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16 Cards in this Set

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  • Back

What four variables are used to desrcibe the state of all gas samples?





n, moles of gas

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

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

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

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


R: universal gas constant 8.314 (J/ K x mol)

How can you implement density into the ideal gas law?

@STP, what is the volume of one mole of gas?


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

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

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


P1V1 = P2V2

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^

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^

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

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


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

Kinetic Molecular Theory

-explains why gases act as they do

Gas Laws (previously theorized) only describe

Gaseous Molecular Behavior


- 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


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


Effusion: gas moves through small whole via Pressure

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