Transcript Document

Gases
Chapter 8
5.1
Physical Characteristics of Gases
•
Gases assume the volume and shape of their containers.
•
Gases are the most compressible state of matter.
•
Gases will mix evenly and completely when confined to
the same container.
•
Gases have much lower densities than liquids and solids.
5.1
Force
Pressure = Area
Units of Pressure
1 pascal (Pa) = 1 N/m2
1 atm = 760 mmHg = 760 torr
1 atm = 101,325 Pa
Barometer
5.2
10 miles
4 miles
Sea level
0.2 atm
0.5 atm
1 atm
5.2
5.2
As P (h) increases
V decreases
5.3
Boyle’s Law
P a 1/V
P x V = constant
P1 x V1 = P2 x V2
5.3
A sample of chlorine gas occupies a volume of 946 mL
at a pressure of 726 mmHg. What is the pressure of
the gas (in mmHg) if the volume is reduced at constant
temperature to 154 mL?
P1 x V1 = P2 x V2
P2 =
P1 = 726 mmHg
P2 = ?
V1 = 946 mL
V2 = 154 mL
P1 x V1
V2
726 mmHg x 946 mL
=
= 4460 mmHg
154 mL
5.3
As T increases
V increases
5.3
Variation of gas volume with temperature
at constant pressure.
Charles’ Law
VaT
V = constant x T
V1/T1 = V2/T2
Temperature must be
in Kelvin
T (K) = t (0C) + 273.15
5.3
A sample of carbon monoxide gas occupies 3.20 L at
125 0C. At what temperature will the gas occupy a
volume of 1.54 L if the pressure remains constant?
V1/T1 = V2/T2
T2 =
V1 = 3.20 L
V2 = 1.54 L
T1 = 398.15 K
T2 = ?
V2 x T1
V1
=
1.54 L x 398.15 K
3.20 L
= 192 K
5.3
Avogadro’s Law
V a number of moles (n)
V = constant x n
Constant temperature
Constant pressure
V1/n1 = V2/n2
5.3
Ammonia burns in oxygen to form nitric oxide (NO)
and water vapor. How many volumes of NO are
obtained from one volume of ammonia at the same
temperature and pressure?
4NH3 + 5O2
1 mole NH3
4NO + 6H2O
1 mole NO
At constant T and P
1 volume NH3
1 volume NO
5.3
Ideal Gas Equation
Boyle’s law: V a 1 (at constant n and T)
P
Charles’ law: V a T (at constant n and P)
Avogadro’s law: V a n (at constant P and T)
Va
nT
P
V = constant x
nT
P
=R
nT
P
R is the gas constant
PV = nRT
5.4
The conditions 0 0C and 1 atm are called standard
temperature and pressure (STP).
Experiments show that at STP, 1 mole of an ideal
gas occupies 22.414 L.
PV = nRT
(1 atm)(22.414L)
PV
R=
=
nT
(1 mol)(273.15 K)
R = 0.082057 L • atm / (mol • K)
5.4
What is the volume (in liters) occupied by 49.8 g of HCl
at STP?
T = 0 0C = 273.15 K
P = 1 atm
PV = nRT
nRT
V=
P
1 mol HCl
n = 49.8 g x
= 1.37 mol
36.45 g HCl
1.37 mol x 0.0821
V=
L•atm
mol•K
x 273.15 K
1 atm
V = 30.6 L
5.4
Argon is an inert gas used in light bulbs to retard the
vaporization of the filament. A certain light bulb
containing argon at 1.20 atm and 18 0C is heated to
85 0C at constant volume. What is the final pressure of
argon in the light bulb (in atm)?
PV = nRT
n, V and R are constant
nR
P
=
= constant
T
V
P1
P2
=
T1
T2
P1 = 1.20 atm
T1 = 291 K
P2 = ?
T2 = 358 K
T2
= 1.20 atm x 358 K = 1.48 atm
P2 = P1 x
291 K
T1
5.4
Gas Stoichiometry
What is the volume of CO2 produced at 370 C and 1.00
atm when 5.60 g of glucose are used up in the reaction:
C6H12O6 (s) + 6O2 (g)
6CO2 (g) + 6H2O (l)
g C6H12O6
mol C6H12O6
5.60 g C6H12O6 x
6 mol CO2
1 mol C6H12O6
x
= 0.187 mol CO2
180 g C6H12O6
1 mol C6H12O6
V=
nRT
=
P
mol CO2
L•atm
0.187 mol x 0.0821
mol•K
1.00 atm
V CO2
x 310 K
= 4.76 L
5.5
Dalton’s Law of Partial Pressures
V and T
are
constant
P1
P2
Ptotal = P1 + P2
5.6
Gas Laws

Dalton’s law of partial pressures: the total
pressure, PT, of a mixture of gases is the sum of
the partial pressures of each individual gas
– Problem: to a tank containing N2 at 2.0 atm and O2 at
1.0 atm we add an unknown quantity of CO2 until the
total pressure in the tank is 4.6 atm. What is the partial
pressure of CO2?
– Solution:
Kinetic Molecular Theory of Gases
1. A gas is composed of molecules that are separated from
each other by distances far greater than their own
dimensions. The molecules can be considered to be points;
that is, they possess mass but have negligible volume.
2. Gas molecules are in constant motion in random directions.
Collisions among molecules are perfectly elastic.
3. Gas molecules exert neither attractive nor repulsive forces
on one another.
4. The average kinetic energy of the molecules is proportional
to the temperature of the gas in kelvins. Any two gases at
the same temperature will have the same average kinetic
energy
5.7
Kinetic Molecular Theory

Ideal gas: the six assumptions of the K MT give
us an idealized picture of the particles of a gas and
their interactions with one another
 Real gases
– their atoms or molecules do occupy some volume
– there are forces of attraction between their atoms or
molecules

In reality, no gases are ideal
– at pressures above 1 to 2 atm and temperatures well
above their boiling points, most real gases behave in
much the same way as predicted by the KMT
Intermolecular Forces

The strength of attractive forces between
molecules determines whether any sample of
matter is a gas, liquid, or solid
– the forces of attraction between molecules of most
gases are so small that they can be ignored
– when T decreases or P increases or both, the forces of
attraction become important to the point that they cause
condensation (gases to liquids) and ultimately
solidification (liquids to solids)
– in order to understand the properties of liquids and
solids, we must look at the nature of these
intermolecular forces of attraction
Intermolecular Forces

Three types of intermolecular forces
– their origin is electrostatic; that is, the attraction
between positive and negative dipoles
– the strengths of covalent bonds are also shown
for comparison
Intermolecu lar forces
of attraction
betw een molecules
Covalent bond s
Lond on disp ersion forces 0.01 - 2 k cal/mol
D ipole-dipole interactions 1-6 k cal/mol
Hyd rogen bondin g
2 - 10 k cal/mol
Sin gle, doub le, and
triple covalen t bonds
70 - 200 kcal/mol
London Dispersion Forces

London dispersion forces are the attraction
between temporary induced dipoles
London Dispersion Forces
– LDF exist between all atoms and molecules
– only forces for nonpolar particles
– they range in strength from 0.01 to 2 kcal/mol
depending on mass, size, and shape
– in general, their strength increases as the mass
and number of electrons in a molecule increases
– even though these forces are very weak, they
contribute significantly to the attractive forces
between large molecules because they act over
large surface areas
Dipole-Dipole Interactions

Dipole-dipole interactions; the electrostatic
attraction between positive and negative dipoles
– butane is a nonpolar molecule; the only interactions
between butane molecules are London forces
– acetone is a polar molecule; its molecules are held
together in the liquid state by dipole-dipole interactions
Hydrogen Bonds

Hydrogen bond: a force of attraction between the
partial positive charge on a hydrogen and the
partial negative charge on a nearby O or N
Hydrogen Bonds
– the strength of hydrogen bonds ranges from 2
–
–
–
–
to 10 kcal/mol
that in liquid water is 5.0 kcal/mol
by comparison, the strength of an O-H covalent
bond in a water molecule is 119 kcal/mol
the relatively high boiling point of water is due
to hydrogen bonding between water molecules
hydrogen bonds are not restricted to water; they
form whenever there is are O-H or N-H groups
Liquids
– as pressure increases attractions between
molecules become more important
– in liquids, there is very little space between
molecules - liquids are difficult to compress
– the density of liquids is much greater than that of
gases because the same mass occupies a much
smaller volume
– the position of molecules in a liquid is random
and there is irregular space between them into
which other molecules can slide; this causes
liquids to be fluid
Surface Tension

Surface tension:
– molecules in the interior of a liquid have equal
intermolecular forces in all directions
– molecules at the liquid-gas interface experience
a greater attraction toward the interior of the
liquid than toward the gas phase above it
– therefore, there is preferential pull of molecules
on the surface toward the interior of the liquid
– this preferential pull crowds the molecules on
the surface, and creates a thin elastic skin-like
layer
Surface Tension
Evaporation/Condensation

Liquids evaporate
– in a liquid there is a distribution of kinetic
energies (KE) among its molecules
– if a molecule at the surface is moving slowly
(has a low KE), it cannot escape from the liquid
because of the attractions of neighboring
molecules
– if, however, it is moving more rapidly (has a
high KE) and moving upward, it can escape the
liquid and enter the gaseous space above it
Evaporation/Condensation
– if the container is open, molecules escape
– if the container is closed, molecules remain in
–
–
–
–
the air space above the liquid
Equilibrium and vapor pressure
vapor pressure is a function of temperature
vapor pressure increases with temperature until
it equals the atmospheric pressure
boiling point: the temperature at which the
vapor pressure of a liquid equals the
atmospheric pressure
Evaporation/Condensation
– normal boiling point:
the temperature at
which the vapor
pressure of a liquid
equals atmospheric
pressure
Boiling Point

Table 5.3
– boiling points of covalent compounds depend
primarily on two factors: (1) the nature and
strength of intermolecular forces and (2)
molecular size and shape
Boiling Points

Intermolecular forces
– consider CH4 (MW 16, bp -164°C) and H2O (MW 18,
bp 100°C). The difference in boiling points between
them is due to the greater strength of hydrogen bonding
in water compared with the much weaker London
dispersion forces in methane
– consider methane, CH4 (MW 16, bp -164°C), and
hexane C6H14 (MW 86, bp 69°C). Because of its larger
surface area, London dispersion forces are stronger
between hexane molecules than between methane
molecules
Boiling Points

Molecular shape
– when molecules are similar in every way except shape,
the strength of London forces determines boiling point
CH3
CH3 CH2 CH2 CH2 CH3
CH3 -C-CH3
CH3
Pentane
2,2-D imeth ylpropane
(bp 36.2°C)
(bp 9.5°C)
– both have the same molecular weight
– 2,2-dimethylpropane is roughly spherical while pentane
is a linear molecule
– pentane has the higher boiling point because it has the
larger surface area
Solids
– when liquids are cooled, their molecules come
so close together and attractive forces between
them become so strong that random motion
stops and a solid is formed
– crystallization (solidification): formation of a
solid from a liquid
Types of Solids
Type
Made u p of
Characteris tics
Ion ic
ions in a crystal
lattice
high melting points N aCl, K2S O 4
Molecular
molecules in
a crystal lattice
low melting points
Polymeric
giant molecu les;
can be crystalline,
semicrystalline, or
amorph ou s
low melting points rubb er,
or cannot be melted; plas tics,
soft or hard
proteins
N etw ork
a very large n umb er very hard; very
of atoms connected high melting or
by covalent bonds can not be melted
diamond,
qu artz
Amorp hous
randomly arran ged
atoms or molecu les
soot,
tar,
glass
mos tly soft, can b e
made to flow , but
no meltin g p oin t
Examples
ice, as pirin
Phase Changes

Phase: any part of a system that looks
uniform throughout
– examples: solid water (ice), liquid water, and
gaseous water (steam)

Phase change: a change from one physical
state (gas, liquid, or solid) to another
Phase Changes

The heating
curve of ice
Phase Changes

Calculation of energy required to heat 1.0
gram of solid water from -20°C to 120°C
En ergy
Ph ys ical chan ge
(cal/g)
w arm ice from -20°C to 0°C
9.6
melt ice; temp = 0°C
80
w armwater from 0°C to 100°C
100
boil w ater; temp = 100°C
540
w arm s team from 100°C to 120°C 9.6
*SH = s pecific heat
Basis for Calcu lation of
En ergy Req uired *
S H of ice = 0.48 cal/g• °C
Heat of fusion of ice = 80 cal/g
S H of liquid water = 1.0 cal/g• °C
Heat of vaporization = 540 cal/g
S H of steam = 0.48 cal/g• °C
Phase Changes

All phase
changes for
any substance
can be shown
on a phase
diagram
Chapter 5
Gases, Liquids, and Solids
End
Chapter 5