Chemistry (Ideal Gases)

GREASE THOSE WHEELS (CHEMISTRY)

#3 IDEAL GASES

• state the basic assumptions of the kinetic theory as applied to an ideal gas
• explain qualitatively in terms of intermolecular forces and molecular size:
• the conditions necessary for a gas to approach ideal behaviour.
• the limitations of ideality at very high pressures and very low temperatures.
• state and use the general gas equation pV = nRT in calculations, including the determination 
• of Mr

------------------------------------------------------------------------------------------------------------------------

Before we start off with this component of our syllabus, you should have some background knowledge about gases...

• Gases exert a pressure on the container walls they are contained in.
• Gases have very weak forces of attraction between them which allows their molecules to move about randomly, independent of each other. Which we condense into constant random motion.
• Gases can be compressed, i.e, they can be forced to consume a smaller volume upon the application of pressure.
• At constant temperature the product of pressure and volume of a gas is constant for all values of pressure and volume thus the important, Boyle's law, P1V1 = P2V2. [Remember you can use any units for calculations involving the Boyle's law unless and until both values of pressure have the same units. (the same is true for volume)]
• At constant pressure, the volume of a gas is proportional to its temperature thus Charles's Law, V=kT [T is always used in Kelvins where the relationship between degree Celsius and Kelvin Scale being: Temperature in degree Celsius + 273 = Temperature in Kelvins while any identical units for volume may be used]
• The above mentioned gas laws can be combined to give P1V1/T1 = P2V2/T2.
• The reference point for discussing changes in gas behaviour is the s.t.p (standard temperature and pressure) where temperature equals absolute zero \ 273 Kelvins and pressure equals 1 atm or 101325 Pa and volume taken as 22.4 dm3 in the light of Avogadro's law which states that all gases have the same number of molecule at equal volumes when temperature and pressure are kept constant, i.e, V = kn where n is the number of moles of the gas.

------------------------------------------------------------------------------------------------------------------------

Combining all the previously stated laws and relationships, we get the ideal gas equation.

"PV = nRT"

Where...

• P = Pressure = in Pascals (1KPa = 1000 Pascals)
• V = Volume = in metre cubed (1 m3 = 10^-3 dm3 = 10^-6 cm3)
• R = Constant = 8.71 J/mol.K
• n = number of moles of gas = in mols
• T = Temperature = in kelvin (Temperature in degree Celsius + 273 = Temperature in Kelvins)

Now you will be asking that what is an Ideal Gas hmm...well researchers from a bygone era laid down these assumptions for an ideal gases.

1. The size of a gas molecules is negligibly small and they are in constant random motion in a straight line.
2. The volume occupied by a gas molecules is negligibly small when compared to the volume of the container and gas molecules have large intermolecular distances.
3. The collisions between the walls of the container and the gas molecules and within the gas molecules are perfectly elastic, i.e, there is no net loss or gain of energy.
4. The attractive and repulsive forces between gas molecules are so weak that they should not be considered.
5. The average kinetic energy of gas molecules is proportional to the absolute temperature of the gas.

But all gases do not behave identically and usually deviate from being an ideal gas. Why? Lets see..

Non-ideal gas behaviors is especially noticeable at low temperature and high pressure. We see this in the light of two main faulty assumptions of the ideal gas theory (see text in red above)...

• At high pressures the volume of the container decreases and so the gas molecules become a considerable fraction of it so at high pressures the product PV is higher than expected for ideal gases.
• At low temperatures the average kinetic energy is way to less and the attractive forces between the gas molecules become dominant and the gas molecules strike the walls with decreasing vigor and the product PV gets lower than expected for ideal gases.

Remember! For an ideal gas the plot of PV against Pressure is a straight line and that gases with a lower Mr get more close to showing ideal gas behaviour, i.e, hydrogen is more ideal than ammonia.

------------------------------------------------------------------------------------------------------------------------

Thats all for this small but interesting topic.