Transcript Chemistry B11 Chapter 5 Gases, Liquids, and Solids Gases T↑ move faster Kinetic energy ↑

Chemistry B11
Chapter 5
Gases, Liquids, and Solids
Gases
T↑
move faster
Kinetic energy ↑
Gases
Physical state of matter depends on:
Attractive forces
Brings molecules together
Kinetic energy
Keeps molecules apart
Gases
Gas
Liquid
Solid
High kinetic energy
(move fast)
Low attractive forces
Medium kinetic energy
(move slow)
medium attractive forces
Low kinetic energy
(move slower)
High attractive forces
Physical Changes
Melting
Boiling
Change of states
Ideal Gases
Kinetic molecular theory:
1.
2.
3.
4.
5.
6.
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 = 0°C (273 K)
P = 1 atm
Pressure (P)
Pressure (P) =
Force (F)
Area (A)
F: constant
A↓
P↑
Atmosphere (atm)
Millimeters of mercury (mm Hg)
torr
in. Hg
Pascal
A: constant
F↑
P↑
1 atm = 760 mm Hg
= 760 torr
= 101,325 pascals
= 29.92 in. Hg
Pressure (P)
Hg
barometer
manometer
atmospheric pressure
pressure of gas in a container
Gases
Boyle’s law:
m,T: constant
P 1/α V
PV = a constant
P1V1 = P2V2
P2 =
P1V1
V2
V2 =
P1V1
P2
Gases
Boyle’s law:
Gases
Charles’s law:
m,P: constant
V
T
T α V
V1
T1
V2 =
V1T2
T1
=
= a constant
V2
T2
T2 =
T1V2
V1
Gases
Charles’s law:
Gases
Gay-Lussac’s law:
m,V: constant
P
T
P α T
P1
T1
P2 =
P1T2
T1
=
= a constant
P2
T2
T2 =
T1P2
P1
Gases
Pressure (P)
Gay-Lussac’s law:
Gases
combined gas law:
PV
= a constant
T
P1V1
T1
=
P2V2
T2
Avogadro’s law:
T1
P1
P1 = P2
V1
T2
P2
V2
V1 = V2
n1 = n2
T1 = T2
n = number of molecules
Gases
Ideal gas law:
n: number of moles (mol)
R: universal gas constant
V: volume (L)
P: pressure (atm)
T: temperature (K)
PV = nRT
Standard Temperature and Pressure (STP)
T = 0°C (273 K)
P = 1.000 atm
R=
PV
nT
1 mole → V = 22.4 L
(1.000 atm) (22.4 L)
=
(1 mol) (273 K)
= 0.0821
L.atm
mol.K
Gases
Dalton’s law of partial pressures:
PT = P1 + P2 + P3 + …
Intermolecular Forces
London dispersion forces
Intermolecular
Forces
Dipole-dipole interaction
<
Ionic bonds
Covalent bonds
Intramolecular
(Bonding) Forces
Hydrogen bonding
London dispersion forces
Attractive forces between all molecules
Only forces between nonpolar covalent molecules
_
He
He
2+ _
2+ _
_
He
δ-
_
_
δ+
2+
Original Temporary
Dipole
No Polarity
He
δ-
_
_
δ+
2+
Original Temporary
Dipole
He
δ-
_
_
δ+
2+
Induced Temporary
Dipole
He
+
_
2+
_
London dispersion forces
He:
T↓
Kinetic energy ↓
Move slower
T = -240°C (1 atm) → liquid
Attractive forces
become more important
liquid
Dipole-Dipole Interactions
Attractive forces between two polar molecules
stronger than London dispersion forces
boiling point ↑
Hydrogen bonding
Between H bonded to O, N, or F (high electronegativity) → δ+
and a nearby O, N, or F → δhydrogen
bond
H
-
+
O
H
H

O
H
H 2O
(a)
Stronger than dipole-dipole interactions & London dispersion forces
Hydrogen bonding
δ+
CH3COOH
Acetic acid
δ-
H-bonding in our body
H-bond
H-bond
Protein (α-helix)
DNA
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.
Evaporation
normal boiling point: the temperature at which a liquid boils under
a pressure of 1 atm.
1 atm. = 760 mm Hg
CH3OH
H 2O
Boiling point
Factors that affect boiling point:
1. Intermolecular forces:
London dispersion forces < Dipole-Dipole interactions < Hydrogen bonds
2. Number of sites for intermolecular interaction (surface area):
Larger surface areas (more electrons)  more sites for L.D.F  b.p.
CH3-CH2-CH2-CH2-CH3
>
CH3-CH2-CH3
3. Molecular shape: With the same molecular weight.
linear CH3-CH2-CH2-CH2-CH3 > spherical
CH3
CH3-C- CH3
CH3
Solids
Network solids (network crystals)
Amorphous solids
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.
Dry ice (solid CO2)
Heating Curve
Heat added (cal)
during the phase changes, the temperature stays constant.
Phase diagram