Gases, Properties of
Gases, Properties of
Gases, Properties of
A gas is one of the four phases of matter, with the other phases being plasma, liquid, and solid. These four phases are distinct with regards to certain properties. The fundamental physical properties of a gas are related to its temperature, pressure and volume. These properties can be described and predicted by a set of equations, which are known as the gas laws. While these laws were originally based on mathematical interpretations for an ideal or perfect gas, modern atomic and kinetic theory of gases has led to a modified expression that more accurately reflects the properties of real gases. The term gas is considered to have been first used by Flemish chemist Jan Baptist van Helmont (1577-1644) as a Dutch translation of the Greek word chaos.
Current understanding of gas properties came as a result of study of the interaction between volume, pressure and temperature. English physicist and chemist Robert Boyle (1627-1691) was the first to describe the relationship between the volume and pressure of a gas. In 1660, he learned that if an enclosed amount of a gas is compressed to half its original volume while the temperature is kept constant, the pressure will double. He expressed this law mathematically as PV = constant, where P stands for pressure, V stands for volume, and the value of the constant depends on the temperature and the amount of gas present. This expression is known as Boyle’s law.
The second fundamental property of gasses was defined by French physicist Jacques Charles (1746-1823) in 1787. He found that the temperature and volume of a gas are directly related. Charles observed that a number of gases expanded equally as heat was applied and the pressure was kept constant. This can be expressed mathematically as
His ideas were expanded upon in research by others in the field, most notably French chemist Joseph Gay-Lussac (1778-1850), who also studied the thermal expansion of gases. Even though Charles did not publish the results of his work, the volume/temperature relationship become known as Charles’s law.
The third property of gases was described by Gay-Lussac who, in addition to his work with volume and temperature, researched the connection between
|Properties of gases. (Thomson Gale.)|
|Properties of gases|
|Name||Formula||% Content in Atm||Color||Odor||Toxicity|
|*Contact with air will immediately convert nitric oxide to nitogen dioxide—oxidation|
|Carbon monoxide||CO||—||Colorless||Odorless||Very toxic|
|Chlorine||Cl2||—||Pale green||Irritating||Very toxic|
|Nitrogen dioxide||NO2||—||Red brown||Irritating||Very toxic|
|Nitric oxide||NO||—||Colorless||Odorless||Very toxic*|
|Comparison of gases (pressure and percentages). (Thomson Gale.)|
|Properties of gases: Pressure|
|Name of gas partial||Pressure in air (mm Hg)||Percent content in air|
|*—This is the equilibrium vapor pressure at 20°C.|
pressure and temperature. In 1802, he formed an additional law:
These three laws can be combined into one generalized equation that expresses the interrelation between pressure, temperature, and volume. This equation, called the ideal gas law, is written as PV = nRT where the R is the gas constant, which has been determined experimentally to be equal to 0.082 liter-atmospheres per Kelvin-moles. The symbol n stands for the number of moles of gas. This expression can be used to predict the behavior of most gasses at moderate temperatures and pressures.
While the ideal gas law works very well in predicting gas properties at normal conditions, it does not accurately represent what happens under extreme conditions. Neither does it account for the fact that real gases can undergo phase change to a liquid form. Modern atomic theory helps explain these discrepancies.
It describes molecules as having a certain freedom of motion in space. Molecules in a solid material are arranged in a regular lattice such that their freedom is restricted to small vibrations about lattice sites. Gas molecules, on the other hand, have no macroscopic spatial order and they can move about their containers at random. The motion of these particles can be described by the branch of physics known as classical mechanics. The study of this particulate motion is known as the kinetic theory of gases. It states that
|Comparison of gases in liquid and solid states. (Thomson Gale.)|
|Properties of gases: Various states|
|State||Volume form shape compressibility||Arrangement and closeness of particles||Motion of particles||Attraction between particles||Boiling point|
|Gas||No definite volume, form or shape. Compressible||Random; far apart||Fast||Little to none||Lower than room temperature|
|Liquid||Has a definite volume, but no definite form or shape. Non-compressive tendency||Random; close||Moderate||Moderate||Higher than room temperature|
|Solid||Definite volume, has own shape or form. Non-compressible||Definite; close||Slow||Strong||Much higher than room temperature|
the volume of a gas is defined by the position distribution of its molecules. In other words, the volume represents the available amount of space in which a molecule can move. The temperature of the gas is proportional to the average kinetic energy of the molecules, or to the square of the average velocity of the molecules. The pressure of a gas, which can be measured with gauges placed on the container walls, is a function of the particle momentum, which is the product of the mass of the particles and their speed.
Atomic theory was used to modify the ideal gas law to take into account the interaction between gas molecules on an atomic level. This can be done by factoring in a set of experimental parameters that describe this interaction. The resultant variation of the ideal gas law equation is known as the van der Waals equation of state (which was derived by Dutch scientist Johannes Diderik van der Waals [1837–1923] in 1873): (P + a/V2) (V - b) = RT, where a and b are adjustable parameters determined by measuring intramolecular forces. According to this expression, a strong repulsive force comes into play when molecules are situated very close to one another. This force becomes mildly attractive when the molecules are at moderate distances, and its effect is not measurable at all at greater distances. Van der Waals forces help
explain how a gas can under go a change from a gas to a liquid state. At low temperatures (reduced molecular motion) and at high pressures or reduced volumes (reduced intermolecular spacing), the molecules in a gas come under the influence of one another’s attractive force and they undergo a phase transition to a
Ideal gas law —The mathematical expression that predicts the behavior of a perfect gas.
Kinetic theory of gases —The physical principles that describe how gas molecules interact.
Van der Waals forces —Weak atomic forces that affect gas molecules when they are in close proximity.
liquid. This modified gas law can be used to predict many secondary gas properties including transport properties such as thermal conductivity, the coefficient of diffusion, and viscosity.
Science continues to explore the basic properties of gases. For example, superconductivity, the study of electricity at very low temperatures, relies on super cooled gases like nitrogen to lower the temperature of materials to a point at which they gain special electrical properties. Furthermore, gas analysis techniques have been developed based on the discovery that the speed of sound through a given gas is a function of its temperature. These techniques rely on recently developed ultrasonic technology to analyze two-component gas mixtures that vary by as little as one percent.
Brock, William Hodson. Norton History of Chemistry. New York: Norton, 2000.
Carey, Francis A. Organic Chemistry. Dubuque, IA: McGraw-Hill, 2006.
Hoffman, Robert V. Organic Chemistry: An Intermediate Text. Hoboken, NJ: Wiley-Interscience, 2004.
Tro, Nivaldo J. Introductory Chemistry. Upper Saddle River, NJ: Pearson Education, 2006.