Gas
Gas is a state of matter with neither fixed volume nor fixed shape. It is a compressible form of fluid, in contrast to a liquid. A pure gas consists of individual atoms, or molecules. Pure gases can also be mixed together such as in the air. What distinguishes gases from liquids and solids is the vast separation of the individual gas particles. This separation can make some gases invisible to the human observer.
The gaseous state of matter occurs between the liquid and plasma states, the latter of which provides the upper-temperature boundary for gases. Bounding the lower end of the temperature scale lie degenerative quantum gases which are gaining increasing attention.
High-density atomic gases super-cooled to 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
The only chemical elements that are stable diatomic homonuclear molecular gases at STP are hydrogen, nitrogen, oxygen, and two halogens: fluorine and chlorine. When grouped with the monatomic noble gases - helium, neon, argon, krypton, xenon, and radon - these gases are referred to as "elemental gases".Etymology
The word gas was first used by the early 17th-century Brabantian or Southern Netherlandish chemist Jan Baptist van Helmont. He identified carbon dioxide, the first known gas other than air. Van Helmont's word appears to have been simply a phonetic transcription of the Ancient Greek word χάος – the g in Dutch being pronounced like ch in "loch" – 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.An alternative story is that Van Helmont's term was derived from "gahst, which signifies a ghost or spirit". That story is given no credence by the editors of the Oxford English Dictionary. 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.
Physical characteristics
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 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.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. A comparison of boiling points for compounds formed by ionic and covalent bonds leads us to this conclusion.
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
When observing 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 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 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 or their surroundings. 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. 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,, named to honor his work in this field.
There are many mathematical tools available for analyzing gas properties. Boyle's lab equipment allowed the use of just a simple calculation 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. However, 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. As gases are subjected to extreme conditions, tools to interpret them become more complex, from the Euler equations for inviscid flow to the Navier–Stokes equations that fully account for viscous effects. This advanced math, including statistics and multivariable calculus, adapted to the conditions of the gas system in question, makes it possible to solve 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 situation, the gas is no longer behaving ideally.
Pressure
In a container, gas pressure is 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. The force imparted by a gas particle into the container during this collision is the change in momentum of the particle. During a collision only the normal component of velocity changes. A particle traveling parallel to the wall does not change its momentum. Therefore, the average force on a surface must be the average change in linear momentum from all of these gas particle collisions.Pressure is the sum of all the normal components of force exerted by the particles impacting the walls of the container divided by the surface area of the wall.
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 which make up the system. In statistical mechanics, temperature is the measure of the average kinetic energy stored in a molecule. The methods of storing this energy are dictated by the degrees of freedom of the molecule itself. Thermal energy added to a gas or liquid 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 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
The symbol used to represent specific volume in equations is "v" with SI units of cubic meters per kilogram.The symbol used to represent volume in equations is "V" with SI units of cubic meters.
When performing a thermodynamic analysis, it is typical to speak of intensive and extensive properties. Properties which depend on the amount of gas are called extensive properties, while properties that do not depend on the amount of gas are called intensive properties. Specific volume is an example of an intensive property because it is the ratio of volume occupied by a unit of mass of a gas that is identical throughout a system at equilibrium. 1000 atoms a gas occupy the same space as any other 1000 atoms for any given temperature and pressure. This concept is easier to visualize for solids such as iron which are incompressible compared to gases. However, volume itself --- not specific --- is an extensive property.