A substance is in the aeriform state when it does not have a defined shape (it is therefore a fluid) or a defined volume, as it tends to expand, completely filling the container that contains it. Any substance above its boiling point can, therefore, be defined as aeriform. An aeriform can also exist below the boiling point, vaporizing from its liquid.
Formally, gas is aeriform that is at a temperature higher than its critical temperature (temperature above which a substance cannot exist in a liquid state), and this means that it cannot be condensed in any way by simple compression that is, it cannot be brought to the liquid state by compressing it at a constant temperature.
When, on the other hand, they are at a temperature lower than their critical temperature, the aeriforms are said to be in a vapor state; a vapor can, therefore, become liquid when it is sufficiently compressed at a constant temperature.
Aeriforms have high compressibility and, in the absence of external forces, can expand indefinitely. Between the molecules of aeriforms only very weak intermolecular forces act; interactions become important only at the moment of collisions. Aeriforms are distinguished in gases and vapors: the former are aeriforms that at room temperature cannot be condensed by simple compression (they must therefore undergo a cooling below their critical temperature) while the latter condense if they are sufficiently compressed. If we report in a pressure-volume diagram the states of an aeriform at different temperatures, we build experimentally as many curves, compression isotherms (continuous lines), which together constitute the so-called Andrews diagram.
At relatively low temperatures (e.g. 0 °C) it is found that, as pressure increases, the volume of the aeriform decreases until it begins to condense (dew point). An attempt to increase the pressure of the aeriform only leads to a decrease in its volume (the pressure remains constant and in the isotherm there is a landing) which continues until it is all condensed (all liquid point). At this point the substance is in the liquid state and even very large increases in pressure do not lead to a noticeable decrease in volume because liquids are practically incompressible and the isotherm rises almost vertically. The same form have the other isotherms that are built for increasing temperatures, with the difference that the length of the landing is decreasing until it is reduced to a point (critical point).
Above the critical temperature there is no more condensation; as the compression of the aeriform increases, it behaves in a similar way to a perfect gas; for high enough temperatures the isotherms get closer and closer to equilateral hyperbolas. The bell-shaped curve formed by the set of dew points, all liquid and the critical point is called Andrews or Mathias curve and it, together with the isotherm passing through the critical point (critical isotherm), divides the pressure-volume plane into four zones. In the zone above the critical isotherm the aeriform cannot be condensed by simple compression: this is the gas zone. In the right zone between critical isotherm and Andrews curve the aeriform is a vapor. In the area below Andrews curve the aeriform is a vapor in presence of its liquid, a saturated vapor. In the area to the left between the critical isotherm and the Andrews curve there is only liquid.
In reality all aeriforms can be gas or vapor depending on whether they are above or below the critical temperature, which is characteristic for each of them; however we assign, as mentioned above, the name of vapor to aeriforms whose critical temperature is higher than ambient temperature and the name of gas to those whose critical temperature is lower. Nitrogen, which has a critical temperature equal to -117.1 ºC is therefore a gas, while ammonia, with a critical temperature equal to 132.4 ºC, is a vapor.