Chapter 3 - The Gaseous State Flashcards
What are 4 simple gas laws?
1) At constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure
2) At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature (in kelvin).
3) At constant temperature and pressure, the volume of a gas is directly proportional to the number of moles of gas present. (Avogadro’s law)
4) at constant temperature and volume, the pressure of a gas is also directly proportional to the number of moles present.
Define ideal gas.
Ideal gas is a hypothetical gas which obeys the ideal gas equation PV=nRT at all values of pressure, volume and temperature.
What is the combined gas equation?
It is when the amount of gas (n) is kept constant, such that P1V1/T1=P2V2/T2, where 1 relates to the gas in its initial condition, and 2 relates to the gas in financial condition. It can be used with any unit.
Define partial pressure.
In a mixture of gases, the pressure exerted by any individual gas on the sides of the container in known as the partial pressure of the gas.
Define Dalton’s law of partial pressure.
It states that the total pressure of a mixture of gases is the sum of the partial pressures of the constituent gases.
What is the relationship between partial pressures of gases and mole fraction/amount of gas?
In a gas mixture, the partial pressure of any individual gas is directly proportional to its mole fraction/amount of gas in the mixture.
Define mole fraction.
It is the ratio of the amount (no. of moles) of an individual gas to the total amount of gas in a mixture of gases.
What are 3 assumptions about ideal gases?
1) The gas PARTICLES have negligible volume compared to the volume of the container. (Gas particles are so small compared to space between them that we can assume the particles themselves have negligible volume)
2) The intermolecular forces of attraction between gas particles are negligible.
3) Collisions between gas particles, and their collisions with the walls of the container are perfectly elastic (no net loss or gain of KE during collision)
What are 2 conditions needed for real gases to behave ideally? Why?
LP and/or HT
1) Low pressures: gaseous molecules would be relatively far apart. The volume of the molecules themselves is negligible compared to the volume of the container. Thus, real gas molecules at low pressure can be approximated to have negligible volume. Also, intermolecular forces are weak as the particles are far apart.
2) High temperatures: gas particles would have enough KE to overcome intermolecular forces, which can thus be considered insignificant.
Why would real gases not obey the ideal gas law?
At high pressures and low temperatures, real gases would deviate from ideal behaviour, especially molecules with stronger intermolecular forces (consider polarity: non polar molecules with large electron clouds&polar molecules with stronger intermolecular forces)
Define mole fraction.
It is the ratio of the amount (no. of moles) of an individual gas to the total amount of gas in a mixture of gases.
What are 3 assumptions about ideal gases?
1) The gas PARTICLES have negligible volume compared to the volume of the container. (Gas particles are so small compared to space between them that we can assume the particles themselves have negligible volume)
2) The intermolecular forces of attraction between gas particles are negligible.
3) Collisions between gas particles, and their collisions with the walls of the container are perfectly elastic (no net loss or gain of KE during collision)
What are 2 conditions needed for real gases to behave ideally? Why?
LP and/or HT
1) Low pressures: gaseous molecules would be relatively far apart. The volume of the molecules themselves is negligible compared to the volume of the container. Thus, real gas molecules at low pressure can be approximated to have negligible volume. Also, intermolecular forces are weak as the particles are far apart.
2) High temperatures: gas particles would have enough KE to overcome intermolecular forces, which can thus be considered insignificant.
Why would real gases not obey the ideal gas law?
At high pressures and low temperatures, real gases would deviate from ideal behaviour, especially molecules with stronger intermolecular forces (consider polarity: non polar molecules with large electron clouds&polar molecules with stronger intermolecular forces)
Why are gaseous particles no longer ideal at high pressures?
At high pressures, the molecules take up a large portion of the volume of the container, resulting in a considerably small space in which the molecules can move. The volume becomes an increasingly significant proportion of the volume of the container. Thus, it is no longer valid to assume that its volume is negligible, so the gas deviates from ideal behaviour.
