chemistry - the gaseous state Flashcards
(28 cards)
what is the definition of avogardo’s law?
equal volumes of all gases, under the same temperature and pressure, contain the same number of particles
what is the definition of one mole?
one mole contains exactly 6.02 x 10^23 elementary entities
what is the formula for the relative molecular mass
average mass of one molecule of a substance divided by 1/12 the mass of one atom of C
what is the gas pressure a gauge of?
- the frequency of collisions between gas particles and the walls of container
- the force of collisions between gas particles and the walls of container
what is the definition of molar volume?
molar volume of any gas is the volume occupied by 1 mole of the gas at a specified temperature and pressure. (no need to memorise values)
what does boyle’s law state?
at constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure
what does charles’ law state?
at constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature
what does avogardo’s law state?
at constant temperature and pressure, the volume and pressure of a gas is directly proportion to the number of molecules present
what is the ideal gas equation?
PV = nRT, where R is the molar gas constant, and n is the number of moles of the gas
what is the combined gas equation?
P1V1/T1 = P2V2/T2, where P1, V1 and T1 relate to the gas in its initial condition, and P2, V2 and T2 relate to the gas in its final condition
when is the combined gas equation useful?
the combined gas equation is useful for changes to gas systems where the amount of the gas n, is kept constant
what is dalton’s law of partial pressures?
if a mixture of non-reacting gases is confined in a container, each gas occupies the volume of the container and exerts its own pressure on the walls of the container as if it alone were present
what is the partial pressure of each gaseous component directly dependent on?
the number of moles of each gas in the mixture
using the concept of partial pressures, what is the total pressure of all the gaseous components?
ntotal(RT/V)
what is the mole fraction of an individual gas?
na/ntotal, where a is the individual gas. therefore, the partial pressure of an individual gas is directly proportional to its mole fraction in the mixture
what are the following assumptions about ideal gases?
- the size of the gas particles is so small compared to the space between them, that we can assume that the particles themselves have negligible volume
- the intermolecular forces of attraction between gas particles are negligible
- collisions between gas particles, and their collisions with the walls of the container, are perfectly elastic, meaning there is no net loss or gain of KE during collison
why are real gases seen to approach ideal gas behaviour at low pressures?
at low pressures, the gaseous molecules are relatively far apart, and 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. intermolecular forces are negligible as the particles are far apart
why are real gases seen to approach ideal gas behaviour at high temperatures?
at high temperatures, gas particles have enough kinetic energy to overcome intermolecular forces, which can thus be considered insignificant
under which two conditions would the assumptions about ideal gases remain invalid?
- the gas particles have negligible volume compared to the volume of the container
- the intermolecular forces of attraction between gas particles are negligible. molecules with a larger electron cloud and hence stronger dispersion forces, would deviate more from ideal behaviour. molecules with other intermolecular forces like pdpd and hydrogen also need to be considered
why would ideal gases never condense into liquids?
liquefaction is a key property of real gases that is not predicted by the kinetic molecular theory of gases, as it requires the action of intermolecular forces in order to occur
why can gas particles no longer be considered to have negligible volume compared to the volume of the container at high pressures?
at high pressures, the volume of the container decreases, and the molecules are pushed closely together and take up a significant portion of the container volume, resulting in less space in which the molecules can move. thus, it is no longer valid to assume that its volume is negligible compared to the container volume
why do intermolecular forces between gas particles become significant at low temperatures?
as temperature is lowered, the KE of the gas particles decreases, causing them to move more slowly and intermolecular forces to become more significant, causing collisions to become inelastic. eventually, it reaches a point where the particles can no longer overcome the intermolecular forces, at which point real gases liquefy
what is the relationship between average kinetic energy of gas particles and the absolute temperature?
the average kinetic energy of the gas particles is directly proportional to the absolute temperature. so at a particular temperature, all types of gaseous particles have the same average kinetic energy, but the individual molecules are moving at different speeds
what is the boltzmann distribution curve?
the speed distribution of the molecules shifts toward higher speeds and becomes less sharply peaked as the temperature of the gas sample is increased. even at low temperatures, a small number of molecules have high speed and KE, and this number increases with temperature, while the number of molecules with low speed and KE becomes smalller
