The Laws of Gases Help

Avogadro’s Law

In 1811, Italian physicist Amedeo Avogadro noted that when you have equal volumes of gases at the same temperatures and pressures, the gases will have equal numbers of molecules. Avogadro found that this direct relationship between the number of molecules and the volume of the sample could be written in the following way: V/n (number of molecules) = k (a constant)

V ₁ / n ₁ = V ₂ / n ₂

Temperature and pressure are unchanged and constant in Avogadro’s law.

Avogadro’s number ( L ) is a constant number of atoms, ions, or molecules in a sample. It is equal to the number of atoms in 12 grams of carbon-12 or 6.022 × 10 ²³ .

To give you an idea of the unthinkable number of molecules contained in Avogadro’s number, take a bunch of hazelnuts and cover the United States. To equal Avogadro’s number, you would cover it with a layer over 100 kilometers (about 70 miles) deep.

Boyle’s Law

In 1662, an Irish chemist named Robert Boyle tried to figure out how gas is affected by outside factors. To test this, he bent a glass tube into a hook shape and sealed one end. Into this tube he poured mercury. Boyle discovered that the mercury pushed ahead of it a small volume of air to the end that could not escape out of the tube. Since mercury is so heavy and dense, Boyle found that the more he poured into the tube, the harder it pushed against the trapped air at the end of the tube. After adding enough mercury to push (or compress) the trapped air into ½ of its space (or volume) he realized that the more pressure pushing against a trapped volume of gas, the more it was compressed if the temperature stays the same. Boyle described pressure affect in the following way.

In Boyle’s law , when temperature is held constant, a volume of gas is inversely proportional to the pressure; V α 1/P

With further tests, Boyle found that by doubling the pressure, the volume of the gas was reduced by ½. When the pressure was tripled, the volume of gas was squashed to ¹ / ₃ of its original volume.

Boyle decided to multiply both sides of the equation by pressure ( P ) to get rid of the inverse 1/ P . Then, if you know two volumes and want to figure out the pressures or have two pressures and want to find the volumes, then the formula below will give the unknown values:

P ₁ V ₁ = P ₂ V ₂

Examples

Charles’ Law

The second of the ideal gas laws has to do with the affect of changing the temperatures of gases. This different angle of research was done in 1787 by French physicist Jacques Charles. Charles was said to be very interested in the hot air ballooning that was sweeping France as a huge sport at the time. He was known as one of the best balloonists in France. In fact, Charles was the first to use helium to inflate a balloon capable of carrying passengers.

Charles’ scientific nature caused him to use ballooning as a way to test his ideas about gases and temperature. He found that the more a gas was heated, the more its volume increased. He described this with the following equation:

V ₁ / T ₁ = V ₂ / T ₂

V ₂ = V ₁ × T ₂ / T ₁

If we think of what is happening with the atoms of the gas, it is easy to remember Charles’ law. Temperature (heat) provides energy to a sample. When atoms are energized (heated), they move around a lot. Like a happy puppy that can’t stay in one place, heated atoms get crazy wild and hit the sides of their containers harder and more often causing it to expand.

In Charles’ law , when pressure is held constant, a volume of gas is directly proportional to the Kelvin temperature; V α T

The kinetic energy (KE) increases, but the mass stays the same, so the velocity has to increase.

Kinetic energy = ½ m v ²