gas
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What is a gas in terms of states of matter?
A gas is one of the three fundamental states of matter, distinctly different from the liquid and solid states. Gases lack a definite size and shape, and they will expand to fill any closed container.
Gases consist of a large number of molecules moving randomly in all directions, colliding with each other and the walls of their container. The pressure exerted by a gas results from the countless impacts of these molecules on the container walls. This behavior is described by the kinetic theory of gases, where properties like heat conductivity, viscosity, and diffusion are attributed to the motion of the molecules.
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How do gas molecules move compared to solids and liquids?
Gases lack a definite size or shape, and their properties depend on the container's volume, not its shape. Gas molecules move randomly and are characterized by high average speeds and large intermolecular separation compared to solids and liquids.
In gases, the ratio of intermolecular separation to molecular diameter is approximately 10 under normal conditions. The average molecular speed in a gas is greater than the speed of sound in the gas.
In contrast, solids have both definite size and shape, while liquids have a definite size but adapt their shape to the container. The molecules in liquids are considered to be touching each other.
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What are some common examples of gases?
Gases are a state of matter characterized by the lack of defined size or shape. They fill any closed container, with properties dependent on the container's volume, not its shape.
Examples of gases include:
* Components of Air Air is a mixture of gases, including nitrogen, oxygen, argon, neon, helium, methane, krypton, hydrogen, nitrous oxide, and xenon.
* Noble Gases Noble gases such as helium, neon, argon, krypton, xenon, and radon are known for their chemical inertness and are used in various applications, including lighting and providing unreactive environments.
* Natural Gas Natural gas, primarily methane, serves as a fuel for power generation, industrial, domestic, and commercial uses, and as a feedstock for chemical products.
* Combustion Gases Gases such as carbon monoxide and nitrogen oxides are produced from burning fossil fuels.
* Gaseous Lubricants Gases like air and steam can be used as lubricants in equipment.6 Britannica SourcesThis answer is created from Britannica articles using AI. AI can make mistakes, so verify using Britannica articles. -
How does gas fill the shape and volume of its container?
Gases lack a definite size or shape, and they'll fill any closed container entirely, with properties depending on the container's volume, not its shape.
Gases consist of numerous molecules moving randomly, colliding with each other and the container walls. This constant motion and collision cause gases to expand and occupy all available space. The pressure a gas exerts results from the countless impacts of these molecules on the container walls.
The kinetic theory of gases explains gas behavior based on the motion of these particles. This model assumes gas molecules move randomly and are separated by large distances, undergoing elastic collisions with each other and the container.
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What happens to a gas when you change its temperature or pressure?
Gases are affected by changes in temperature and pressure. The ideal gas law describes the relationship between pressure, volume, and temperature of a gas.
- Boyle's Law At a constant temperature, the pressure of a gas varies inversely with its volume.
- Charles's Law At a constant pressure, the volume of a gas is directly proportional to its absolute temperature.
- Ideal Gas Law Combines Boyle's and Charles's laws into PV = nRT, where P is pressure, V is volume, n is the number of moles, R is the universal gas constant, and T is the absolute temperature.
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How are gases important in our everyday lives and the environment?
Gases are vital to both our daily lives and the environment. Biogenic gases, which are produced by living organisms, include oxygen, nitrogen, water vapor, carbon dioxide, and methane.
These gases affect Earth's radiation budget, the thermodynamics of the atmosphere, and wind systems. The atmosphere is largely transparent to sunlight, but gases like water vapor, carbon dioxide, and methane absorb infrared radiation, maintaining Earth's temperature. Human activities, such as burning fossil fuels, have increased carbon dioxide levels, contributing to global warming.
Industrially, gas turbines are used for power generation and to drive compressors. Oxygen is used in steel production, sewage treatment, and medical applications. However, gases like sulfur dioxide from burning fossil fuels can cause air pollution and acid rain, harming the environment and human health.
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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.
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.
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.
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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.