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Gases. I. Physical Properties. Real vs. Ideal Gases:. Ideal gas = an imaginary gas that conforms perfectly to all the assumptions of the kinetic-molecular theory We will assume that the gases used for the gas law problems are ideal gases . Real Gases vs. Ideal Gases:. - PowerPoint PPT Presentation
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I. Physical Properties
Gases
Real vs. Ideal Gases: Ideal gas = an imaginary gas that
conforms perfectly to all the assumptions of the kinetic-molecular theory• We will assume that the gases
used for the gas law problems are ideal gases.
Real Gases vs. Ideal Gases:Real gas = a gas that does not
behave completely according to the assumptions of the kinetic-molecular theory.
All real gases deviate to some degree from ideal gas behavior. However, most real gases behave nearly ideally when their particles are sufficiently far apart and have sufficiently high kinetic energy.
Causes of non-ideal behavior:The kinetic-molecular theory is
more likely to hold true for gases whose particles have NO attraction for each other.
Causes of non-ideal behavior:
1. High pressure (low volume): • Space taken up by gas particles
becomes significant• Intermolecular forces are more
significant between gas particles that are closer together
Causes of non-ideal behavior:
2. Low temperature: •Gas particles move slower so intermolecular forces become more important
Kinetic Molecular Theory (KMT) = the idea that particles of matter are always in motion and that this motion has consequences.
theory developed in the late 19th century to account for the behavior of the atoms and molecules that make up matter
KMT (Kinetic Molecular Theory)
based on the idea that particles in all forms of matter are always in motion and that this motion has consequences
KMT (Kinetic Molecular Theory)
SOLID LIQUID GAS
can be used to explain the properties of solids, liquids, and gases in terms of the energy of particles and the forces that act between them
KMT (Kinetic Molecular Theory)
B. Kinetic Molecular Theory – Ideal gas
KMT describing particles in an IDEAL gas:• have no volume.• have elastic collisions. • are in constant, random motion.• don’t attract or repel each other.• have an avg. KE directly related
to Kelvin temperature.
C. Kinetic Molecular Theory - Real Gases
KMT describing particles in a REAL gas:• have their own volume• attract each other
Gas behavior is most ideal…• at low pressures• at high temperatures• in nonpolar atoms/molecules
D. Characteristics of Gases- using KMT
Gases expand to fill any container.KMT - gas particles move rapidly
in all directions without significant attraction or repulsion between particles
D. Characteristics of Gases- using KMT
Gases are fluids (like liquids). KMT - No significant attraction or
repulsion between gas particles; glide past each other
D. Characteristics of Gases- using KMT
Gases have very low densities.• KMT - particles are so much
farther apart in the gas state
sodium in the solid state:
sodium in the liquid state:
sodium in the gas state:
Gases can be compressed.• KMT - gas particles are far apart
from one anther with room to be “squished” together
D. Characteristics of Gases- using KMT
Gases undergo diffusion & effusion.• KMT – gas particles move in
continuous, rapid, random motion
D. Characteristics of Gases- using KMT
Effusion
E. Temperature
ºF
ºC
K
-459 32 212
-273 0 100
0 273 373
32FC 95 K = ºC + 273
Always use Kelvin temperature when working with gases.
F. Pressure
areaforcepressure
Which shoes create the most pressure?
F. PressureWhy do gases exert
pressure? • Gas particles exert a
pressure on any surface with which they collide!More collisions =
increase in pressure!
F. PressureBarometer = measures
atmospheric pressureThe height of the Hg in
the tube depends on the pressure• The pressure of the
atmosphere is proportional to the height of the Hg column, so the height of the Hg can be used to measure atmospheric pressure! Mercury Barometer
F. PressurePressure UNITS
101.3 kPa (kilopascal)
1 atm
760 mm Hg
760 torr
G. STP
Standard Temperature & Pressure
0°C (exact) 1 atm (exact)
273 K 101.3 kPa
760 mm Hg (exact)
760 torr (exact)
STP
II. The Gas Laws
Gases
GAS LAW PROBLEMS- MUST USE KELVIN
K = ºC + 273
A. Boyle’s Law
P
V
PV = k
Volume (mL)
Pressure (torr)
P·V (mL·torr)
10.0 760.0 7.60 x 103 20.0 379.6 7.59 x 103 30.0 253.2 7.60 x 103 40.0 191.0 7.64 x 103
A. Boyle’s LawThe pressure and volume
of a gas are INVERSELY related• at constant mass & temp
P
V
P1V1 = P2V2
A. Boyle’s Law•Real life application: When you breathe, your diaphragm moves downward, increasing the volume of the lungs. This causes the pressure inside the lungs to be less than the outside pressure so air rushes in.
A. Boyle’s Law•Ex: Halving the volume leads to twice the rate of collisions and a doubling of the pressure.
A. Boyle’s Law
• As the volume increases, the pressure decreases!
kTV
V
T
B. Charles’ LawVolume
(mL) Temperature
(K) V/T
(mL/K) 40.0 273.2 0.146 44.0 298.2 0.148 47.7 323.2 0.148 51.3 348.2 0.147
V1_ = V2__
T1 T2
V
T
B. Charles’ LawThe volume and
temperature (K) of a gas are DIRECTLY related • at constant mass &
pressure
B. Charles’ LawReal life application: Bread dough
rises because yeast produces carbon dioxide. When placed in the oven, the heat causes the gas to expand, and the bread rises even further.
B. Charles’ LawAs temperature decreases, the
volume of the gas decreases
Liquid nitrogen’s temp. is about 63K or -210 ºC or -346 ºF!
B. Charles’ LawAs
temperature decreases, the volume of the gas decreases
B. Charles’ Law
NUMBER OF PARTICLES & PRESSURE ARE CONSTANT!!!
kTP
P
T
C. Guy-Lussac’s LawTemperature
(K) Pressure
(torr) P/T
(torr/K) 248 691.6 2.79 273 760.0 2.78 298 828.4 2.78 373 1,041.2 2.79
_P1_ = P2__
T1 T2
P
T
C. Guy-Lussac’s LawThe pressure and
temperature (K) of a gas are DIRECTLY RELATED • at constant mass &
volume
C. Guy-Lussac’s LawWhen the temp. of a gas increases
(KE increases) and gas particles move faster and hit container walls more frequently and collisions are more forceful
C. Guy-Lussac’s LawReal life application: The air
pressure inside a tire increases on a hot summer day.
= kBoyles PV PGuy Lussac’sT
VCharles’ T
PVT
D. Combined Gas Law
P1V1
T1=
P2V2
T2
E. EXAMPLE ProblemsYou need your calculators!
C. Johannesson
GIVEN:
V1 = 473 cm3
T1 = 36°C = 309KV2 = ?T2 = 94°C = 367K
WORK:
Gas Law Problems1.) A gas occupies 473 cm3 at 36°C. Find its volume at
94°C. CHARLES’ LAW V1_ = V2__
T1 T2
T V
V2= (473 cm3)(367 K)
(309 K)
V2 = 562 cm3
V2 = V1T2
T1
C. Johannesson
GIVEN:
V1 = 100. mLP1 = 150. kPaV2 = ?P2 = 200. kPa
WORK:V2 = P1V1
P2
Gas Law Problems2.) A gas occupies 100. mL at 150. kPa. Find its
volume at 200. kPa.
BOYLE’S LAW P1V1 = P2V2
P V
V2 = (150.kPa)(100.mL)
200.kPa
V2 = 75.0 mL
C. Johannesson
GIVEN:V1 = 7.84 cm3
P1 = 71.8 kPaT1 = 25°C = 298 KV2 = ?P2 = 101.3 kPaT2 = 273 K
WORK:V2 = P1V1T2
P2T1
V2 = (71.8 kPa)(7.84 cm3)(273 K)
(101.3 kPa) (298 K)
V2 = 5.09 cm3
Gas Law Problems3.) A gas occupies 7.84 cm3 at 71.8 kPa & 25°C. Find its
volume at STP.
P T V
COMBINED GAS LAW P1V1 = __P2V2
T1 T2
C. Johannesson
GIVEN:
P1 = 765 torrT1 = 23°C = 296KP2 = 560. torrT2 = ?
WORK:
T2 = P2T1
P1
Gas Law Problems4.) A gas’ pressure is 765 torr at 23°C. At what
temperature will the pressure be 560. torr?
P T
T2 = (560. torr)(296K)
765 torrT2 = 217 K
GUY LUSSAC’S LAW P1_ = P2__
T1 T2
III. Standard Molar Volume,
Gas Densities & Molar Mass
Gases
A. Standard Molar VolumeThe volume occupied by one mole
of a gas at STP • 22.41410 L/mol or about 22.4
L/mol In other words, one mole of any
ideal gas at STP will occupy 22.4 L
B. Example Problems-standard molar volume
1.) A chemical reaction is expected to produce 0.0680 mol of oxygen gas. What volume of gas in L will be occupied by this gas sample at STP?
0.068 mol O2 22.4 L O2
1 mol O2
=
1.52 L O2
B. Example Problems-standard molar volume
1.) A chemical reaction is expected to produce 0.0680 mol of oxygen gas. What volume of gas in L will be occupied by this gas sample at STP?
0.068 mol O2 22.4 L O2
1 mol O2
=
1.52 L O2
B. Example Problems-standard molar volume
2.) A chemical reaction produced 98.0 mL of SO2 at STP. What mass (in grams) of the gas was produced?
98.0 mL SO2 1 L SO2 1 mol O2
=
0.280 g SO2
22.4 L SO21000 mL SO2 1 mol SO2
64.064 g SO2
C. DensityThe ratio of an object’s mass to its
volume.D = M V The volume (and density) of a gas
will change when pressure and temperature change. • Densities are usually given in g/L at
STP
C. DensityFor an ideal gas:
Density (at STP) = molar mass standard molar volume
D. Example Problems- density
1.) What is the density of CO2 in g/L at STP?
44.009 g CO2 1 mol CO2
1 mol CO2 22.4 L CO2
=
1.96 g/L CO2
D. Example Problems- density2.) What is the molar mass of a gas
whose density at STP is 2.08 g/L.
2.08 g 22.4 L1 L 1 mol =
46.6 g/mol
IV. More Gas Laws: Ideal Gas
Law & Avogadro's Law
Gases
knV
V
n
A. Avogadro’s LawEqual volumes of gases
contain equal numbers of moles• at constant temp &
pressure
PVTVn
PVnT
B. Ideal Gas Law
= kUNIVERSAL GAS
CONSTANTR=0.0821
Latm/molKR=8.315
dm3kPa/molK
= R
You don’t need to memorize these values!
• Boyle’s, Charles, and Avogadro’s laws are contained within the ideal gas law!
• Merge the Combined Gas Law with Avogadro’s Law:
B. Ideal Gas Law
GAS CONSTANTR=0.0821
Latm/molKR=8.315
dm3kPa/molK
PV=nRT
You don’t need to memorize these values!
P= pressureV = volume n = molesT = temperature (in K)R = the ideal gas constant
C. Johannesson
GIVEN:P = ? atmn = 0.412 molT = 16°C = 289 KV = 3.25 LR = 0.0821Latm/molK
WORK:P = nRT
VP =(0.412 mol)(0.0821Latm/molK) (289K)
3.25L
P = 3.01 atm
C. Ideal Gas Law Problems1.) Calculate the pressure in atmospheres of
0.412 mol of He at 16°C & occupying 3.25 L.
PV = nRT
C. Johannesson
GIVEN:
V = ?n = 85 gT = 25°C = 298 KP = 104.5 kPaR = 8.315 dm3kPa/molK
C. Ideal Gas Law Problems2.) Find the volume of 85 g of O2 at 25°C
and 104.5 kPa. PV = nRT
= 2.7 mol
WORK:
85 g O2 1 mol O2 = 2.7 mol 31.998 g O2 O2
V = nRT PV =(2.7 mol)(8.315 dm3kPa/molK) (298K)
104.5 kPa
V = 64 dm3
C. Johannesson
GIVEN:
V = 1.00 Ln = ?T = 28°C = 301 KP = 98.7 kPaR = 8.314 LkPa/molK
C. Ideal Gas Law Problems3.) At 28C and 98.7 kPa, 1.00 L of an unidentified gas has
a mass of 5.16 g. Calculate the number of moles of gas present and the molar mass of the gas. PV = nRT
WORK:
n = PV RT
(98.7 kPa) (1.00L)(8.314 LkPa/molK) (301 K)
n = 0.0394 mol
m = 5.16g 0.0394 mol
=
= 131 g/mol
Gases
IV. Two More Laws: Dalton’s Law & Graham’s Law
A. Dalton’s Law The total pressure of a mixture
of gases equals the sum of the partial pressures of the individual gases.
Ptotal = P1 + P2 + P3 ...
A. Dalton’s Law
Ptotal = P1 + P2 + P3 ...
A. Dalton’s Law
When a H2 gas is collected by water displacement, the gas in the collection bottle is actually a mixture of H2 AND water vapor.
C. Johannesson
GIVEN:PH2 = ?Ptotal = 94.4 kPaPH2O = 2.72 kPa
WORK:Ptotal = PH2 + PH2O
94.4 kPa = PH2 + 2.72 kPa
PH2 = 91.7 kPa
A. Dalton’s Law1.) Hydrogen gas is collected over water at
22.5°C. Find the pressure of the dry gas if the atmospheric pressure is 94.4 kPa.
Look up water-vapor pressure on gas formula sheet for
22.5°C.
Sig Figs: Round to least number of decimal places.
The total pressure in the collection bottle is equal to atmospheric pressure and is a mixture of H2 and water vapor.
C. Johannesson
GIVEN:Pgas = ?Ptotal = 742.0 torrPH2O = 42.2 torr
WORK:Ptotal = Pgas + PH2O
742.0 torr = PH2 + 42.2 torr
Pgas = 699.8 torr
2.) A gas is collected over water at a temp of 35.0°C when the barometric pressure is 742.0 torr. What is the partial pressure of the dry gas?
Look up water-vapor pressure on gas formula sheet for
35.0°C.
Sig Figs: Round to least number of decimal places.
A. Dalton’s Law
The total pressure in the collection bottle is equal to barometric pressure and is a mixture of the “gas” and water vapor.
B. Graham’s LawDiffusion =
• Spreading of gas molecules throughout a container until evenly distributed.
Effusion =• Passing of gas molecules
through a tiny opening in a container
B. Graham’s Law
KE = ½mv2
Speed of diffusion/effusion• Kinetic energy is determined by
the temperature of the gas.• At the same temp & KE, heavier
molecules move more slowly.Larger m smaller v
B. Graham’s LawKE = kinetic energym= molar massv = velocity
KE = ½mv2
C. Johannesson
B. Graham’s LawGraham’s Law = Rate of effusion of
a gas is inversely related to the square root of its molar mass.• The equation shows the ratio of
Gas A’s speed to Gas B’s speed.
A
B
B
A
mm
vv
1.) Determine the relative rate of effusion for krypton and bromine.
1.381
Kr effuses 1.381 times faster than Br2.
Kr
Br
Br
Kr
mm
vv 2
2
A
B
B
A
mm
vv
g/mol83.80 g/mol159.80
C. Graham’s Law Problems
The first gas is “Gas A” and the second gas is “Gas B”. Relative rate mean find the ratio “vA/vB”.
C. Johannesson
2.) A molecule of oxygen gas has an average speed of 12.3 m/s at a given temp and pressure. What is the average speed of hydrogen molecules at the same conditions?
A
B
B
A
mm
vv
2
2
2
2
H
O
O
H
mm
vv
g/mol 2.02g/mol32.00
m/s 12.3vH 2
C. Graham’s Law Problems
3.980m/s 12.3
vH 2
m/s49.0 vH 2
Put the gas with the unknown
speed as “Gas A”.
C. Johannesson
3.) An unknown gas diffuses 4.0 times faster than O2. Find its molar mass.
Amg/mol32.00 16
A
B
B
A
mm
vv
A
O
O
A
mm
vv 2
2
Amg/mol32.00 4.0
16g/mol32.00 mA
2
Amg/mol32.00 4.0
g/mol2.0
C. Graham’s Law Problems
The first gas is “Gas A” and the second gas is “Gas B”. The ratio “vA/vB” is 4.0.
Square both sides to get rid of the square
root sign.
Gases VI. Gas Stoichiometry at Non-STP Conditions
A. Gas Stoichiometry Moles Liters of a Gas:
• STP - use 22.4 L/mol • Non-STP - use ideal gas law
Non-STP• Given liters of gas?
start with ideal gas law• Looking for liters of gas?
start with stoichiometry conv.
B. Volume Mass Stoich.Thinking:Gas volume A moles A moles B mass B
If gas A is at STP:
? L “A” 1 mol “A” ? mol “B” molar mass(g) “B” 22.4 L “A” ? mol “A” 1 mol “B”
Mole Ratio
B. Volume Mass Stoich.If gas “A” is NOT at STP:
Use PV = nRT to find moles of “A”
? mol “A” ? mol “B” molar mass (g) “B”
? mol “A” 1 mol “B”
C. Johannesson
WORK:n = PV RTn= (97.3 kPa) (15.0 L) (8.315dm3kPa/molK) (294K)
n = 0.597 mol O2
B. Volume Mass1.) How many grams of Al2O3 are formed from
15.0 L of O2 at 97.3 kPa & 21°C?
GIVEN:P = 97.3 kPaV = 15.0 Ln = ?T = 21°C = 294 KR = 8.315 dm3kPa/molK
4 Al + 3 O2 2 Al2O3 15.0 L
non-STP ? gGiven liters: Start with
Ideal Gas Law and calculate moles of O2.
NEXT
2 mol Al2O3
3 mol O2
B. Volume Mass1.) How many grams of Al2O3 are formed
from 15.0 L of O2 at 97.3 kPa & 21°C?
0.597mol O2 = 40.6 g Al2O3
4 Al + 3 O2 2 Al2O3
101.96 g Al2O3
1 molAl2O3
15.0Lnon-STP
? gUse stoich to convert moles of O2 to grams Al2O3.
1 mol SO2
22.4L SO2
B. Volume Mass2.) What mass of sulfur is required to
produce 12.6 L of sulfur dioxide at STP according to the equation:
S8 (s) + 8 O2 (g) 8 SO2 (g)
12.6 L SO2 = 18.0 g
S8
256.582 g S8
1 molS8
12.6L? g
1 mol S8
8 molSO2
C. Mass VolumeThinking:Mass A moles A moles B gas volume B
If gas “B” is at STP:
? g “A” 1 mol “A” ? mol “B” 22.4 L “B” molar mass (g) “A” ? mol “A” 1 mol “B”
C. Mass Volume If gas “B” is NOT at STP:
? g “A” 1 mol “A” ? mol “B” molar mass (g) “A” ? mol “A”
• Then use PV = nRT to find volume of “B”
C. Mass Volume1.) What volume of chlorine gas at STP is needed to react
completely with 10.4 g of sodium to form NaCl?
2 Na + Cl2 2 NaCl 10.4 g ? L
10.4 g Na
22.990 g Na
22.4 L Cl21 mol Cl2
2 mol Na 1 mol Cl2
1 mol Na
= 5. 07 L Cl2
1 molCaCO3
100.09g CaCO3
C. Mass Volume2.) What volume of CO2 forms from 5.25 g
of CaCO3 at 103 kPa & 25ºC?
5.25 gCaCO3 = 1.26 mol CO2
CaCO3 CaO + CO2
1 molCO2
1 molCaCO3
5.25 g ? Lnon-STPLooking for liters: Start with stoich
and calculate moles of CO2.
Plug this into the Ideal Gas Law to find liters.
C. Johannesson
WORK:V = nRT PV = (1.26mol)(8.315dm3kPa/molK) (298K)
(103 kPa)
V = 30.3 dm3 CO2
C. Mass Volume
2.) What volume of CO2 forms from 5.25 g of CaCO3 at 103 kPa & 25ºC?
GIVEN:P = 103 kPaV = ?n = 1.26 molT = 25°C = 298 KR = 8.315 dm3kPa/molK