Ch. 8: The Gas Phase (Complete) Flashcards
why are gases classified as fluids?
because they can flow and take on the shapes of their containers
defn + aka + options: phases
the three different physical forms that matter can exist
aka: states
options: gas, liquid, solid
what are four characteristics of gas particles?
- the atoms or molecules in gas move rapidly
- they are far apart from each other
- only very weak intermolecular forces exist between gas particles
- easily, but not infinitely, compressible
what are the four variables that allow us to describe the state of a gaseous sample?
pressure (P)
volume (V)
temperature (T)
and number of moles (n)
what are the 4 common units of gas pressures + what are their conversions?
1 atm = 760 mmHg = 760 torr = 101.325 kPa
what is the SI unit for pressure?
pascal (Pa)
defn + unit: sphygmomanometer
medical devices that measure blood pressure
unit: mmHg
explain why the mercury rises in a barometer (4)
- atmospheric pressure creates a downward force on the pool of mercury at the base of the barometer while the mercury in the column exerts an opposing force (its weight) based on its density
- the weight of the mercury creates a vacuum in the top of the tube
- when the external air exerts a higher force than the weight of the mercury in the column, the column rises
- when the external air exerts a lower force than the weight of the mercury, the column falls
how can a reading be obtained on a barometer?
by measuring the height of the mercury column (in mm), which will be directly proportional to the atmospheric pressure being applied
is atmospheric pressure the only external pressure that can exert this force?
no! for example a clinical blood pressure cuff creates a force that is opposed by the person’s systolic and diastolic arterial blood pressure
what units are typically used for volume of gas and temperature?
volume: liters, milliliters
temperature: kelvin, celsius sometimes
defn: standard temperature and pressure (STP)
many processes involving gases take place under these conditions
273 K (0 C) and 1 atm
defn: ideal gas
a hypothetical gas with molecules that have no intermolecular forces and occupy no volume
under what circumstances are real gases similar to ideal gases? different from ideal gases?
real gases deviate from this ideal behavior at HIGH pressures, LOW volumes, and LOW temperatures
many COMPRESSED real gases demonstrate behavior that is close to ideal
why do real gases deviate from gas laws at high pressures and low temperatures?
because of intermolecular forces or volume effects
eqn: ideal gas law
PV = nRT
P = pressure
V = volume
n = number of moles
T = temperature
R = ideal gas constant
what are the two values for the ideal gas constant, R? (don’t need to memorize, but recognize which to use)
what are the two circumstances that the ideal gas law is used for?
- determine the missing term when given all of the others
- calculate the change in a term while holding two of the others constant
what is the ideal gas law most commonly used to solve for?
volume and pressure at any given temperature and number of moles
what are two things you can use the ideal gas law to help solve for that you might not initially think of?
- gas density
- molar mass
defn + symbol: density
the ratio of the mass per unit volume of a substance
what unit is typically used for densities of gases?
g/L
how do we rearrange the ideal gas law to calculate the density of the gas?
a mole of an ideal gas at STP occupies what volume?
22.4 L
equation: combined gas law
where the subscripts 1 and 2 refer to the two states of the gas (at STP and at the conditions of actual temp and pressure, for example)
how do we experimentally calculate the molar mass of a gas using the equation derived from the ideal gas law? (5)
- the pressure and temperature of a gas contained in a bulb of a given volume are measured, and the mass of the bulb with the sample is measured
- the bulb is evacuated (the gas is removed) and the mass of the empty bulb is determined
- the mass of the bulb with the sample minus the mass of the evacuated bulb gives the mass of the sample
- the density of the sample is determined by dividing the mass of the sample by the volume of the bulb
- this gives the density at the given temperature and pressure
defn: Avogadro’s principle
all gases at a constant temperature and pressure occupy volumes that are directly proportional to the number of moles of gas present
equal amounts of all gases at the same temperature and pressure will occupy equal volumes
eqn + summary: Avogadro’s principle
as the number of moles of gas increases, the volume increases in direct proportion
eqn + summary + layman’s term: Boyle’s law
for a given gaseous sample held at constant temperature, the volume of the gas is inversely proportional to its pressure
isothermal compression
eqn + summary + layman’s term: Charles’s law
at constant pressure, the volume of a gas is proportional to its absolute temperature, expressed in kelvin
isobaric expansion
eqn + layman’s term: Gay-Lussac’s Law
isovolumetric heating
why is the pressure exerted by each gas in a mixture of 2 or more gases that do not chemically interact, equal to the pressure that the gas would exert if it were the only one in the container?
when two or more gases that do not chemically interact are found in one vessel, each gas will behave independently of the others
that is, each gas will behave as if it were the only gas in the container
defn: partial pressure
the pressure exerted by each individual gas
defn + eqn: Dalton’s law of partial pressures
the total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individual components
eqn: the partial pressure of a gas is related to its mole fraction
defn: vapor pressure
the pressure exerted by evaporated particles above the surface of a liquid
process + eqn: Henry’s Law
where [A] is the concentration of A in solution
kH is Henry’s constant
PA is the partial pressure of A
what does the value of Henry’s constant depend on?
the identity of the gas
does kinetic molecular theory apply to ideal gases or real gases?
they were developed in reference to ideal gases, but can be applied to real gases with reasonable accuracy
what is the difference between the goals of kinetic molecular theory and the gas laws?
kinetic molecular theory EXPLAINS the behavior of gases, while the other laws just DESCRIBE the behavior of gases
what are the 5 assumptions of the kinetic molecular theory?
- Gases are made up of particles with volumes that are negligible compared to the container volume
- Gas atoms or molecules exhibit no intermolecular attractions or repulsions
- gas particles are in continuous, random motion, undergoing collisions with other particles and the container walls
- collisions between any two gas particles (or between particles and the container walls) are elastic, meaning that there is conservation of both momentum and kinetic energy
- the average kinetic energy of gas particles is proportional to the absolute temperature of the gas (in kelvin), and it is the same for all gases at a given temperature, irrespective of chemical identity or atomic mass
what is a good tangible way to imagine gas particles as?
as little rubber balls bouncing off each other and off the walls of the container
eqn + statement: kinetic molecular theory of gases
the average kinetic energy of a gas particle is proportional to the absolute temperature of the gas
where kB is the boltzmann constant
value + func: boltzmann constant
serves as a bridge between the macrosopic and microscopic behaviors of gases (that is, as a bridge between the behavior of the gas as a whole and the individual gas molecules)
defn + eqn + symbol: root-mean-square speed
a way to find an average speed of the molecules of a gas by finding the average kinetic energy per particle and then calculate the speed to which it corresponds
R is the ideal gas constant
T is the temperature
M is the moral mass
func: Maxwell-Boltzmann distribution curve
shows the distribution of gas particle speeds at a given temperature
at higher temperatures, do gas molecules move at higher or lower speeds?
higher speeds
defn: diffusion
the movement of molecules from high concentration to low concentration through a medium
do heavier gases or lighter gases diffuse more slowly? why?
heavier gases diffuse more slowly than lighter ones because of their differing average speeds
because all gas particles have the same average kinetic energy at the same temperature, it must be true that particles with greater mass travel at slower average speed
defn + eqn + example: Graham’s law
under isothermal and isobaric conditions, the rates at which two gases diffuse are inversely proportional to the square roots of their molar masses
where r1 and r2 are the diffusion rates of gas 1 and gas 2, respectively, and M1 and M2 are the molar masses of gas 1 and gas 2
example: a gas that has a molar mass four times that of another gas will travel half as fast as the lighter gas
defn: effusion
the flow of gas particles under pressure from one component to another through a small opening
eqn: effusion
the same as Graham’s law!
what is another commonality between effusion and diffusion?
they will both be slower for larger molecules
what happens as the pressure of a gas increases?
the particles are pushed closer and closer together
what happens as the condensation pressure for a given temperature is approached?
intermolecular attraction forces become more and more significant, until the gas condenses into a liquid
what happens to the volume of a gas at moderately high pressure? at extremely high pressure?
moderately high pressure: a gas’s volume is less than would be predicted by the ideal gas law due to intermolecular attraction
extremely high pressure: the size of the particles becomes relatively large compared to the distance between them, this causes the gas to take up a larger volume than would be predicted by the ideal gas law (the ideal gas law assumes that a gas can be compressed to take up zero volume)
what happens as the temperature of a gas decreases?
the average speed of the gas molecules decreases and the attractive intermolecular forces becomes increasingly significant
what happens as the condensation temperature is approached for a given pressure?
intermolecular attractions eventually cause the gas to condense to a liquid state
intermolecular attraction causes the gas to have a smaller volume than that would be predicted by the ideal gas law
what happens to a gas’s behavior the closer it is to its boiling point/condensation point?
the less ideal it acts
what happens to gases at extremely low temperatures? why?
gases will occupy more space than predicted by the ideal gas law because the particles cannot be compressed to zero volume
func + eqn: van der Waals equation of state
one of several gas equations that attempt to correct for the deviations from ideality that occur when a gas does not closely follow the ideal gas law
where a and b are physical constants experimentally determined for each gas
what do the a and b terms in the van der Waals equation correct for?
the a term corrects for the attractive forces between molecules and as such will be smaller for gases that are small and less polarizable, larger for gases that are larger and more polarizable, and largest for polar molecules
the b term corrects for the volume of the molecules themselves (larger molecules have larger values of b)
which is usually larger, a or b?
a values are generally much larger than b
mnemonic for remembering a and b of the van der Waals equation
a = van der Waals term for the Attractive forces
b = van der Waals term for Big particles
