Ch5.1 – Gas Pressure
“Fill a tube with water, then invert it into a bowl of its own liquid,
what happens?”
vacuum
Patm
“The water will remain at a height where its weight creates a pressure
equal to the pressure exerted from the atmosphere”
vacuum
Patm
Barometer – device to measure atmospheric pressure.
Standard atmospheric pressure:
760 mm Hg (760 torr)
1 atm
vacuum
101,325 Pascals (14.7 psi)
Patm
Barometer – device to measure atmospheric pressure.
Standard atmospheric pressure:
760 mm Hg (760 torr)
1 atm
vacuum
101,325 Pascals (14.7 psi)
Patm
Ex1) Convert 49 torr to atms, and then to Pascals
Ideal gases – hypothetical substance.
Most gases behave ideally (obey the gas laws)
at low pressures and high temps.
vs
Kinetic Theory (for gases)
1. All matter is made of particles.
2. Particles in constant motion.
3. All collisions are perfectly elastic. (No energy lost)
Ideal Gas Law
P.V = n.R.T
R is the ideal gas law constant.
L ⋅ atm
R = 0.08206
K ⋅ mol
Ex2) A sample of hydrogen gas has a volume of 8.56L at a temp of 0oC
and a pressure of 1.5atm. Calculate the moles of H2 present.
The ideal gas law can be altered into the other gas laws:
P1.V1 = n1.R.T1
1. Boyle’s Law:
(moles and temp
held constant)
2. Charles’ Law:
(moles and pressure
held constant)
3. Guy Lussac’s Law:
(moles and volume
held constant)
4. Avagadro’s Law:
(pressure and temp
held constant)
5. Combine Gas Law:
(moles only
held constant)
P2.V2 = n2.R.T2
P1 ⋅ V1 P2 ⋅ V2
=
n1 ⋅ T1 n2 ⋅ T2
1. Boyle’s Law:
(moles and temp
P1.V1 = P2.V2
held constant)
2. Charles’ Law:
V1 V2
=
(moles and pressure
T1 T2
held constant)
3. Guy Lussac’s Law:
P1 P2
=
(moles and volume
T1 T2
held constant)
4. Avagadro’s Law:
V1 V2
=
(pressure and temp
n1 n2
held constant)
5. Combine Gas Law: P1 ⋅ V1
P2 ⋅ V2
=
(moles only
T1
T2
held constant)
P.V = k
V = b.T
P = c.T
V = a.n
k, b, c, a are
constants
Don’t copy, just listen…
Ex3) The following data was collected for 1 mol of NH3 gas
at a constant temp of 0oC. Calculate Boyle’s Law constant.
Exp
1
2
3
4
5
6
Pressure (atm) Volume (L)
0.1300
172.1
0.2500
89.28
0.3000
74.35
0.5000
44.49
0.7500
29.55
1.000
22.08
22.5
P.V 22.4
(L.atm) 22.3
22.2
22.1
.2 .4 .6 .8 1.0
P (atm)
P.V (L.atm)
Ex3) The following data was collected for 1 mol of NH3 gas
at a constant temp of 0oC. Calculate Boyle’s Law constant.
Exp
1
2
3
4
5
6
Pressure (atm) Volume (L)
0.1300
172.1
0.2500
89.28
0.3000
74.35
0.5000
44.49
0.7500
29.55
1.000
22.08
22.5
P.V 22.4
(L.atm) 22.3
22.2
22.1
.2 .4 .6 .8 1.0
P (atm)
P.V (L.atm)
22.4
22.3
22.3
22.25
22.16
22.08
We know molar volume of a gas
should be 22.4 Liters…
… what experiment
was most accurate?
Ex4) Suppose a 12.2L sample of oxygen gas contains 0.50mol
at a pressure of 1 atm and a temp of 25oC.
If this O2 were converted to ozone, O3, at the same temp and pressure,
what would be the volume of ozone?
Ch5 HW#1 p232+ 23,27a,29,30,32,35
Ch5 HW#1 p232+ 23,27a,29,30,32,35
23. Freon-12 (CF2Cl2) is commonly know as the refrigerant in central
home air conditioners. The system is initially charged to a pressure
of 4.8 atm. Express this pressure in each of the following units:
a) mm Hg
b) torr
c) Pa
d) psi (1 atm = 14.7 psi)
27. If the flask is open to the atmosphere,
the mercury levels are equal.
For each of the following situations
where a gas is contained in the flask,
calculate the pressure in the flask
in torr, atmospheres, and pascals.
a)
27. If the flask is open to the atmosphere,
the mercury levels are equal.
For each of the following situations
where a gas is contained in the flask,
calculate the pressure in the flask
in torr, atmospheres, and pascals.
a)
Flask is 118 torr less than atm = 642 torr…
29. An aerosol can contains 400. mL of compressed gas at 5.20 atm
pressure. When all the gas is sprayed into a large plastic bag, the bag
inflates to a volume of 2.14 L. What is the pressure inside the plastic bag?
Assume temperature is constant.
V1 = .400L
P1 = 5.20atm
T1 = const
V2 = 2.14L
P2 = ?
T2 = const
P1 ⋅ V1 P2 ⋅ V2
=
T1
T2
30. A balloon is filled to a volume of 7.00x102 mL at a temperature
of 20.0°C. The balloon is then cooled at a constant pressure to a
temperature of 1.00x102 K. What is the final volume of the balloon?
V1 = .700L
P1 = const
T1 = 293K
V2 = ?
P2 = const
T2 = 100.K
P1 ⋅ V1 P2 ⋅ V2
=
T1
T2
32. Consider the following chemical equation. 2NO2(g) N2O4(g)
If 25.0 mL of NO2 gas is completely converted to N2O4 gas under the
same conditions, what volume will the N2O4 occupy?
2NO2(g) N2O4(g)
mole ratio:
volume:
2
25ml
:
V1 V2
=
n1 n2
1
?
35. The steel reaction vessel of the bomb calorimeter, which has a volume
of 75.0 mL, is charged with oxygen gas to a pressure of 145 atm of 22˚C.
Calculate the mole of oxygen in the reaction vessel.
P.V = n.R.T
P ⋅⋅V
V
n=
=
R ⋅T
Ch5.3 Gas Stoichiometry
Ex1) What volume does 1 mol of gas occupy at STP?
STP stands for:
Standard Temp and Pressure
0oC and 1 atm
- used as common
reference conditions
Ex1) What volume does 1 mol of gas occupy at STP?
STP stands for:
Standard Temp and Pressure
nRT
V=
0oC and 1 atm
P
- used as common
reference conditions
L ⋅ atm
1.000mol 0.08206
273.2 K
mol ⋅ K
= 22.42 L
1atm
22.42 L is the molar volume of an ideal gas at STP
Ex2) A sample of nitrogen gas has a volume of 1.75L.
How many moles of N2 are present?
Ex3) Quicklime, CaO, is produced by the thermal decomposition of
calcium carbonate. Calculate the volume of CO2 produced at STP
from the decomp of 152g CaCO3.
Ex4) A sample of methane gas having a volume of 2.80L
at 25oC and 1.65atm was mixed with a sample of oxygen gas
having a volume of 35.0L at 31oC and 1.25atm.
the mix was ignited. What volume of CO2 formed
at a pressure of 2.50atm and temp of 125oC?
Ch5 HW#2 p233+ 38,43,47,48,49,50,52
Ch5 HW#2 p233+ 38,43,47,48,49,50,52
38. What volume is occupied by 2.0 g of oxygen at 25˚C and a pressure
of 775 mm Hg?
2.0g O 2 1mol O 2
= 0.0625 mol O 2
32.0g O 2
775mm Hg
1atm
= 1.02 atm
760mm Hg
P.V = n.R.T
43. An ideal gas is contained in a cylinder with a volume of 5.0х102 mL.
at a temperature of 30˚C and a pressure is 710. torr. The gas is then
compressed to a volume of 25 mL, and the temperature is raised to
820˚C. What is the new pressure of the gas?
V1 = 0.50L
P1 = 710 torr
T1 = 303K
V2 = 0.025L
P2 = ?
T2 = 1093K
P1 ⋅ V1 P2 ⋅ V2
=
T1
T2
!
47. What mass of helium is required to fill a 1.5-L balloon of STP?
(mass can be found from moles.)
V ⋅P
n=
R ⋅T
48. A student adds 4.00g of dry ice (solid CO2) to an empty balloon.
What will be the volume of the balloon at STP after all the dry ice
sublimes (converts to gaseous CO2)?
4.00g CO 2 1mol CO 2
= 0.0909mol CO 2
44.0g CO 2
49. Calculate the volume of O2, at STP, required for the complete
combustion of 125 g octane (C8H18) to CO2 and H2O.
C8H18 + O2 CO2 + H2O
125g
?L
50. The method used by Joseph Priestley to obtain oxygen made us
of the thermal decomposition of mercuric oxide:
What volume of oxygen gas, measured at 30°C and 725 torr, can be
produced from the complete decomposition of 4.10g mercuric oxide?
2HgO(s) → 2Hg(l) + O2(g)
4.10g
V=? (@STP)
52. Air bags are activated when a severe impact causes a steel ball
to compress a spring and electrically ignite a detonator cap. This
causes a sodium azide (NaN3) to decompose explosively according
to the following reaction:
What mass of NaN3(s) must be
reacted in order to inflate an air bag to 70.0 L at STP?
2NaN3(s) → 2Na(s) + 3N2(g)
?g
70.0 L
Ch5.3 – Molar Mass of a Gas
PV = nRT
PV = nRT
nRT
P=
→
V
(molarmmass )RT → (Vm )RT
dRT
→
V
molar mass
molar mass
dRT
molar mass =
P
Ex1) The density of a gas was measured at 1.50atm and 27oC,
and found to be 1.95g/L. Calc the molar mass of the gas.
54. Urea (H2NCONH2) is used extensively as a nitrogen source in fertilizers.
It is produced commercially from the reaction of ammonia and carbon dioxide:
2NH3(g) + CO2(g) → H2NCONH2(s) + H2O(g)
Ammonia gas at 223˚C and 90 atm flows into a reactor at a rate of 500 L/min.
Carbon dioxide at 223˚C and 45 atm flows into the reactor at a rate
of 600 L/min. What mass of urea is produced per minute by this reaction
assuming 100% yield?
(Solve on next slide)
Ch5 HW#3 p234 39,53,54,55,57,59 + Lab5.1 Pre-Lab HW
2NH3 + CO2 → H2NCONH2 + H2O
T=223˚C
T=223˚C
P=90atm
P=45atm
rate=500L/min rate=600L/min
? g/min
Ch5 HW#3 p234 39,53,54,55,57,59 + Lab5.1 Pre-Lab HW
Ch5 HW#3 p234 53,54,55,57,59 +
Lab5.1 Pre-Lab HW (check for it today, go over tomorrow)
39. A particular balloon is designed by its manufacturer to be inflated to a
volume of no more than 2.5 L. If the balloon is filled with 2.0 L of helium
at sea level, is released, and rises to a altitude at which the atmoshperic
pressure is only 500. mm Hg, will the balloon burst? (temp is constant.)
V1 = 2L
P1 = 760mm Hg
T1 = const
V2 = ?
P2 = 500mm Hg
T2 = const
P1 ⋅ V1 P2 ⋅ V2
=
T1
T2
P1.V1 = P2.V2
53.Consider the reaction between 50.0 mL of liquid methyl alcohol,
CH3OH (density = 0.85 g/mL), and 22.8 L of O2 at 27˚C and a pressure
of 2.00 atm. The products of the reaction are CO2(g) and H2O(g).
Calculate the number of moles of H2O formed if the reaction goes to
completion.
CH3OH + O2 CO2 + H2O
V=50ml
d=0.85g/ml
V=22.8L
P=2atm
T=300K
?mol
(Solve on next slide)
53.
2CH3OH + 3O2 2CO2 + 4H2O
V=50ml
d=0.85g/ml
V=22.8L
P=2atm
T=300K
?mol
54. Urea (H2NCONH2) is used extensively as a nitrogen source in
fertilizers. It is produced commercially from the reaction of ammonia and
carbon dioxide: 2NH3(g) + CO2(g) → H2NCONH2(s) + H2O(g)
Ammonia gas at 223˚C and 90 atm flows into a reactor at a rate
of 500 L/min. Carbon dioxide at 223˚C and 45 atm flows into the reactor
at a rate of 600 L/min. What mass of urea is produced per minute by this
reaction assuming 100% yield?
(Solved in class)
55. Hydrogen cyanide is prepared commercially by the reaction of
methane, CH4(g), ammonia, NH3(g), and oxygen, O2(g), at high
temperature. The other product is gaseous water.
a. Write a chemical equation for the reaction.
CH4 + NH3 + O2 b. What volume of HCN(g) can be obtained from 20 L CH4(g),
20 L NH3(g), and 20 L O2(g)? The volumes are all measured
at the same temperature and pressure.
57. A gas consisting of only carbon and hydrogen has an empirical
formula of CH2. The gas has a density of 1.65 g/L at 27oC and 734 torr.
Determine the molar mass and molecular formula of the gas.
57. A gas consisting of only carbon and hydrogen has an empirical
formula of CH2. The gas has a density of 1.65 g/L at 27oC and 734 torr.
Determine the molar mass and molecular formula of the gas.
empirical formula: CH2
molecular formula: CXH2X
d=1.65g/L
T=300K
P=734torr
dRT
molar mass =
P
59. Silicon tetrachloride (SiCl4) and trichlorosilane (SiHCl3) are both
starting materials for the production of electronics-grade silicon. Calculate
the densities of pure SiCl4 and pure SiHCl3 vapor at 85oC and 758 torr.
T = 85oC 358K
P = 758torr 0.997atm
SiCl4 :
SiHCl3 :
dRT
molar mass =
P
Ch5.4 Dalton’s Law of Partial Pressures
For a mixture of gases in a container, the total pressure exerted
is the sum of the pressures that each gas would exert
if it were alone.
Ptotal = P1 + P2 + P3 + …
Ex1) In a scuba tank, 46L of He and 12L of O2 are placed in a tank
at 25oC and 1atm. the tank has a volume of 5.0L.
Calc the partial pressure of each gas and the total pressure
in the tank at 25oC.
Mole Fraction: the ratio of the number of moles of a given component
in a mixture to the total number of moles in the mixture.
Χ1 =
n1
ntotal
It follows that the mole fraction can also be found by comparing
partial pressures:
P1
Χ1 =
Ptotal
Ex2) The partial pressure of oxygen was observed to be 156 torr in
a sample of air, when the total pressure was743 torr.
Calc the mole fraction.
Ex3) The mole fraction of nitrogen in air is 0.7808.
Calc the partial pressure of N2 in air when
the atmospheric pressure is 760 torr.
Ex4) Solid potassium chlorate was heated in a test tube and decomposed
by the following reaction:
2KClO3(s) 2KCl(s) + 3O2(g)
If 2.20g of KClO3(s) was decomposed, what was the theoretical yield
of O2(g)?
If 0.650L of gas was collected by water displacement,
at 22oC and 754 torr, what was the experimental yield of O2(g)?
The vapor pressure of water at 22oC is 21 torr.
Ch5 HW#4 p234+ 41,61,63,65,67
Ch5 HW#4 p234+ 41,61,63,65,67 Go over Lab5.1 today
41. A compressed gas cylinder, at 13.7 MPa and 23˚C, is in a room
where a fire raises the temperature to 450˚C.
What is the new pressure in the cylinder.
V1 = const
P1 = 13.7 MPa
T1 = 296K
V2 = const
P2 = ?
T2 = 723K
P1 ⋅ V1 P2 ⋅ V2
=
T1
T2
61. A piece of solid carbon dioxide with a mass of 7.8 g is placed in a
4.0 L otherwise empty container at 27o C. What is the pressure in the
container after all the carbon dioxide vaporizes? If 7.8 g solid carbon
dioxide were placed in the container but it already container air at 740
torr, what would be the partial pressure of the carbon dioxide and the
total pressure after the carbon dioxide has vaporized?
CO2:
m = 7.8g
V = 4.0L
T = 27o C = 300K
P=?
63. Consider the flasks diagrammed below, what are the partial pressures
of H2 and N2 after the stopcock between the two flasks is opened?
(Assume the final volume is 3 L) What is the final pressure in torr?
1.00 L N2
0.200 atm
2.00 L H2
475 torr
65. The partial pressure of CH4(g) is 0.175 atm and that of O2(g)
is 0.250 atm in a mixture of the two gasses.
a. What is the mole fraction of each gas in the mixture?
P1
Χ1 =
Ptotal
b. If the mixture occupies a volume of 10.5 L at 65o C, calculate
the total number of moles of gas in the mixture
V ⋅P
n=
R ⋅T
c. Calculate the number of grams of each gas in the mixture.
67. A sample of nitrogen gas was collected over water at 20oC and
a total pressure of 1.00 atm. A total volume of 2.50x102 mL was
collected. What mass of nitrogen was collected?
(At 20o C the vapor pressure of water is 17.5 torr)
T = 293K
Vtotal = 0.250L
mN2 = ?
Ptotal = PN2 + PH20
1atm = PN2 + 0.023atm
PN2 = 0.977atm
Lab 5.1 Pre-Lab Questions
2NaHCO3 Na2CO3 + H2O + CO2
1. Classify the type of reaction represented by the above chemical equation.
2. Predict the physical state of each product in the reaction.
3. How might the H2O affect the results of what we are trying to achieve in this lab?
H2O, CO2
NaHCO3
Ch5.5 – Kinetic Molecular Theory of Gases
1. Particles are so small compared to the distances between them,
that the volume of the individual particle is negligible
Oxygen trivias: (Gas Trivias)
1. Diameter of O2 molecule:
0.339 nm = .000000000339 meters!
(Very small)
2. Average distance traveled between collisions:
106 nm (Mean free path)
Travels over 300x its own diameter
before it hits something else.
(Mostly empty space)
1. Particles are so small compared to the distances between them,
that the volume of the individual particle is negligible
Oxygen trivias: (Gas Trivias)
1. Diameter of O2 molecule:
0.339 nm = .000000000339 meters!
(Very small)
2. Average distance traveled between collisions:
106 nm (Mean free path)
Travels over 300x its own diameter
before it hits something else.
(Mostly empty space)
2. The particles are in constant motion. The collisions of the particles
with the walls of the container are the cause of pressure.
3. Average velocity of 02:
443m/s
(~1000mph!)
(Very fast)
4. Frequency of collisions:
4.5 billon collisions per second
(A lot of collisions)
1. Particles are so small compared to the distances between them,
that the volume of the individual particle is negligible
(Gas Trivias)
1. (Very small)
2. (Mostly empty space)
2. The particles are in constant motion. The collisions of the particles
with the walls of the container are the cause of pressure.
3. (Very fast)
4. (A lot of collisions)
3. The particles are assumed to exert no forces on each other.
Don’t attract, don’t repel.
4. The average kinetic energy of the gas particles is assumed to be
directly proportional to their temperature in Kelvins.
KMT proves the concepts of the Ideal Gas Law
PV = nRT
1. Boyle’s Law:
P.V = nRT
VP
.
V.
P
KMT proves the concepts of the Ideal Gas Law
PV = nRT
1. Boyle’s Law:
P.V = nRT
VP
V.
.
P
2. Guy Lussac:
P = nR
T V
P
T
Fire
P
T
PV = nRT
1. Boyle’s Law:
P.V = nRT
VP
V.
.
P
2. Guy Lussac:
P = nR
T V
Fire
P
T
Fire
V
T
P
T
3. Charles Law:
V nR
=
T P
V
T
PV = nRT
4. Volume and moles:
V RT
=
n P
V
n
V
n
Kinetic Theory also explains Vapor Pressure (Ch10)
Liquids that vaporize easily are volatile
(low intermolecular forces).
Kinetic Theory also explains Vapor Pressure (Ch10)
Liquids that vaporize easily are volatile
(low intermolecular forces).
Vapor pressure increases significantly with temperature.
Kinetic Theory also explains Vapor Pressure (Ch10)
Liquids that vaporize easily are volatile
(low intermolecular forces).
Vapor pressure increases significantly with temperature.
When the rate of evaporation
equals the rate of condensation,
the system reaches equilibrium.
Temperature (in Kelvin) is a measure of the average kinetic energy
of the particles.
Kinetic Energy Formula: KE = ½m.u2
KEave = ½m.(uave)2
KE avg = 32 RT
u rms
3RT
=
M
Joules
K ⋅mol
M is the molar mass in kilograms
R =8.31
Ex1) Calculate the root mean square velocity for the atoms in
a sample of helium gas at 25oC.
Ch5 HW#5 p231+ 17,18,20,22,71,73,75
Ch5 HW#5 p231+ 17,18,20,22,71,73,75
17. Show how Boyle’s law and Charles’s law are special cases of the ideal
gas law.
V1 ⋅ P1
V2 ⋅ P2
=n=
R ⋅ T1
R ⋅ T2
18. Using the postulate of the kinetic molecular theory, give molecular
interpretations of Boyle’s law and Charles’s law.
VP
.
V.
P
V
T
Fire
V
T
20. At room temperature water is a liquid with a molar volume of 18mL.
At 150oC and I atm pressure, water is a gas and has a molar volume
of over 30L. Explain the large difference in molar volume.
293K
423K
22. Consider a sample of gas molecules for the following questions.
a. How is the average kinetic energy of the gas molecules related to
temperature?
b. How is the average velocity of the gas molecules related to
temperature?
c. How is average velocity of the gas molecules related to
the molar mass of the gas at constant temperature?
22. Consider a sample of gas molecules for the following questions.
a. How is the average kinetic energy of the gas molecules related to
temperature?
b. How is the average velocity of the gas molecules related to
temperature?
c. How is average velocity of the gas molecules related to
the molar mass of the gas at constant temperature?
Temperature (in Kelvin) is a measure of the average kinetic energy
of the particles.
KEave = ½m.(uave)2
71. Calculate the kinetic energy of the CH4 molecules in the sample
of CH4 gas at 217 K and at 546 K.
Temperature (in Kelvin) is a measure of the average kinetic energy
of the particles.
Kinetic Energy Formula: KE = ½m.u2
KEave = ½m.(uave)2
73. Calculate the root mean square velocity of the CH4 molecules in
a sample of CH4 gas at 273 K and at 546 K.
u rms
3RT
=
M
Joules
K ⋅mol
M is the molar mass in kilograms
R =8.31
75. Do all the molecules of a 1-mol sample of CH4(g) have the same energy
at 273 K?
KE avg = 32 RT
Ch5.6 – Effusion and Diffusion, and Real Gases
Diffusion – describes the mixing of gases. (ammonia with air)
Ex1) I have an 80cm long tube, as shown. When HCl and NH3 meet…
Where will they meet?
Effusion – describes the passage of a gas thru a tiny orifice.
Graham’s Law of Effusion
Rate of effusion for gas 1
=
Rate of effusion for gas 2
M2
M1
(M is molar mass in g/mol)
Real Gases
If asked on a test the difference…
In a real gas,
1. There may be interparticle interactions.
2. The volume of each particle isnt zero.
3. Approaches Ideal when low pressures and hi temps.
Ch5 HW#6 p237+ 78,79,81,85a
Ch5 HW#6 p237+ 78,79,81,85a
78. Consider a 1.0-L container of neon gas at STP. Will the average kinetic
energy, average velocity, and frequency of collisions of gas molecules
with the walls of the container increase, decrease, or remain the same
under each of the following conditions?
a. the temperature is increased to 100o C.
b. The temperature is decreased to -50o C.
c. The volume is decreased to 0.5L.
d. The number of moles of the neon is doubled.
79. Consider three identical flasks filled with different gases.
Flask A: CO at 760 torr and 0oC m = 26.0 g/mol
Flask B: N2 at 250 torr and 0oC m = 28.0 g/mol
Flask C: H2 at 100 torr and 0oC m = 2.0 g/mol
a. in which flask will the molecules have the greatest average kinetic
energy?
KEave = ½m.(uave)2
b. In which flask will the molecules have the greatest average velocity?
81. The effusion rate of an unknown gas is measured and found to be
31.50mL/min. Under identical experiment conditions,
the effusion rate of O2 is found to be 30.50mL/min.
If the choices are CH4, CO,NO,CO2, and NO2 what is the identity
of the unknown gas?
O2 : 32.0g/mol
Rate of effusion for gas 1
=
Rate of effusion for gas 2
CH4 : 16.0g/mol
CO : 28.0g/mol
NO : 30.0g/mol
CO2 : 44.0g/mol
NO2 : 46.0g/mol
M2
M1
(M is molar mass in g/mol)
85. Calculate the pressure exerted by 0.5000 mol N2 in a 1.0000-L
container at 25o C using the ideal gas law.
n = 0.5000 mol
V = 1.0000-L
T = 25o C = 298K
P=?
PV = nRT
Ch5 Rev p230+ 1,2,3,4,5,6,12,34a,b,58,60,62,68 Handout Midterm
1. Consider the following apparatus: a test tube covered with a Pre-lab HW
nonpermeable elastic membrane inside a container that is closed Ch5 test
with a cork. A syringe goes through the cork.
a) As you push down the syringe, how does the membrane covering
the test tube change?
b) You stop pushing the syringe but continue to hold it down.
In a few seconds what happens to the membrane?
2. Which of the following statements is the best explanation of how this
barometer works?
a) Air pressure outside the tube causes the mercury to
move in the tube until the air pressure inside and
outside the tube is equal.
b) Air pressure inside the tube causes the mercury to
move in the tube until the air pressure inside and
outside the tube is equal.
c) Air pressure outside the tube counterbalances
the weight of the mercury in the tube.
d) Capillary action of the mercury causes the mercury
to go up the tube.
3. The barometer bellow shows the level of mercury at a given atmosphere
pressure. Fill all the other barometers with mercury for the same
atmospheric pressure. Explain your answer.
4. As you increase the temperature of a gas in a sealed, rigid container,
what happens to the density of the gas? Would the results be the same
if you did the same experiment in a container with a piston at constant
pressure?
5. A diagram in a chemistry book shows a magnified view of a flask of air
as shown. What do you suppose is between the dots (the dots represent air
molecules)?
6. If you put a drinking straw in water, place your finger over the opening,
and lift the straw out of the water, some water stays in the straw. Explain.
12. If you have any two gases in different containers that are the same size
at the same pressure and same temperature, what is true about the moles
of each gas? Why is this true?
34. Complete the following table for an ideal gas.
58. A compound has the empirical formula CHCl.
A 256-ml flask, at 373 K and 750. torr, contains 0.800 g
of the gaseous compound. Give the molecular formula.
60. Calculate the density of ammonia gas at 27°C and 635 torr.
62. A mixture of 1.00 g H2 and 1.00 g He is placed in a 1.00-L container
at 27°C. Calculate the partial pressure of each gas and the total pressure.
68. Helium is collected over water at 25°C and 1.00 atm total pressure.
What total volume of gas must be collected to obtain 0.586 g of helium?
(At 25°C the vapor pressure of water is 23.8 torr.)
FRQ. A sample of pure, gaseous hydrocarbon is introduced into a previously
evacuated rigid 1.00L vessel. The pressure of the gas is 0.200 atm
at a temperature of 127oC.
a. Calculate the number of moles of the hydrocarbon in the vessel.
b. O2(g) is introduced into the same vessel containing the hydrocarbon.
After the addition of the O2(g), the total pressure of the gas mixture in the vessel
is 1.40 atm at 127oC. Calculate the partial pressure of O2(g) in the vessel.
(FRQ cont) The mixture of the hydrocarbon and oxygen is sparked so that
a complete reaction occurs, producing CO2(g) and H2O(g). The partial pressures
of these gases at 127oC 0.600 atm for CO2(g) and 0.800 atm for H2O(g).
There is O2(g) remaining in the container after the reaction is complete.
c. Use the partial pressures of CO2(g) and H2O(g) to calculate the partial pressure
of the O2(g) consumed in the combustion.
d. On the basis of your answers above, write the balanced chemical equation
for the combustion reaction AND determine the formula of the hydrocarbon.
Extra Examples?
Ex) 1.53 L sample of gaseous sulfur dioxide is at a pressure of
5.6 x 103 Pa. If the pressure is changed 1.5 x 104 Pa,
while at a constant temp, what will be the new volume?
20
P 15
(x103Pa)10
5
1 2 3 4 5 6
V (L)
Ex) A sample of gas at 15oC and 1 atm has a volume of 2.58L.
What volume will this gas occupy at 38oC and 1 atm?
Ex) 3.5L of ammonia gas at 1.68atm. The gas is compressed to
1.35L at a constant temp. Show how the ideal gas law can be used
to find the final pressure.
Ex) A sample of methane gas has a volume of 3.8L at 5oC.
It is heated to 86oC, at constant pressure. Calc new volume.
Ex) Diborane gas (B2H6) at a pressure of 345 torr and –15oC,
occupies a volume of 3.48L. If temp raised to 36oC
and pressure of 468 torr, calc new volume.
Ex) A sample containing 0.35mol argon gas at a temp of 13oC
and a pressure of 568 torr is heated to 56oC and a pressure of 897 torr.
Calc the change in volume.