Gas is aeriform whose temperature is higher than the critical temperature; as a result, gases cannot be liquefied without first being cooled, unlike vapors. Gas is a fluid that has no volume of its own (tends to occupy all the volume at its disposal) and that is easily compressible. A gas is one of the four fundamental states of matter, in which, the atoms or molecules are far apart due to they are not bounded at all, meaning, they do not have any attractive forces but only repulsive forces. The term “gas” attributed in 1620 by the chemist J.B. van Helmont to substances that are in the gaseous state and therefore have no proper volume.
The interaction of gas particles in the presence of electric and gravitational fields are considered negligible. Due to that, they can occupy a large volume. They do not have their shape or volume but assume the shape and the volume of the container. Finally, gas particles spread apart or diffuse in order to homogeneously distribute themselves throughout any container. Some typical examples are oxygen, hydrogen, and helium at room temperature.
The gaseous state, like any other state of aggregation, depends on the temperature and pressure conditions and is not characteristic of certain substances. Saying that a substance, for example, air is a gas, it only means that it is such in the ordinary conditions of temperature and pressure, by varying these conditions it can instead manifest itself as a liquid or even as a solid.
The only chemical elements that are stable diatomic homonuclear molecules at standard conditions for temperature and pressure are hydrogen (H2), nitrogen (N2), oxygen (O2), and two halogens: fluorine (F2) and chlorine (Cl2). When the pressure is changed and is higher or lower, or when the temperature is changed and is higher or lower, then the element may exist in a different form such as in liquid form or solid form. When grouped together with the monatomic noble gases – helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) – these gases are called “elemental gases”.
Physical characteristics of a gas
The physical properties or macroscopic characteristics of a gas are:
From the microscopic point of view, however, the properties of gases are:
- kinetic theory of particles: the model of the kinetic theory of gases describes a gas as a large number of identical submicroscopic particles (atoms or molecules) with the same mass, all of which are in constant, rapid, random motion. Their size is assumed to be much smaller than the average distance between the particles. The particles undergo random elastic collisions between themselves and with the enclosing walls of the container. The basic version of the model describes the ideal gas and considers no other interactions between the particles and, thus, the nature of kinetic energy transfers during collisions is strictly thermal.
- brownian motion: is the mathematical model used to describe the random motion of particles suspended in a fluid (a liquid or a gas) resulting from their collision with the fast-moving molecules in the fluid.
- intermolecular forces: momentary attractions (or repulsions) between particles affect gas dynamics; the name given to these intermolecular forces is van der Waals force. These forces play a key role in determining physical properties of a gas such as viscosity and flow rate. Ignoring these forces in certain conditions allows a real gas to be treated like an ideal gas. This assumption allows the use of ideal gas laws which greatly simplifies calculations.
Compared to the other states of matter, gases have low density and low viscosity. Since gas molecules can move freely within a container, their mass is normally characterized by density. 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.
In a gas (at standard temperature), the attractive forces existing between the molecules are not such as to keep the molecules bound together; when instead the gas is brought to a very low temperature, it happens that the short-range forces end up prevailing on the tendency of the molecules to remain independent one from the other.
Under normal conditions of pressure (1 atm) and temperature (25 °C), the molecules are practically free from each other, and this is the reason why a gas always tends to occupy the entire volume at its disposal.
To characterize a gas we need different parameters, unlike liquids or solids, for which even a single parameter is often sufficient: for example, if we have 1 liter of water, there is no possibility of confusion, as it is permissible to neglect the phenomenon of “cubic expansion” of liquids (in relation to sudden changes in pressure or temperature), so that a precise volume will always correspond to 1 liter of water. Same thing for solids: to study a substance in the solid state it is generally not necessary to specify under which experimental conditions we conduct our analysis.
Different is the case of gases because the “quantity” for gas is something entirely different from the volume that contains it. Given a certain mass (m) of a gas, or a certain quantity of such gas, it is necessary to use other parameters to conduct further analyzes: the pressure (P), the volume (V) of the vessel containing the gas and the temperature (T) at which the gas is located.
Of these four quantities (P, V, T, m), each can be expressed as a function of the other three: while in solids and liquids the dependence of “V” and “m” from “P” and “T”, it is not possible to neglect this dependency for gases (it is correct to speak of a specific volume only if we specify in what conditions of pressure and temperature we consider it). This is the only way we can get the mass of gas available.