Gases, Liquids, and Solids

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Chemistry B11
Chapter 5 Gases, Liquids, and Solids
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Gases
move faster Kinetic energy ?
T ?
3
Gases
Physical state of matter depends on
Attractive forces
Kinetic energy
Keeps molecules apart
Brings molecules together
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Gases
Gas
High kinetic energy (move fast)
Low attractive forces
Liquid
Medium kinetic energy (move slow)
medium attractive forces
Solid
Low kinetic energy (move slower)
High attractive forces
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Physical Changes
Melting
Boiling
Change of states
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Ideal Gases
Kinetic molecular theory

• Particles move in straight lines, randomly.
• Kinetic energy of particles depends on
  temperature.
• Particles collide and change direction (they may
  exchange
• kinetic energies).
• Gas particles have no volume.
• No attractive forces between gas particles.
• More collision greater pressure.

In reality, There is no ideal gas (all gases are
real).
At STP (Standard Temperature and Pressure) we
can consider them as ideal.
T 0C (273 K) P 1 atm
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Pressure (P)
F constant
A constant
P ?
P ?
A ?
F ?
Atmosphere (atm) Millimeters of mercury (mm
Hg) torr in. Hg Pascal
1 atm 760 mm Hg 760 torr
101,325 pascals 29.92 in. Hg
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Pressure (P)
barometer atmospheric pressure
manometer pressure of gas in a container
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Gases
m,T constant
Boyles law
P 1/a V
PV a constant
P1V1 P2V2
P1V1
P1V1
P2
V2
V2
P2
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Gases
Boyles law
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Gases
m,P constant
Charless law
V
T a V
a constant
T
V1
V2

T2
T1
V1T2
T1V2
V2
T2
T1
V1
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Gases
Charless law
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Gases
m,V constant
Gay-Lussacs law
P
P a T
a constant
T
P1
P2

T2
T1
P1T2
T1P2
P2
T2
T1
P1
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Gases
Gay-Lussacs law
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Gases
combined gas law
P2V2
P1V1
PV

a constant
T
T2
T1
Avogadros law
T1
P1
P1 P2
V1
n1 n2 n number of molecules
V1 V2
T2
P2
T1 T2
V2
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Gases
Ideal gas law
n number of moles (mol) R universal gas
constant V volume (L) P pressure (atm) T
temperature (K)
PV nRT
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Gases
Daltons law of partial pressures
PT P1 P2 P3
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Intermolecular Forces
London dispersion forces
Intermolecular Forces
lt
Dipole-dipole interaction
Ionic bonds Covalent bonds
Intramolecular (Bonding) Forces
Hydrogen bonding
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London dispersion forces
Attractive forces between all molecules
Only forces between nonpolar covalent molecules
He
He
He
He
d-
d
_
_
_

_
_
_
_
2
2
_
2
2
Original Temporary Dipole
No Polarity
He
He
d-
d
d-
d
_
_
2
_
2
_
Original Temporary Dipole
Induced Temporary Dipole
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London dispersion forces
He T -240C (1 atm) ? liquid
Kinetic energy ? Move slower
Attractive forces become more important
T ?
liquid
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Dipole-Dipole Interactions
Attractive forces between two polar molecules
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Hydrogen bonding
Between H bonded to O, N, or F (high
electronegativity) ? d and a nearby O, N, or F ?
d-
H2O
Stronger than dipole-dipole interactions London
dispersion forces
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Hydrogen bonding
d
CH3COOH Acetic acid
d-
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H-bonding in our body
Protein (a-helix)
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Evaporation
equilibrium
Vapor pressure the pressure of a gas in
equilibrium with its liquid form in a closed
container.
Boiling point the temperature at which the vapor
pressure of a liquid is equal to the atmospheric
pressure.
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Evaporation
normal boiling point the temperature at which a
liquid boils under a pressure of 1 atm.
1 atm.
760 mm Hg
Diethyl ether
CH3OH
H2O
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Boiling point
Factors that affect boiling point
1. Intermolecular forces London dispersion
forces lt Dipole-Dipole interactions lt Hydrogen
bonds
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Solids
Network solids (network crystals)
Amorphous solids
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Solids
Solidification (Crystallization) change phase
from liquid to solid.
Fusion (melting) change phase from solid to
liquid.
Sublimation change phase from solid directly
into the vapor.
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Heating Curve
during the phase changes, the temperature stays
constant.
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Phase diagram