Chapter 14 Liquids and Solids
Phase changes and
temperature
 Normally when heat is added the
temperature goes up.
 However when you hit a phase change
point (melting/freezing,
boiling/condensation)…
 The temperature stays constant when
heat is added (until the phase change is
complete).
Why should you…
 Turn the heat down once the water is
boiling?
 Recipes will always tell you to do this.
 Heat the water to a boil. Add spaghetti,
and turn the heat down.
 Won’t your spaghetti cook faster if you
turn the heat up?
 No
 The water can only get to 100o C
 Increasing the heat would increase how
fast it boils off, but that water leaves.
So a graph would look like…
Time vs. Temperature of water
under constant heat
boiling point
temperature
gas
100oC
0oC
melting
point
liquid
solid
time
Changes in phase require
energy
 It takes more energy to completely turn
water at 100o C into steam than it does to
take the same water from 0o C to 100o C.
 It actually takes 10x more energy to
convert 100o C water to steam than it
does to heat 0o C water to 100o C water.
 Steam has a much higher heat energy
content than 100o water.
 This is why steam burns are much worse
than water burns (scalding).
Phase Diagram graphs
 Phase changes normally occur with a
temperature change.
 However a change in pressure can also
force a phase change.
 Like the butane in a Bic lighter.
 It is a liquid inside (higher pressure), but
once released it is a gas (lower pressure).
 No temperature change caused this
Terminology
 Triple point is the point where the
substance can exist in all three phases of
matter. It is the meeting point of all three
phases
 Critical point is the temperature where no
matter the pressure, the substance will
always be a gas.
Phase Diagram Graph of H2O
Water is odd since the liquid is more dense than the
solid. This line normally veers the other way.
Critical point
Liquid
Solid
Pressure
normal melting point
101 kPa
Normal boiling point
Triple point
Gas
0o C
100o C
Temperature
“Normal” Phase Diagram
Liquid
Solid
Pressure
Triple point
Gas
Temperature
Critical point
Why is water more dense than
ice?
 Intermolecular forces- forces of
attraction between molecules that
forces them to come together to form
solids or liquids.
 Intermolecular Forces are collectively
called Van der Waals Forces.
 Don’t confuse these with bonds which
are intramolecular forces or forces that
hold a molecule together.
The bonds
holding
hydrogen and
oxygen
together are
intramolecular
forces
The forces between these two
water molecules are
intermolecular forces.
Phase changes
 When intermolecular forces are strong
enough to hold particles in place you
have a solid.
 As you increase the amount of energy
in the particles, they break free of Van
der Waals forces and start to move
around some. This is a liquid.
 When the atoms break free of all
significant intermolecular forces they
become a gas.
Dipole-Dipole Attraction
 There are several intermolecular forces
that we are not discussing.
 One specific intermolecular force is
dipole-dipole attraction.
 Remember we said some molecules
have a dipole moment or positive and
negative ends.
 A dipole-dipole attraction is when the
molecules arrange themselves so that
the opposite ends face each other.
Before Dipole-Dipole Attraction
Cl
H
Cl
H
H
Cl
Dipole-Dipole Attraction
Cl
H
Cl
H
Cl
H
Now the negative side (chlorine) is next
to the positive side (hydrogen)
A really strong dipole-dipole force
 A strong dipole-dipole force occurs when
you have a molecules that have hydrogen
bonding with nitrogen, oxygen or fluorine.
 This is called hydrogen bonding.
 The name is a misnomer, it is not an
intramolecular force (regular bond), it is an
intermolecular force.
 It is much weaker than a regular bond, but
stronger than the average intermolecular
force.
Hydrogen bonding in water
Why is liquid water more dense…
 Hydrogen bonding.
 In solid water, the molecules can’t
rearrange themselves.
 In liquid water, they are capable of
moving around.
 Normally random movement would
increase the spaces between molecules,
but with hydrogen bonding the molecules
“purposefully” move to a position where
they can be pulled in closer.
London Dispersion Forces
 ~A short lived induced dipole-dipole
attraction between atoms that don’t
normally have a dipole moment.
 An orbital is an area of probability of an
electron.
 The electron does not have a uniform
motion, at least we don’t think it does.
We don’t know what the motion of an
electron is
Probability
 Flip a coin 2 times, should you get 1 heads
and 1 tails?
 Not really. There is a chance you will but,
but it is only the highest probability.
 You have a 50% chance of getting 1 heads
1 tails, 25% chance of getting 2 heads and
25% chance of getting 2 tails.
 Apply that to our atom…
Helium
Positive nucleus with 2
electrons in a 1 s orbital.
It is neutral because the
negative electrons cancel
out the positive charges.
Imagine a line cutting the orbital in half.
What is the probability the electron is on
either side?
50/50, but just like the coin flip should we
always expect to find 1 electron on either
side? No. However, if we don’t…
Still Helium
 -  + None on
2 electrons
this side.
on this side.
This side is
This side is
now negative.
now positive.
Put this atom near another atom...
 -  + The positive side will
+
attract the electrons,
increasing the
chance of poles
forming again.
It not only forces another atom to have poles, but the
“new” atom forces the original to keeps its poles.
London Forces
 This force is random and short lived, as the
electrons do constantly move, and will
eventually end the dipole moment.
 It is also fairly weak.
 You can tell it is really weak in helium because
it stays a gas until -268.9o C.
 Larger atoms or molecules (with more
electrons) have stronger London forces.
 With more electrons it is easier for the atom or
molecule to have its electrons unbalanced and
stay that way for an extended period of time.
 Iodine (I2) is a solid at room temperature.
Evaporation and Vaporization
Evaporation
 Evaporation is a change in phase from
liquid to gas, but is not the same as
vaporization!
 Vaporization requires you to heat the
substance to its boiling point.
 Evaporation can happen at much lower
temperatures.
 Volatility- A measure of how easily a
liquid evaporates.
Is the vapor above the boiling
point?
 No it is not! (water vapor is not +100o C)
 It is possible to get matter in a phase that its
temperature does not agree with.
 It is like a solution (dissolved water in air)
 It is also possible to get liquids above or
below their freezing points. (supercooled or
superheated liquids)
Evaporation works like this
Liquids have molecules moving around in them
Temperature is the average kinetic
energy (which depends on the speed)
If they are
of these molecules.
moving
fast
Some are moving
enough,
at
just
faster than others!
the right angle,
some will escape
Molecules are
the surface of the
held in by
liquid
intermolecular
and turn into a gas.
forces.
These evaporated!
Why are they a “gas”
 Intermolecular forces determine whether
something is a solid, liquid or gas.
 In order to have intermolecular forces you
need to have multiple particles.
 The ones that escaped aren’t next to any
other particles.
 Since they have almost no intermoleluar
forces they have to be a gas.
Where did they go?
 They are in the air around the liquid.
 They are called vapors, anything that naturally
is a solid or liquid under standard conditions
that is currently a gas at standard conditions.
 If enough of them get together they will
condense and reform a liquid.
 As more of the molecules evaporate and fill the
air around the liquid, the chance that some of
them may condense increases.
 Provided the vapors can’t escape, the liquid will
reach a state where the rate of condensation
and evaporation equal each other.
Evaporation works like this
Liquids have molecules moving around in them
Temperature is the average
kinetic energy (which depends onIf they are
moving
fast
the speed) of these molecules.
enough, at just
Some are moving
the right angle,
faster than others!
some will escape
the surface of
Molecules are
the
held in by
liquid
intermolecular
and turn into a gas.
forces.
These evaporated!
Why are they a “gas”
 Intermolecular forces determine whether
something is a solid liquid or gas.
 In order to have intermolecular forces you
need to have multiple particles.
 The ones that escaped aren’t next to any
other particles.
 Since they have almost no intermoleluar
forces they have to be a gas.
Where did they go?
 They are in the air around the liquid.
 They are called vapors, anything that naturally
is a solid or liquid under standard conditions
that is currently a gas at standard conditions.
 If enough of them get together they will
condense and reform a liquid.
 As more of the molecules evaporate and fill the
air around the liquid, the chance that some of
them may condense increases.
 Provided the vapors can’t escape, the liquid will
reach a state where the rate of condensation
and evaporation equal each other.
Vapor Pressure
 Vapor Pressure of a substance is the
pressure of the vapor required for the rate of
evaporation and condensation to be the
same.
 At this pressure the substance will reach a
dynamic equilibrium.
 Dynamic means changing, equilibrium
means staying the same.
 At a molecular level, constantly molecules
are evaporating and condensing. However,
since these cancel out, there is not net
change.
After a liquid evaporates
 The remaining liquid is cooler.
 This is because the molecules with the most
kinetic energy (heat) escaped.
 Water has a “cooling” effect because it
evaporates.
 Sweat cools your body by evaporation.
 Provided it is not humid out.
 Humidity is a measure of the amount of water
vapor present in the air.
Muggy (humid) weather
 In humid weather, the water vapor in the
air is closer to its vapor pressure.
 Less net water can evaporate, and cool
you off.
 The rate of evaporation hasn’t changed,
but more water vapor is condensing than
normally.
Increasing Evaporation
 Intermolecular forces play a big part.
 Low molecular forces mean the substance
will easily evaporate. These substances are
volatile.
 Evaporation occurs at the surface of a liquid
so increasing the surface area will increase
the rate of evaporation.
 Allow evaporated vapors to escape so it
can’t reach vapor pressure.
 Heat the substance to increase kinetic
energy.
Why do fans/wind feel cool?
 The majority of the water vapor from your
sweat is directly around you.
 A fan or wind pushes air from somewhere
else over to you, and the air that was around
you somewhere else.
 The water vapor that evaporated can’t
condense back on you.
 This only works if it isn’t extremely humid
out.
 If it is extremely humid the air from
somewhere else contains a lot of water
vapor that will condense on you.
Vaporization or Boiling
 Evaporation occurs at the surface of a liquid.
 As you continually heat a liquid, the particles
inside move faster.
 Eventually the particles move so quickly,
that they break free of all intermolecular
forces and form gas pockets inside of liquid.
 These are always less dense than the liquid
so the float to the surface and escape.
 This is vaporization or boiling.
Evaporation and Vaporization
Evaporation occurs at the surface
If I get the
substance hot
enough
Evaporation and Vaporization
This is vaporization or boiling
I can force gas
bubbles to form in
the middle of the
liquid
Types of Solids
Solids
 Crystalline Solids- have a regular repeating
arrangement of their particles.
 Salts, Sugars, Metals
 Amorphous Solids- have no regular
repeating arrangement of their molecules
 Common glass, several polymers.
Crystalline Structure
Amorphous
Amorphous solids
 Amorphous solids, due to a lack of
arrangement of molecules, will actually flow,
slowly.
 If you look at very old windows, you will find
there is more glass at the bottom than at the
top. That is because the glass flowed down.
 You can also see the same effect with Silly
putty.
Making solids…
 Technically, anything can be made
amorphous.
 A rapid cooling from liquid to solid makes it
amorphous. The particles just don’t have
time to arrange themselves in a pattern.
 A slower cooling or heat treatment can
make some amorphous solids crystalline.
Safety Glass
 Cars don’t use common glass for their
windshield because it breaks into dangerous
shard when it breaks.
 Instead they use a heat strengthened glass,
one that is slowly cooled to a solid to allow
for a better arrangement of molecules, so
that when it breaks it breaks into less
dangerous “dice”.
Glass
Safety Glass
Back to crystalline solids
 Crystalline solids can be made up of 3
different things
 Ionic Solids –made of ions
 Molecular Solids- made of molecules held
together by covalent bonds
 Atomic Solids- Made of atoms
Ionic Compounds
 Ionic Compounds have very high melting
points.
 Sodium Chloride melts at 801oC
 That is because every single negative
particle is attracted to every single positive
particle and vice versa.
 This is in essence a very strong
intermolecular force.
Ionic Solids
 Ionic solids are brittle. When they break
their crystal structure shows, as they break
into similar shapes.
 NaCl breaks into
CaCl2 into cubes
spheres.
Conduction of electricity
 Electricity is a flow of electrons
 Anything that allows electrons to easily pass
through will be a good conductor of
electricity.
 While solids, electrons can only jump from
ion to ion.
 This is a very slow process so solid ionic
compounds are not good conductors.
Melts and solutions
 If you melt an ionic compound, then the ions
can move. Electrons can now easily move
through the substance.
 If you dissolve an ionic compound, the ions
are also free to move.
 Therefore, liquid ionic compounds and ionic
solutions are good conductors.
Molecular Compounds
 Molecular Compounds have much lower
melting points.
 Several are liquids (water) or gases (carbon
dioxide) at room temperature.
 Molecular compounds are not good
conductors of electricity.
Atomic Solids/Elements
 Solid nonmetals and metalloids commonly
form very large molecules.
 A diamond (any size) could actually be
viewed as one molecule of all carbon.
 These solids are called network solids.
 They have high melting points and don’t
conduct electricity.
Allotopes of Carbon
Nonmetal Gases
 Noble gases and diatomic elements
(except bromine, and iodine)
 These all have only London dispersion
forces.
 These are very weak intermolecular
forces.
 They all have very low melting points,
obviously since they are gases.
 None are good conductors
Bromine and Iodine
 These act the same as the other diatomic
elements but since the atoms are larger the
London dispersion forces are greater.
 That is why they are a liquid (bromine) or a
solid (iodine) at room temperature.
Metals
 Metals have high melting points and are
good conductors of electricity.
 Metals are held together by metallic bonds.
 Similar to ionic bonds these are somewhere
in between intramolecular forces and
intermolecular forces.
Metallic Bonding
 Bonds between metals
 Metallic bonds only occur with the same
metal not with other metals.
 Ca can bond with other Ca atoms, but not
Ba.
Metallic Bond
 In metallic bonds the valence electrons
become community property, traveling
anywhere they want to throughout the
element.
 This “Sea of Electrons” is why metals are
such good conductors of electricity and
heat.
Model of Metallic Bonds
Calcium has 2 valence electrons
Ca
Ca
Ca
Ca
Ca
Ca
All of the
electrons
move like
this.
The “sea of electrons” is kind of like
bees (valence electrons) swarming
around a few flowers (rest of the atoms).
Properties
 The nuclei inside the “sea of electrons” are
movable without breaking the structure.
 This is why metals are malleable and ductile.
 Electrons can easily move through so they
are great conductors of electricity.
 Heat is the speed of the particles. If I heat
up electrons at one end they quickly hit the
slower moving ones and speed them up. So
the whole material gets hot. That is why they
conduct heat.
Alloys
 ~a substance that is mixture of elements
and has metallic properties.
 Alloys are mixtures so they can be
separated without chemical reactions
 Steel is an alloy. It is made of iron and 0.21.5% carbon.
 The carbon makes it harder, stronger, and
less malleable than normal iron.
 More carbon makes it stronger.
Interstitial Alloy
 Steel is an interstitial alloy because the
carbon atoms fit into the “holes” between the
iron atoms in the crystal structure.
Substitutional Alloy
 A substitutional alloy is when a metal atom
of similar size replaces the host metal.
 Brass (copper and zinc), sterling silver
(silver and copper), white gold (gold,
palladium, silver, and copper) are all
substitutional alloys.
 This changes the
properties of the metal.
Both substitutional and interstitial alloys
 Stainless Steel is iron and carbon
(interstitial) mixed with chromium and nickel
(substitutional).
 It resists corrosion.
 Slightly changing the presence of any of
these drastically changes the properties of
the final metal.