Transcript Ch. 11: Gases

Ch. 11: Gases
Dr. Namphol Sinkaset
Chem 152: Introduction to
General Chemistry
I. Chapter Outline
I. Introduction
II. Kinetic-Molecular Theory of Gases
III. Pressure
IV. Individual Gas Laws
V. The Combined Gas Law
VI. The Ideal Gas Law
VII. Partial Pressures
VIII. Gases and Stoichiometry
I. The Unique Gas Phase
• Physical properties of a gas are nearly
independent of its chemical identity!
• Gas behavior is markedly different than
solid or liquid behavior.
• We look at a theory that explains why
gas behavior is universal and then at
the origin of equations that allow us to
do gas calculations.
II. Kinetic Molecular Theory
• This simple theory is very successful in
explaining the physical behavior of most
gases under “normal” conditions.
• Kinetic molecular theory can be
summarized into 4 statements.
II. Kinetic Molecular Theory
1) A gas is a collection of particles in
constant, straight line motion.
2) Gas particles do not attract nor repel
one another. They collide with each
other and the walls of the container.
3) There is a lot of space between gas
particles.
4) The average KE is proportional to the
kelvin temperature of the gas.
II. Kinetic Molecular Theory
II. Application of KM Theory
• KM theory predicts properties of gases
well. For example:
• Compressibility. Gases can be
compressed because of the amount of
space between particles.
• Assume shape and volume of container.
Gas particles are in constant motion
and have no interaction with each other.
III. Pressure
• Pressure is simply a force exerted over a
surface area.
III. Origin of Gas Pressure
• Gas pressure is
the result of the
cumulative force
of many collisions
between gas
particles and
container walls.
III. Pressure Imbalance
III. Atmospheric Pressure
• Patm is simply the weight
of the earth’s
atmosphere pulled
down by gravity.
• Barometers are used to
monitor daily changes
in Patm.
• Torricelli barometer was
invented in 1643.
III. Units of Pressure
• For historic reasons, we have units such
as torr and mm Hg. (Why?)
• The derived SI unit for pressure is N/m2,
known as the pascal (Pa).
• Note that 1 atm = 760 mm Hg = 760 torr =
101.325 kPa.
• Pounds per square inch, psi, is an
everyday unit. 1 atm = 14.7 psi.
III. Sample Problem
• Perform the pressure unit conversions
below.
 Convert 575 torr to atm.
 Convert 2.17 atm to mm Hg.
IV. Gas Laws
• A sample of gas can be physically
described by its pressure (P),
temperature (T), volume (V), and
amount of moles (n).
• If you know any 3 of these variables,
you know the 4th (via calculation).
• We look at the history of how the ideal
gas law was formulated.
IV. Pressure and Volume
IV. Boyle’s Law
• At constant temperature and constant
amount of gas, the volume of a gas and
its pressure are inversely proportional.
1
V
P
IV. Boyle’s Law and KM
Theory
IV. Volume and Temperature
IV. Absolute Zero
• The graph shows an extrapolation to
zero volume for the gas.
• Of course, zero volume is impossible,
so the corresponding temperature is
known as absolute zero.
• Absolute zero is the coldest possible
temperature; 0 K = -273.15 °C.
IV. Charles’s Law
• For constant pressure and constant
moles of gas, the volume of a gas and
its kelvin temperature are directly
proportional.
V T
IV. Charles’s Law and KM
Theory
V. The Combined Gas Law
• Boyle’s and
Charles’s Laws can
be combined into a
convenient form.
• The equation holds
only when amount
of gas remains
constant.
V. Sample Problem
• What’s the final pressure of a sample of
N2 with a volume of 952 m3 at 745 torr
and 25 °C if it’s heated to 62 °C with a
final volume of 1150 m3?
V. Sample Problem
• A sample of N2 has a volume of 880 mL
and a pressure of 740 torr. What
pressure will change the volume to 870
mL at the same temperature?
VI. Combined Gas Law to
Ideal Gas Law
• The combined gas law is actual very
close to the ideal gas law.
• The only quantity missing is the moles
of gas.
• We need one more gas law derive the
ideal gas law from the combined gas
law.
VI. Volume and Moles
VI. Avogadro’s Law
• At constant temperature and pressure,
the volume of a gas and the amount of
moles of gas are directly proportional.
V n
VI. Avogadro’s Law and KM
Theory
VI. The Ideal Gas Law
• The ideal gas law is
a combination of the
combined gas law
and Avogadro’s
Law.
R = 0.082058 L atm/K mole
VI. Sample Problem
• What volume, in mL, does a 0.245 g
sample of N2 occupy at 21 °C and 750
torr?
VI. What Is An Ideal Gas?
• The ideal gas law works best when
gases are acting ideally.
• To be an ideal gas, (1) the volume of
the gas particles must be small relative
to space between them and (2) the
forces between the gas particles are not
significant.
• Gases behave nonideally at low
temperature and high pressure.
VI. Ideal Vs. Nonideal
VII. Mixtures of Gases
• According to KM theory, each gas in a
mixture of gases acts independently of
the others.
• Each individual gas pressure is called a
partial pressure.
• Dalton’s Law of Partial Pressures: the
sum of all partial pressures equals the
total pressure.
 Ptotal = P1 + P2 + P3 + … + Pn
VII. Gas Collection Over Water
• When gas is collected over water, the
total pressure is a sum of Pgas and PH2O.
• Dalton’s Law of Partial Pressure is used
to calculate the pressure of the gas by
itself.
 Ptotal = Pgas + PH2O
• Partial pressures of water are tabulated
as vapor pressures.
VII. Collecting H2 Over Water
VII. Water Vapor Pressures
VIII. Gases and Stoichiometry
• Gases can be either products or
reactants in a reaction, so they can be
involved in stoichiometry problems.
• Mole relationships allow gas
calculations via the ideal gas equation.
• Note that at STP (0 °C and 1 atm) 1
mole of gas occupies 22.4 L.
VIII. Sample Problem
• How many mL of HCl(g) forms at 725
mm Hg and 32.3 °C when 0.117 kg of
NaCl reacts with excess H2SO4?
H2SO4(aq) + 2NaCl(s)  Na2SO4(aq) + 2HCl(g)
VIII. Sample Problem
• How many liters of oxygen at STP are
needed to form 100.0 g of water
according to the reaction below?
2H2(g) + O2(g)  2H2O(g)