Gas

Gas

I

INTRODUCTION

Gas, one of the three ordinary states of matter. The other two ordinary states of matter are solid and liquid. Both solids and liquids are made up of particles that touch one another. The attraction between the particles of a solid is so strong that the particles hold rigidly together. This rigidity gives solids a definite shape and volume. The attraction between particles in a liquid is great enough to hold the particles near each other but too weak to prevent the particles from sliding around. Liquids have a definite volume but take the shape of their container. The particles that make up a gas, however, are completely separated from one another. Empty space accounts for more than 99 percent of the total volume of air, for example. Because gas particles are separated, the attractive forces between them are extremely small and are insufficient to hold gases in a definite shape or volume. Gases expand freely to fill their containers.

II

PROPERTIES OF GASES

The characteristics or properties of gases vary widely. Some gases are transparent, some have a strong smell, some dissolve in water, and some react violently with almost any substance. Other gases exhibit exactly the opposite properties. The chemical structure of gases also varies greatly.

A

Color

A number of gases have a characteristic color. For example, fluorine gas appears green, chlorine appears yellow-green, and nitrogen dioxide (a component of smog) appears red-brown. The majority of gases, however, are colorless.

B

Odor

Many gases, including nitrogen, oxygen, and hydrogen, are odorless. Ammonia, however, has a sharp, pungent odor. Because fuel gases such as methane, propane, and butane are odorless, an intensely odorous sulfur compound is added to them to ensure early detection should these gases leak from their containers.

C

Solubility

Some gases, such as carbon dioxide, dissolve well in water. Many others, including nitrogen, hydrogen, and oxygen, are only slightly soluble in water. The solubility of any gas decreases as the temperature of the gas increases, and it increases as the pressure increases.

D

Chemical Reactivity

Some gases can react with other substances to form new chemical compounds. Oxygen, for example, reacts with iron to form rust. The chemical reactivity of gases varies widely. Oxygen, chlorine, and fluorine are extremely reactive gases. In fact, fluorine will react with almost any other substance; even water and glass will burn in a fluorine atmosphere. At the other extreme are the noble gases, which are generally considered inert (unreactive). Neon, a noble gas, is not known to react with any other substance.

E

Structure

Gas particles are the smallest units into which a gas can be divided without changing the chemical properties of the gas. These particles can either be single atoms or molecules (combinations of atoms). The noble gases, such as neon and helium, are composed of individual atoms. Other gases, including carbon dioxide (CO₂), methane (CH₄), and ammonia (NH₃), contain atoms of more than one element chemically bound together in molecules. Some gases that contain only a single element, such as hydrogen, oxygen, and nitrogen, are also composed of molecules. The oxygen in Earth’s atmosphere, for example, consists mostly of oxygen molecules (O₂) rather than individual oxygen atoms (O).

III

GAS LAWS

From the 17th to the 19th century, scientists noticed that gases respond to changes in temperature, pressure, and volume in predictable ways. Scientists established four laws that govern the behavior of gases: Boyle’s law, Charles’s law, Dalton’s law, and Avogadro’s law. These four gas laws can be combined and expressed as a single equation known as the combined ideal gas equation.

A

Boyle's Law

The smaller the volume a given amount of gas is squeezed into, the greater the pressure the gas exerts on the walls of its container. Boyle's law, a mathematical equation that more precisely describes this relationship, states that at constant temperature, the volume of a given quantity of gas varies inversely with the pressure exerted on it. Mathematically, this relationship can be expressed: V is proportional to k (1/P)where V is volume, k is a constant, and P is pressure. Boyle’s law asserts that if the pressure on a given amount of gas is doubled, its volume will decrease by one-half (as long as the temperature of the gas remains unchanged). Conversely, if the pressure is decreased by one-half, the volume will double.

B

Charles's Law

Raising the temperature of a gas causes the gas to fill a greater volume as long as pressure remains constant. Gases expand at a constant rate as temperature increases, and the rate of expansion is similar for all gases. Charles's law (also called Gay-Lussac’s law) describes the relationship between volume and temperature of an enclosed gas. The law says that at constant pressure, the volume of a fixed number of particles of gas is directly proportional to the absolute (Kelvin) temperature, mathematically expressed as: V = k Twhere T is temperature on the Kelvin scale (see Temperature: Temperature Scales). If the temperature of a given amount of gas is doubled, for example, its volume will also double (as long as pressure remains unchanged).

C

Dalton's Law

Dalton’s law states that in a mixture of different gases, such as air, the sum of the partial pressures of all the gases equals the total pressure. The partial pressure of a gas is the pressure that gas would exert if it was the only gas present. This law, expressed mathematically, is: Ptotal = P₁ + P₂ + P₃ + ...where each subscripted P value is the partial pressure of a different gas.

Air, for example, consists mostly of nitrogen and oxygen, with small quantities of argon, water vapor, and carbon dioxide. The partial pressure of nitrogen accounts for 78 percent of the total pressure exerted by Earth’s atmosphere; oxygen accounts for 21 percent; and argon accounts for 0.9 percent. Together, these three gases account for 99.9 percent of air pressure. Water vapor, carbon dioxide, and all the other trace gases present in the atmosphere combined contribute only a tenth of a percent.

D

Avogadro's Law

The behavior of gases described by Boyle’s, Charles’s, and Dalton’s laws is nearly the same for all gases. Avogadro’s law states that under identical conditions of temperature and pressure, equal volumes of different gases contain equal numbers of particles (atoms or molecules).

The temperature 0ºC (32ºF) and the pressure equal to the pressure Earth’s atmosphere exerts at sea level are called standard temperature and pressure (STP). According to Avogadro’s law, 1 cubic meter of oxygen at STP contains the same number of particles as 1 cubic meter of nitrogen at STP. Restated, Avogadro's law says that one mole of any gas at STP occupies a volume of 22.4 liters. A mole is 6.02 × 10²³ basic particles (atoms or molecules) of a substance. The extremely large number 6.02 × 10²³ is called Avogadro's number.

E

Ideal Gas Equation

The gas laws can be combined as a more general expression called the ideal gas equation or ideal gas law: PV = nRT

In this equation, n represents the number of moles of a gas. The constant R on the right-hand side of the equation is a universal constant and has a value of 8.31447 J/mol·K. This single equation can predict the behavior of a gas even if multiple conditions are changed simultaneously. If both the pressure and volume of a gas double, for example, its temperature will increase by a factor of four.

IV

UNDERSTANDING GAS BEHAVIOR

The gas laws were discovered empirically—that is, scientists performed experiments, observed the behavior of gases, and came up with equations that fit that behavior. Scientists such as Boyle and Charles could describe how gases behave, but they did not know the reasons underlying gas behavior. The behaviors they observed can now be explained by the kinetic theory of gases (also called the kinetic-molecular theory). Kinetic theory holds that gases are collections of tiny particles flying around randomly, bumping into each other and into the walls of any container enclosing them. Pressure is really a macroscopic (large-scale) reflection of how hard and how often these microscopic particles strike the walls of their container. Temperature is a manifestation of how fast the particles are moving.

According to kinetic theory, gases exert pressure because their particles have kinetic energy and constantly move around and collide with the walls of the container holding the gas. Kinetic energy is energy of motion, and is related to temperature. Raising the temperature of a gas raises the kinetic energy of its particles. A gas particle has kinetic energy in proportion to its speed: the faster it is moving, the greater its kinetic energy. In mathematical terms, the kinetic energy of a gas particle is equal to ½mv², where m is the particle's mass and v is its velocity. The more particles that strike a given area at any instant of time and the harder they hit, the greater the gas pressure. So the pressure that a gas exerts depends on the number of particles in the sample, the volume of its container, and the temperature of the gas.

A

Real Gases and Ideal Gases

A gas that obeyed the ideal gas equation exactly under any conditions would be an ideal gas, but no actual gas perfectly conforms to the equation at all temperatures and pressures. Under the conditions of high temperatures and low pressures present over much of Earth’s surface, however, most real gases behave as ideal gases. Gases with boiling points below –173ºC (-279ºF), such as hydrogen, oxygen, nitrogen, and the noble gases, come closest to being ideal gases. Gases with relatively high boiling points, such as carbon dioxide, obey the gas laws only approximately.

B

Van der Waals Equation

The ideal gas equation assumes that there are no attractive forces between particles of a gas. If that assumption were true, an ideal gas would remain in the gaseous state under all conditions. Ideal gases, therefore, could never become liquids or solids, no matter how much they were cooled or compressed. Very small attractive forces between gas particles do exist, however, and gases do liquefy if cooled sufficiently. The ideal gas equation has no provision for a change of state from gas to liquid. According to the equation, an ideal gas would simply become denser and denser at lower temperatures and higher pressures without ever liquefying. Because all real gases can in fact be converted to liquids, Dutch physicist Johannes van der Waals came up with an adjusted version of the ideal gas equation (PV = nRT): (P + a/V²) (V - b) = nRT

The adjustable parameters a and b are determined from experimental measurements carried out on actual gases. Their values account for the strength of attractive forces between real gas particles and for particle size, factors that are different for different gases.

Van der Waals equation can be explained in terms of interactions between gas particles. Gas particles strongly repel each other at close range and mildly attract each other at intermediate range. The ideal gas equation must be modified slightly when these attractive and repulsive forces are considered. The mutual repulsion between particles, for example, prevents neighboring particles from getting too close. The ideal gas equation assumes that particles are free to move anywhere within the volume (V) of a container, so this fraction of unavailable space around each particle requires an adjustment to that volume (V – b).

C

Diffusion and Effusion

If not constrained in some way, gases expand to fill all available space. They also mix with other gases if no barriers keep them separate. Diffusion is the movement of a gas into a space or the mixing of one gas with another. When room deodorants or perfumes are released in one part of a room, they diffuse and one can soon detect the odor in all parts of the room. Effusion describes the escape of a gas through a tiny hole. If gases are placed in a container with porous walls, such as in a balloon, the particles effuse through its walls, causing the volume to gradually decrease. Both diffusion and effusion occur because of the vast amount of space between gas particles and the kinetic energy of the gas particles. The constant, rapid, random motion of gas particles makes them spread out rapidly in all directions and distribute themselves uniformly throughout any container.

Different gases comprise particles with different masses. In 1832 British chemist Thomas Graham proposed that the rates of effusion and diffusion of gases are inversely proportional to the square roots of the masses of their particles. In other words, gases made up of smaller particles effuse and diffuse faster than gases made up of larger particles. This principle is now known as Graham’s law. Since the average kinetic energy of a gas particle, ½mv², is the same for all gases at the same temperature, less massive gas particles must travel faster than heavier particles. Because they move faster, atoms of lighter gases such as helium effuse through the tiny openings of porous balloon walls, for example, more quickly than the heavier molecules of gases such as oxygen or nitrogen.

V

COMMON USES OF GASES

Improved understanding of the properties of gases has led to their large-scale exploitation for industrial and consumer applications. A few common examples may serve to illustrate the importance of gases for human beings.

Oxygen is perhaps the most familiar gas, since it occurs in the atmosphere and animals require it to survive. Animals, including humans, take in oxygen from the air when they breathe. Fires require oxygen to burn; familiar items such as candles, gas stoves, and fireplaces will not work without oxygen. Oxygen is also used as a fuel in rockets.

Carbon dioxide is another familiar gas. Some of the oxygen that animals breathe in is combined with carbon to produce carbon dioxide that is subsequently exhaled. The bubbles in soda and beer are actually bubbles of carbon dioxide. The gas is dissolved under pressure in flavored solutions to produce many kinds of carbonated beverages.

Helium is used to fill party balloons, airships, and weather balloons because it is much less dense than air. Helium and other gases such as argon, nitrogen, krypton, and xenon do not react with most elements. Because they do not react, they provide inert working environments for the chemical and electronics industries, for metallurgical processes, and for high-temperature welding. Such environments make possible the production of ultrapure silicon and germanium semiconductors, for example, items that might otherwise be marred by exposure to air. Incandescent light bulbs often contain argon because their extremely hot filaments would quickly react and disintegrate if exposed to air.

Liquefied gases are often used as coolants because gases liquefy at very low temperatures. Most cooling systems work by transforming a coolant substance from gas to liquid and back again. Liquid nitrogen is used as a freezing agent in the food processing industry. Liquid helium produces the extremely low temperatures that the superconductors and electromagnets used in particle accelerators and nuclear fusion research need to function.

The fuel gases methane, propane, and butane can be burned with the oxygen in air to provide energy. Natural gas, used to power stoves, heating systems, clothes dryers, and hot-water heaters in many homes, is mostly composed of methane. These gases, or their derivatives, are sometimes also used as aerosol propellants.

VI

HISTORY OF GAS RESEARCH

The intensive investigation of gases dates back only a few hundred years, but scientists have made great progress in understanding gases. During the 18th and 19th centuries scientists realized that air comprises more than one gas, and conducted experiments to determine how gases respond to changes in temperature, pressure, and volume. Although the understanding of gases solidified by the early 20th century, scientists continue to conduct extensive research to find more applications for gases.

A

Early Observations

Until the 17th century the only well-known gas was air. Alchemists, ancient and medieval experimenters who attempted to find ways to produce gold or silver from base metals, frequently produced “airs” and “vapors” in their experiments. These gases, however, were usually ignored rather than studied.

In 1640 the Belgian physician Jan van Helmont noted that the gas produced by burning wood resembled air but did not behave quite like air. He subsequently coined the word gas from the Flemish pronunciation of the word for chaos. Today we call the gas he produced carbon dioxide.

In 1643 Evangelista Torricelli, an Italian mathematician and physicist, conducted laboratory experiments to show that the gases in air exerted pressure and that air supported a column of mercury 76 cm (30 in) high. In so doing, he invented the barometer.

B

Major Advances in Understanding Gases

During the mid- to late 1600s, several experimenters observed that most combustible objects disappear upon burning, and that the remaining soot or ash was much lighter than the original substance. Rust, however, was observed to be heavier than the original metal. Gases became a subject of study in part because the prevalent theories could not explain these phenomena. This research led to an understanding of how gases behave and why they behave in those ways.

The English chemist Robert Boyle spent much of his career studying the behavior of gases. In 1662 he devised an air pump that he used in the first attempt to precisely measure the relationship between volume and pressure in gases. As he studied what he called “the spring of air”—the pressure with which a compressed gas sample pushes back on the walls of its container—Boyle observed an indirect relationship between volume and pressure. As volume decreases, pressure increases. French physicist Edme Mariotte, who discovered the same relationship independently in 1680, noted that temperature must be held constant for the relationship to remain valid.

As early as 1738, the Swiss scientist Daniel Bernoulli applied Sir Isaac Newton's laws of motion to gas molecules. He envisioned gaseous molecules in ceaseless motion, exerting pressure when they struck the walls of their container.

In 1766 the British chemist Henry Cavendish systematically investigated the properties of hydrogen, the gas produced when an acid reacts with a metal. He discovered that hydrogen, when burned in air, produces water.

In 1772 British physician Daniel Rutherford isolated nitrogen from air and discovered that objects would not burn in it. Four years later British chemist Joseph Priestley discovered that objects burned more brightly and rapidly in oxygen, another component of air. Based on Priestley's discoveries, French chemist Antoine Lavoisier postulated that air was a mixture of oxygen and nitrogen and that only one-fifth of air was oxygen. He proposed that oxygen was the part of air that combined chemically with burning or rusting materials. Lavoisier also consolidated discoveries made in connection with gases into the law of conservation of mass. The law of conservation of mass states that, in a chemical reaction, the total amount of matter remains constant. The law of conservation of mass served as the cornerstone of 19th-century chemistry.

During the late 18th and early 19th centuries, two French scientists, Jacques Charles and Joseph Gay-Lussac, studied the expansion of gases as the gases were heated. The two scientists observed that raising the temperature of a gas caused the gas to fill a greater volume as long as pressure remained constant. In 1807, while studying the difference between moist air and dry air, British chemist and physicist John Dalton discovered Dalton’s law, the law of partial pressure. In 1811 Italian chemist Amedeo Avogadro proposed an explanation for the fact that different gases seem to uniformly obey Boyle’s and Charles’s laws. He hypothesized that equal volumes of different gases contain equal numbers of particles (atoms or molecules) when the gases are at the same temperature and pressure.

British physicist James Maxwell and Austrian physicist Ludwig Boltzmann were able to derive Boyle's law mathematically, rather than experimentally, based on three assumptions: gases are vast numbers of randomly moving particles, there are no attractive forces between gas molecules, and gas molecules have no size. Their work contributed to the kinetic theory of gases. In 1857 German physicist Rudolf Clausius, building upon their work, published a theory that explained Boyle's, Charles’s, Dalton's, and Avogadro's observations in terms that are now known as kinetic theory.

Argon was the first of the noble gas family of elements to be identified. Lord Rayleigh, an English physicist, and Sir William Ramsay, a British chemist, discovered it in 1894. The name argon, which means “inactive” or “lazy,” was chosen because argon did not react with magnesium in the way that nitrogen did. French astronomer Pierre Janssen observed the spectral lines of helium, another noble gas, in the Sun in 1868, but helium was not discovered on Earth until 1895. In 1898 Ramsay and Morris William Travers succeeded in separating neon, krypton, and xenon from air. German chemist Friedrich Dorn discovered radon in 1900. Ramsay later determined that helium and radon were products of the radioactive decay of radium.

Dutch physicist Johannes van der Waals conducted research on gases during the late 19th and early 20th centuries. This research led him to van der Waals equation, a corrected version of the ideal gas equation. Physicists and chemists continued to refine the ideal gas equation throughout the 20th century.

C

Current Research

One of the most interesting areas of current gas research is cryogenics, the study of changes in the properties of matter at very low temperatures. Gases exhibit some of the most interesting changes in behavior at low temperatures, and they are also used to achieve the very low temperatures that cryogenic research requires. Cryogenic research could result in advances including better superconductors, more powerful computers, and faster communication.

Gases also have a number of potentially beneficial environmental applications. Safe, economical storage and large-scale production of hydrogen could provide an efficient, environmentally sound energy source because when hydrogen burns in air it produces energy with only pure water as a by-product. Direct injection of gases such as oxygen into streams and sewage systems could neutralize wastes by oxidation, preventing the wastes from reacting with the environment. Greater understanding of chemical reactions in the atmosphere could determine courses of action that would prevent the destruction of the ozone layer and reduce production of pollutant gases.

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