Chapter 12 States of Matter
Gases
Kinetic-molecular theory:
• Describes behavior of gases in terms of particles in
motion
• Kinetic means “to move”
• Objects in motion have kinetic energy
• Proposed by Boltzmann & Maxwell
• Makes many assumptions
Gases
Assumptions of Kinetic-molecular theory:
(1) Particle Size:
• Gases are small particles separated from each other
by empty space
• Volume of particle is small compared with volume
of empty space
• Particles are far apart, have no attractive or repulsive
forces between gas particles.
Gases
Assumptions of Kinetic-molecular theory:
(2) Particle Motion:
• Gas particles are in constant random motion
• Move in straight line until collision w/ walls of
container or another gas particle
• Collisions are elastic (no lost kinetic energy)
Gases
Assumptions of Kinetic-molecular theory:
(3) Particle Energy:
• Kinetic Energy (KE) of particle represented by :
1
• KE = mν2
2
• m = mass
• ν = velocity (speed & direction)
• All particles have same mass but different velocities;
do not have same kinetic energy
• Temperature: measure of average kinetic energy of
particles
Diffusion of gases
Diffusion: describes movement of one material through
another
• Smell food cooking in kitchen while watching TV in
living room
• Rate of diffusion depends on mass of particles involved
• lighter particles diffuse more rapidly
Graham’s Law of Effusion
Effusion: similar to diffusion; gas escapes through a tiny
hole.
Graham’s Law of Effusion: relationship between effusion
rates and molar mass of gas
Rate of Effusion =
1
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠
Graham’s Law of Effusion
Use it to relate rates of diffusion between different gases
since heavier particles diffuse more slowly.
𝑅𝑎𝑡𝑒 𝑨
𝑅𝑎𝑡𝑒 𝑩
=
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑩
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑨
Practice Problem
Find the ratio of diffusion rates between ammonia (NH3)
and hydrogen chloride (HCl).
molar mass NH3: 17.0 g/mol
molar mass HCl: 36.5 g/mol
Practice Problem
Find the ratio of diffusion rates between ammonia (NH3)
and hydrogen chloride (HCl). (molar mass NH3: 17.0 g/mol
molar mass HCl: 36.5 g/mol)
𝑅𝑎𝑡𝑒 𝑵𝑯𝟑
𝑅𝑎𝑡𝑒 𝑯𝑪𝒍
=
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑯𝑪𝒍
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑵𝑯𝟑
Practice Problem
Find the ratio of diffusion rates between ammonia (NH3)
and hydrogen chloride (HCl). (molar mass NH3: 17.0 g/mol
molar mass HCl: 36.5 g/mol)
𝑅𝑎𝑡𝑒 𝑵𝑯𝟑
𝑅𝑎𝑡𝑒 𝑯𝑪𝒍
* Plug & Chug
𝑔
=
36.5𝑚𝑜𝑙𝑯𝑪𝒍
𝑔
17.0𝑚𝑜𝑙𝑵𝑯𝟑
Practice Problem
Find the ratio of diffusion rates between ammonia (NH3)
and hydrogen chloride (HCl).
𝑅𝑎𝑡𝑒 𝑵𝑯𝟑
𝑅𝑎𝑡𝑒 𝑯𝑪𝒍
𝑔
=
36.5𝑚𝑜𝑙𝑯𝑪𝒍
𝑔
17.0𝑚𝑜𝑙𝑵𝑯𝟑
= 2.15
= 1.47
* NH3 diffuses ~1.5 times as fast as HCl
Practice Problem
Calculate the ratio of effusion rates for nitrogen (N2) and
neon (Ne).
Practice Problem
Calculate the ratio of effusion rates for nitrogen (N2) and
neon (Ne).
𝑅𝑎𝑡𝑒 𝑵𝟐
𝑅𝑎𝑡𝑒 𝑵𝒆
=
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑵𝒆
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑵𝟐
Practice Problem
Calculate the ratio of effusion rates for nitrogen (N2) and
neon (Ne).
𝑅𝑎𝑡𝑒 𝑵𝟐
𝑅𝑎𝑡𝑒 𝑵𝒆
=
20.18 𝑔/𝑚𝑜𝑙𝑵𝒆
28.02 𝑔/𝑚𝑜𝑙 𝑵𝟐
Practice Problem
Calculate the ratio of effusion rates for nitrogen (N2) and
neon (Ne).
𝑅𝑎𝑡𝑒 𝑵𝟐
𝑅𝑎𝑡𝑒 𝑵𝒆
=
20.18 𝑔/𝑚𝑜𝑙𝑵𝒆
28.02 𝑔/𝑚𝑜𝑙 𝑵𝟐
= 0.849
Gas Pressure
Pressure: force per unit area
• Gas particles exert force when they collide with walls of
containers.
• Individual gas particles has little mass; exerts little
pressure
• But, 1022 gas particles in a liter container; pressure can
be substantial
• How do we measure air pressure?
Gas Pressure
How do we measure air pressure?
• A Barometer!
• Increase in air pressure: Hg rises
• Decrease in air pressure: Hg falls
Units of Pressure
(1) Pascal (Pa): SI unit of pressure
• Derived unit from kilogram, meter, second
• One Pascal is equal to force of one newton per
square meter: 1 Pa = 1 N/m2
(2) Pounds per square inch (psi)
(3) Millimeter of mercury (mm Hg)
(4) Torr (1 torr = 1 mm Hg)
(4) Atmosphere (atm): represents air pressure
1 atm = 760 mm Hg = 760 torr
Units of Pressure
Comparison of Pressure Units
Unit
kilopascal (kPa)
Millimeters of Hg
(mm Hg)
Compared w/ 1 atm
Compared w/ 1 kPa
1 atm = 101.3 kPa
1 atm = 760 mm Hg
1 kPa = 7.501 mm Hg
torr
1 atm = 760 torr
1 kPa = 7.501 psi
Pounds per square
inch (psi)
1 atm = 14.7 psi
1 kPa = 0.145 psi
Atmosphere (atm)
1 kPa = 0.009 atm
Dalton’s Law of Partial Pressure
• Dalton studied properties of gases & found that each
gas in a mixture exerts pressure independently of other
gases present.
• Dalton’s Law of Partial Pressure: total pressure of
mixture of gases is equal to sum of the pressures of all
the gases in mixture.
Ptotal = P1 + P2 + P3 + ….Pn
Dalton’s Law of Partial Pressure
Example: A mixture of oxygen, carbon dioxide, and
nitrogen has a total pressure of 0.97 atm. What is the partial
pressure of O2, if the partial pressure of CO2 is 0.70 atm
and the partial pressure of N2 is 0.12 atm?
Dalton’s Law of Partial Pressure
Example: A mixture of oxygen, carbon dioxide, and
nitrogen has a total pressure of 0.97 atm. What is the partial
pressure of O2, if the partial pressure of CO2 is 0.70 atm
and the partial pressure of N2 is 0.12 atm?
Ptotal = P1 + P2 + P3 + ….Pn
Dalton’s Law of Partial Pressure
Example: A mixture of oxygen, carbon dioxide, and
nitrogen has a total pressure of 0.97 atm. What is the partial
pressure of O2, if the partial pressure of CO2 is 0.70 atm
and the partial pressure of N2 is 0.12 atm?
Ptotal = PO2 + PN2 + PCO2
Dalton’s Law of Partial Pressure
Example: A mixture of oxygen, carbon dioxide, and
nitrogen has a total pressure of 0.97 atm. What is the partial
pressure of O2, if the partial pressure of CO2 is 0.70 atm
and the partial pressure of N2 is 0.12 atm?
Ptotal = PO2 + PN2 + PCO2
0.97 atm = PO2 + 0.12 atm + 0.70 atm
Dalton’s Law of Partial Pressure
Example: A mixture of oxygen, carbon dioxide, and
nitrogen has a total pressure of 0.97 atm. What is the partial
pressure of O2, if the partial pressure of CO2 is 0.70 atm
and the partial pressure of N2 is 0.12 atm?
Ptotal = PO2 + PN2 + PCO2
0.97 atm = PO2 + 0.12 atm + 0.70 atm
PO2 = 0.97 atm – 0.12 atm – 0.70 atm
PO2 = 0.15 atm
Practice Problem
What is the partial pressure of hydrogen gas in a mixture of
hydrogen and helium if the total pressure is 600 mm Hg
and the partial pressure of helium is 439 mm Hg?
13.2 Forces of Attraction
If all particles of matter have the same average kinetic
energy (at room temp), why are some materials gases while
others are liquids or solids?
Forces of Attraction
The answer is the attractive forces between particles.
Two types of attractive forces:
1) Intramolecular forces:
- Attractive forces that hold particles together in ionic,
covalent, and metallic bonds
- “intra” means “within”
2) Intermolecular forces:
- Attractive forces that hold together water molecules
in drop of water, two different particles
- “inter” means “between”
Intramolecular Forces
Comparison of Intramolecular Forces
Force
Ionic
Covalent
Metallic
Model
+ - + - + - +
+
+
+
+
+
+
+
+
Basis of
Attraction
Example
cations & anions
NaCl
positive nuclei &
shared electrons
H2
metal cations &
mobile electrons
Fe
Intermolecular Forces
Three types of intermolecular forces:
1) Dispersion Forces
2) Dipole-dipole Forces
3) Hydrogen Bonds
Dispersion Forces
• Dispersion Forces: weak forces that result from
temporary shifts in the density of electrons in electron
cloud.
• Sometimes called London forces after GermanAmerican physicist who first described them, Fritz
London.
Dispersion Forces
• Oxygen molecules (O2) are nonpolar because electrons
are equally distributed between equally electronegative
oxygen atoms.
• But, electrons in the electron cloud are in constant
motion.
• When two nonpolar molecules collide the electron cloud
of one molecule repels the electron cloud of the other
molecule.
• The electron density around one nucleus is greater in
one region of each cloud, and forms temporary dipole.
Dispersion Forces
• Temporary dipole forms in each molecule.
δ-
δ+
temporary dipole
δ-
δ+
temporary dipole
• Temporary dipoles close together cause weak dispersion
force between oppositely charged regions of dipole.
Dispersion Forces
• Dispersion forces exist between all particles, but only
play a significant role when no stronger forces of
attraction acting on particles.
• Dominant force of attraction between identical
nonpolar molecules!!
• Weakest forces of attraction.
• Forces have noticeable effect with increasing number of
electrons involved (bigger dipole and hence larger
forces)…explains why fluorine and chlorine are gases,
bromine is a liquid, and iodine is solid at room temp.
Dipole-dipole Forces
• Dipole-dipole forces: attractions between oppositely
charged regions of polar molecules.
• Polar molecules contain permanent dipoles; some
regions of a polar molecule are always partially negative
and some regions are always negative.
• Typically stronger than dispersion forces as long as the
molecules being compared have similar masses.
Dipole-dipole Forces
• Neighboring polar molecules orient themselves so that
oppositely charged regions line up.
• In HCl, the partially positive hydrogen atom is attracted
to the partially negative chlorine atom.
δ-
δ+
δδ+
δ+ δ -
δ+
δ-
δδ+
δ+
δ-
Hydrogen Bonds
• Hydrogen bond: is a special type of dipole-dipole
attraction that occurs between molecules containing a
hydrogen atom bonded to a small, highly electronegative atom
with at least one lone pair.
• Examples: hydrogen is bonded to fluorine (HF), oxygen
(H2O), or nitrogen (NH3).
• Hydrogen bonds explain why water is liquid at room
temperature, while other molecules of comparable
masses are gases. (H2O vs. CH4)
Hydrogen bonds
• Water molecules
δ+
δ-
δ - δ+
δ+
δ+
δ+
δδ+
δ+
δ-
δ+ δ -
δ+
δ+
δ-
δ+
δ+
δ+ δ -
δ+
δ+
δδ+
Intermolecular Forces
Comparison of Intermolecular Forces
Force
Type of
molecule
Dispersion nonpolar
Strength of
Force
Example
weak
H2
Dipoledipole
polar, permanent
dipole
Stronger than
dispersion
HCl
Hydrogen
bonds
polar, contains
H-F, H-O, or H-N
typically
strongest
NH3
Practice Problem
Determine the type of intermolecular forces in each
molecule.
1) F2
2) H2O
3) HI
4) HF
13.3 Liquids: Fluidity
• Fluidity: ability to flow
• Both liquids and gases flow
• Liquids can diffuse through
another liquid (food
coloring in water)
• Liquids diffuse more slowly
than gases at the same
temperature
Liquids: Fluidity
Which will diffuse more
quickly
1) Food coloring in cold
water?
2) Food coloring in hot
water?
Liquids: Viscosity
• Viscosity: measure of the amount of resistance of a
liquid to flow.
• What is more viscous… maple syrup OR milk?
Liquids: Viscosity
• Particles in liquid have attractive forces that slow
movement as they flow past one another.
• Viscosity of liquid is determined by size and shape of
particle
• Bigger particles have higher viscosity
• Longer particles have higher viscosity compared to
shorter molecules
Liquids: Viscosity
• Viscosity is affected by temperature.
• What happens when you put oil into a hot pan?
• Increase in temp = increase in kinetic energy
• Higher temperature = Decrease in Viscosity
Liquids: Surface Tension
• Intermolecular forces: attractions between particles
• Intermolecular forces not equal for all particles in a liquid.
• Particles in middle of liquid attracted to particles above
and to side of them.
• Particles on surface have no attractions from above to
balance attractions below
Liquids: Surface Tension
• Surface tension: measurement of inward pull by the
particles in the interior.
• Allows spiders to “walk on water”
Liquids: Capillary Action
• Water in graduated cylinder has meniscus due to attractive
forces between water molecules and glass molecules
(silicon dioxide)
• Water molecules more attracted to silicon than other
water molecules
• If narrow tubes, water will be
drawn upward
• Called “Capillary Action”
Solids: Density
• Density of solids
13.4 Phase Changes
• Phase changes of water: solid → liquid → gas
• These phase changes require energy
• Energy flows from higher temperature to lower
temperature
Phase Changes
gas
solid
melting
freezing
liquid
Phase Changes
• Phase changes that require energy:
• Melting
• Vaporization
• Evaporation
• Sublimation
Phase Changes: Melting
• Melting: solid particles absorb enough heat to change to
the liquid phase
Phase Changes: Vaporization
• Molecules have more kinetic
energy w/ increasing temp
• Vaporization: process where a
liquid changes to a gas or vapor
• Evaporation: vaporization
occurs only at surface of liquid
• Your body controls its temp by
evaporation
Vapor Pressure
• Evaporation in an open container (a) or evaporation in a
closed container (b).
• Vapor pressure: pressure exerted by vapor over liquid
Boiling Point
• Boiling Point: temperature at which vapor pressure of a
liquid equals the external or atmospheric pressure
• At boiling point, molecules have enough energy to
vaporize
Sublimation
• Sublimation: process which solid changes directly to a
gas without first becoming a liquid
• “Dry Ice” is solid carbon dioxide
Phase Changes
• Phase changes that release energy:
• Condensation
• Deposition
• Freezing
Phase Changes: Condensation
• Water vapor molecules lose
energy, velocity is reduced,
interact with other water
molecules
• Condensation: change
from the vapor to the liquid
phase
• Condensation is the reverse
of vaporization
Phase Changes: Deposition
• Deposition: process by
which substance changes
from a gas or vapor to solid
without becoming liquid.
• Reverse of sublimation
• Formation of snowflakes
(energy released as crystals
form)
Phase Changes: Freezing
• Freezing Point:
temperature at which liquid
is converted to crystalline
solid
• Heat is removed from
water, lose kinetic energy,
molecules become fixed or
“frozen” into set positions
Phase Diagram
• Phase Diagram: graph of pressure versus temperature;
shows what phase a substance exists under different
temperatures and pressures
• Triple point: point on
graph that represents
temp and pressure
where all 3 phases exist