4 The Gaseous State Flashcards
State the basic assumptions of the kinetic theory as applied to an ideal gas. How is a real gas different from an ideal gas? Explain quantitively in terms of intermolecular forces and molecular size, the (i) conditions necessary for gas to approach ideal behaviour (ii) limitations of ideality at very high pressures and very low temperatures
a. An ideal gas consists of particles or negligible volume as the size of gas particles is negligible compared to the volume of the container. (Gas particles of a real gas have finite size and volume)
b. The gas particles exert negligible attractive forces on one another. (There are FOA between the gas particles in a real gas and some attractive forces are stronger than orders e.g. hydrogen bonding, id-id etc)
c. The gas particles are in constant random motion, colliding with one another and with the walls of the container. Between the collisions, they move in straight lines.
d. Collisions between the gas particles are perfectly elastic (no loss of overall K.E. upon collision) though K.E. can be transferred between colliding particles. (Collisions between gas particles of a real gas may not be elastic)
e. The average K.E. of the particles in a gas is constant at a constant temperature as the average kinetic energy is proportional to the absolute temperature. At the same temperature, all the gas particles have the same kinetic energy.
f. Obeys the ideal gas equation under all conditions of pressure and temperature. (Real gas approaches ideal gas behaviour at low pressure and high temperature.)
(i) Under low pressure, the gas particles are far apart and the volume occupied by them is negligible compared to the volume of the container and the intermolecular FOA are also negligible as the gas particles are far apart. At high temperature, the gas particles have large kinetic energy to overcome the attractive forces so the intermolecular FOA between them are negligible.
(ii) However, under high pressure, the gas particles are much closer together and the gas occupies a smaller volume. As such, the volume of the gas particles is not negligible as compared to the volume of the container and the intermolecular FOA between the closely spaced gas particles also become significant. At low temperature, the gas particles possess less kinetic energy and the intermolecular FOA become significant (KE insufficient to overcome FOA). Hence, real gas deviates from ideal gas behaviour at high pressure and low temperature.
How do you determine Mr using the general gas equation? State the units of m, M and Mr.
pV = nRT = m/M(RT) where m is the mass of the gas (in g) and M is the molar mass of the gas (in g mol^-1)
Note that while molar mass M has the same numerical value as the relative molecular mass, Mr. M has units and Mr does not.
State the general gas equation and the units of each term. State the value of 1 atm, 1 bar, 1 dm^3 and 1cm^3 in m^3.
pV = nRT p = pressure exerted by the gas (Pa or N m^-2) V = volume occupied by the gas (m^3) R = molar gas constant (8.31 J K^-1 mol ^-1) T= temperature of the gas (K) n = amount of the gas (mol)
1 atm = 101 325 Pa
1 bar = 10^5 Pa
1 dm^3 = 1000cm^-3 = 10^-3 m^3
1cm^3 = 10^-6 m^3
State Gay-Lussac’s Law. What equation expresses this law?
Gay-Lussac’s Law states that for a fixed mass of gas at constant volume, the pressure of the gas is directly proportional to its absolute temperature. At constant V and n, This law is expressed as p1/T1 = p2/T2 as p/T is a constant.
State Charles’ Law. What equation expresses this law?
Charles’ Law states that for a fixed mass of gas at constant pressure, the volume of gas is directly proportional to its absolute temperature. At constant p and n, V1/T1 = V2/T2 as V/T is a constant.
State Boyle’s Law. What equation expresses this law?
Boyle’s Law states that for a fixed mass of gas at a constant temperature, the volume of the gas is inversely proportional to the pressure. At constant T and n, p1V1 = p2V2 as pV is a constant.
How do you determine the density of a gas using the general gas equation? State the units of density.
pV = m/M(RT) p = m/V(RT/M) where m/V is the density (in g dm^-3)
What is the value of pV/RT for 1 mole of an ideal gas?
1
Which 2 factors explain why real gases experience positive and negative deviations as pressure increases? What affects these 2 factors and why?
- Intermolecular attractive forces
- Intermolecular repulsive forces
The molecular sizes affect the intermolecular attractive forces and intermolecular repulsive forces as the greater the molecular size, the larger the electron cloud which results in stronger intermolecular attractive and repulsive forces.
Which 2 factors explain why real gases experience positive and negative deviations as pressure increases? What affects these 2 factors and why?
- Intermolecular attractive forces
- Intermolecular repulsive forces
The molecular sizes affect the intermolecular attractive forces and intermolecular repulsive forces as the greater the molecular size, the larger the electron cloud which results in stronger intermolecular attractive and repulsive forces. (occupy more significant volume compared to the volume of the container)
The strength of the intermolecular FOA also affects the extent of deviation as the deviation is greatest for polar molecules and smallest for non-polar molecules. (Pg 15 of notes)
Define partial pressure.
Under the same physical conditions, each gas exerts its own pressure that is equal to the pressure a gas exerts when it occupies the container alone, this pressure is known as the partial pressure of the gas. (In a mixture of gases, each gas behave as if it is the only gas present.)
State Dalton’s Law of partial pressure. Define mole fraction. What is the relationship between mole fraction, the partial pressure and the total pressure of the mixture?
Dalton’s Law of partial pressure states that the total pressure of a mixture of non-reacting gases is equal to the sum of the partial pressures of the individual gases in the mixture.
Mole fraction of a component in a non-reacting gas mixture is the ratio of the number of moles of A to the total number of moles of gaseous components present in the mixture.
Mole fraction times the total pressure of the mixture = Partial pressure of the gas.
Define vapour pressure,
Pressure exerted by the gaseous particles on the liquid surface and on the walls of the container.
Define dynamic equilibrium.
Equilibrium in which the number of particles leaving the surface is exactly balanced by the number rejoining it. In this equilibrium at a particular temperature, there will be a fixed number of gaseous particles in the space above the liquid.
Define saturated vapour pressure.
Vapour pressure exerted by the gaseous particles on the liquid surface and on the walls of the container at the temperature where dynamic equilibrium is set up.
