Chapter 11.
States of Matter
States of Matter
State is Determined by:
Chemical Identity
Temperature
Pressure
States of Elements
Kinetic Molecular Theory
of Matter
The Kinetic Molecular Theory of Matter is
an explanation of the behavior of matter,
based on the idea that the particles that
make up the matter are always in motion.
The particles can be atoms, molecules,
or ions.
Kinetic Molecular Theory
of Matter
1. Matter is composed of tiny particles; the
size of the particles is fixed for each substance.
2. The particles are in constant random motion and therefore possess kinetic energy
(energy of motion, which can be transferred by collisions).
Kinetic Molecular Theory
of Matter
3. The particles interact with each other by
means of electrostatic attractions and
repulsions, and therefore possess
potential energy (stored energy, possessed by matter as a result of its position, condition, and or composition).
Kinetic Molecular Theory
of Matter
4. The kinetic energy (velocity) of the
particles increases as temperature is
increased.
5. The particles in a system transfer energy
by means of elastic collisions (collisions
in which all energy transfer results in
motion, not deformation).
Kinetic Molecular Theory
of Matter
No kinetic energy is lost in elastic collisions.
Kinetic energy is transformed to work and/or
heat in inelastic collisions.
Kinetic energy is a disruptive force between
particles; it makes them more independent
of each other.
Potential energy is a cohesive or attractive
force between particles.
Comparison of States
Some Definitions:
Density is the ratio of mass to volume.
Compressibility is a measure of volume
change resulting from a pressure change.
Thermal Expansion is a measure of volume
change resulting from temperature change.
Properties of States
Solids
Definite Volume, Definite Shape
Density is High: 1.0 – 20 g/cm3
Compressibility is Low
Thermal Expansion is Low: 0.01% per C
The Solid State
In solids, cohesive forces predominate over
kinetic energy. Particles are usually held
together in a regular array, and vibrate
about fixed positions.
Electrostatic attractions between particles keep
them close together in fixed positions.
The particles fill 50-70% of the space
available, the rest is “void volume.”
Properties of States
Liquids
Definite Volume, Indefinite Shape
Density is Fairly High: 0.5 – 15 g/mL
Compressibility is Low
Thermal Expansion is Fairly Low: 0.1% per C
The Liquid State
In liquids, cohesive forces are balanced by
kinetic energy. Particles move freely about
each other but do not separate.
Electrostatic attractions between particles keep
them close together, but able to move relative to one another.
The particles fill about 50% of the available
space.
Properties of States
Gases
Indefinite Volume, Indefinite Shape
Density is Low: 0.2 – 10 g/L
Compressibility is High
Thermal Expansion is Moderate: 0.3% per C
The Gaseous State
In gases, cohesive forces are overcome by
kinetic energy. Particles move independently of each other.
Electrostatic attractions between particles are
very weak, do not hold them together.
The particles fill about 1% of the space
available.
Changes in State
Most substances can exist in any phase,
solid, liquid, or gas.
The phase at which the substance exists
depends on its temperature and the
applied pressure.
A phase diagram is a graph showing the
phase behavior of a given substance.
Phase Diagram for Water
Phase Diagram for CO2
Changes of State
Exothermic:
Release Heat Energy
H is negative
Endothermic
Absorb Heat Energy
H is positive
Changes of State
Freezing point
the temperature at which the
liquid and solid phases of a
substance are in equilibrium
Boiling point
the temperature at which the
liquid and vapor phases of a
substance are in equilibrium.
Equilibrium
is a state in which two opposing processes occur at equal
rates.
Energy and Heat
Energy is the capacity to do work.
Energy can exist in different forms, and
can change in form:
Heat
Light
Electrical
Mechanical
Heat Energy
First Law of Thermodynamics
Energy is neither created nor
destroyed, just changed in
form and/or transferred.
Energy Units
The joule (J) is the base unit for energy
or work (force x distance).
1 J = 1 kgm2/sec2
4.18 J = 1 calorie
1 calorie is the amount of heat energy
required to raise the temperature of
1 g of water by 1C.
Specific Heat
Specific heat is the amount of heat energy
required to raise the temperature of 1.00
gram of a substance by 1.00C.
It takes 4.18 J to raise the temperature of
1.00 g of water by 1.00C.
4.18 J/gC is the specific heat of water.
Specific Heat
4.18 J/gC very high!
Metals have specific heats of
0.1 to 1.0 J/gC;
They conduct heat.
Most nonmetallic materials are
insulators, with specific heats
of 1 to 2 J/gC.
Specific Heat
How much heat is absorbed if I heat 100 g
of water from 25.0C (room temperature)
to 100.0C (boiling point of water)?
How much heat is given off if I cool 1.00 lb
(454 g) of iron metal from 100.0C to
25.0C? Specific heat of iron is
0.444 J/gC.
Specific Heat
What is the specific heat of a rock?
Its mass is 125 g. I heat it to
100.0C in a boiling water bath,
then drop it into 100.0 g of water
that's at 20.0C. The water
temperature rises to 30.0C.
Energy and Changes of State
A heating or cooling curve is a graph
showing the amount of energy required
to change the temperature or phase of
a given amount of a substance.
Heating and Cooling Curves
Energy and Changes of State
In an endothermic phase change:
heat energy is absorbed by a substance
its particles gain kinetic energy
forces between particles are overcome
The substance goes into a less ordered state.
It melts, boils, or sublimes.
Energy and Changes of State
In an exothermic phase change:
heat energy is given off by a substance
its particles lose kinetic energy
forces between particles can act
The substance goes into a more ordered state.
It freezes, condenses, or deposits.
Energy and Changes of State
During an endothermic phase change, all
the energy being supplied to the substance
is used to disrupt forces between particles.
No temperature change is observed.
The temperature of a substance that is
present in two phases will remain constant
until all of one phase is consumed.
Energy and Changes of State
During an exothermic phase change, all
the energy being released by the substance
is allowing intermolecular forces to bring
particles to a more ordered state.
No temperature change is observed.
The temperature of a substance that is
present in two phases will remain constant
until all of one phase is consumed.
Heats of Fusion and
Vaporization
The amount of heat required to cause a phase
change in a given material is a physical property of that material.
Heat of fusion, Hf = heat to convert one gram
of substance from solid to liquid at its melting
point. Units: J/g
Heat of vaporization, Hv = heat to convert
one gram of substance from liquid to gas at
its boiling point. Units: J/g
Heats of Fusion and
Vaporization
How much energy will be consumed
when 150 g of water at 100C is
boiled to steam, also at 100C?
Hv water = 2260 J/g
How much energy will be released if 250
g of water freezes to ice, all at 0C?
Hf water = 334 J/g
Heating and Cooling Curve
Calculations
How much energy will be consumed when
200 g of ice at -5.0C is converted to
steam at 120.0C?
Specific Heats, J/gC
Ice: 2.09 Water: 4.18
Steam: 2.03
Evaporation and Boiling
What is really going on when a substance
goes from the liquid to the gas phase or
vice-versa?
In evaporation, some particles (atoms or
molecules) have enough kinetic energy to
overcome cohesive forces and escape
from the surface of the liquid.
Evaporation and Boiling
What is really going on when a substance
goes from the liquid to the gas phase or
vice-versa?
In boiling, many particles have enough
kinetic energy to enter the gas phase.
Bubbles of gas form in the bulk liquid.
Evaporation and Boiling
Evaporation and Boiling
Evaporation and Boiling
If a liquid is put in a closed container, molecules of the liquid will escape into the gas
(or vapor) phase.
The amount of vapor will depend on:
Temperature
Higher T increases amount of gas
Cohesive forces between molecules
Stronger forces decrease amount of gas
Evaporation and Boiling
Device for measuring
the vapor pressure
of a liquid.
Evaporation and Boiling
Vapor Pressure of Water as a function of Temperature
1
Vapor Pressure,
Atmospheres
0.9
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0
0
10
20
30
40
50
60
70
Temperature, degrees C
80
90
100
Evaporation and Boiling
A liquid boils when its vapor pressure equals
the external pressure on the liquid.
The normal boiling point of a liquid is the
temperature at which it boils under
atmospheric pressure.
The boiling point of a liquid will increase or
decrease with changes in applied pressure.
Evaporation and Boiling
Evaporation is conversion from liquid to vapor
at temperatures below the boiling point.
Vapor is usually used for the gas phase of
compounds that are liquid at room temperature and pressure.
Gas is the term used for compounds that are
not liquid at room temperature and pressure.
Evaporation and Boiling
Evaporation and Boiling
Intermolecular Forces
Intermolecular forces are attractive forces
that act between molecules (also atoms
and ions). There are three types:
London Dispersion Forces
Dipole – Dipole Interactions
Hydrogen Bonds
Intermolecular forces are hierarchical and
additive.
Dipole-dipole forces are electrostatic forces
that occur between polar molecules.
Hydrogen bonds are especially strong dipoledipole forces that occur in molecules with
these bonds:
F-H O-H N-H
Hydrogen Bonds
London dispersion forces are weak, induced,
temporary dipole-dipole interactions.
These are the only forces between nonpolar
molecules.
They are strongest between large molecules
and atoms.
London Dispersion Forces
Intermolecular Forces
London dispersion forces
1 – 10 kJ/mol
Dipole-dipole forces
3 – 4 kJ/mol
Hydrogen bonds
10 – 40 kJ/mol
Single covalent bonds
150 – 550 kJ/mol
IM Forces and Boiling Points
Types of Solids
There are two types of solids:
Crystalline solids are characterized by a
regular three-dimensional arrangement of
the atoms, molecules, or ions that are make
up the substance. A crystal lattice is the
regular arrangement of these particles.
Amorphous solids are characterized by a
random, nonrepetitive three-dimensional
arrangement of the atoms, molecules, or
ions that make up the substance.
Types of Solids
Types of Crystalline solids:
Ionic
Network
Molecular
Metallic
Types of Amorphous solids:
Molecular
Network
Ionic solids are crystalline solids composed
of ions. The ions are arranged to maximize
interactions between unlike charges and
minimize interactions between like charges.
Crystalline molecular solids are composed
of molecules that are placed in a regular
array to maximize intermolecular forces.
Crystalline network solids are crystalline
solids in which the atoms are held in a
regular array by covalent bonds.
Metallic solids are made of metal atoms,
usually in a closely packed array. Electrons move freely among the atoms.
Amorphous molecular solids are composed
of large molecules (polymers and plastics)
that exist as random coils. Think spaghetti
but longer! They will melt and dissolve.
Amorphous network solids are usually polymers and plastics that have been crosslinked to
form covalent bonds between spaghetti
strands. They will not melt or dissolve.