Properties and Composition of Air: Exploring the Gas Laws
Properties and Composition of Air
The initial properties observed in gases were transparency, compressibility, expansibility, and diffusion capacity.
Air Pressure
The first property of air quantitatively studied was pressure, a crucial measurement for establishing gas laws.
Chemical Composition of Air
The Earth’s atmosphere is a mixture of gases. Four gases constitute 99% of the total volume and are often referred to as the major components of air due to their relatively stable presence. The concentration of most atmospheric gases is subject to variations based on season and altitude.
Ideal Gas Laws
Boyle’s Law
The compressibility of air is a proven fact. In a J-shaped glass tube closed at one end, a quantity of mercury is poured, trapping a volume of air at the closed end. Adding successive amounts of mercury compresses the trapped air. After each addition of mercury, the total volume of trapped air and the pressure exerted upon it are measured.
PV = constant
Boyle could not control the temperature during his experiment but assumed it remained constant. At a constant temperature, the pressure and volume of a fixed amount of gas are inversely proportional.
Charles’s and Gay-Lussac’s Law
Following Boyle’s work, sustained efforts were made to discover the effect of temperature changes on the pressure and volume of a gas. Repeating Boyle’s experiment at different temperatures yielded different curves called isotherms. These curves indicate that temperature influences pressure for the same volume, or volume for the same pressure.
If pressure remains constant, volume increases with increasing temperature; the gas expands. The relationship between volume and temperature was determined by Charles.
In a sealed tube with a small amount of mercury enclosing a given mass of air at the closed end, as the temperature increases, the mercury rises, increasing the volume of trapped air. The results can be represented in a volume-temperature diagram, resulting in a line that intersects the x-axis in the first quadrant. If extended to negative temperatures, the line intersects the axis at -273°C.
V/T = constant
For a fixed amount of gas kept at constant pressure, volume and absolute temperature are directly proportional. This law was established by Gay-Lussac in 1808 and confirmed by Dalton.
Boyle’s and Charles’s laws can be combined into an equation relating pressure, volume, and temperature for a fixed mass of gas. Consider a mass of gas enclosed in a balloon undergoing three consecutive states characterized by volume, temperature, and pressure values. This law is useful for determining the volume of a gas under specific pressure and temperature conditions, knowing the volume occupied by the same gas under different pressure and temperature conditions.
Avogadro’s Hypothesis
Avogadro’s hypothesis, proposed in 1811, states that equal volumes of all gases, measured under the same pressure and temperature conditions, contain the same number of molecules. The scale used to measure the amount of gas is the number of molecules, incorporating the concept of the mole.
If the combined law’s constant term is replaced with the number of moles of gas, we obtain the general equation for ideal gases:
PV = nRT
In this equation, ‘r’ is the ideal gas constant, whose value depends on the units used to measure pressure and volume. When volume is measured in liters, pressure in atmospheres, and temperature in Kelvin:
R = 0.082 atm·L/mol·K
Dalton’s Law of Partial Pressures
Often, the relationship between pressure, volume, and temperature of a gas mixture is needed. This requires understanding how the total pressure of the mixture relates to its individual components. Each gas exerts pressure within the mixture, called partial pressure. In 1801, Dalton formulated a law, known as Dalton’s Law of Partial Pressures, which states that the total pressure of a gas mixture is the sum of the pressures each gas would exert if it were alone.
Graham’s Law of Diffusion
When a second gas is introduced into a container with an existing gas, the two gases eventually form a homogeneous mixture. This gradual mixing process is known as diffusion and was studied by Thomas Graham in 1832. He established a law stating that under the same pressure and temperature conditions, the diffusion rates of two gases are inversely proportional to the square root of their molar masses. This is known as Graham’s Law of Diffusion. This equation is also used to describe the phenomenon of effusion.