Gas

• 1. Introduction

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  Gas is the only state of matter that holds neither a fixed volume nor a fixed shape, filling whatever container surrounds it. A pure gas can be single atoms, as in a noble gas like neon, or whole molecules, as in oxygen or carbon dioxide. Mix several pure gases together and you get something familiar: the air. What sets gas apart from liquids and solids is the vast separation between its individual particles. That separation can be so wide that some gases vanish from human sight entirely. Yet this almost-nothing has been weighed, named, and bottled. Who first gave it a name? How did glass tubes filled with mercury reveal its hidden laws? And why does a gas that behaves perfectly in a jet engine misbehave during a spacecraft's fiery return to Earth? The answers begin with a single word coined in the early 17th century.

• 2. Naming The Invisible

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  Jan Baptist van Helmont, a chemist from Brabant in the Southern Netherlands, gave the world the word "gas" in the early 17th century. He also identified carbon dioxide, the first gas known to be something other than air. Van Helmont's coinage appears to be a phonetic transcription of the Ancient Greek word chaos. The Dutch g is pronounced like the ch in "loch," a sound called a voiceless velar fricative. In choosing it, Van Helmont followed an alchemical usage first attested in the works of Paracelsus, for whom chaos meant something elusive and formless. A rival explanation traces the word to "gahst," or geist, signifying a ghost or spirit, but the editors of the Oxford English Dictionary give that story no credence. The French-American historian Jacques Barzun offered yet another idea: that Van Helmont borrowed the German word Gascht, the froth thrown up by fermentation. From that single ambiguous syllable grew a whole science of measurement.

• 3. Four Numbers

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  Pressure, volume, number of particles, and temperature are the four properties through which gas is understood, since most gases resist direct observation. Chemists count the particles in units called moles. These four characteristics were measured again and again by scientists including Robert Boyle, Jacques Charles, John Dalton, Joseph Gay-Lussac, and Amedeo Avogadro, across many gases and many settings. Pressure, written p or P, has SI units of pascals and is the average force per unit area a gas exerts on its container. Picture the particles flying in straight lines until they strike the walls. Only the component of velocity perpendicular to the wall changes during a collision, so a particle skimming parallel to the surface imparts nothing. Temperature, written T in kelvins, tracks the motion of the particles themselves. The speed of a gas particle is proportional to its absolute temperature, which is why a balloon cooled with extremely cold nitrogen visibly shrinks as its trapped particles slow. In statistical mechanics, temperature is the measure of the average kinetic energy stored in a molecule. Heat a gas and its molecules occupy a wider range of speeds, described by the Maxwell-Boltzmann distribution.

• 4. Inside The Particle Storm

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  Under a powerful enough microscope, a gas would reveal a swarm of particles with no definite shape or volume, moving in more or less random motion. Each particle changes direction only when it collides with another particle or with the container wall. The kinetic theory of gases describes this world by assuming those collisions are perfectly elastic. That same theory ignores the forces of attraction and repulsion that real molecules feel for one another. Starting from a sealed box at absolute zero, where atoms have no thermal energy at all, adding heat lifts the particles above their zero-point energy and sets them moving. As more energy enters, average speed rises and collisions grow more frequent, both between particles and against the walls. The measurable pressure is simply the sum of countless tiny forces from those wall collisions. The motions themselves are richer than straight-line travel. Kinetic energy in real molecules combines translation, rotation, and vibration, which together produce a finite number of microstates that physicists group into an ensemble. Counting those microstates through a partition function is the key tool that links the microscopic world to measurable quantities like heat capacity, internal energy, enthalpy, and entropy.

• 5. When Molecules Touch

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  Van der Waals forces, the attractions and repulsions between molecules in close proximity, are the real reason a true gas defies simple math. They shape nearly every physical property of a fluid, including viscosity, flow rate, and gas dynamics. Ignore them and a real gas can be treated like an ideal one, which greatly simplifies the calculation. The Lennard-Jones potential models how those forces shift with distance, and it is one of the most studied of all interatomic potentials. So central is it that the model system earned its own nickname, Lennard-Jonesium. The potential splits into two parts: a long-range pull from the London dispersion force, and a short-range shove from electron-electron exchange, tied to the Pauli exclusion principle. Distance decides which force wins. Slow-moving molecules at low temperature linger near each other long enough for attraction to take hold. Compress that cold gas into a small volume and collisions multiply until repulsion dominates. The result is a clear rule. At low temperatures and low pressures, a real gas occupies less volume than the ideal gas law predicts; at high temperatures and high pressures, it occupies more.

• 6. Models Of An Ideal World

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  An equation of state is a mathematical model meant to predict the state properties of a gas, and no single one works for all gases under all conditions. The three most discussed models are the perfect gas, the ideal gas, and the real gas, each with its own assumptions and each widening the range of conditions it can handle. The ideal gas law reads PV equals nRT, where n is the amount of gas in moles, R is the universal gas constant of 8.314 joules per mole-kelvin, and T is temperature. This is sometimes called the chemist's version, since it counts molecules directly. A second form, favored by gas dynamicists, swaps in density and a specific gas constant for modeling accelerating flows without chemical reactions. An ideal gas is a simplified real gas whose compressibility factor Z is fixed at one. Inside the combustion chamber of a jet engine, that approximation works well. But push toward the upper end of engine temperatures, near 1300 kelvin in a combustor, and complex fuel particles store energy in rotations and vibrations that bend their specific heats away from the simple cases. Past roughly double that temperature, electronic excitation and dissociation begin, and the gas starts its transition into plasma.

• 7. Laws Written In Mercury

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  In 1662, Robert Boyle ran a series of experiments with a J-shaped glass tube sealed at one end. Mercury trapped a fixed quantity of air in the short end, and as he poured in more, the gas volume shrank. He read the pressure from the difference in mercury levels between the short sealed end and the long open one. Boyle found that at constant temperature, pressure varies inversely with volume: halve the volume and the pressure doubles. Multiply pressure by volume and the product stays a constant, the relationship now named Boyle's law. Other measurers followed across the decades. In 1787, the French physicist and balloon pioneer Jacques Charles found that oxygen, nitrogen, hydrogen, carbon dioxide, and air all expand to the same extent over the same 80-kelvin interval, showing volume rises in direct proportion to temperature. In 1802, Joseph Louis Gay-Lussac published more extensive experiments and credited Charles by naming the law in his honor; in 1809 he added his own law linking pressure to temperature. John Dalton, in 1801, published the law of partial pressures, finding the pressure of a mix of non-reactive gases equals the sum of each gas alone. Avogadro's law arrived in 1811, when Amedeo Avogadro verified that equal volumes of pure gases hold the same number of particles.

• 8. Where Ideal Gas Fails

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  The 1990 eruption of Mount Redoubt sent up gases of exactly the kind the ideal gas law cannot describe, formed under extreme geological conditions. Real gas effects gather wherever pressure climbs and intermolecular forces grow too strong to ignore. Engineers list the corrections plainly: compressibility, with Z allowed to drift from one; heat capacities that vary with temperature; van der Waals forces; non-equilibrium effects; and molecular dissociation with changing composition. The Space Shuttle's re-entry is the textbook case, where extreme temperatures and pressures strip the ideal model of any meaning. There the surrounding gas no longer behaves ideally, and the math must climb from the Euler equations for inviscid flow up to the Navier-Stokes equations that fully account for viscous effects. Other phenomena ride along with that complexity. Particles stick to a moving surface to form a boundary layer, visible as the thickening seen along the leading edge of a delta wing. When that layer separates, it reshapes the entire flow, as in a stalling airfoil. Some gases resist liquefaction altogether: a permanent gas has a critical temperature below normal habitable temperatures and so cannot be liquefied by pressure within that range, a fact that still governs how such gases are stored and transported at high pressure today.