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| What are the units used to express gas pressures? How do you convert between them? |
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Definition
| 1 atm (atmosphere) = 760 mmHg = 760 torr |
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Term
| What is standard temperature and pressure (STP)? |
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Definition
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Term
| What are standard state conditions? |
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Definition
| 298 K (25 C), 1 atm and 1 M concentrations |
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Term
| How much volume does one mole of an ideal gas occupy at STP? |
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Definition
| 22.4 L, this is the same for all gases |
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Term
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Definition
PV = nRT
P is pressure (atm)
V is volume (L)
n is numer of moles
T is temperature (K)
R is the ideal gas constant = 8.21 x 10-2 (L•atm)/(mol•K) |
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Definition
| A hypothetical gas with molecules that have no intermolecular forces and occupy no volume; real gases deviate from this ideal at high pressures(low volumes) and low temperatures |
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Term
| How is density incorporated into the ideal gas law? |
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Definition
PV = nRT
where n = m (mass)/M (molar mass)
Therefore, PV = mRT/M
and ρ = m/V = PM/RT
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Definition
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Definition
n/V = K or n1/V1 = n2/V2
all gases at a constant temperature and pressure occupy volumes that are directly proportional to the number of moles of gas present; as teh number of moles of gas increases the volume increases in direct proportion |
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Term
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Definition
for a given gaseous sample held at constant temperature (isothermal conditions), the volume of the gas is inversely proportional to its pressure
PV = k or P1V1 = P2V2 |
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Definition
at constant pressure, the volume of a gas is proportional to its absolute temperature, in kelvins
V/T = k or V1/T1 = V2/T2 |
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Definition
constant volume, pressure and temperature are proportional
P/T = k or P1/T1 = P2/T2 |
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Term
| Dalton's law of partial pressures |
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Definition
| individual gas components of a mixture of gases will exert individual pressures in proportion to their mole fractions; the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases |
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Term
| How to determine partial pressure of a gas |
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Definition
PA = XAPT
where XA= moles of gas A/total moles of gas |
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Term
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Definition
the amount of gas dissolved in solution is directly proportional to the partial pressure of that gas at the surface of a solution
[A]1/P1 = [A]2/P2 = kH
[A] is the concentration of A in solution
P is the partial pressuer of A
kH is Henry's constant which depends on the identity of the gas |
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Term
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Definition
- gas particles have negligible volume
- gas particles do not have intermolecular attractions or repulsions
- gas particles undergo random collisions with each other and the walls of the container
- collisions between gas particles and the walls are elastic
- the average kinetic energy of the gas particles is directly proportional to temperature
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Term
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Definition
gases with lower molar masses will diffuse or effuse faster than gases with higher molar masses at the same temperature
v = √(T/mw) |
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Term
| How do real gases deviate from ideal behavior? |
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Definition
- at mdoerately high pressures, low volumes or low temperatures, real gases will occupy less volume than predicted by the ideal gas law because the particles have intermolecular attractions
- at extremely high pressures, low volumes, or low temperatures, real gases will occupy more volume than predicted by the ideal gas law because they particles occupy physical space
- the van der Waals equation of state is used to correct teh ideal gas law for intermolecular attractions and molecular volume
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