Principles of Chemistry I - Test 3: Gases, Nomenclature, and Quantitative Relationships

Pressure

force/area

Gas pressure

force exerted when gas molecules strike the area of a container wall

vacuum

opposite of pressure; a deficit of molecules

SI Pressure unit

Pascals (Pa); 1 Pa = 1N/m^2

1 barr =

10^5 Pa

1 atm =

101.325 kPa, 760 torr, 760 mmHg, 14 lb/in^2

barometer

-an instrument that measures atmospheric pressure
-the pressure at the bottom of the Hg tube is balanced by atmospheric pressure on the dish
-h is proportional to P

Most common pressure units

torr and atm

Hydrostatic pressure

pressure exerted by a fluid due to gravity

hydrostatic pressure equation

p=rho(h)(g)

wind

created by air moving from high to low pressure

manometer

-measures gas trapped in a container
-delta h is proportional to pressure

Why do hot air balloons rise?

-due to buoyant force
-hot air is less dense than normal air (less gas mlx)
-causes balloon to rise

the atmosphere

exerts a large pressure on objects on earth's surface

Atmospheric pressure analogy

like a bowling ball on a thumbnail

exosphere

>700 km

thermosphere

80-700 km

mesosphere

50-80 km

stratosphere

12-50 km

troposphere

Air density and temperature in troposphere

increase with distance from the earth

Other units

ppm, ppb, and ppt

Selenium

protects the body against an excess of heavy metals

Where does the jet stream exist?

stratosphere

state function

-depends on initial and final volume
-path independent
-denoted by a capital letter

Collisions of an ideal gas

must be completely elastic

ideal gas

hypothetical construct that real gases approximate under certain conditions

ideal gas law

-relates P, V, n, T
-most gases are ideal at room temp
-accurate for low P and moderate T

first hot air balloon

caught on fire due to use of hydrogen gas

Luer

end of syringe with a specific length and taper

Absolute zero

zero kelvin, -273.15 C

Celsius and Kelvin intervals

are equal

Charles' Law

-volume is proportional to temperature
-P and n are held constant
-V1/T1 = V2/T2

Charles' Law origin

extrapolates so P=0 @ T=-273 C

Charles' Law Graph (V and T)

linear

Boyle's Law

-volume varies inversely with pressure
-T and n are constant
-P1V1 = P2V2

Graph of P vs. V

parabolic

Graph of 1/P vs. V

linear

What prevents containers from collapsing?

collisions from gas molecules inside the container counteract the atmospheric pressure

STP

T=273.15 K
P=1 atm

standard molar volume

22.4 L

Gay-Lussac's Law

-pressure and temperature are proportional
-V and n are constant
-P1/T1 = P2/T2

Gay-Lussac's Law graph (P and T)

linear

Gay-Lussac's law synonym

Amontons' Law

Ideal Gas Law

PV=nRT

Combined Gas Law

PV/nT = R = PV/nT

Antoine Lavoisier

-father of modern chemistry
-made quantitative strides by studying gases

Gas density

d=MP/RT

Molar Mass Equation

M=m/n

Pressure in a Mixture of Gases

-individual gases do not affect each other's pressures
-each gas exerts the same pressure it would alone

Partial Pressure

pressure exerted by an individual gas in a mixture

Dalton's Law of Partial Pressures

Pt = Pa + Pb + Pc...

Mole Fraction (X)

-unit of concentration
-# of moles of component/ total # of moles

Mole Fraction equation

Pa = Xa(Pt) when Xa = na/nt

Collection of Gases over Water

-bottle is inverted in water dish
-reaction runs that produces a gas
-gas bubbles into bottle
-gas pressure = atmospheric pressure when the water level is equal in and out of the bottle

saturation of gas with vapor

Ptotal = Pgas + Pvapor

Vapor Pressure of Water

pressure exerted by water vapor in equilibrium with liquid water in closed container; depends on temperature

boiling point

the point where gas and liquid phase are in equilibrium

Mixing of Ideal gases

ideal gases are completely miscible and form homogeneous mixtures

lighter gas particles....

move faster

Mean Free Path

average distance a molecule travels between collisions; increases with decreasing V

Diffusion

-molecule put in closed container disperse quickly
-results in a equal concentration of gases

Rate of Diffusion

amount of gas passing thru area/unit of time

Effusion

escape of gas molecules through a tiny hole

Graham's Law of effusion

the rate of effusion is inversely proportional to molar mass

Graham's Law equation

ratea/rateb = SQrt(Mb)/SQrt(Ma)

compression

increases temperature

expansion

decreases temperature

pressure of chemical reactions in the lab

constant

maxwell-boltzmann distribution

distribution curve broadens with higher average

V total in ideal gas =

Vspace + Vparticles(0)

molecular velocity is proportional to

molecular mass

compression of liquids and solids

extremely difficult

Kinetic Molecular Theory

a simple microscopic model that effectively explains the gas laws

First Article of KMT

a gas consists of tiny molecules in random, rapid motion; molecules change direction after collision

Second Article of KMT

Gas molecules are small compared to the distances between them, making them compressible and miscible

Third Article of KMT

Gas molecules exert pressure thru collisions between molecules and molecule walls

Fourth Article of KMT

Gas molecules has small attraction/repulsion for each other

Fifth Article of KMT

Gas molecules make elastic collisions with each other and do not lose energy thru collision

Kinetic Energy depends on

temperature and n

kinetic energy and temperature

directly proportional

Peak on Maxwell Curve

most probable particle speed

Vrms on maxwell curve

after the peak when the graph switches from increasing to decreasing

Speed distribution flattens...

as temperature increases

Vrms

square root of average squares of speed where n= number of particles

Average KE =

1/2 Mvrms

KMT and Amonton's Law

-KE increases with T
-more collisions = higher P

KMT and Boyle's Law

-V decreases, decreasing wall area
-more frequent collision = higher P

KMT and Avogadro's Law

-V increase = n increase
-more mlxs require greater space to reduce collision and keep P constant

KMT and Dalton's Law

-gas molecules of one gas bombard walls regardless of other gases
-Ptotal = total bombardment from all gases

KMT and effusion

effusion depends on the average speed of molecules

Vrms a / Vrms b

SqrtMb/SqrtMa

low Pressure deviation

negative

high Pressure deviation

positive

Compressibility Factor (Z)

actual volume of one mole of gas divided by molar volume of an ideal gas at the same temperature and pressure

When Z=1

the gas is ideal

Gases at High pressure

-molecules begin to crowd and volume of the molecules becomes appreciable
-attractive force becomes more signficant

Gases at low temperature

begin to stick together and condense

Van der Waals equation

P = nRT = (P +an^2/V^2)(V-nb)

boiling point of polar molecules

high