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16 Cards in this Set
- Front
- Back
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What four variables are used to desrcibe the state of all gas samples? |
PV=nRT Pressure Volume Temperature n, moles of gas |
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Unit conversions for gas pressure what is 1 atmosphere equal to in mmHg? |
usually expressed in atmospheres (atm) or milimeters of mercury (mmHg) - - equivalent to Torr or Pascal (Pa) 1 atm= 760 mmHg = 760 Torr = 101.32kPa |
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What medical device measures blood pressure? - how does this relate to the Barometer? - How does mercury rise in a Barometer?
- what creates a downward force on the pool of mercury at the base of a barometer? - how does this affect the mercury in the column? |
Sphygomomanometer measures BP - utilizes Barometer concept concept: the atmospheric pressure creates downward force onto pool mercury at base of barometer 1. Mercury exerts force back(weight) - based on its density 2. if the F(air) > F(Hg) then the Hg will rise if Force of air is less than down Hg Force, the column will fall *Height is proportional to the pressure applied - can get systolic, diastolic pressures |
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Standard Temperature and Pressure: STP What is the volume measured in? what is the temperature constant at? and given as.. what is the pressure constant at? |
STP: for gases T= 273k, or 0 degree Celsius P= 1atm **remember standard start is different - 298k (25C) instead and 1M, 1 atm |
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Ideal Gas Do ideal gases occupy volume? What about inter molecular forces? What is the idea gas law |
Ideal gases... occupy NO volume and have No inter-molecular forces Ideal gas law: shows relationship w/ four variables PV=nRT R: universal gas constant 8.314 (J/ K x mol) |
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How can you implement density into the ideal gas law? @STP, what is the volume of one mole of gas? |
PV=nRT n=mol mol= mass/ molar mass p= density = mass(g) / volume (L) -- rearrange PV=nRT to p= (n/V)= P x molar mass Volume = 22.4L @ STP - can use density@STP to find molar mass mm= d(stp) x 22.4 L/mol |
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Avogadros principle w/ Gases - constant temperature - constant pressure - How will these affect the volume of all gases? |
@ constant pressure and temperature Volumes are directly proportional to mol gas -equal amounts of all gases @same temp, P will have equal volumes n1/ V1 = n2/ V2 *given moles and initial volume, can calc. the new volume w/ addition of gas, solve for V2 |
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Boyles Law @Isothermal conditions... How does the volume relate to the gas? what does isothermal mean? |
Isothermal: constant Temperature Boyles Law: @constant T, - the Pressure is inverse Proportion to Volume *derivation of the ideal gas law PV= k(constant) * if pressure goes up, volume has to go down vise versa OR P1V1 = P2V2 |
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Charle's Law @isobaric conditions - how does Volume relate to Temperature? Hint:use the PV=nRT equation and set P as constant(k) |
Charles' Law @constant pressure(isobaric) -moles, n, would be also constant Volume Gas is Proportional to Temperature or k= V/ T .. better seen as T= V/k if T inc^ then V inc^ |
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Gay-Lussac's Law @isovolumetric conditions how does Pressure relate to Volume? -hint: use PV=nRT, set constants |
@ constant Volume(isovolumetric) n and V are constant in PV=nRT P/ T = k... or P=T x k if T inc^, then P inc^ |
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Dalton's Law of Partial pressures |
Dalton's law of partial pressures Ptot= P(A) + P(B) + P(C)..... Partial pressure: pressure exerted by each individual gas * if gases do not react, - gases will act as if they are only gas in container |
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Henry's Law of .. vapor pressure @ surface of liquids describe the relationship between solubility of gases....and pressure How does concentration relate? |
Law: w/ applied pressure, [gas] would increase or decrease [A] = Kh x Pa * solubility and pressure directly related (concentration) Inc^ Pressure --> Inc^ solubility ex: increase partial pressure of oxygen raises, amount dissolved in blood also elevates Vapor pressure: pressure exerted by evaporated particles above surface of liquid Concentration: is solubility *remember: Evaporation: dynamic process, req. molecules at surface gain enough energy to escape into gas phase |
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Kinetic Molecular Theory -explains why gases act as they do Gas Laws (previously theorized) only describe |
Gaseous Molecular Behavior Assumptions: - particles of gases' volumes negligible vs container - gases exhibit no intermolecular forces -Gases move constantly, collide w/ wall and each other - Collisions: are elastic-- conservation of momentum and kinetic energy - average KE of each gas is proportional to temperature(absolute) - will be same for all gases@same temperature KE= 1/2 mv^2 = 3/2 Kb T ** speed of gas directly related to Temp. - Kb: constant=1.38x10^-23 *@same temp, KE same for all gases **larger molecules will move slower |
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Diffusion: how does kinetic energy relate to diffusion? Graham's Law - How do rate of two gases change with different masses? |
diffusion: movement molecules down concentration gradient - high to low concentration * heavier gases diffuse more slowly( diff. KE) r1/ r1 = Sqrt (M2/ M1) *if mass is x4 as big, the rate will be half as fast |
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Effusion: |
Effusion: gas moves through small whole via Pressure |
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what will cause a deviation from IDEAL gases? |
High Pressure and Low temp - more intermolecular forces w/ molecules so close together, gases will actually expand |
