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Gas
Gas is a state of matter with neither fixed volume nor fixed shape. It is a compressible form of fluid. A pure gas consists of individual atoms (e.g. a noble gas like neon), or molecules (e.g. oxygen (O2) or carbon dioxide). 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.
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".
The word gas was first used by the early 17th-century Flemish 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 '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'.
An alternative story is that Van Helmont's term was derived from "gahst (or geist), 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.
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. 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.
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Gas
Gas is a state of matter with neither fixed volume nor fixed shape. It is a compressible form of fluid. A pure gas consists of individual atoms (e.g. a noble gas like neon), or molecules (e.g. oxygen (O2) or carbon dioxide). 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.
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".
The word gas was first used by the early 17th-century Flemish 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 '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'.
An alternative story is that Van Helmont's term was derived from "gahst (or geist), 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.
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. 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.