4]
High-density atomic gases super-cooled to very low temperatures are classified by their statistical behavior as either Bose gases or Fermi gases. For a comprehensive listing of these exotic states of matter, see list of states of matter.
Elemental gases[edit]
The only chemical elements that are stable diatomic homonuclear molecular gases at STP are hydrogen (H2), nitrogen (N2), oxygen (O2), and two halogens: fluorine (F2) and chlorine (Cl2). When grouped with the monatomic noble gases – helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) – these gases are referred to as "elemental gases".
Etymology[edit]
The word gas was first used by the early 17th-century Flemish chemist Jan Baptist van Helmont.[5] He identified carbon dioxide, the first known gas other than air.[6] 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" (voiceless velar fricative, /x/) – in which case Van Helmont simply was following the established alchemical usage first attested in the works of Paracelsus. According to Paracelsus's terminology, chaos meant something like 'ultra-rarefied water'.[7]
An alternative story is that Van Helmont's term was derived from "gahst (or geist), which signifies a ghost or spirit".[8] That story is given no credence by the editors of the Oxford English Dictionary.[9] In contrast, the French-American historian Jacques Barzun speculated that Van Helmont had borrowed the word from the German Gäscht, meaning the froth resulting from fermentation.[10]
Physical characteristics[edit]
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 § Ideal and perfect gas 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 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.[11][12] A comparison of boiling points for compounds formed by ionic and covalent bonds leads us to this conclusion.[13]
Compared to the other states of matter, gases have low density and viscosity. 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.
Macroscopic view of gases[edit]
Shuttle imagery of re-entry phase
See also: Gas kinetics
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 view.
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[14] 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 behaving ideally.
Pressure[edit]
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.[15] 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.
Temperature[edit]
Air balloon shrinks after submersion in liquid nitrogen
Main article: Thermodynamic temperature 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.[16] In statistical mechanics, temperature is the measure of the average kinetic energy stored in a molecule (also known as the thermal energy). The methods of storing this energy are dictated by the degrees of freedom of the molecule itself (energy modes). Thermal (kinetic) energy added to a gas or liquid (an endothermic process) produces translational, rotational, and vibrational motion. In contrast, a solid can only increase its internal energy by exciting additional vibrational modes, as the crystal lattice structure prevents both translational and rotational motion. These heated gas molecules have a greater speed range (wider distribution of speeds) with a higher average or mean speed. The variance of this distribution is due to the speeds of individual particles constantly varying, due to repeated 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.
Specific volume[edit]
Main article: Specific volume
The symbol used to represent specific volume in equations is "v" with SI units of cubic meters per kilogram.
See also: Gas volume
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.[17] 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. However, volume itself --- not specific --- is an extensive property.
Density[edit]
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 specific volume.
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.
Microscopic view of gases[edit]
Gas-phase particles (atoms, molecules, or ions) move around freely in the absence of an applied electric field.If one could observe a gas under a powerful microscope, one would see a collection of particles without any definite shape or volume that are in more or less random motion. These gas particles only change direction when they collide with another particle or with the sides of the container. This microscopic view of gas is well-described by statistical mechanics, but it can be described by many different theories. The kinetic theory of gases, which makes the assumption that these collisions are perfectly elastic, does not account for intermolecular forces of attraction and repulsion.
Kinetic theory of gases[edit]
Main articles: Kinetic theory of gases and Maxwell–Boltzmann distribution
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,[18] 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 kinetic theory of gases can help explain how the system (the collection of gas particles being considered) responds to changes in temperature, with a corresponding change in kinetic energy.
For example: Imagine you have a sealed container of a fixed-size (a constant volume), containing a fixed-number of gas particles; starting from absolute zero (the theoretical temperature at which atoms or molecules have no thermal energy, i.e. are not moving or vibrating), you begin to add energy to the system by heating the container, so that energy transfers to the particles inside. Once their internal energy is above zero-point energy, meaning their kinetic energy (also known as thermal energy) is non-zero, the gas particles will begin to move around the container. As the box is further heated (as more energy is added), the individual particles increase their average speed as the system's total internal energy increases. The higher average-speed of all the particles leads to a greater rate at which collisions happen (i.e. greater number of collisions per unit of time), between particles and the container, as well as between the particles themselves.
The macroscopic, measurable quantity of pressure, is the direct result of these microscopic particle collisions with the surface, over which, individual molecules exert a small force, each contributing to the total force applied within a specific area. (Read "Pressure" in the above section "Macroscopic view of gases".)
Likewise, the macroscopically measurable quantity of temperature, is a quantification of the overall amount of motion, or kinetic energy that the particles exhibit. (Read "Temperature" in the above section "Macroscopic view of gases".)
Thermal motion and statistical mechanics[edit]
In the kinetic theory of gases, kinetic energy is assumed to purely consist of linear translations according to a speed distribution of particles in the system. However, in real gases and other real substances, the motions which define the kinetic energy of a system (which collectively determine the temperature), are much more complex than simple linear translation due to the more complex structure of molecules, compared to single atoms which act similarly to point-masses. In real thermodynamic systems, quantum phenomena play a large role in determining thermal motions. The random, thermal motions (kinetic energy) in molecules is a combination of a finite set of possible motions including translation, rotation, and vibration. This finite range of possible motions, along with the finite set of molecules in the system, leads to a finite number of microstates within the system; we call the set of all microstates an ensemble. Specific to atomic or molecular systems, we could potentially have three different kinds of ensemble, depending on the situation: microcanonical ensemble, canonical ensemble, or grand canonical ensemble. Specific combinations of microstates within an ensemble are how we truly define macrostate of the system (temperature, pressure, energy, etc.). In order to do that, we must first count all microstates though use of a partition function. The use of statistical mechanics and the partition function is an important tool throughout all of physical chemistry, because it is the key to connection between the microscopic states of a system and the macroscopic variables which we can measure, such as temperature, pressure, heat capacity, internal energy, enthalpy, and entropy, just to name a few. (Read: Partition function Meaning and significance)
Using the partition function to find the energy of a molecule, or system of molecules, can sometimes be approximated by the Equipartition theorem, which greatly-simplifies calculation. However, this method assumes all molecular degrees of freedom are equally populated, and therefore equally utilized for storing energy within the molecule. It would imply that internal energy changes linearly with temperature, which is not the case. This ignores the fact that heat capacity changes with temperature, due to certain degrees of freedom being unreachable (a.k.a. "frozen out") at lower temperatures. As internal energy of molecules increases, so does the ability to store energy within additional degrees of freedom. As more degrees of freedom become available to hold energy, this causes the molar heat capacity of the substance to increase.[19]Random motion of gas particles results in diffusion.
Brownian motion[edit]
Main article: Brownian motionBrownian 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 different from 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 examined.
Intermolecular forces - the primary difference between Real and Ideal gases[edit]
Main articles: van der Waals force, Intermolecular force, and Lennard-Jones potential
Forces between two or more molecules or atoms, either attractive or repulsive, are called intermolecular forces. Intermolecular forces are experienced by molecules when they are within physical proximity of one another. These forces are very important for properly modeling molecular systems, as to accurately predict the microscopic behavior of molecules in any system, and therefore, are necessary for accurately predicting the physical properties of gases (and liquids) across wide variations in physical conditions.
Arising from the study of physical chemistry, one of the most prominent intermolecular forces throughout physics, are van der Waals forces. Van der Waals forces play a key role in determining nearly all physical properties of fluids such as viscosity, flow rate, and gas dynamics (see physical characteristics section). The van der Waals interactions between gas molecules, is the reason why modeling a "real gas" is more mathematically difficult than an "ideal gas". Ignoring these proximity-dependent forces allows a real gas to be treated like an ideal gas, which greatly simplifies calculation.
Isothermal curves depicting the non-ideality of a real gas. The changes in volume (depicted by Z, compressibility factor) which occur as the pressure is varied. The compressibility factor Z, is equal to the ratio Z = PV/nRT. An ideal gas, with compressibility factor Z = 1, is described by the horizontal line where the y-axis is equal to 1. Non-ideality can be described as the deviation of a gas above or below Z = 1.
The intermolecular attractions and repulsions between two gas molecules are dependent on the amount of distance between them. The combined attractions and repulsions are well-modelled by the Lennard-Jones potential, which is one of the most extensively studied of all interatomic potentials describing the potential energy of molecular systems. The Lennard-Jones potential between molecules can be broken down into two separate components: a long-distance attraction due to the London dispersion force, and a short-range repulsion due to electron-electron exchange interaction (which is related to the Pauli exclusion principle).
When two molecules are relatively distant (meaning they have a high potential energy), they experience a weak attracting force, causing them to move toward each other, lowering their potential energy. However, if the molecules are too far away, then they would not experience attractive force of any significance. Additionally, if the molecules get too close then they will collide, and experience a very high repulsive force (modelled by Hard spheres) which is a much stronger force than the attractions, so that any attraction due to proximity is disregarded.
As two molecules approach each other, from a distance that is neither too-far, nor too-close, their attraction increases as the magnitude of their potential energy increases (becoming more negative), and lowers their total internal energy.[20] The attraction causing the molecules to get closer, can only happen if the molecules remain in proximity for the duration of time it takes to physically move closer. Therefore, the attractive forces are strongest when the molecules move at low speeds. This means that the attraction between molecules is significant when gas temperatures is low. However, if you were to isothermally compress this cold gas into a small volume, forcing the molecules into close proximity, and raising the pressure, the repulsions will begin to dominate over the attractions, as the rate at which collisions are happening will increase significantly. Therefore, at low temperatures, and low pressures, attraction is the dominant intermolecular interaction.
If two molecules are moving at high speeds, in arbitrary directions, along non-intersecting paths, then they will not spend enough time in proximity to be affected by the attractive London-dispersion force. If the two molecules collide, they are moving too fast and their kinetic energy will be much greater than any attractive potential energy, so they will only experience repulsion upon colliding. Thus, attractions between molecules can be neglected at high temperatures due to high speeds. At high temperatures, and high pressures, repulsion is the dominant intermolecular interaction.
Accounting for the above stated effects which cause these attractions and repulsions, real gases, delineate from the ideal gas model by the following generalization:[21]
At low temperatures, and low pressures, the volume occupied by a real gas, is less than the volume predicted by the ideal gas law.
At high temperatures, and high pressures, the volume occupied by a real gas, is greater than the volume predicted by the ideal gas law.
Mathematical models[edit]
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.[22] Each successive model expands the temperature range of coverage to which it applies.
Ideal and perfect gas[edit]
Main article: Perfect gas
The equation of state for an ideal or perfect gas is the ideal gas law and reads
P
V
=
n
R
T
,
{\displaystyle PV=nRT,}
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
P
=
ρ
R
s
T
,
{\displaystyle P=\rho R_{s}T,}
where
R
s
{\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 science.
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.[23] It may also be useful to keep the elementary reactions and chemical dissociations for calculating emissions.
Real gas[edit]
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).[24] 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 variable composition.
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.
Permanent gas[edit]
Permanent gas is a term used for a gas which has a critical temperature below the range of normal human-habitable temperatures and therefore cannot be liquefied by pressure within this range. Historically such gases were thought to be impossible to liquefy and would therefore permanently remain in the gaseous state. The term is relevant to ambient temperature storage and transport of gases at high pressure.[25]
Historical research[edit]
See also: Gas laws
See also: Timeline of fluid and continuum mechanics
Boyle's law[edit]
Boyle's equipment
Main article: Boyle's law
Boyle's law was perhaps the first expression of an equation of state. In 1662 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.[26] 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:
P
V
=
k
{\displaystyle PV=k}
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:
P
1
V
1
=
P
2
V
2
.
{\displaystyle \qquad P_{1}V_{1}=P_{2}V_{2}.}
Charles's law[edit]
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:
V
1
T
1
=
V
2
T
2
{\displaystyle {\frac {V_{1}}{T_{1}}}={\frac {V_{2}}{T_{2}}}}
Gay-Lussac's law[edit]
Main article: Gay-Lussac's law
In 1802, Joseph Louis Gay-Lussac published results of similar, though more extensive experiments.[27] 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.
P
1
T
1
=
P
2
T
2
{\displaystyle {\frac {P_{1}}{T_{1}}}={\frac {P_{2}}{T_{2}}}\,}
Avogadro's law[edit]
Main article: Avogadro's law
In 1811, 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 physical constant that bears his name (the Avogadro constant) is the number of atoms per mole 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 amount of substance in the volume. This gives rise to the molar volume of a gas, which at STP is 22.4 dm3/mol (liters per mole). The relation is given by
V
1
n
1
=
V
2
n
2
,
{\displaystyle {\frac {V_{1}}{n_{1}}}={\frac {V_{2}}{n_{2}}},}
where n is the amount of substance of gas (the number of molecules divided by the Avogadro constant).
Dalton's law[edit]
Dalton's notation.
Main article: Dalton's law
In 1801, 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 species as:
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 following:[28]
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).
Special topics[edit]
Compressibility[edit]
Compressibility factors for air.
Main article: Compressibility factor
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.
Reynolds number[edit]
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.
Viscosity[edit]
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).
Turbulence[edit]
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.
Boundary layer[edit]
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[edit]
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 grouping.
Thermodynamic equilibrium[edit]
Main article: Thermodynamic equilibrium
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.
Phase transitions of matter (vte)
ToFrom
Solid
Liquid
Gas
Plasma
Solid
Melting
Sublimation
Liquid
Freezing
Vaporization
Gas
Deposition
Condensation
Ionization
Plasma
Recombination
See also[edit]
Greenhouse gas
List of gases
Natural gas
Volcanic gas
Breathing gas
Wind
Notes[edit]
^ "Gas". Merriam-Webster. 7 August 2023.
^ 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).
^ Zelevinsky, Tanya (2009-11-09). "—just right for forming a Bose-Einstein condensate". Physics. 2 (20): 94. arXiv:0910.0634. doi:10.1103/PhysRevLett.103.200401. PMID 20365964. S2CID 14321276.
^ "Quantum Gas Microscope Offers Glimpse Of Quirky Ultracold Atoms". ScienceDaily. Retrieved 2023-02-06.
^ Helmont, Jan Baptist Van (1652). Ortus medicine, id est initial physicae inaudita... authore Joanne Baptista Van Helmont,... (in Latin). apud L. Elzevirium. 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 "chaos" ...)
^ 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.
^ ""gas, n.1 and adj."". OED Online. Oxford University Press. June 2021. There is probably no foundation in the idea (found from the 18th cent. onwards, e.g. in J. Priestley On Air (1774) Introd. 3) that van Helmont modelled gas on Dutch geest spirit, or any of its cognates
^ Barzun, Jacques (2000). For Dawn to Decadence: 500 Years of Western Cultural Life. New York: HarperCollins Publishers. p. 199.
^ 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. p. 67.
^ Anderson, p.501
^ J. Clerk Maxwell (1904). Theory of Heat. Mineola: Dover Publications. pp. 319–20. ISBN 978-0-486-41735-6.
^ See pages 137–8 of Society, Cornell (1907).
^ Kenneth Wark (1977). Thermodynamics (3 ed.). McGraw-Hill. p. 12. ISBN 978-0-07-068280-1.
^ For assumptions of kinetic theory see McPherson, pp.60–61
^ Jeschke, Gunnar (26 November 2020). "Canonical Ensemble". Archived from the original on 2021-05-20.
^ "Lennard-Jones Potential - Chemistry LibreTexts". 2020-08-22. Archived from the original on 2020-08-22. Retrieved 2021-05-20.
^ "14.11: Real and Ideal Gases - Chemistry LibreTexts". 2021-02-06. Archived from the original on 2021-02-06. Retrieved 2021-05-20.
^ Anderson, pp. 289–291
^ John, p.205
^ John, pp. 247–56
^ "Permanent gas". www.oxfordreference.com. Oxford University Press. Retrieved 3 April 2021.
^ McPherson, pp.52–55
^ McPherson, pp.55–60
^ John P. Millington (1906). John Dalton. pp. 72, 77–78.
References[edit]
Anderson, John D. (1984). Fundamentals of Aerodynamics. McGraw-Hill Higher Education. ISBN 978-0-07-001656-9.
John, James (1984). Gas Dynamics. Allyn and Bacon. ISBN 978-0-205-08014-4.
McPherson, William; Henderson, William (1917). An Elementary study of chemistry.
Further reading[edit]
Look up gas in Wiktionary, the free dictionary.
Wikimedia Commons has media related to Gases.
Philip Hill and Carl Peterson. Mechanics and Thermodynamics of Propulsion: Second Edition Addison-Wesley, 1992. ISBN 0-201-14659-2
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 2008.
Lewes, Vivian Byam; Lunge, Georg (1911). "Gas" . Encyclopædia Britannica. Vol. 11 (11th ed.). p. 481–493.
vteStates of matter (list)State
Solid
Liquid
Gas / Vapor
Supercritical fluid
Plasma
Low energy
Bose–Einstein condensate
Fermionic condensate
Degenerate matter
Quantum Hall
Rydberg matter
Strange matter
Superfluid
Supersolid
Photonic molecule
High energy
QCD matter
Quark–gluon plasma
Color-glass condensate
Other states
Colloid
Crystal
Liquid crystal
Time crystal
Quantum spin liquid
Exotic matter
Programmable matter
Dark matter
Antimatter
Magnetically ordered
Antiferromagnet
Ferrimagnet
Ferromagnet
String-net liquid
Superglass
Transitions
Boiling
Boiling point
Condensation
Critical line
Critical point
Crystallization
Deposition
Evaporation
Flash evaporation
Freezing
Chemical ionization
Ionization
Lambda point
Melting
Melting point
Recombination
Regelation
Saturated fluid
Sublimation
Supercooling
Triple point
Vaporization
Vitrification
Quantities
Enthalpy of fusion
Enthalpy of sublimation
Enthalpy of vaporization
Latent heat
Latent internal energy
Trouton's rule
Volatility
Concepts
Baryonic matter
Binodal
Compressed fluid
Cooling curve
Equation of state
Leidenfrost effect
Macroscopic quantum phenomena
Mpemba effect
Order and disorder (physics)
Spinodal
Superconductivity
Superheated vapor
Superheating
Thermo-dielectric effect
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From Wikipedia, the free encyclopedia
Boiling liquid oxygen
This is a list of gases at standard conditions, which means substances that boil or sublime at or below 25 °C (77 °F) and 1 atm pressure and are reasonably stable.
List[edit]
This list is sorted by boiling point of gases in ascending order, but can be sorted on different values. "sub" and "triple" refer to the sublimation point and the triple point, which are given in the case of a substance that sublimes at 1 atm; "dec" refers to decomposition. "~" means approximately.
Name
Formula
Boiling pt (°C)
Melting pt (°C)
Molecular weight
CAS No
Helium-3
3He
−269.96
N/A
3
14762-55-1
Helium-4
4He
−268.928
N/A
4
7440-59-7
Hydrogen
H2
−252.879
−259.16
2
1333-74-0
Deuterium[1]
D2
−249.49
−254.43
4
7782-39-0
Tritium[2]
T2
−248.12
−254.54
6
10028-17-8
Neon
Ne
−246.046
−248.59
20
7440-01-9
Nitrogen
N2
−195.795
−210.0
28
7727-37-9
Carbon monoxide
CO
−191.5
−205.02
28
630-08-0
Fluorine
F2
−188.11
−219.67
38
7782-41-4
Argon
Ar
−185.848
−189.34
40
7440-37-1
Oxygen
O2
−182.962
−218.79
32
7782-44-7
Methane
CH4
−161.5
−182.50
16
74-82-8
Krypton
Kr
−153.415
−157.37
84
7439-90-9
Nitric oxide
NO
−151.74
−163.6
30
10102-43-9
Oxygen difluoride
F2O
−144.3
−223.8
54
7783-41-7
Tetrafluoromethane
CF4
−127.8
−183.6
88
75-73-0
Nitrogen trifluoride
NF3
−128.74
−206.79
71
7783-54-2
Silane
SiH4
−111.9
−185
32
7803-62-5
trans-Dinitrogen difluoride
N2F2
−111.45
−172
66
13776-62-0
Ozone
O3
−111.35
−193
48
10028-15-6
Xenon
Xe
−108.099
−111.75
131
7440-63-3
cis-Dinitrogen difluoride
N2F2
−105.75
−195
66
13812-43-6
Ethylene
CH2=CH2
−103.7
−169.2
28
74-85-1
Phosphorus trifluoride
PF3
−101.8
−151.5
88
7783-55-3
Chlorine monofluoride
ClF
−101.1
−155.6
54.5
7790-89-8
Boron trifluoride
BF3
−99.9
−126.8
68
7637-07-2
Fluorosilane
SiH3F
−98.6
50
13537-33-2
Trifluorosilane
SiHF3
−95
−131
86
13465-71-9
Trifluoromethyl hypofluorite[3]
CF3OF
−95
−215
104
373-91-1
Diborane
B2H6
−92.49
−164.85
28
19287-45-7
3,3-Difluorodiazirine[4]
CF2N2
−91.3
78
693-85-6
Acetylene
CH≡CH
−84.7
−81.5
26
74-86-2
Ethane
CH3CH3
−88.5
−182.8
30
74-84-0
Germane
GeH4
−88.1
−165
77
7782-65-2
Nitrous oxide
N2O
−88.48
−90.8
44
10024-97-2
Phosphine
PH3
−87.75
−133.8
34
7803-51-2
Trifluoramine oxide
NOF3
−87.5
−161
87
13847-65-9
Tetrafluorosilane
SiF4
−86
−90.2
104
7783-61-1
Trifluoronitrosomethane
CF3NO
−85
−196.6
99
334-99-6
Azidotrifluoromethane
CF3N3
−85
−152
111
3802-95-7
Hydrogen chloride
HCl
−85
−114.17
36.5
7647-01-0
1,1-Difluoroethene
CF2=CH2
−85.5
−144
64
75-38-7
Phosphorus pentafluoride
PF5
−84.6
−93.8
126
7647-19-0
Carbonyl fluoride
COF2
−84.5
−111.2
66
353-50-4
Trifluoromethane
CHF3
−82.1
−155.2
70
75-46-7
Chlorotrifluoromethane
CClF3
−81.5
−181
104.5
75-72-9
Bis(difluoroboryl)methane[3]
BF2CF2BF2
−81.4 ?
148
55124-14-6
Trifluoroisocyanomethane
CF3NC
−80
95
105879-13-8
Difluoromethylborane
CH3BF2
−78.5
64
373-64-8
Carbon dioxide
CO2
−78.464 sub
−56.561 triple
44
124-38-9
Fluoromethane
CH3F
−78.4
−137.8
34
593-53-3
Hexafluoroethane
CF3CF3
−78.1
−100.015
138
76-16-4
Pentafluoromethanamine
CF3NF2
−78
−130
121
335-01-3
Difluorosilane
SiH2F2
−77.8
−122
68
13824-36-7
Tetrafluoroethene
CF2=CF2
−76
−131.14
100
116-14-3
Fluoroacetylene
FCCH
−74
−196
44
2713-09-9
Tetrafluorohydrazine
N2F4
−74
−164.5
104
10036-47-2
Nitryl fluoride
NO2F
−72.4
−166
65
10022-50-1
Fluoroethene
CH2CHF
−72
−160.5
46
75-02-5
Chlorotrifluorosilane
SiClF3
−70
−138
120.5
14049-36-6
Trifluoroacetonitrile
CF3CN
−68.8
95
353-85-5
Chlorodifluoroamine
NClF2
−67
−195
87.5
13637-87-1
Hydrogen bromide
HBr
−66.38
−86.80
81
10035-10-6
Difluorophosphine[5]
PHF2
−65
−124
70
14984-74-8
Borane carbonyl
BH3CO
−64
−137
42
13205-44-2
Fluoroperoxytrifluoromethane[6]
CF3OOF
−64
120
Bis(fluoroxy)difluoromethane[7]
CF2(OF)2
−64
120
16282-67-0
Sulfur hexafluoride
SF6
−63.8
−49.596 triple
146
2551-62-4
Tetrafluorooxirane[citation needed]
C2F4O
−63.5
116
694-17-7
Arsine
AsH3
−62.5
−166
78
7784-42-1
Thiocarbonyl fluoride[8]
CSF2
−62.1
−163.5
82
420-32-6
Radon
Rn
−61.7
−71
222
10043-92-2
Difluorocyanamide[3][9]
NF2CN
−61
−196
78
7127-18-6
Nitrosyl fluoride
ONF
−59.9
−132.5
49
7789-25-5
Hydrogen sulfide
H2S
−59.55
−85.5
34
7783-06-4
Trifluoroacetyl fluoride[10]
CF3COF
−59
−159.5
116
354-34-7
Hexafluorodimethyl ether[11]
CF3OCF3
−59
154
333-36-8
Bromotrifluoromethane
CBrF3
−57.75
−167.78
149
75-63-8
Difluoroaminooxyperfluoromethane[3][12]
CF3ONF2
−57.63
137
4217-93-0
Methylsilane
CH3SiH3
−57.5
−156.5
46
992-94-9
Dioxygen difluoride
F2O2
−57 dec
−163.5
70
7783-44-0
Sulfuryl fluoride
SO2F2
−55.4
−135.8
102
2699-79-8
Dichlorofluorosilane
SiHCl2F
−54.3
119
19382-74-2
trans-1,2-Difluoroethene
CHF=CHF
−53.1
64
1630-78-0
Trifluoroethene
CF2=CHF
−53
82
359-11-5
Arsenic pentafluoride
AsF5
−52.8
−79.8
170
7784-36-3
Phosphorothioic trifluoride
PSF3
−52.25
−148.8
120
2404-52-6
Difluoromethane
CH2F2
−52
−136
52
75-10-5
Difluorocarbamyl fluoride[13][14]
F2NCOF
−52
−152.2
99
2368-32-3
Stannane
SnH4
−51.8
−146
123
2406-52-2
Tetrafluoropropyne[15]
CF3C≡CF
−50.39
112
20174-11-2
Carbonyl sulfide
OCS
−50.2
−138.8
60
463-58-1
Pentafluoroethyl hypofluorite[16]
C2F5OF
−50
154
Chlorodifluorosilane[17]
SiHClF2
−50~
102.5
80003-43-6
Digallane
Ga2H6
−50~
145.494
12140-58-8
Ethenone
CH2=C=O
−49.7
−151
42
463-51-4
Thionyl tetrafluoride
SOF4
−48.5
−99.6
124
13709-54-1
3,3,3-Trifluoro-1-propyne
CF3CCH
−48.3
94
661-54-1
Pentafluoroethane
CF3CHF2
−48.1
−100.6
120
354-33-6
Propene
C3H6
−47.6
−185.2
42
115-07-1
Chlorodifluorophosphine
PClF2
−47.3
−164.8
104.5
14335-40-1
Carbonyl chloride fluoride
COClF
−47.2
−148
82.5
353-49-1
1,1,1-Trifluoroethane
CF3CH3
−47
−111.8
84
420-46-2
Trifluoromethyl hypochlorite[3]
CF3OCl
−47
−164
120.5
22082-78-6
Perchloryl fluoride
ClO3F
−46.75
−147
102.5
7616-94-6
Selenium hexafluoride
SeF6
−46.6 sub
−34.6 triple
193
7783-79-1
Cyanogen fluoride
FCN
−46
−82
45
1495-50-7
Fluorine nitrate
FNO3
−46
−175
81
7789-26-6
Pentafluoronitrosoethane[18]
C2F5NO
−45.7
137
354-72-3
Difluoromethylene dihypofluorite[19]
CF2(OF)2
−45.8
−142
120
16282-67-0
cis-1,2-Difluoroethene
CHF=CHF
−45
64
1630-77-9
1,1-Difluoropropene[3]
CH3CH=CF2
−44
78
430-63-7
Dimethylfluoroborane[citation needed]
(CH3)2BF
−44
60
353-46-8
Fluoro(trifluoromethyl)silane[20]
CF3SiH2F
−44
118
Thionyl fluoride
SOF2
−43.8
−110.5
86
7783-42-8
Phosphorus chloride tetrafluoride
PClF4
−43.4
−132
142.5
13498-11-8
Methyldiborane
CH3B2H5
−43
42
23777-55-1
Difluoro(trifluoromethyl)phosphine[21]
CF3PF2
−43
138
1112-04-5
N,N,1,1-Tetrafluoromethylamine[22]
CHF2NF2
−43
103
24708-53-0
Propane
C3H8
−42.25
−187.7
44
74-98-6
Trifluoro(trifluoromethyl)silane[23]
CF3SiF3
−42
154
335-06-8
Bromotrifluorosilane
BrSiF3
−41.7
−70.5
169
14049-39-9
Hydrogen selenide
H2Se
−41.25
−65.73
81
7783-07-5
Chlorodifluoromethane
CHClF2
−40.7
−175.42
86.5
75-45-6
Sulfur tetrafluoride
SF4
−40.45
−125
108
7783-60-0
Hexafluorodiazomethane cis[24]
CF3NNCF3
−40
−127
166
73513-59-4
Phosphoryl fluoride
POF3
−39.7 sub
−39.1 triple
104
13478-20-1
Chloropentafluoroethane
CF3CClF2
−39.1
−99
135.5
76-15-3
Tetrafluoro(trifluoromethyl)phosphorane[25]
CF3PF4
−39
−113
176
1184-81-2
tetrafluorophosphorane[26]
PHF4
−39.0
−100
108
13659-66-0
Tellurium hexafluoride
TeF6
−38.9
−37.6 triple
242
7783-80-4
Vinyldifluoroborane[27]
CH2=CHBF2
−38.8
−133.4
76
(Trifluoromethyl)silane
CF3SiH3
−38.3
−124
100
10112-11-5
Heptafluoroethanamine[3]
C2F5NF2
−38.1
−183
171
354-80-3
Tetrafluoroallene[28]
CF2=C=CF2
−38
112
461-68-7
Hexafluorooxetane[29]
C3F6O
−38
166
425-82-1
Trifluoromethanethiol[30]
CF3SH
−37.99
−157.11
102
1493-15-8
Fluoroethane
CH3CH2F
−37.7
−143.2
48
353-36-6
Bis(trifluoromethyl)peroxide
CF3OOCF3
−37
170
927-84-4
Pentafluoropropanenitrile[31]
C2F5CN
−37
145
422-04-8
Perfluorodimethylamine[32]
(CF3)2NF
−37
171
359-62-6
Octafluoropropane
C3F8
−36.8
−147.7
188
76-19-7
Germanium tetrafluoride
GeF4
−36.5
−15 triple
149
7783-58-6
Cyclopropene
C3H4
−36
40
2781-85-3
Trifluoromethyl fluoroformate[33]
CF3OC(O)F
−36
−120
132
3299-24-9
Trifluoromethyl isocyanate[34]
CF3NCO
−36
111
460-49-1
Tetrafluoro-1,2-diazetidine
C2F4N2H2
−36
130
Hydrogen iodide
HI
−35.5
−50.76
128
10034-85-2
Pentafluorosulfur hypofluorite
SOF6
−35.1
−86
162
15179-32-5
Difluoromethoxy(trifluoro)methane
CF3OCHF2
−35.0
−157
136
3822-68-2
Propadiene
CH2=C=CH2
−34.8
−136
40
463-49-0
Chlorine
Cl2
−34.04
−101.5
71
7782-50-5
Trifluoromethyl fluoroformate[3][35]
CF3OC(O)F
−34
132
3299-24-9
Diboron tetrafluoride
B2F4
−34
−56
98
13965-73-6
Ammonia
NH3
−33.33
−77.73
17
7664-41-7
Hexafluorocyclopropane[36]
-CF2CF2CF2-
−33
−80
150
931-91-9
Trifluoronitromethane[37]
CF3NO2
−32
115
335-02-4
Dichlorodifluorosilane
SiCl2F2
−32
−44
137
18356-71-3
(Difluoroamino)difluoroacetonitrile[38]
NF2CF2CN
−32
128
5131-88-4
Hexafluoromethanediamine[3][39]
(NF2)2CF2
−31.9
−161.9
154
4394-93-8
Bis(trifluoromethyl)diazene trans[3]
CF3NNCF3
−31.1
166
372-63-4
Cyclopropane
C3H6
−31
−127.6
42
75-19-4
Chlorosilane
SiH3Cl
−30.4
−118
66.5
13465-78-6
Hexafluoropropylene
CF2=CFCF3
−30.2
−156.6
150
116-15-4
Chloroacetylene
HCCCl
−30
−126
60.5
593-63-5
Methyltrifluorosilane
CH3SiF3
−30
−73
100
373-74-0
Fluorine azide[40]
FN3
−30
−139
61.019
14986-60-8
Dichlorodifluoromethane
CCl2F2
−29.8
−157.7
121
75-71-8
2,3,3,3-Tetrafluoropropene[41]
CF3CF=CH2
−29.5
−152.2
114
754-12-1
Tetrafluorodiaziridine[3]
CF4N2
−29
116
17224-09-8
fluoroxypentafluoroselenium[42]
F5SeOF
−29
209
[43]
Perfluorooxetane
C3OF6
−28.6
−117
166
425-82-1
Chlorotrifluoroethene
CClF=CF2
−28.3
−158.14
116.5
79-38-9
Methyldifluorophosphine
CH3PF2
−28
−110
84
753-59-3
Perfluoroacetone
CF3COCF3
−27.4
−125.45
166
684-16-2
Trifluoro(trifluoromethyl)oxirane
C2OF3CF3
−27.4
−144
166
428-59-1
Thiazyl trifluoride
N≡SF3
−27.1
−72.6
103
15930-75-3
Trifluoroacetyl chloride
CF3COCl
−27
−146
132.5
354-32-5
3,3,3-Trifluoropropene
CF3CH=CH2
−27
96
677-21-4
Phosphonium chloride
PH4Cl
−27 sub
70.5
24567-53-1
Formyl fluoride
HCOF
−26.5
−142.2
48
1493-02-3
1,1,1,2-Tetrafluoroethane
CF3CH2F
−26.1
−103.296
102
811-97-2
Trifluoromethyl trifluorovinyl ether[3]
CF3OCF=CF2
−26
166
5930-63-2
Methyl trifluoromethyl ether
CF3OCH3
−25.2
−149.1
100
421-14-7
Bis(trifluoromethyl)nitroxide[44]
(CF3)2NO
−25
−70
168
2154-71-4
Sulfur cyanide pentafluoride[45]
SF5CN
−25
−107
153
1512-13-6[46]
Dimethyl ether
CH3OCH3
−24.8
−141.49
46
115-10-6
1,1,1,4,4,4-Hexafluoro-2-butyne
CF3C≡CCF3
−24.6
−117.4
162
692-50-2
1-Chloro-1-fluoroethene[3]
CClF=CH2
−24.1
80.5
2317-91-1
1,1-Difluoroethane
CHF2CH3
−24.05
−118.6
66
75-37-6
2-Fluoropropene[47]
CH3CF=CH2
−24
60
1184-60-7
Borirane
C2H4BH
−24
−129
40
39517-80-1
Chloromethane
CH3Cl
−23.8
−97.4
50.5
74-87-3
Trifluoronitrosoethylene[48]
CF2=CFNO
−23.7
111
2713-04-4
Pentafluoro(trifluoromethoxy)ethane[49]
C2F5OCF3
−23.6
204
665-16-7
1,1-Difluorocyclopropane[50]
CF2CH2CH2
−23.5
78
558-29-2
Propyne or methylacetylene
CH3CCH
−23.2
−103.0
40
74-99-7
Diazomethane
CH2N2
−23
−145
42
334-88-3
Methylgermane
CH3GeH3
−23
−158
91
1449-65-6
Difluoramine [de]
NHF2
−23
−116
53
10405-27-3
Prop-1-en-1-one or methylketene
CH3CH=CO
−23
−80
56
6004-44-0
Vinylsilane
CH2=CHSiH3
−22.8
58
7291-09-0
Trifluoroiodomethane
CF3I
−22.5
−110
196
2314-97-8
Ethynylsilane
HC≡CSiH3
−22.5
56
1066-27-9
Hexafluoro-1,3-dioxolane[51]
c-CF2OCF2OCF2-
−22.1
182.02
21297-65-4
Chloromethane sulfonyl chloride[3]
CH2ClS(O)(O)Cl
−22
149
3518-65-8
Trifluoromethyl peroxychloride[3]
CF3OOCl
−22
−132
136.5
32755-26-3
Carbonyl selenide
COSe
−21.7
−124.4
107
1603-84-5
Trifluoromethanesulfonyl fluoride
CF3SOF
−21.7
136
335-05-7
Chlorine trifluoride dioxide
ClO2F3
−21.6
−81.2
124.5
38680-84-1
Carbonyl bromide fluoride
COBrF
−21
127
753-56-0
Bromopentafluoroethane
C2BrF5
−21
199
354-55-2
Cyanogen
NCCN
−21
−27.83
52
460-19-5
Methoxysilane
CH3OSiH3
−21
−98.5
62
2171-96-2
1,1,3,3,3-Pentafluoropropene
CF2=CHCF3
−21
-153
132
690-27-7
Carbonyl bromide fluoride[3]
CBrFO
−20.6
127
753-56-0
Trifluoromethylsulfur pentafluoride[3]
CF3SF5
−20.4
−87
196
373-80-8
Chlorotrifluorogermane
GeClF3
−20.3
−66.2
165.5
14188-40-0
Trimethylborane
(CH3)3B
−20.2
−159.93
56
593-90-8
Dimethylsilane
(CH3)2SiH2
−20
−150
60
1111-74-6
1,1,2,2-Tetrafluoroethane
CHF2CHF2
−20
−89
66
359-35-3
Formaldehyde
H2CO
−19.1
−92
30
50-00-0
Hexafluorodisilane
SiF3SiF3
−19.1
−18.7 triple
170
13830-68-7
Sulfur chloride pentafluoride
SClF5
−19.05
−64
158.5
13780-57-9
1-Chloro-2,2-difluoroethene
CHCl=CF2
−18.8
−138.5
98.5
359-10-4
E-1,2,3,3,3-Pentafluoropropene
CFH=CFCF3
−18.7
132
5595-10-8
1,1,1,2,2-Pentafluoropropane
CF3CF2CH3
−18
133
1814-88-6
Hexafluoropropene
CF2=CFCF3
−18
−153
150
116-15-4
Fluoral[52]
CF3CHO
−18
98
75-90-1
2-Chloro-1,1-difluoroethylene[53]
CF2=CHCl
−17.7
−138.5
98
359-10-4
Difluoroamino sulfur pentafluoride[54]
NF2SF5
−17.5
179
13693-10-2
Stibine
SbH3
−17
−88
125
7803-52-3
1,1,2,2,3,3,3-Heptafluoropropane[55]
CF2HCF2CF3
−17
−148.5
170
2252-84-8
1,1,1,2,3,3,3-Heptafluoropropane
CF3CHFCF3
−16.34
−126.8
170
431-89-0
Phosphorus(III) bromide difluoride
PBrF2
−16.1
−133.8
149
15597-40-7
Methylphosphine
CH3PH2
−16
48
593-54-4
N,N-Difluoromethanamine[3][56]
CH3NF2
−16
−114.8
67
753-58-2
Fluorine perchlorate
FOClO3
−16
−167.3
118.5
10049-03-3
Bis(trifluoromethyl) trioxide[57]
CF3OOOCF3
−16
−138
186
1,3,3,3-Tetrafluoropropene[3]
CF3CH=CHF
−16
−104.53
114
1645-83-6
1-Trifluoromethyl-1,2,2-trifluorocyclopropane[50]
CF3C3H2F3
−15.8
152
Disiloxane
(SiH3)2O
−15.2
−144
78
13597-73-4
cis-1-Fluoropropene
CH3CH=CHF
−15
60
19184-10-2
trans-1-Fluoropropene
CH3CH=CHF
−15?
60
20327-65-5
Nitryl chloride
NO2Cl
−15
−145
81.5
13444-90-1
Chlorazide
ClN3
−15
−100
77.47
13973-88-1
Disilane
Si2H6
−14.8
−129.4
62
1590-87-0
Z-1,2,3,3,3-Pentafluoropropene
CHF=CFCF3
−14.7
132
5528-43-8
Bromodifluoromethane
CHBrF2
−14.6
−145
131
1511-62-2
Chloroethene
CH2=CHCl
−13.8
−153.84
62.5
75-01-4
Monoethylsilane[58]
CH3CH2SiH3
−13.7
−180
60
2814-79-1
Chlorine pentafluoride
ClF5
−13.1
−103
130.5
13637-63-3
Perfluorocyclopropene[59]
-CF=CFCF2-
−13
−60
112
19721-29-0
1,1,1-Trifluoropropane
CF3CH2CH3
−13
98
421-07-8
1-Chloro-1,1,2,2-tetrafluoroethane
CClF2CHF2
−13
−117
135.5
354-25-6
Carboimidic difluoride
CF2NH
−13 dec
−90
65
2712-98-3
Plumbane
PbH4
−13
211
15875-18-0
Methyl nitrite
CH3NO2
−12
−16
61
624-91-9
Trifluoromethylarsine[60]
CF3AsH2
−12
146
420-42-8
1-Chloro-1,2,2,2-tetrafluoroethane
CHClFCF3
−11.96
−199.15
136.5
2837-89-0
Isobutane
(CH3)2CHCH2CH3
−11.7
−159.42
58
75-28-5
Trifluoromethoxy sulfur pentafluoride[61]
CF3OSF5
−11
−143
212
1873-23-0
Thiothionyl fluoride
SSF2
−10.6
−164.6
102
101947-30-2
Sulfur dioxide
SO2
−10.05
−75.5
64
7446-09-5
Pentafluorocyclopropane[59]
-CHFCF2CF2-
−10
−10
132
872-58-2
2-Fluoropropane
CH3CHFCH3
−10
62
420-26-8
Pentafluoroethyl hypochlorite[62]
C2F5OCl
−10±
170.5
22675-67-8
Fluoroformyl sulfurpentafluoride[63]
SF5C(O)F
−10
174
Trifluoromethyl fluoroformyl peroxide[64]
CF3OOC(O)F
−10~
148
16118-40-4
Perfluorodimethoxymethane
CF3OCF2OCF3
−10
−161
220
53772-78-4
1-Chloro-1,1-difluoroethane
CClF2CH3
−9.6
−130.8
100.5
75-68-3
Chlorofluoromethane
CH2ClF
−9.1
−133.0
68.5
593-70-4
Pentafluoroethyl isocyanate[65]
C2F5NCO
−9
157
356-74-1
Bis(trifluoromethyl)chloramine[66]
(CF3)2NCl
−9
187.5
Selenium dioxydifluoride
SeO2F2
−8.4
−99.5
149
14984-81-7
Fluoro(trifluoro-methyl)diazine[3]
CF4N2O
−7.63
132
815-10-1
Isobutene
(CH3)2C=CH2
−7.0
−140.7
56
115-11-7
3-Fluoropropene
CH2=CHCH2F
−7
60
818-92-8
Bis(trifluromethyl)amine[3]
(CF3)2NH
−7
153
371-77-7
Ethoxytrifluorosilane[67]
CH3H2OSiF3
−7
−122
118
460-55-9
Trifluoromethylsulfur trifluoride[68]
CF3SF3
−7
−110
158
374-10-7
Perfluoro-2-methyl-1,2-oxazetidine[69]
(CF3-N)CF2CF2O
−6.8
199
Tris(trifluoromethyl)-amine[32]
(CF3)3N
−6.5
221
432-03-1
Methylamine
CH3NH2
−6.4
−93.42
31
74-89-5
1-Butene
CH2=CHCH2CH3
−6.3
−185.33
56
106-98-9
Diphosphorus tetrafluoride
P2F4
−6.2
−86.5
138
13824-74-3
Chloryl fluoride
ClO2F
−6
−115
86.5
13637-83-7
Trifluoromethyl iminosulfur difluoride[70]
CF3N=SF2
−6
153
1512-14-7
Perfluorocyclobutane
-CF2CF2CF2CF2-
−5.91
−40.16
200
115-25-3
Perfluoro-2-butene
CF3CCF=CF3
−5.9
200
360-89-4 trans
Nitrosyl chloride
ONCl
−5.55
−59.4
65.5
2696-92-6
Difluorocarbamoylchloride
CClF2NO
−5.5
115.5
16847-30-6
Hexafluoro 1,2-dioxolane[71]
-CF2CF2CF2OO-
−5
−115.5
182.02
2,3,3,3-tetrafluoro-2-(trifluoromethyl)-propanenitrile[72]
N≡CCF(CF3)2
−4.7
−118
195.038
42532-60-5
1,3-Butadiene
CH2=CHCH=CH2
−4.6
−108.9
54
106-99-0
Ethyltrifluorosilane[73]
CH3CH2SiF3
−4.4
−105
114
353-89-9
Difluoro-N-fluoromethanimine
F2C=NF
−4
83
338-66-9
1,1-Dimethyldiborane
(CH3)2B(μ-H)2BH2
−4
−150.2
56
16924-32-6
trans-1-Chloro-2-fluoroethene[74]
CHCl=CHF
−4
80.5
2268-32-8
Bromochlorodifluoromethane
CF2ClBr
−3.7
−159.5
165.5
353-59-3
N-Nitroso-bis(trifluoromethyl)amine[75]
ONN(CF3)2
−3.5
182
Trifluoromethyl 1,1,2,2-tetrafluoroethyl ether[51]
CF3OCF2CF2H
−3.3
−141
186
2356-61-8
1-Fluoropropane[3]
CH2FCH2CH3
−3.2
−159
62
460-13-9
3-Fluoropropene[47]
CH2CHCH2F
−3
60
818-92-8
Dimethylperoxide
CH3OOCH3
−3
−100
62
690-02-8
Trifluoromethyl thionitrite[76]
CF3SNO
−3
131
Dichlorodifluorogermane
GeCl2F2
−2.8
−51.8
182
24422-21-7
Bromotrifluoroethene
CBrF=CF2
−2.5
161
598-73-2
Trifluoromethane sulfinyl fluoride[3]
CF3SOF
−2.5
136
812-12-4
Perfluorobutane
C4F10
−2.1
−129
238
355-25-9
Hydrogen telluride
H2Te
−2
−49
130
7783-09-7
1-Chloroheptafluoropropane[77]
CF3CF2CF2Cl
−2
204.5
422-86-6
2-Chloroheptafluoropropane[77]
CF3CFClCF3
−2
204.5
76-18-6
Bis(trifluoromethyl)selenium[78]
(CF3)2Se
−2
217
371-79-9
Trifluoromethyl sulfinyl fluoride[68]
CF3S(O)F
−1.6
136
812-12-4
1,1,1,2,2,3-Hexafluoropropane
CF3CF2CFH2
−1.44
−98.38
152
677-56-5
1,1,1,3,3,3-Hexafluoropropane
CF3CH2CF3
−1.4
−93.6
152
690-39-1
Pentafluoroguanidine[79]
CF5N3
−1
149
10051-06-6
1,1,2,2-Tetrafluoropropane[80]
CHF2CF2CH3
−0.8
−121.1
116
40723-63-5
Heptafluoronitrosopropane[18]
C3F7NO
−0.7
199
354-72-3
Trifluoromethanesulfenylchloride[3]
CF3SCl
−0.6
136.5
421-17-0
1,1,1,2-Tetrafluoropropane
CF3CHFCH3
−1-0
116
421-48-7
1,1,2,2,3,3-Hexafluoropropane[3]
CHF2CF2CHF2
−0.3
−98.38
152
680-00-2
Butane
C4H10
0
−140
58
106-97-8
2,2-Difluoropropane
CH3CF2CH3
0
−104.8
80
420-45-1
Perfluoroisobutane
C4F10
0
238
354-92-7
Nitrosyl bromide
NOBr
0
−56
110
13444-87-6
Xenon tetroxide
XeO4
0 dec
−35.9
195
12340-14-6
Trifluoromethylsulfonyl hypofluorite[81]
CF3SO2OF
0
−87
168
Trifluoromethyl chloroformate[82]
CF3OC(O)Cl
0
148.5
23213-83-4
Decafluorodiethyl ether perfluoro ether[83]
CF3CF2OCF2CF3
0
254
358-21-4
Perfluorocyclobutanone[84]
-CF2CF2CF2C(O)-
0
178
699-35-4
Trifluoromethyl peroxonitrate[3]
CF3OON2
0.4
129
50311-48-3
Thiazyl fluoride[85]
NSF
0.4
−89
65.07
18820-63-8
Perfluorotetrahydrofuran
-OCF2CF2CF2CF2-
0.6
−85
216
773-14-8
Tetrafluoro(trifluoromethylimino)-λ6-sulfane[86]
F4S=NCF3
0.75±0.25
191
trans-2-Butene
CH3CH=CHCH3
0.9
−43.3
56
624-64-6
Methylcyclopropane
CH3CHCH2CH2
1
−177.2
56
594-11-6
Bis(trifluoromethyl)phosphine[87]
(CF3)2PH
1
170
Oxalyl fluoride[13]
CFOCFO
1±1
−12.42
94
359-40-0
Methylstannane
CH3SnH3
1.4
137
1631-78-3
Azomethane
CH3N=NCH3
1.5
−78
58
503-28-6
1,1,2,3,3-Pentafluoropropene
CF2=CFCHF2
1.5
−101.2
132
433-66-9
Bromosilane
SiH3Br
1.9
−94
111
13465-73-1
Methylarsine
CH3AsH2
2
−143
92
593-52-2
Hexafluorocyclobutene
C4F6
2
−60
162
697-11-0
Chlorine monoxide
Cl2O
2.2
−120.6
87
7791-21-1
Cyclobutene
C4H6
2.5
54
822-35-5
Difluorodimethylsilane
(CH3)2SiF2
2.5
−87.5
96
353-66-2
1,1,1-Trifluoroazomethane[3]
CF3NNCH3
2.5
112
690-21-1
dichorotrifluorophosphorane[26]
PCl2F3
2.5
−124
159
13659-65-9
Trimethylamine
(CH3)3N
2.8
−117.1
59
75-50-3
1,1-Dichloro-1,2,2,2-tetrafluoroethane
CCl2CF3
3
−56.6
152
374-07-2
Sulfur bromide pentafluoride
SBrF5
3.1
−79
207
15607-89-3
1,1,2-Trifluoroethane
CHF2CH2F
3.5
−84
84
430-66-0
1,2-Dichloro-1,1,2,2-tetrafluoroethane
CClF2CClF2
3.6
−92.52
171
76-14-2
cis-2-Butene
CH3CH=CHCH3
3.72
−138.9
56
590-18-1
Phosphorus dihydride trifluoride[26]
PH2F3
3.9
−52
90
Bromomethane
CH3Br
4
−93.66
95
74-83-9
1,2-Dimethyldiborane
[(CH3)BH2]2
4
−124.9
56
17156-88-6
Selenium chloride pentafluoride
SeClF5
4.5
−19
209.5
34979-62-9
1,1,4,4-Tetrafluoro-1,3-butadiene[88]
CF2=CFCF=CF2
4.5±
162
407-70-5
Trifluoromethyl phosphorodifluoridate[81]
CF3OP(O)F2
4.6
−96.2
170
39125-43-4
Bromoacetylene
C2HBr
4.7
105
593-61-3
Iodine heptafluoride
IF7
4.8
6.5 triple
250
16921-96-3
Dimethylchloroborane[3]
(CH3)2BCl
4.9
−39.9
76.5
1803-36-7
Perfluoro-1-butene[3]
CF3CF2CF=CF2
5
200
357-26-6
Sulfur pentafluoride cyanate[89]
F5SOCN
5
−60
169
Pentafluorosulfanyl cyanate[86]
F5SOCN
5-5.5
169
1,1,2,3,4,4-Hexafluoro-1,3-butadiene
CF2CFCFCF2
5.4
−132
162
685-63-2
Bis(difluoromethyl) ether
CHF2OCHF2
5.5
118
1691-17-4
Methyl pentafluoroethyl ether
CH3OC2F5
5.6
140
22410-44-2
Tris(difluoroamine)fluoromethane[3]
(NF2)3CF
5.6
−136.9
187
14362-68-6
Perfluoro ethyl methyl ether
C2F5OCF3
5.61
204
1-Bromo-2,2-difluoro-ethylene[90]
CHBr=CF2
5.7
143
Perfluoro-1,3-butadiene[91]
CF2=CFCF=CF2
5.8
−132
162
685-63-2
Methanethiol[3]
CH3SH
5.95
−123
48
74-93-1
1-Buten-3-yne
CH2CHC≡CH
6
54
689-97-4
Methoxyethane
CH3OC2H5
6
−113
60
540-67-0
Methyl vinyl ether
CH3OCH=CH2
6
−122
58
107-25-5
1,1,1-Trifluoro-2-chloroethane
CF3CH2Cl
6.1
−105.5
118.5
75-88-7
1,1,1,2,3,3-Hexafluoropropane
CF3CH2CF3
6.2
152
431-63-0
Phosphorothioic chloride difluoride
PSClF2
6.3
−155.2
136.5
2524-02-9
Perfluoro-2-methoxypropionylfluoride[84]
CF3OCF(CF3)C(O)F
5-8
232
Trimethylsilane
(CH3)3SiH
6.7
−153.9
74
993-07-7
Carbon suboxide
OCCCO
6.8
−111.3
68
504-64-3
2-Chloropentafluoropropene[55]
CF3CCl=CF2
6.8
166.5
2804-50-4
dimethylgermane[92]
CH3GeH2CH3
7.0
104.72
1449-64-5
Perfluoroisobutene
(CF3)2C=CF2
7
−130
200
382-21-8
Pentafluoroethyl trifluorovinyl ether
CF3CF2OCFCF2
7
216
10493-43-3
1,1-Difluoropropane
CHF2CH2CH3
7-8
80
430-61-5
1,1,1,2,4,4,4-Heptafluoro-2-butene[69]
CF3CF=CHCF3
7-8
182
Chloromethylsilane
CH3ClSi
7
−135
78.5
993-00-0
Fluoromethyldifluoroborane[93][94]
CH2FBF2
7
−47
82
Nitryl cyanide[95]
NCNO2
7
−85
72
105879-05-8
Silylgermane[96]
GeH3SiH3
7.0
−119.7
107
13768-63-3
Phosphorus(V) dichloride trifluoride
PCl2F3
7.1
−125
159
13454-99-4
Sulfuryl chloride fluoride
SO2ClF
7.1
−124.7
118.5
13637-84-8
Dimethylamine
(CH3)2NH
7.3
−93
45
124-40-3
3-Chloropentafluoropropene[55]
CF2ClCF=CF2
7.4
166.5
79-47-0
Phosgene
COCl2
7.5
−127.77
99
75-44-5
Chloropentafluoroacetone
CClF2COCF3
7.8
−133
182.5
79-53-8
1-Butyne
CH3CH2C≡CH
8.08
−125.7
54
107-00-6
Dichlorosilane
SiH2Cl2
8.3
−122
101
4109-96-0
trans-1,1,1,4,4,4-hexafluoro-2-butene[97]
CF3CH=CHCF3
8.5
164
407-60-3
2-Bromo-1,1,1,2-tetrafluoroethane[69]
CF3CHFBr
8.65
181
Methyl chlorosilane[3]
CH3SiH2Cl
8.7
−134.1
80.5
993-00-0
Pentafluorosulfanyl hypochlorite[citation needed]
SF5OCl
8.9
178.5
Dichlorofluoromethane
CHCl2F
8.92
−135
103
75-43-4
ethylgermane[98]
CH3CH2GeH3
9.2
104.66
Neopentane
(CH3)4C
9.5
−16.5
72
463-82-1
Trifluoromethylperchlorate
CF3OClO3
9.5
168.5
52003-45-9
1,3-Butadiyne
HC≡CC≡CH
10
−35
50
460-12-8
N-Nitroso-O,N-bis(trifluoromethyl)-hydroxylamine or O-Nitroso-bis(trifluoromethy1)hydroxylamine[99][100]
CF3(CF3O)NNO or (CF3)2NONO
10
198
367-54-4
Ethylene oxide
CH2OCH2
10.4
−112.46
44
75-21-8
1,2-Difluoroethane[3]
CH2FCH2F
10.5
−118.6
66
624-72-6
1,2-Butadiene
CH3CH=C=CH2
11
−136.20
54
590-19-2
Dichloromethylborane
CH3BCl2
11
97
7318-78-7
Chlorine dioxide
ClO2
11
−59
103
10049-04-4
2-Chloro-2,3,3,3-tetrafluoropropanoyl fluoride[101]
CF3CFClC(O)F
11
182.5
28627-00-1
Methyl trifluorovinyl ether[102]
CH3OCF=CF2
11
112
3823-94-7
Trifluoromethyl hydroperoxide[3][103]
CF3OOH
11.3
102
16156-36-8
Methyl trifluoromethyl sulfide[78]
CH3SCF3
11.5
116
421-16-9
Chlorine trifluoride
ClF3
11.75
−76.34
128
7790-91-2
1-bromoheptafluoropropane[77]
CF3CF2CF2Br
12
249
422-85-5
Tert-butylfluoride[104]
(CH3)3CF
12
76
353-61-7
2-Fluoro-2-methylpropane
CH3(CH3)CFCH3
12.1
76
353-61-7
Trichlorofluorosilane
SiCl3F
12.25
153.5
14965-52-7
Chloroethane
CH3CH2Cl
12.27
−138
64.5
75-00-3
Pentafluoroiodoethane[105]
CF3CF2I
12.5
−92
246
354-64-3
Cyclobutane
C4H8
12.5
−90.7
56
287-23-0
1,1-difluoro-N-(pentafluoro-λ6-sulfanyl)methanimine[86]
F5SN=CF2
12.5±0.5
191
2-diazo-1,1,1,3,3,3-hexafluoropropane[106]
(CF3)2CN2
12.5±0.5
178
684-23-1
Silylphosphine[107]
SiH3PH2
12.7
68
14616-47-8
Cyanogen chloride
ClCN
13
−6.55
61.5
506-77-4
trans-1-Bromo-1,2-difluoroethylene[90]
CBrF=CFH
13
143
358-99-6
Trifluoromethyl phosphine[3]
CF3PH2
13.1
102
420-52-0
2,2,2-Trifluorodiazoethane[108]
CF3CHNN
13.2
91
371-67-5
2-Chloro-1,1,1,2-tetrafluoropropane HCFC-244bb[109]
CF3CClFCH3
13.23
150.5
421-73-8
Pentafluoroethyl sulfur pentafluoride[110]
C2F5SF5
13.5
246
Phosphorus(III) dichloride fluoride
PCl2F
13.85
−144
121
15597-63-4
2-Chloro-1,1,1,3,3,3-hexafluoropropane[55]
CF3CHClCF3
14
−120.8
186.5
51346-64-6
2-Chloro-3,3,3-trifluoroprop-1-ene[74]
CH2=CClCF3
14
130.5
2730-62-3
Difluoro(difluorochloromethyl)amine
CClF2NF2
14.14
137.5
13880-71-2
Nitrosyl bromide
ONBr
14.5
110
13444-87-6
Hexafluoroisobutylene[3]
(CF3)2C=CH2
14.5
164
382-10-5
N,N-Difluoroethylamine[56]
CH3CH2NF2
14.9
−150.3
81
758-18-9
2-(Pentafluorothio)-3,3-difluorooxaziridine[111]
SF5(-NCF2O-)
14.0
207
73002-62-7
Disulfur difluoride
FSSF
15
−133
102
13709-35-8
1,1,1,3,3-Pentafluoropropane[3]
CF3CH2CHF2
15.14
−102.10
134
460-73-1
1,1,1,2,2,3,3,4,4-Nonafluorobutane[3]
CF3CF2CF2CF2H
12
220
375-17-7
cis-1-Chloro-2-fluoroethene[74]
CHCl=CHF
15
80.5
2268-31-7
(Z)-1-chloro-2,3,3,3-tetrafluoropropene[74]
CHCl=CFCF3
15
148.5
111512-60-8
Trifluoromethyl phosphorodifluoroperoxoate[81]
CF3OOP(O)F2
15.5
−88.6
167
39125-42-3
(Trifluoroacetyl)sulfur pentafluoride[63]
CF3C(O)SF5
15.6
−112
224
82390-51-0
Vinyl bromide
CH2=CHBr
15.8
−137.8
107
593-60-2
Bis(fluorocarbonyl) peroxide[112][13]
CF(O)OOCFO
15.9
−42.5
126
692-74-0
1-Chloro-1,1,2-trifluoroethane
CClF2CH2F
16
118.5
421-04-5
cis-1-Bromo-1,2-difluoroethylene[90]
CBrF=CHF
16
143
1-Fluoro-2-methylpropane[104]
CH2FCHCH3CH3
16
76
359-00-2
Difluoromethyl 1,1,2-trifluoroethyl ether
CHF2OCF2CH2F
16.14
150
69948-24-9
1-Chloro-1-fluoroethane[3]
CHClFCH3
16.15
82.5
1615-75-4
Fluorotrimethylsilane
(CH3)3SiF
16.4
−74.3
92
420-56-4
Bis(trifluoromethyl)nitramine[66]
(CF3)2NNO2
16.4
198
Hexafluoroacetone imine[113]
CF3C(=NH)CF3
16.5±
−47
165
1645-75-6
Dichloro(trifluoromethyl)amine[3]
CF3NCl2
16.6
154
13880-73-4
Ethylamine
CH3CH2NH2
16.6
−81
45
75-04-7
1,1,1-Trifluorobutane[114]
CF3CH2CH2CH3
16.74
−114.79
112
460-34-4
Bismuthine
BiH3
17
−67
212
18288-22-7
2,2,3,3-Tetrafluorobutane[69]
CH3CF2CF2CH3
17
130
tris(trifluoromethyl)phosphine[115]
(CF3)3P
17
238
432-04-2
Tungsten hexafluoride
WF6
17.1
1.9
294
7783-82-6
1-Chloro-1,2,2-trifluoroethane
CHClFCHF2
17.3
99.5
431-07-2
Bis (trifluoromethyl) carbamyl fluoride
(CF3)2NC(O)F
17.5±2.5
199
Ethyl nitrite
C2H5NO2
17.5
75
109-95-5
Tetraborane(10)
B4H10
18
−120
54
18283-93-7
trifluoro-(sulfinylamino)methane[116]
CF3N=S=O
18
131
10564-49-5
F-2,3-dihydro-1,4-dioxin[117]
-CF2CF2OCF=CFO-
18.5
158
Bromofluoromethane
CH2BrF
19 (CRC=23)
113
373-52-4
Bis(trifluoromethyl)arsine[60]
(CF3)2AsH
19
214
1,1-Dichloro-2,2-difluoroethene
CCl2=CF2
19
−116
135
79-35-6
Perfluorovinylsulphur pentafluoride[118]
CF2=CFSF5
19
208
1186-51-2
Difluorotris(trifluoromethyl)phosphorane[115]
(CF3)3PF2
19
276
661-45-0
trans-1-Chloro-3,3,3-trifluoropropene[74]
CHCl=CHCF3
19
130.5
102687-65-0
Fluoroformic acid anhydride[119]
FC(O)OC(O)F
19.2
−46.2
110
177036-04-3
Trifluorovinyl isocyanate[120]
CF2=CFNCO
19.5
123
41594-57-4
Hydrogen fluoride
HF
20
−83.36
20
7664-39-3
Bromine monofluoride
BrF
20 dec
−33
99
13863-59-7
Perbromyl fluoride
BrO3F
20 dec
−110
147
25251-03-0
1-Chloro-1,1,2,2-tetrafluoropropane[121]
CF2ClCF2CH3
20
150.5
421-75-0
Perfluorooxaspiro[2.3]hexane[84]
C5F8O
18-21
228
pentafluoroethylsulfinyl fluoride[68]
C2F5S(O)F
20
186
20621-31-2
3-Methyl-1-butene
CH2=CHCH(CH3)2
20.1
−168.41
70
563-45-1
Acetaldehyde
CH3CHO
20.2
−123.37
44
75-07-0
Chlorotetrafluoro(trifluoromethyl)sulfur
CF3SClF4
20.2
212.5
42179-04-4
trans-Bis(trifluoromethyl)sulfur tetrafluoride[122]
CF3SF4CF3
20.5
246
42179-02-2
Decafluorocyclopentane[3]
C5F10
20.5
250
376-77-2
1,1-Dimethylcyclopropane
(CH3)2CCH2CH2
21
−109.0
70
1630-94-0
Acetyl fluoride
CH3C(O)F
21
−84
62
557-99-3
Perfluoro-N-methyloxazolidine
CF3-NCF2OCF2CF2-
21
237
Bis(trifluoromethyl)cyanoamine[66]
(CF3)2NCN
21
178
Bis(trifluoromethyl)sulfur difluoride[78]
(CF3)2SF2
21
198
30341-38-9
Methaneselenol[123]
CH3SeH
21
95
6486-05-1
Perfluoroethyldimethylamine[32]
C2F5(CF3)2N
21±1
271
815-28-1
cis-1,2-Dichloro-1,2-difluoroethene
CClF=CClF
21.1
−119.6
133
598-88-9
Trifluoromethyl trifluoromethanesulfonate[124]
CF3SO2OCF3
21.2
−108.2
218
3582-05-6
Dinitrogen tetroxide
N2O4
21.15
−9.3
92
10544-72-6
Difluoroiodomethane
CHF2I
21.5
−122
178
1493-03-4
Trifluoromethylcyclopropane[50]
C5H6F3CH3
21.6
110
381-74-8
1,1,1-Trifluoroacetone
CF3C(O)CH3
21.9
−78
112
421-50-1
trans-1,2-Dichloro-1,2-difluoroethene
CClF=CClF
22
−93.3
133
27156-03-2
Heptafluoroisopropyl hypochlorite[62]
(CF3)2CFOCl
22
220.5
22675-68-9
Perfluoroazoethane[34]
C2F5NNC2F5
22
266
1,2,2-Trifluoropropane[125]
CH2FCF2CH3
22
98
811-94-9
Bis(trifluoromethyl)bromamine[66]
(CF3)2NBr
22
232
Bis(trifluoromethyl)sulfde[126]
(CF3)2S
22.2
170
371-78-8
1,1,1,3,3-Pentafluorobutane
CF3CH3CF2CH3
22.6
−34.1
149
406-58-6
octafluoro-1,4-dioxane[127]
(-CF2CF2OCF2CF2O-)
22.75±0.25
232
32981-22-9
Cyanic acid
HNCO
23
−86
43
420-05-3
2-Chloropropene
CH3CCl=CH2
23
−137.4
76.5
557-98-2
1,2,2,2-Tetrafluoroethyl difluoromethyl ether
CF3CHFOCHF2
23
168
57041-67-5
1,1,1,2,3-Pentafluoropropane[128]
CF3CHFCH2F
23
134
431-31-2
2,2,3,3,4,4,5-Heptafluoro oxolane[129]
-CF2CF2CF2CHFO-
23
179
Pentafluoroethyl iminosulfur difluoride[70]
CF3CF2N=SF2
23±1
203
Methoxyacetylene[130]
CH3OC≡CH
23±0.5
56
6443-91-0
Carbonyl fluoride iodide
COFI
23.4
174
1495-48-3
1,2,2,2-Tetrafluoroethyl difluoromethyl ether[3]
CF3CHFOCHF2
23.4
168
57041-67-5
Propylsilane[58]
CH3CH2CH2SiH3
23.5±
78
13154-66-0
Trichlorofluoromethane
CCl3F
23.77
−110.48
137.5
75-69-4
1-Chloro-1,3,3,3-tetrafluoropropene[55]
CF3CH=CFCl
24
148.5
460-71-9
1,1,2,2-Tetrafluoro-1-nitro-2-nitrosoethane[131]
NO2CF2CF2NO
24.2
176
679-08-3
Germyl methyl ether[132]
GeH3OCH3
24.3
107
5910-93-0
Dibromodifluoromethane[3]
CBr2F2
24.45
−141.5
210
75-61-6
Heptafluoro-N-propyl isocyanate[18]
C3F7NCO
24.5
211
87050-96-2
Hexafluorobut-2-yne[69]
CF3C≡CCF3
24.6
−117.4
162
692-50-2
1,1,1,4,4,4-Hexafluorobutane[133]
CF3CH2CH2CF3
24.6
166
407-59-0
Chloroheptafluorocyclobutane
C4ClF7
25
−39.1
216.5
377-41-3
Dimethylphosphine
(CH3)2PH
25
62
676-59-5
Ethyl phosphine[134]
CH3CH2PH2
25
62
593-68-0
Octafluorocyclopentene[135]
C5F8
25
212
559-40-0
2-Fluorobutane[136]
CH3CHFCH2CH3
25
−121
76
359-01-3
Known as gas[edit]
The following list has substances known to be gases, but with an unknown boiling point.
Fluoroamine
Trifluoromethyl trifluoroethyl trioxide CF3OOOCF2CF3 boils between 10 and 20°[137]
Bis-trifluoromethyl carbonate boils between −10 and +10°[35] possibly +12, freezing −60°[138]
Difluorodioxirane boils between −80 and −90°.[139]
Difluoroaminosulfinyl fluoride F2NS(O)F is a gas but decomposes over several hours[140]
Trifluoromethylsulfinyl chloride CF3S(O)Cl[68]
Nitrosyl cyanide ?−20° blue-green gas 4343-68-4[141]
Thiazyl chloride NSCl greenish yellow gas; trimerises.[85]
Possible[edit]
This list includes substances that may be gases. However reliable references are not available.
cis-1-Fluoro-1-propene
trans-1-Chloropropene ?
cis-1-Chloropropene ?
Perfluoro-1,2-butadiene[142]
Perfluoro-1,2,3-butatriene −5[143] polymerizes[144]
Perfluoropent-2-ene
Perfluoropent-1-ene 29-30°[145]
Trifluoromethanesulfenylfluoride CF3SF
Difluorocarbamyl fluoride F2NCOF −52°
N-Sulfinyltrifluoromethaneamine CF3NSO 18°
(Chlorofluoromethyl)silane 373-67-1 274.37 K (1.22 °C)[3]
difluoromethylsilane 420-34-8 237.56 K (−35.59 °C)[3]
trifluoromethyl sulfenic trifloromethyl ester[146]
pentafluoro(penta-fluorethoxy)sulfur 900001-56-6 15°
ethenol 557-75-5 10.5° =vinyl alcohol (tautomerizes)
1,1,1,2,2,3,4,4,4-nonafluorobutane 2-10° melt −129°[55]
trans-2H-Heptafluoro-2-butene
pentafluoroethylhypochlorite around −10°[62]
trifluoromethyl pentafluoroethyl sulfide 6° 33547-10-3 [3]
1,1,1-Trifluoro-N-(trifluoromethoxy)methanamine 671-63-6 0.6°[3]
1-chloro-1,1,2,2,3,3-hexafluoropropane 422-55-9 16.7[3]
1-chloro-1,1,2,3,3,3-hexafluoropropane 359-58-0 17.15[3]
2-chloro-1,1,1,2,3,3-hexafluoropropane 51346-64-6 16.7°[3]
3-chloro-1,1,1,2,2,3-hexafluoropropane 422-57-1 16.7°[3]
trifluormethyl 1,2,2,2-tetrafluoroethyl ether 2356-62-9 11°[3]
2-chloro-1,1,1,3,3-pentafluoropropane HFC-235da 134251-06-2 8°[3]
1,1,2,3,3-pentafluoropropane 24270-66-4 −3.77
2,2,3,3,4,5,5-heptafluoro oxolane[129]
(Heptafluoropropyl)carbonimidic difluoride 378-00-7
Pentafluoroethyl carbonimidic difluoride 428-71-7
(Trifluoromethyl)carbonimidic difluoride 371-71-1 CF3N=CF2
[[perfluoro[n-methyl-(propylenamine)]]] 680-23-9
Perfluoro-N,N-dimethylvinylamine 13821-49-3
[[3,3,4-trifluoro-2,4-bis-trifluoromethyl-[1,2]oxazetidine]] 714-52-3
bis(trifluoromethyl) 2,2-difluoro-vinylamine 13747-23-4
bis(trifluoromethyl) 1,2-difluoro-vinylamine 13747-24-5
1,1,2-Trifluoro-3-(trifluoromethyl)cyclopropane 2967-53-5
bis(trifluoromethyl) 2-fluoro-vinylamine 25211-47-6
2-Fluoro-1,3-butadiene 381-61-3
trifluormethylcyclopropane 381-74-8
cis-1-fluoro-1-butene 66675-34-1
trans-1-fluoro-1-butene 66675-35-2
2-Fluoro-1-butene
3-Fluoro-1-butene
trans-1-fluoro-2-butene
cis-2-fluoro-2-butene
trans-2-fluoro-2-butene
1-fluoro-2-methyl-1-propene
3-fluoro-2-methyl-1-propene
perfluoro-2-methyl-1,3-butadiene 384-04-3
1,1,3,4,4,5,5,5-octafluoro-1,2-pentadiene 21972-01-0
Near misses[edit]
This list includes substances that boil just above standard condition temperatures. Numbers are boiling temperatures in °C.
1,1,2,2,3-Pentafluoropropane 25–26 °C[147][3]
Dimethoxyborane 25.9 °C
1,4-Pentadiene 25.9 °C
2-Bromo-1,1,1-trifluoroethane 26 °C
1,2-Difluoroethane 26 °C
Hydrogen cyanide 26 °C
trimethylgermane 26.2 °C[92]
1,H-Pentafluorocyclobut-1-ene[135]
1,H:2,H-hexafluorocyclobutane[135]
Tetramethylsilane 26.7 °C
Chlorosyl trifluoride 27 °C
2,2-Dichloro-1,1,1-trifluoroethane 27.8 °C
Perfluoroethyl 2,2,2-trifluoroethyl ether 27.89 °C
Perfluoroethyl ethyl ether 28 °C
perfluorocyclopentadiene C5F6 28 °C[148]
2-Butyne 29 °C
Digermane 29 °C
Perfluoroisopropyl methyl ether 29 °C
Trifluoromethanesulfonyl chloride 29–32 °C[149]
Perfluoropentane 29.2 °C
Rhenium(VI) fluoride 33.8 °C
Chlorodimethylsilane 34.7 °C
1,2-Difluoropropane 43 °C
1,3-Difluoropropane 40-42 °C
Dimethylarsine 36 °C
Spiro[2.2]pentane 39 °C
Ruthenium(VIII) oxide 40 °C
Nickel carbonyl 42.1 °C
Trimethylphosphine 43 °C
Unstable substances[edit]
Gallane liquid decomposes at 0 °C.
Nitroxyl and diazene are simple nitrogen compounds known to be gases but they are too unstable and short lived to be condensed.
Methanetellurol CH3TeH 25284-83-7 unstable at room temperature.[150]
Sulfur pentafluoride isocyanide isomerises to sulfur pentafluoride cyanide.[151]
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^ Clifford, A. F.; El-Shamy, H. K.; Emeléus, H. J.; Haszeldine, R. N. (1953). "483. Organometallic and organometalloidal fluorine compounds. Part VIII. The electrochemical fluorination of dimethyl sulphide and carbon disulphide". J. Chem. Soc.: 2372–2375. doi:10.1039/JR9530002372.
^ Amouroux, David; Donard, Olivier F. X. (1996-07-01). "Maritime emission of selenium to the atmosphere in Eastern Mediterranean seas". Geophysical Research Letters. 23 (14): 1777–1780. Bibcode:1996GeoRL..23.1777A. doi:10.1029/96GL01271.
^ Kobayashi, Yoshiro; Yoshida, Tsutomu; Kumadaki, Itsumaro (January 1979). "Trifluoromethyl trifluoromethanesulfonate (CF3SO2OCF3)". Tetrahedron Letters. 20 (40): 3865–3866. doi:10.1016/S0040-4039(01)95545-5.
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^ Gibbon, G. A.; Wang, Jin Tsai; Van Dyke, Charles H. (November 1967). "Preparation and properties of germyl methyl ether and germylmethyl methyl ether". Inorganic Chemistry. 6 (11): 1989–1994. doi:10.1021/ic50057a012. ISSN 0020-1669.
^ Baker, Max T; Ruzicka, Jan A; Tinker, John H (April 1999). "One step synthesis of 1,1,1,4,4,4-hexafluorobutane from succinonitrile". Journal of Fluorine Chemistry. 94 (2): 123–126. doi:10.1016/S0022-1139(98)00311-X.
^ Wagner, Ross I.; Freeman, LeVern D.; Goldwhite, H.; Rowsell, D. G. (March 1967). "Phosphiran". Journal of the American Chemical Society. 89 (5): 1102–1104. doi:10.1021/ja00981a013.
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^ Thrasher, Joseph S.; Madappat, Krishnan V. (September 1989). "Sulfur-Pentafluoride Isocyanide, SF5NC". Angewandte Chemie International Edition in English. 28 (9): 1256–1258. doi:10.1002/anie.198912561.
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Physics & Mathematics
Properties of Matter: Gases
References
By Mary BagleyContributions from Daisy Dobrijevic published 16 February 2022
Gases will fill a container of any size or shape evenly.
Gases are a state of matter with no fixed shape or volume.
(Image credit: MirageC via Getty Images)
Jump to:
Measurable properties of gases
Boyle's law
Charles' law (Gay-Lussac's law)
Avogadro's number
Ideal gas constant
Ideal gas law
Additional resources
Bibliography
Gas is a state of matter that has no fixed shape and no fixed volume. Gases have a lower density than other states of matter, such as solids and liquids. There is a great deal of empty space between particles, which have a lot of kinetic energy and aren’t particularly attracted to one another. Gas particles move very fast and collide with one another, causing them to diffuse, or spread out until they are evenly distributed throughout the volume of the container. According to the educational website Lumen Learning gas can only be contained by either being fully surrounded by a container or held together by gravity.When more gas particles enter a container, there is less space for the particles to spread out, and they become compressed. The particles exert more force on the interior volume of the container. This force is called pressure. There are several units used to express pressure. Some of the most common are atmospheres (atm), pounds per square inch (psi), millimeters of mercury (mmHg) and pascals (Pa). The units relate to one another this way: 1 atm = 14.7 psi = 760 mmHg = 101.3 kPa (1,000 pascals).Related: Greenhouse gases: Causes, sources and environmental effects A gas can be converted to a liquid through compression at a suitable temperature, according to Purdue University. But if the critical temperature is reached, the vapor cannot be liquified regardless of how much pressure is applied. Critical pressure is the pressure needed to liquefy a gas at its critical temperature.Examples of critical temperatures and pressure of different substances according to Engineering Toolbox Measurable properties of gasesBesides pressure, denoted in equations as P, gases have other measurable properties: temperature (T), volume (V) and number of particles, which is expressed in a mole number (n or mol). In work involving gas temperature, the Kelvin scale is often used. Because temperature and pressure vary from place to place, scientists use a standard reference point, called standard temperature and pressure (STP), in calculations and equations. Standard temperature is the freezing point of water — 32 degrees Fahrenheit (0 degrees Celsius, or 273.15 Kelvin). Standard pressure is one atmosphere (atm) — the pressure exerted by the atmosphere on Earth at sea level. Gas lawsTemperature, pressure, amount and volume of a gas are interdependent, and many scientists have developed laws to describe the relationships among them. Boyle's lawChemist Robert Boyle stated that if the temperature is held constant, volume and pressure have an inverse relationship; that is, as volume increases, pressure decreases. This is known as Boyle’s law. (Image credit: GeorgiosArt via Getty Images)Named after Robert Boyle, who first stated it in 1662. Boyle's law states that if the temperature is held constant, volume and pressure have an inverse relationship; that is, as volume increases, pressure decreases, according to the University of California, Davis' ChemWiki.Increasing the amount of space available will allow the gas particles to spread farther apart, but this reduces the number of particles available to collide with the container, so pressure decreases. Decreasing the volume of the container forces the particles to collide more often, so the pressure is increased. A good example of this is when you fill a tire with air. As more air goes in, the gas molecules get packed together, reducing their volume. As long as the temperature stays the same, the pressure increases.Charles' law (Gay-Lussac's law)In 1802, Joseph Louis Gay-Lussac, a French chemist and physicist referenced data gathered by his countryman, Jacque Charles, in a paper describing the direct relationship between the temperature and volume of a gas kept at a constant pressure. Most texts refer to this as Charles' law, but a few call it Gay-Lussac's law, or even the Charles Gay-Lussac law. This law states that the volume and temperature of a gas have a direct relationship: As temperature increases, volume increases when pressure is held constant. Heating a gas increases the kinetic energy of the particles, causing the gas to expand. In order to keep the pressure constant, the volume of the container must be increased when a gas is heated. This law explains why it is an important safety rule that you should never heat a closed container. Increasing temperature without increasing the volume available to accommodate the expanding gas means that pressure builds up inside the container and may cause it to explode. The law also explains why a turkey thermometer pops out when the turkey is done: The volume of air trapped under the plunger increases as the temperature inside the turkey climbs.Joseph Louis Gay-Lussac collects air samples at different heights with Jean-Baptiste Biot in 1804. (Image credit: Luisa Vallon Fumi via Getty Images)Avogadro's numberIn 1811, Italian scientist Amedeo Avogadro proposed the idea that equal volumes of gas at the same temperature and pressure will have an equal number of particles, regardless of their chemical nature and physical properties. Ideal gas constantThe kinetic energy per unit of temperature of one mole of a gas is a constant value, sometimes referred to as the Regnault constant, named after the French chemist Henri Victor Regnault. It is abbreviated by the letter R. Regnault studied the thermal properties of matter and discovered that Boyle's law was not perfect. When the temperature of a substance nears its boiling point, the expansion of the gas particles is not exactly uniform. Ideal gas lawAvogadro's Number, the ideal gas constant, and both Boyle's and Charles' laws combine to describe a theoretical ideal gas in which all particle collisions are absolutely equal. The laws come very close to describing the behavior of most gases, but there are very tiny mathematical deviations due to differences in actual particle size and tiny intermolecular forces in real gases. Nevertheless, these important laws are often combined into one equation known as the ideal gas law. Using this law, you can find the value of any of the other variables — pressure, volume, number or temperature — if you know the value of the other three. Additional resources Learn more about supercritical fluids and their uses with this article from SciMed. For quick children-friendly facts about gases head over to the educational website Love My Science. Discover more examples of gases with this informative material from the educational website Science Notes. BibliographyKnez, Željko, et al. "Industrial applications of supercritical fluids: A review." Energy 77 (2014): 235-243. Silberberg, Martin. Principles of general chemistry. McGraw-Hill Education, 2012. Levy, Sharona T., and Uri Wilensky. "Gas laws and beyond: Strategies in exploring models of the dynamics of change in the gaseous state." annual meeting of the National Association for Research in Science Teaching, San Francisco, CA. 2006.
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Definition and Examples of Gas in Chemistry
Definition and Examples of Gas in Chemistry
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Gas Definition and Examples in Chemistry
Chemistry Glossary Definition of Gas
Water vapor is the gas state of water.
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By
Anne Marie Helmenstine, Ph.D.
Anne Marie Helmenstine, Ph.D.
Chemistry Expert
Ph.D., Biomedical Sciences, University of Tennessee at Knoxville
B.A., Physics and Mathematics, Hastings College
Dr. Helmenstine holds a Ph.D. in biomedical sciences and is a science writer, educator, and consultant. She has taught science courses at the high school, college, and graduate levels.
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Updated on May 04, 2019
A gas is defined as a state of matter consisting of particles that have neither a defined volume nor defined shape. It is one of the four fundamental states of matter, along with solids, liquids, and plasma. Under ordinary conditions, the gas state is between the liquid and plasma states. A gas may consist of atoms of one element (e.g., H2, Ar) or of compounds (e.g., HCl, CO2) or mixtures (e.g., air, natural gas).
Examples of Gases
Whether or not a substance is a gas depends on its temperature and pressure. Examples of gases at standard temperature and pressure include:
air (a mixture of gases)
chlorine at room temperature and pressure
ozone
oxygen
hydrogen
water vapor or steam
List of the Elemental Gases
There are 11 elemental gases (12 if you count ozone). Five are homonuclear molecules, while six are monatomic:
H2 - hydrogenN2 - nitrogenO2 - oxygen (plus O3 is ozone)F2 - fluorineCl2 - chlorineHe - heliumNe - neonAr - argonKr - kryptonXe - xenonRn - radon
Except for hydrogen, which is at the top left side of the periodic table, elemental gases are on the right side of the table.
Properties of Gases
Particles in a gas are widely separated from each other. At low temperature and ordinary pressure, they resemble an "ideal gas" in which the interaction between the particles is negligible and collisions between them are completely elastic. At higher pressures, intermolecular bonds between gas particles have a greater effect on the properties. Because of the space between atoms or molecules, most gases are transparent. A few are faintly colored, such as chlorine and fluorine. Gases tend not to react as much as other states of matter to electric and gravitational fields. Compared with liquids and solids, gases have low viscosity and low density.
Origin of the Word "Gas"
The word "gas" was coined by 17th-century Flemish chemist J.B. van Helmont. There are two theories about the origin of the word. One is that it is Helmont's phonetic transcription of the Greek word Chaos, with the g in Dutch pronounced like the ch in chaos. Paracelsus's alchemical use of "chaos" referred to rarified water. The other theory is that van Helmont took the word from geist or gahst, which means spirit or ghost.
Gas vs Plasma
A gas may contain electrically charged atoms or molecules called ions. In fact, it's common for regions of a gas to contain random, transient charged regions because of van der Waals forces. Ions of like charge repel each other, while ions of opposite charge attract each other. If the fluid consists entirely of charged particles or if the particles are permanently charged, the state of matter is a plasma rather than a gas.
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10.1: Characteristics of Gases - Chemistry LibreTexts
10.1: Characteristics of Gases - Chemistry LibreTexts
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Learning ObjectivesSummary
Learning Objectives
To describe the characteristics of a gas.
The three common phases (or states) of matter are gases, liquids, and solids. Gases have the lowest density of the three, are highly compressible, and completely fill any container in which they are placed. Gases behave this way because their intermolecular forces are relatively weak, so their molecules are constantly moving independently of the other molecules present. Solids, in contrast, are relatively dense, rigid, and incompressible because their intermolecular forces are so strong that the molecules are essentially locked in place. Liquids are relatively dense and incompressible, like solids, but they flow readily to adapt to the shape of their containers, like gases. We can therefore conclude that the sum of the intermolecular forces in liquids are between those of gases and solids. Figure \(\PageIndex{1}\) compares the three states of matter and illustrates the differences at the molecular level.
Figure \(\PageIndex{1}\): A Diatomic Substance (O2) in the Solid, Liquid, and Gaseous States: (a) Solid O2 has a fixed volume and shape, and the molecules are packed tightly together. (b) Liquid O2 conforms to the shape of its container but has a fixed volume; it contains relatively densely packed molecules. (c) Gaseous O2 fills its container completely—regardless of the container’s size or shape—and consists of widely separated molecules.
The state of a given substance depends strongly on conditions. For example, H2O is commonly found in all three states: solid ice, liquid water, and water vapor (its gaseous form). Under most conditions, we encounter water as the liquid that is essential for life; we drink it, cook with it, and bathe in it. When the temperature is cold enough to transform the liquid to ice, we can ski or skate on it, pack it into a snowball or snow cone, and even build dwellings with it. Water vapor (the term vapor refers to the gaseous form of a substance that is a liquid or a solid under normal conditions so nitrogen (N2) and oxygen (O2) are referred to as gases, but gaseous water in the atmosphere is called water vapor) is a component of the air we breathe, and it is produced whenever we heat water for cooking food or making coffee or tea. Water vapor at temperatures greater than 100°C is called steam. Steam is used to drive large machinery, including turbines that generate electricity. The properties of the three states of water are summarized in Table 10.1.
Table \(\PageIndex{1}\): Properties of Water at 1.0 atm
Temperature
State
Density (g/cm3)
≤0°C
solid (ice)
0.9167 (at 0.0°C)
0°C–100°C
liquid (water)
0.9997 (at 4.0°C)
≥100°C
vapor (steam)
0.005476 (at 127°C)
The geometric structure and the physical and chemical properties of atoms, ions, and molecules usually do not depend on their physical state; the individual water molecules in ice, liquid water, and steam, for example, are all identical. In contrast, the macroscopic properties of a substance depend strongly on its physical state, which is determined by intermolecular forces and conditions such as temperature and pressure.
Figure \(\PageIndex{2}\) shows the locations in the periodic table of those elements that are commonly found in the gaseous, liquid, and solid states. Except for hydrogen, the elements that occur naturally as gases are on the right side of the periodic table. Of these, all the noble gases (group 18) are monatomic gases, whereas the other gaseous elements are diatomic molecules (H2, N2, O2, F2, and Cl2). Oxygen can also form a second allotrope, the highly reactive triatomic molecule ozone (O3), which is also a gas. In contrast, bromine (as Br2) and mercury (Hg) are liquids under normal conditions (25°C and 1.0 atm, commonly referred to as “room temperature and pressure”). Gallium (Ga), which melts at only 29.76°C, can be converted to a liquid simply by holding a container of it in your hand or keeping it in a non-air-conditioned room on a hot summer day. The rest of the elements are all solids under normal conditions.
Figure \(\PageIndex{2}\): Elements That Occur Naturally as Gases, Liquids, and Solids at 25°C and 1 atm. The noble gases and mercury occur as monatomic species, whereas all other gases and bromine are diatomic molecules.
Purple elements are gaseous elements, green are liquid, and gray and solid. H, N, O, F, Cl, He, Ne, Ar, Kr, Xe, and Rn are all purple. Br and Hg are green. The rest are gray.
All of the gaseous elements (other than the monatomic noble gases) are molecules. Within the same group (1, 15, 16 and 17), the lightest elements are gases. All gaseous substances are characterized by weak interactions between the constituent molecules or atoms.
Summary
Bulk matter can exist in three states: gas, liquid, and solid. Gases have the lowest density of the three, are highly compressible, and fill their containers completely. Elements that exist as gases at room temperature and pressure are clustered on the right side of the periodic table; they occur as either monatomic gases (the noble gases) or diatomic molecules (some halogens, N2, O2).
10.1: Characteristics of Gases is shared under a CC BY-NC-SA 3.0 license and was authored, remixed, and/or curated by LibreTexts.
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A general chemistry Libretexts Textmap organized around the textbook
Chemistry: The Central Science
by Brown, LeMay, Bursten, Murphy, and Woodward
Previously, we focused on the microscopic properties of matter—the properties of individual atoms, ions, and molecules—and how the electronic structures of atoms and ions determine the stoichiometry and three-dimensional geometry of the compounds they form. We will now focus on macroscopic properties—the behavior of aggregates with large numbers of atoms, ions, or molecules. An understanding of macroscopic properties is central to an understanding of chemistry. Why, for example, are many substances gases under normal pressures and temperatures (1.0 atm, 25°C), whereas others are liquids or solids? We will examine each form of matter—gases, liquids, and solids—as well as the nature of the forces, such as hydrogen bonding and electrostatic interactions, that hold molecular and ionic compounds together in these three states.
In this chapter, we explore the relationships among pressure, temperature, volume, and the amount of gases. You will learn how to use these relationships to describe the physical behavior of a sample of both a pure gaseous substance and mixtures of gases. By the end of this chapter, your understanding of the gas laws and the model used to explain the behavior of gases will allow you to explain how straws and hot-air balloons work, why hand pumps cannot be used in wells beyond a certain depth, why helium-filled balloons deflate so rapidly, and how a gas can be liquefied for use in preserving biological tissue.
10.1: Characteristics of GasesBulk matter can exist in three states: gas, liquid, and solid. Gases have the lowest density of the three, are highly compressible, and fill their containers completely. Elements that exist as gases at room temperature and pressure are clustered on the right side of the periodic table; they occur as either monatomic gases (the noble gases) or diatomic molecules (some halogens, N₂, O₂).10.2: PressurePressure is defined as the force exerted per unit area; it can be measured using a barometer or manometer. Four quantities must be known for a complete physical description of a sample of a gas: temperature, volume, amount, and pressure. Pressure is force per unit area of surface; the SI unit for pressure is the pascal (Pa), defined as 1 newton per square meter (N/m²). The pressure exerted by an object is proportional to the force it exerts and inversely proportional to the area.10.3: The Gas LawsThe volume of a gas is inversely proportional to its pressure and directly proportional to its temperature and the amount of gas. Boyle showed that the volume of a sample of a gas is inversely proportional to pressure (Boyle’s law), Charles and Gay-Lussac demonstrated that the volume of a gas is directly proportional to its temperature at constant pressure (Charles’s law), and Avogadro showed that the volume of a gas is directly proportional to the number of moles of gas (Avogadro’s law).10.4: The Ideal Gas EquationThe empirical relationships among the volume, the temperature, the pressure, and the amount of a gas can be combined into the ideal gas law, PV = nRT. The proportionality constant, R, is called the gas constant. The ideal gas law describes the behavior of an ideal gas, a hypothetical substance whose behavior can be explained quantitatively by the ideal gas law and the kinetic molecular theory of gases. Standard temperature and pressure (STP) is 0°C and 1 atm.10.5: Further Applications of the Ideal-Gas EquationsThe relationship between the amounts of products and reactants in a chemical reaction can be expressed in units of moles or masses of pure substances, of volumes of solutions, or of volumes of gaseous substances. The ideal gas law can be used to calculate the volume of gaseous products or reactants as needed. In the laboratory, gases produced in a reaction are often collected by the displacement of water from filled vessels; the amount of gas can be calculated from the volume of water displaced.10.6: Gas Mixtures and Partial PressuresThe pressure exerted by each gas in a gas mixture is independent of the pressure exerted by all other gases present. Consequently, the total pressure exerted by a mixture of gases is the sum of the partial pressures of the components (Dalton’s law of partial pressures). The amount of gas in a mixture may be described by its partial pressure or its mole fraction. In a mixture, the partial pressure of each gas is the product of the total pressure and the mole fraction.10.7: Kinetic-Molecular TheoryThe behavior of ideal gases is explained by the kinetic molecular theory of gases. Molecular motion, which leads to collisions between molecules and the container walls, explains pressure, and the large intermolecular distances in gases explain their high compressibility. Although all gases have the same average kinetic energy at a given temperature, they do not all possess the same root mean square speed. The actual values of speed and kinetic energy are not the same for all gas particles.10.8: Molecular Effusion and DiffusionDiffusion is the gradual mixing of gases to form a sample of uniform composition even in the absence of mechanical agitation. In contrast, effusion is the escape of a gas from a container through a tiny opening into an evacuated space. The rate of effusion of a gas is inversely proportional to the square root of its molar mass (Graham’s law), a relationship that closely approximates the rate of diffusion. As a result, light gases tend to diffuse and effuse much more rapidly than heavier gases.10.9: Real Gases - Deviations from Ideal BehaviorNo real gas exhibits ideal gas behavior, although many real gases approximate it over a range of conditions. Gases most closely approximate ideal gas behavior at high temperatures and low pressures. Deviations from ideal gas law behavior can be described by the van der Waals equation, which includes empirical constants to correct for the actual volume of the gaseous molecules and quantify the reduction in pressure due to intermolecular attractive forces.10.E: ExercisesThese are homework exercises to accompany the Textmap created for "Chemistry: The Central Science" by Brown et al. Complementary General Chemistry question banks can be found for other Textmaps and can be accessed here. In addition to these publicly available questions, access to private problems bank for use in exams and homework is available to faculty only on an individual basis; please contact Delmar Larsen for an account with access permission.10.S: Gases (Summary)
Thumbnail: Motion of gas molecules. The randomized thermal vibrations of fundamental particles such as atoms and molecules—gives a substance its “kinetic temperature.” Here, the size of helium atoms relative to their spacing is shown to scale under 1950 atmospheres of pressure. (CC BY-SA 3.0; Greg L).
10: Gases is shared under a CC BY-NC-SA 3.0 license and was authored, remixed, and/or curated by LibreTexts.
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9.S: Molecular Geometry and Bonding Theories (Summary)
10.1: Characteristics of Gases
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Gas - Simple English Wikipedia, the free encyclopedia
Gas - Simple English Wikipedia, the free encyclopedia
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Gas
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From Simple English Wikipedia, the free encyclopedia
For the fuel "gas", see Gasoline.
An illustration of the random way gas molecules move, without being attached to each other.
A gas is one of the four states of matter. In a gas, the molecules move freely and are not attached to each other. This makes it different from a liquid where the molecules are loosely attached to or touching each other. It is also different from a solid where the molecular bonds are strong and hold the molecules together in one shape.
A gas does not have only one volume like a liquid or solid does. Instead, gas can expand until it fills whatever container it is in.
In a pure gas, each molecule may be made of an individual atom. It may be elemental, where each molecule is made of more than one of the same atom bound together. It may be compounds where molecules are made of many types of atoms together. An example of a monoatomic gas is neon, an example of an elemental gas is hydrogen and an example of a compound gas is carbon dioxide.
A gas mixture contains a mix of any of the above types, for example air which is 78% nitrogen, 21% oxygen, less than 1% argon, around 0.03% carbon dioxide and more other in very small quantities.[1]
Poison gases were used as chemical weapons in WWI.
Physical characteristics[change | change source]
All gases can flow, like liquids. This means the molecules move about independently of each other. Most gases are colourless, like hydrogen.[2] Gas particles will spread about, or diffuse, in order to fill all the space in any container such as a bottle or a room. Compared to liquids and solids, gases have a very low density and viscosity. We cannot directly see most gases since they aren't coloured. However it is possible to measure their density, volume, temperature and pressure.
Pressure[change | change source]
See the main article: Pressure
Pressure is the measure of how much pushing force something is putting on another object. In a gas, this is usually the gas pushing on the container of the object or, if the gas is heavy, something inside the gas. Pressure is measured in pascals. Because of Newton's third law, we can change the pressure of a gas by putting force on the object containing it. For example, squeezing a bottle with air inside pressurises (gives more pressure) to the air inside.
When talking about gas, pressure is often related to the container. A lot of gas in a small container would have very high pressure. A small amount of a gas in a big container would have low pressure.
Gas can create pressure itself when there is a lot of it. The weight of the gas creates pressure on anything underneath it, including other gas. On a planet, this is called atmospheric pressure.
Temperature[change | change source]
See the main article: temperature
The temperature of a gas is how hot or cold it is. In physics it is usually measured in kelvins although degrees Celsius are used more elsewhere. In a gas, the average velocity (how fast they move) of the molecules is related to the temperature. The faster the gas molecules are moving, the more they collide, or smash into each other. These collisions release energy, which in a gas comes in the form of heat. Conversely if the temperature around the gas becomes hotter then the gas particles will convert the thermal energy to kinetic energy, making them move faster and making the gas hotter.[3]
State changes[change | change source]
A gas can go through two different state changes. If the temperature is low enough the gas can condense and turn into a liquid. Sometimes, if the temperature is low enough it can go through deposition, where it changes straight to a solid. Normally a gas must first condense to a liquid, and then freeze to become a solid, but if the temperature is very low it can skip the liquid stage and instantly become solid. Frost on the ground in winter is caused by this. Water vapour (a gas) goes into the air which is very cold, and instantly becomes ice due to deposition.
Related pages[change | change source]
Ideal gas
References[change | change source]
↑ "Composition of Air". mistupid.com. Retrieved 2024-01-18.
↑ "Colours of Gases". Archived from the original on 2009-10-19. Retrieved 2010-01-09.
↑ "Heat and temperature". Archived from the original on 2010-02-10. Retrieved 2010-01-10.
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Order and disorder (physics)
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Category: Gases
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