#### Transcript Slide 1

```Chapter 13
Gases
13.1 Pressure
Objectives:
1) To learn about atmospheric pressure and
the way in which barometers work
2) To learn the various units of pressure.
13.1: Pressure
and soda bottles.
• What happens to a balloon when you
place it in a freezer?
– Why?
13.1 Pressure
• Evangelista Torricelli
invented the
– BAROMETER: device
to measure pressure
– At sea level the height
of the column of
mercury is 760 mm.
Pressure exerted by the atmospheric gases
on the surface of the mercury in the dish
keeps the mercury in the tube.
13.1: Pressure
Weather:
-low pressure STORMS
-high pressure nice weather
13.1: Pressure
Units of Pressure
-mm Hg= torr
Standard atmosphere (atm)
1 std. atm.=1,000 atm=760.0mm Hg=760.0 torr
1 std. atm=101, 325 Pascal
1.000 std. atm=14.69 psi
Measuring pressure
Atmospheric pressure-h
Manometer
Atmospheric pressure + h
13.1: Pressure
The height of mercury in a mercury barometer is
measured to be 732 mm. Represent this
pressure in atm, torr, and pascals.
0.963 atm; 732 torr; 9.76 x 104 Pa
The pressure of a gas is measured to be 2.79 x
105 Pa. Represent this pressure in atm, torr,
and psi.
2.75 atm; 2.10 x 103 torr; 40.4 psi
Homework, Self-Check exercise 13.1
13.2: Pressure and Volume: Boyle’s
Law
• Objectives:
1) To understand the law that relates the
pressure and volume of a gas.
2) To do calculations involving this law.
Robert Boyle: Irish scientist
-studied the relationship
between the pressure of the
trapped gas and its volume
P x V=const.
PV=k
Boyle’s Law
Inversely
proportional
Figure 13.6: Illustration of
Boyle’s law.
P1V1=P2V2
13.2: Pressure and Volume: Boyle’s
Law
•
A sample of neon gas has a pressure of
7.43 atm in a container with a volume of
45.1 L. This sample is transferred to a
container with a volume of 18.4 L. What
is the new pressure of the neon gas?
Assume constant temperature.
18.2 atm
13.2: Pressure and Volume: Boyle’s
Law
•
A steel tank of oxygen gas has a volume
of 2.00L. If all of the oxygen is
transferred to a new tank with a volume
of 5.50 L, the pressure is measured to be
6.75 atm. What was the original
pressure of the oxygen gas?
Assume constant temperature.
18.6 atm
Homework: 7-10 p. 435
13.3: Volume and Temperature:
Charles’s Law
• Objectives:
1) To learn about absolute zero.
2) To learn about the law relating the
volume and temperature of a sample of
gas at constant moles and pressure, and
to do calculations involving that law.
13.3: Volume and Temperature:
Charles’s Law
• Objectives:
1) To learn about absolute zero.
2) To learn about the law relating the
volume and temperature of a sample of
gas at constant moles and pressure, and
to do calculations involving that law.
Jacques Charles
(first solo H balloon
flight)
-showed that the
volume of a gas
(at constant pressure)
increases with the
temperature.
Absolute zero:
point where you
get 0 volume
-273oC
13.3: Volume and Temperature:
Charles’s Law
•
Charles’s Law
V/T= constant
V1 = V2
T1
T2
13.3: Volume and Temperature:
Charles’s Law
•
A 2.45 L sample of nitrogen gas is
collected at 273 K and heated to 325K.
Calculate the volume of the nitrogen gas
at 325 K.
Assume constant pressure.
2.92 L
13.3: Volume and Temperature:
Charles’s Law
•
A sample of methane gas is collected at
285 K and cooled to 245K. At 245 K the
volume of the gas is 75.0 L. Calculate
the volume of the methane gas at 285K.
Assume constant pressure.
87.2 L
13.3: Volume and Temperature:
Charles’s Law
•
Consider a gas with a volume of 5.65 L
at 27 C and 1 atm pressure. At what
temperature will this gas have a volume
of 6.69 L and 1 atm pressure.
82oC (355K)
13.3: Volume and Temperature:
Charles’s Law
•
Consider a gas with a volume of 9.25L at
47oC and 1 atm pressure. At what
temperature does this gas have a
volume of 3.50 L and 1 atm pressure.
-152oC (121K)
13.4: Volume and Moles:
• Objective:
1) To understand the law relating the volume
and the number of moles of a sample of
gas at constant temperature and
pressure, and to do calculations involving
this law.
V1 = V2
n1 n2
13.4: Volume and Moles:
If 2.55 mol of helium gas occupies a volume
of 59.5 L at a particular temperature and
pressure, what volume does 7.83 mol of
helium occupy under the same
conditions?
183 L
13.4: Volume and Moles:
If 4.35 g of neon gas occupies a volume of
15.0 L at a particular temperature and
pressure, what volume does 2.00 g of
neon gas occupy under the same
conditions?
6.90 L
13.5 The Ideal Gas Law
Ideal Gas Law
PV=nRT
R=Universal gas constant (proportionality
constant)
R= 0.08206 L atm/ oK
Based on experimental measurements.
Most gases obey this equation at 1 atm or
lower and 0oC or higher
13.5 The Ideal Gas Law
Ideal Gas Law
When the number of moles and type of gas
are a constant…….
P1V1 = P2V2
T1
T2
13.5 The Ideal Gas Law
A sample of neon gas has a volume of 3.45
L at 25oC and a pressure of 565 torr.
Calculate the number of moles of neon
present in the gas sample.
0.105 mol
13.5 The Ideal Gas Law
A 0.250 mol sample of argon gas has a
volume of 9.00 L at a pressure of 875 mm
Hg. What is the temperature (in oC) of the
gas?
232oC
13.5 The Ideal Gas Law
Consider a sample of helium gas at 23oC
with a volume of 5.60 L at a pressure of
2.45 atm. The pressure is changed to
8.75 atm and the gas is cooled to 15oC.
Calculate the new volume of the gas using
the ideal gas law equation.
1.53 L
13.5 The Ideal Gas Law
Consider a sample of helium gas at 28oC
with a volume of 3.80 L at a pressure of
3.15 atm. The gas expands to a volume of
9.50 L and the gas is heated to 43oC.
Calculate the new pressure of the gas
using the ideal gas law equation.
1.32 atm
13.6: Dalton’s Law of Partial
Pressures
Objectives: To understand the relationship
between the partial and total pressures of a
gas mixture, and to use this relationship in
calculations.
Scuba divers: use helium and oxygen instead of
air. Air contains nitrogen that dissolves in the
blood as a result of high pressure. Nitrogen
bubbles out and the diver gets the bends.
13.6: Dalton’s Law of Partial
Pressures
John Dalton: For a mixture of gases in a container,
the total pressure exerted is the sum of the
partial pressures of the gases present.
Partial pressure: pressure that the gas would exert
if it were alone in the container.
Dalton’s law of partial pressures:
Ptotal=P1 + P2 + P3
Ptotal=ntotal (RT/V)
The pressure doesn’t depend on the
forces amongst the particles.
The volume of the individual gas particles
is not important.
Figure 13.12: The production of
oxygen by thermal decomposition.
13.6: Dalton’s Law of Partial
Pressures
A sample of solid potassium chlorate KClO3, was
heated in a test tube and decomposed
according to the reaction
2KClO3(s) 2KCl(s) + 3O2
The oxygen produced was collected by
displacement of water at 22oC. The resulting
mixture of O2 and H2O vapor had a total
pressure of 754 torr and a volume of 0.650L.
Calculate the partial pressure of O2 in the gas
collected and the number of moles of O2
present. The vapor pressure of water at 22oC
is 21 torr.
13.6: Dalton’s Law of Partial
Pressures
Ptotal=PO2 + PH2O
754=PO2 + 21
PO2= 733 torr
nO2 =PO2V
733/760=0.964 atm
RT
nO2= (0.964 atm)(0.650L)
(0.08206) (295)
13.6: Dalton’s Law of Partial
Pressures
A 5.00 g sample of helium gas is added to a
5.00 g sample of neon in a 2.50 L
container at 27oC. Calculate the partial
pressure of each gas and the total
pressure.
12.3 atm He; 2.44 atm Ne; 14.7 atm total
13.6: Dalton’s Law of Partial
Pressures
A sample of oxygen gas is saturated with
water vapor at 30.0oC. The total
pressure is 753 torr and the vapor
pressure of water at 30.0 C is 31.824
torr. What is the partial pressure of the
oxygen gas in atm?
0.949 atm
Homework: 23-30 and 33-36
13.6: Dalton’s Law of Partial
Pressures
A sample of oxygen gas is saturated with
water vapor at 27oC. The total pressure
is 785 torr and the partial pressure of
oxygen is 758.3 torr. What is the vapor
pressure of water at 27oC?
26.7 torr
13.8: The Kinetic Molecular Theory
of Gases
Objectives: To understand the basic
postulates of the kinetic molecular theory
Postulates of KMT.
13.9: The Implications of the Kinetic
Molecular Theory
Objectives: To understand the term
temperature.
To learn how the kinetic molecular theory
explains the gas laws.
13.9: The Implications of the Kinetic
Molecular Theory
The temperature of a gas: how rapidly, its
individual particles are moving.
High temperatures: move very fast.
Low temperatures: move slower.
As the gas is heated to a higher temperature, the
particles move faster, hitting the walls more often.
Pressure increases with increasing temperature
13.10: Real Gases
Objectives: To describe the properties of real
gases.
As real gases are compressed into smaller and
smaller volumes, the particles of the gas begin to
occupy a significant fraction of the available
volume..
Start to attract to each
other here PV=nRT not true
13.11:Gas Stoichiometry
Objectives:
1) To understand the molar volume of an
ideal gas.
2) To learn the definition of STP
3) To use these concepts and the ideal gas
equation.
13.11:Gas Stoichiometry
For 1 mol of an ideal gas at 0oC (273K) and 1 atm,
the volume will be.
V=nRT/P= (1.00mol)(0.08206)(273) = 22.4L
1 atm.
22.4 L is called the molar volume
Standard temperature and pressure (abbreviated
STP). Contains 1 mol of an ideal gas at STP.
13.11:Gas Stoichiometry
A sample of argon gas has a volume of 3.45
L at STP. What is the mass of the
argon?
6.15 g
13.11:Gas Stoichiometry
A sample of hydrogen gas occupies a
volume of 15.0L at STP. What volume
will this sample occupy at 22oC and 2.50
atm?
6.48 L
13.11:Gas Stoichiometry
When magnesium reacts with hydrochloric
acid, hydrogen gas is produced:
Mg(s) + 2HCl
MgCl2(aq) + H2(g)
Calculate the volume of hydrogen gas
produced at STP by reacting 5.00 g Mg
and an excess of HCl (aq)
4.61 L
13.11:Gas Stoichiometry
When subjected to an electric current, water
decomposes to hydrogen and oxygen
gas: 2H2O(l)
2H2(g) + O2(g)
If 25.0g of water is decomposed, what
volume of oxygen gas is produced at
STP?
15.5 L
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