Liquids and solids slides

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Liquids and Solids

AP Ch i t R id Learning Series
Chemistry Rapid L
i
S i

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Wayne Huang, PhD
Kelly Deters, PhD
Russell Dahl, PhD
Elizabeth James, PhD
Debbie Bilyen, M.A.

© Rapid Learning Inc. All rights reserved. :: http://www.RapidLearningCenter.com

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© Rapid Learning Inc. All rights reserved.

1


Learning Objectives
By completing this tutorial you will learn…
What intermolecular forces are
Properties of liquids
P
ti
f li id
How vapor pressure is affected
by intermolecular forces and
temperature
How solids are structured
How matter changes states
The energy changes during a
phase change

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Concept Map
Previous content
Chemistry

New content
Studies

Matter

Bonding
Structures

Has different
Have different
possibilities of

Solids

One is

States

One is

Liquids

Can change by
g y
Involves breaking
or forming

Phase Changes

Intermolecular
Forces

Boiling and melting point
determined by

Vapor Pressure

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2


Intermolecular
Forces

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Definition: Intra & Intermolecular
Intramolecular Forces –
Chemical bonds within a
molecule.
Intermolecular Forces (IMF) –
Physical attractions between
separate molecules.
t
l
l

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3


London Dispersion Forces
All molecules have London Dispersion Forces.
Due to a temporary “ganging up” of electrons on one side
of the molecule.
Nucleus

δ-

+

-

δ+

-

+
-

-

Electrons

-

As the l t
A th electrons move around the molecule, they can
d th
l
l th
temporarily end up on one side.
This results in a portion of the molecule that has a partially
negative charge and a portion with a partial positive charge.
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London Dispersion Forces Properties
Properties of London Dispersion Forces:
Temporary
p
y

The electrons continue
g
p
moving and “spread out”
again.

Weakest IMF

It’s temporary, and
therefore weak.

All molecules have them

All molecules have
electrons which can “gain
up”.
up”

The larger the molecule,
the larger the force

The more electrons in a
molecule, the greater the
effect of “ganging up”.

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4


Dipole-Dipole Forces
Polar molecules exhibit dipole-dipole forces.
Polar molecules have a permanent partial separation of charges.

δ+

δ-

Polar Molecule

e.g. sugar and water

The opposite charges on separate polar molecules are
attracted to one another.
They are not chemical bonds, just physical attractions
between opposite charges.
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Ion-Dipole Forces
An ion and a polar molecule exhibit ion-dipole forces.
Polar molecules have a permanent partial separation of charges.

δ+

δ-

Polar Molecule

+
Cation

e.g. water and Na+

Ions are attracted to the partially charged regions of a polar
molecule.
They are not chemical bonds, just physical attractions
between opposite charges.
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5


Dipole Forces Properties
Properties of Dipole Forces:
Permanent dipoles
p

Partial separation of
g
charge within the
molecule is permanent.

Stronger IMF

Although the IMF is not
permanent, the ability to
form the IMF is.

Only polar molecules
can exhibit them

A permanent dipole
(separation of charges) is
needed.

The stronger the dipole,
the stronger the force

Stronger dipoles have
greater attraction to other
charges.

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Hydrogen Bonding
Hydrogen bonding is an especially strong case of
dipole-dipole forces.
Hydrogen atoms contain only 1 proton and 1 electron.

- - - O
- - H
H

δ-

δ+

When hydrogen is bonded to a very electronegative atom (N,
O or F), the separation of charges is very large as there are
no other electrons around the hydrogen proton at all.
The hydrogen in this extreme dipole can be attracted to the
lone pairs on an N, O or F atom on another molecule.
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6


Hydrogen Bonding Properties
Properties of Hydrogen Bonds:
Extreme case of dipole
forces

H has no other electrons,
it is very positive when
sharing electrons with a
very electronegative atom.

Strongest IMF

Stronger than typical
dipole forces

Molecules with an H on
an N, O or F can
hydrogen bond

With lone pairs on N, O or
F atoms in another
molecule

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Intermolecular Forces Summary
A summary of the 4 IMF’s:
Type of force

Happens with

Relative
strength

London
Dispersion

All molecules

Weakest

Dipole-dipole

Polar molecules

Medium

Ion dipole
Ion-dipole

Ion & polar molecule

Medium

Hydrogen
bonding

H on an N, O or F with
another N, O or F

Strongest

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7


Liquids

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Properties of Liquids
Some general properties of liquids:
Property

Example

Definite volume

No definite shape

Drink pools out when
poured on a table.

Molecules are free to
move past each other

A drop of food coloring
moves through the liquid
over time.

Not very compressible
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Pouring a drink into a
larger glass doesn’t make
more of the drink.

A full bottle of water (no
air) cannot be
compressed.


8


Definition: Vapor Pressure
Vapor Pressure – Pressure
created above a sample by
t d b
l b
particles evaporating from
the sample and becoming
gas particles.

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Vapor Pressure of a Liquid
Solvent particles on the very top layer of the
sample can evaporate.
Looking down on the top of a solution in a beaker:

Side view

Gas particles
now cause
pressure-vapor
pressure.

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9


Definition: Equilibrium
Equilibrium – Rate of change equals the
rate of the opposite change.
Rate forward = Rate reverse

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Vapor Pressure Equilibrium
Vapor Pressure equilibrium can be achieved in a
closed system.
The rate of evaporation of a liquid is constant.
constant
rate

evaporation

condensation

time

As the liquid evaporates, gas particles are formed which
can then begin to condense down to the liquid form again.
As more gas particles are created, the rate of condensation
increases.
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Equilibrium is reached when rate evaporation = rate
condensation.


10


Volume of the Liquid and Equilibrium
The volume change as equilibrium is established.

Liquid volume
decreases as
initial gas
particles are
formed.

Liquid volume
begins to
increase as gas
particles begin
to re-condense.

Liquid volume is
constant once
equilibrium has
been reached.
It’s less than the
initial volume.

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Vapor Pressure and Temperature
Solvent particles on the very top layer of the
sample can evaporate.
In order to evaporate, the particle must have enough
energy to break the intermolecular forces connecting it
with the other liquid particles.

Higher temperature
means the
average kinetic
energy of the
molecules is
higher.

More molecules
have the
minimum
energy needed
to vaporize.

As temperature
increases,
vapor pressure
increases.

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11


Solids

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Properties of Solids
Some general properties of solids:
Property

Example

Definite volume

A solid cannot “grow or
shrink”

Definite shape

Solids cannot
spontaneously change
shape

Molecules are not free to
move past each other

A dot of ink doesn t travel
doesn’t
across the top of a desk

Not compressible

A piece of ice cannot be
compressed

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12


Types of Solids Structures - Part 1
Solids

Amorphous
Solids

Crystalline
Solids

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Definition: Amorphous and Crystalline

Amorphous Solid – Has a
fair amount of disorder in
the structure.

Crystalline Solid – Has a
highly regular, repeating
arrangement of particles.
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13


Amorphous Solids
Amorphous solids are still rigid—they are still
solid…
But they’re p
y particles are not trapped in a repeating
pp
p
g
pattern as in crystalline solids.

An example is glass.
The particles are trapped in a position before they have a
chance to arrange themselves in a repeating pattern.
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Crystalline Solids
Crystalline structures are lattice structures,
composed of unit cells.
Lattice
Overall crystalline
structure

Unit Cell
Smallest repeating
unit
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14


Types of Unit Cells
These are the most simple unit cells are:

Cubic
8 particles create a cube

Body-centered cubic
A cube with a particle in the center
of the “body” of the cube

Face-centered cubic
A cube with a particle in the center of
each “face” of the cube
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Types of Crystalline
Structures

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15


Types of Solids Structures - the Rest
Solids

Amorphous
Solids

Crystalline
Solids

Atomic
Solids

Metallic

Molecular
Solids

Ionic
Solids

Network

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Definition: Atomic Solids

Atomic solids –
Atoms are the
particles in the
unit cell.

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16


Metallic Solids
Metallic bonding contains metal atoms packed
closely together and bonded to the atom in each
direction equally.

This is called “closest packing”.
packing”
The electrons form a large “pool” that are free to move
throughout the structure.
This allows metals to conduct electricity!
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Network Solids
Network solids can be thought of as one giant
molecule.

Graphite

All of the atoms in the network solids are covalently bonded
to their neighbors.
Examples of network solids include: graphite, diamonds, silica (sand)

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17


Definition: Molecular Solids

Molecular S
Solids – Molecules are
the particles in the unit cell.

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Molecular Solids
Strong covalent forces within the molecule with
weaker forces between the molecules.

An example of a molecular solid is ice.

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18


Definition: Ionic Solids

Ionic Solids – Structure containing
positive and negative ions (cations
and anions).

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Binary Ionic Solids
Binary ionic solids have a variation of the closest
packing structures.

The larger ion is packed in the closest packing structure.
The smaller ion fits in between the holes created by the
larger ion.
This minimizes repulsions from like-charged ions.
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19


Solid Structure Summary
Type

Unit cell

Bonding

There is no
unit cell

Strong intramolecular,
S
weaker intermolecular

Metallic

Metal atoms

Pool of electrons that are free
to move

Network

Non-metal
atoms

Covalent bonding throughout

Molecular

Molecules

Covalent within molecule,
weaker between

Ionic

Ions

Electrostatic attraction
between ions

Crysta
alline

Amorphous

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Solid Structure Properties
General properties of the different structure types
Type
Amorphous

Properties
Disorder in their structure—no repeating
pattern
Excellent conductors of heat and electricity,
malleable and ductile

Network

Brittle, poor conductors of heat and
electricity

Molecular

Strong bonds within but weak between (not
much energy needed to melt, but a lot to
break the chemical bond)

Ionic

Crystal
lline

Metallic

Stable, high melting points, brittle

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20


Phase Changes

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Definition: Phase Changes
Phase Change – Changes between solids,
liquids and gases.
sublimation
melting

Solid

boiling

Liquid

Gas
condensing

freezing

Deposition
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21


IMF’s and Phase Changes
What role to intermolecular forces play in phase
changes?
Break all IMF’s

Break some
IMF’s

Solid

Break rest
of IMF’s

Liquid

Gas
Form some
IMF’s

Form more
IMF’s

Form IMF’s

Breaking IMF’s requires energy

Forming IMF’s releases energy

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Definition: Boiling & Melting Points
Boiling/Condensation Point – When
liquid and gas phases are at
equilibrium with one another.

Melting/Freezing Point – When solid
and liquid are at equilibrium.

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22


How is Boiling Point Determined?
Boiling point is where …
vapor pressure = atmospheric pressure.
When the vapor pressure of the liquid is great enough to
form “bubbles” boiling will occur.
Atmospheric pressure

Bubbles rise (less
dense)…boiling!

As temperature increases, the
vapor pressure of the liquid
increases.
increases
As vapor pressure within the
liquid increases, bubbles can
begin to form—pushing
against atmospheric pressure.

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How is Melting Point Determined? - 1
Melting point is where…
vapor pressure liquid = vapor pressure solid.
Temperature 1: Solid vapor pressure > Liquid vapor pressure

Molecules escape
the solid faster
than they
escape the
liquid.

The solid particles
escape and join
the liquid…but
q
the liquid
molecules
aren’t crossing
over as fast.

After time…end up
with all liquid.
Temperature is
above
melting/freezing
point.

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23


Temperature 2: Solid vapor pressure < Liquid vapor pressure

Molecules escape
the liquid faster
than th
th they
escape the
solid

The liquid
particles
escape and join
the solid…but
solid but
the solid
molecules
aren’t crossing
over as fast

After time…end up
with all solid.
Temperature is
T
t
i
below
melting/freezing
point

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Temperature 3: Solid vapor pressure = Liquid vapor pressure

Molecules escape
the liquid at the
sa e
same pace as
they escape the
solid.

The liquid
particles
escape and join
the solid at the
same rate the
solid particles
escape and join
the liquid.

At equilibrium.
Amount of solid
and liquid don’t
change over
h
time.
Temperature is
freezing point.

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24


Why Do Some Solids Sublime?
Why do some substances (such as dry ice) go
straight from solid to gas?

The intermolecular forces are very weak.

The solid particles have enough energy to break all
of the IMF’s and go straight to a gas, rather than
IMF s
gas
only breaking some of them and going to a liquid.

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Phase Diagrams - 1
Phase diagrams show what state of matter a
substance would exist as at various temperatures
and pressures.
Phase Diagram for H2O
Triple Point
All 3 states exist
together

Pressur (atm)
re

melting

l
s
1
0.006

Critical Point
Point above which it
cannot exist as a liquid

boiling
sublimation
g
0.0099

100

Temperature (°C)
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25


Phase Diagrams - 2
Phase diagrams show what state of matter a
substance would exist as at various temperatures
and pressures.
Phase Diagram for H2O

Pressur (atm)
re

freezing

l
s
1
0.006

condensing
deposition
g
0.0099

100

Temperature (°C)
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Energy of Phase
Changes

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26


Definition: Enthalpy of Fusion

Enthalpy of fusion ( fus) – energy
py
(H
gy
necessary to break enough IMF’s to
turn a solid into a liquid.

Energy released when turning a
liquid into a solid = - Hfus
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Enthalpy of Fusion Calculations
Enthlapy of fusion is used in calculating energy
needed to melt or released when freezing.

ΔH = m × H fus

Example:

H = enthalpy (energy)
m = mass of sample
Hfus = enthalpy of fusion (melting)
(use –Hfus for freezing)

Find the enthalpy of fusion of water if it takes
4175 J to melt 12.5 g of water

4175 J = (12.5 g ) × H fus
Hfus = ?
m = 12.5 g H2O
ΔH = 4175 J

4175 J
= H fus
12.5 g

Hfus = 334 J/g

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27


Definition: Enthalpy of Vaporization

Enthalpy of vaporization (Hvap) –
py
p
(
energy necessary to break the rest of
the IMF’s and turn a liquid into a gas.

Energy released when turning a gas
into a liquid = - Hvap
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Enthalpy of Vaporization Calculations
Enthalpy of vaporization is used in calculating
energy needed to vaporize or released when
condensing.

ΔH = m × H vap

Example:

H = enthalpy (energy)
th l (
)
m = mass of sample
Hvap = enthalpy of vaporization (vaporizing)
(use –Hvap for condensation)

If the heat of vaporization of water is 2287 J/g,
how much energy is released when 15.75 g of
water is condensed?

Hvap = - 2287 J/g
m = 15.75 g H2O
ΔH = ? J

ΔH = (15.75 g ) × ⎛ − 2287 J
⎜
⎝

⎞
g⎟
⎠

ΔH = -36020 J
Energy is released because it’s condensing.
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Solids, Liquids &
The AP Exam

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Liquids & Solids in the Exam
Common Liquid & Solid problems:
Using intermolecular forces to describe properties of
compounds
d
What types of particles are at lattice points in
different types of solids
The properties different types of solid bonding
exhibit

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29


Multiple Choice Questions
Example: Which of the following has the highest melting point?
A.
B.
C.
D.
E.

F
Cl
Br
I
Same

Melting point is largely determined by intermolecular forces.
All of these are pure non-metal elements.
They have London Dispersion Forces only.
The larger the atom, the more electrons it has, the more London
Dispersion Forces it has.
The more London Dispersion Forces it has, the more energy is
needed to break the IMF’s to melt the solid.
Answer: D
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Free Response Questions
Example:

Use chemical bonding and/or intermolecular forces to
explain the following observations:
A. At STP, propane (C3H8) is a gas, while octane (C8H18) is
,p p
(
g ,
(
a liquid.
B. MgO melts at a much higher temperature than CO2
C. Ethanol (CH3CH2OH) dissolves in water much more
easily than ethane (CH3CH3).

These are the two sub-questions you can
answer after this tutorial.

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Answering Free Response Questions
A.

At STP, propane (C3H8) is a gas, while octane (C8H18) is a liquid.
Both molecules only have London Dispersion Forces. Propane
is a larger molecule, therefore it has greater LDF. The greater
g
,
g
g
the IMF’s, the more energy is needed to melt or boil a
substance. C8H18 is larger, and therefore has more IMF’s, and
therefore needs more energy to boil—which it hasn’t gotten at
STP and is therefore a liquid still.

B.

MgO melts at a much higher temperature than CO2
MgO is an ionic compound while CO2 is a non-polar covalent
compound. MgO has London Dispersion Forces and Ion-Ion
interactions while CO2 only has LDF. The more IMF’s, the more
energy is needed to pull apart molecules (melt) and therefore, the
MgO melts at a higher temperature.

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Learning Summary
Matter undergoes
phase changes,
shown by phase
diagrams.

Solids and
liquids are
condensed states
of matters that
have
intermolecular
forces.

Solids have
structures—either
t
t
ith
amorphous or
crystalline (which can
be atomic, molecular
or ionic).

The energy
changes during
phase changes
are governed by
db
enthalpies of
fusion and
vaporization.

Liquids d lid
Li id and solids
have vapor pressures
as particles escape to
the gas form.

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Congratulations
You have successfully completed
the core tutorial

Liquids and Solids

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Chemistry :: Biology :: Physics :: Math

What’s N t
Wh t’ Next …

Step 1: Concepts – Core Tutorial (Just Completed)
Step 2: Practice – Interactive Problem Drill
Step 3: Recap – Super Review Cheat Sheet

Go for it!

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