Chapter 11: States of Matter &
Intermolecular Forces
Essential Question: Sections 1&2
How do particles interact with each
other and how does this affect
their properties?
Do You Remember…?
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Electronegativity
Polarity
Dipole
Ion
Terms to Remember
• Electronegativity: how well an atom pulls a
shared pair of electrons toward itself.
– Think tug of war!
– Higher value = better at pulling electrons
• Polarity: unequal sharing of electrons results in
partial positive/negative charges.
• Dipole: a molecule with partial positive and
partial negative charges.
• Ion: atom with a charge
– Not a partial charge! Electrons have been lost or
gained to give a full charge.
Bonds vs. Intermolecular Forces of
Attraction (IFA)
The H and O atoms in
each water molecule are
held together by a bond.
Each water molecule is
attracted to another
water molecule because
there is a force of
attraction between
them.
Intermolecular Forces of Attraction
• Each person will be given a type of force to
research and summarize.
• After completing the summary, people will be
put into groups of three to explain their
information to others.
Intermolecular Forces of Attraction
(IFA)
• Ionic compounds are held together by attractive
ionic forces between ions.
• Two factors affect strength of ionic forces:
– Size of charge on the ion.
– Ex: NaCl vs. MgCl2; MgCl2 has stronger forces because
Mg+2 has a larger positive charge than Na+.
– Size of the ion itself (radius).
– Ex: NaCl vs. KCl; NaCl has stronger forces because Na+
is smaller than K+.
• Stronger forces = higher boiling/melting point!
Types of IFA
• Three types:
– London Dispersion
– Dipole-Dipole
– Hydrogen Bonding
• In general, the strengths of each are:
London Dispersion < Dipole-Dipole < H Bonding
• Bonds between atoms (covalent & ionic) are
MUCH stronger than any of these
intermolecular forces!
London Dispersion
• Arise from random shifts in electrons in
substances.
• Temporary (instantaneous) dipoles cause
temporary partial charges.
• These temporary dipoles can induce other
temporary dipoles in molecules nearby.
• Partial charges attract molecules, until they
disappear.
• All substances exhibit this type of
intermolecular force (even ionic compounds)
because all substances have electrons!
LDF Continued
• LDF are most noticeable in nonpolar
substances because this is the only
intermolecular force holding molecules
together.
• LDF increases in strength with more
electrons/as molecules get bigger.
• A very large, nonpolar molecule can have very
strong LDF.
Dipole-Dipole
• IFA exhibited in polar covalent
compounds.
• Polar = permanent partial
charges.
– Electronegativity differences!
• Partial positive charges and
partial negative charges attract
in different molecules and hold
them together.
• Stronger partial charges =
stronger dipole-dipole forces.
Hydrogen Bonding
• Some dipole-dipole forces are
REALLY strong and have a
class of their own.
• Hydrogen bonding: O, F, or N
bonded to a H atom causes
VERY strong partial charges to
develop.
• If a molecule has one of these
atoms bonded to a H, it will
exhibit hydrogen bonding.
Polar vs. Nonpolar
• Use electronegativity values to determine
polarity!
• Cutoff points?
– Greater than 2.1 = ionic
– Between 2.1 and 0.5 = polar covalent (partial
charges)
– Less than 0.5 = nonpolar covalent (no partial
charges)
Effects on Properties
• The type and strength of intermolecular forces
of attraction influence MANY properties of
substances. We will examine the following:
– State of matter at room temperature
– Melting and boiling points
– Cohesion
– Adhesion
– Surface tension
– (Also explains why certain liquids do/do not mix
well together)
States of Matter
Solids
• Fixed, rigid position
• Held tightly together by strong intermolecular
forces of attraction
Liquids
• Held loosely together by medium
intermolecular forces of attraction
• Not held tightly enough to prevent motion;
particles slide past each other
Gases
• Particles are extremely far apart, barely held
together due to weak intermolecular forces of
attraction
• Particles move quickly and almost
independently
Summary of States
• What to remember: weakest intermolecular
forces = gases; intermediate forces = liquids;
strongest forces = solids
Melting & Boiling Points
• Remember that when substances melt and
boil their motion increases.
• Stronger intermolecular forces of attraction
means more energy is needed to break these
forces of attraction and allow for more
motion.
• This results in higher MP and BP.
• What to remember: stronger intermolecular
forces = higher MP and BP.
Cohesion vs. Adhesion
• Cohesion: attraction of liquid particles to each
other.
• Adhesion: attraction of liquid particles to solid
surfaces.
• The strength of these two forces determines if
a liquid will spread out over a solid surface or
bead up.
– Lesser adhesion = beading.
– Greater adhesion = liquid spreads over surface
• Strength of cohesion and adhesion depends
on intermolecular forces present.
Example: Water
• Water on wax paper?
– Droplets bead all over the surface
– High cohesion for each other, low adhesion to wax
• Surface of water in a glass?
– Meniscus!
– Stronger adhesion forces pull water molecules up the
sides
– These water molecules pull other water molecules
with them (very cohesive)
– Water rises up until adhesive and cohesive forces are
counterbalanced by gravity (capillary action)
Surface Tension- A Closer Look
• You know what this is! Here’s why it exists:
– Particles below the surface feel attraction
(cohesion) in all directions.
– Particles at the surface only feel attraction
sideways and downward. Creates ‘tension’ at the
surface.
– This is why water bugs can
stand/walk on the surface of
water. There’s enough tension
to support their weight and
prevent them from falling in.
Surface Tension Cont.
• The stronger the intermolecular forces of the
liquid, the stronger the attractions will be.
• Stronger attractions means greater surface
tension.
Review Questions
• Explain how adhesion plays a role in the
formation of a meniscus in a glass of water.
• Describe the motions of particles when water
boils.
• Water has hydrogen bonding intermolecular
forces, and olive oil has London dispersion.
Which would you expect to have a greater
surface tension? Why?
• Which has a higher BP: CO2 or H2S? Why?
Lesson Essential Question: Section 3
• How are enthalpy and entropy involved when
matter changes?
Section 3: Energy & State Changes
• Thermodynamics: studies the effect of heat,
work, and energy on a system.
– Enthalpy and entropy are two important properties.
• Enthalpy vs. entropy
– Enthalpy: total energy of a system, H.
• Depends upon heat energy & kinetic energy of
particles.
– Entropy: disorder in a system, S.
• Not a form of energy!
• Both influence whether a change will occur or
not.
Enthalpy
• Does matter prefer to be at higher or lower
energy?
– Hint: think about the matter around you- is it
constantly undergoing violent, energetic reactions?
Are people constantly running around at high energy
levels?
• Lower energy is favorable: -∆H.
– Less internal energy = greater stability.
– This is why energy must be added to get substances
to melt and vaporize; the particles have greater
kinetic energy!
Determining Enthalpy Changes: ΔH
• The sign of ∆H can be determined by examining
temperature changes.
• If a change is exothermic (releases heat), ΔH is
negative.
• If a change is endothermic (absorbs heat), ΔH is
positive.
• This is true for physical and chemical changes!
You just have to look at temperature changes.
– Remember, you have to focus on the system- what
is actually undergoing the change!
Enthalpy Examples
• When NaOH is dissolved in water, the
temperature of the solution increases from
22.8°C to 24.6°C . (Remember, focus on the
system- the NaOH dissolving.)
– What is ∆T? T increases, so +∆T (for surroundings)
– What is ∆H? Heat is released by system, so -∆H
• When water freezes, what happens to the
temperature?
– It drops! Therefore you have a -∆T (for system).
– What is ∆H? Heat is released by the system,
so -∆H
Enthalpy Examples Cont.
• When hydrochloric acid is added to sodium
hydroxide, the temperature increases.
– Is this reaction exothermic or endothermic?
Exothermic
– What is the sign for ∆H?
-ΔH
– Animation
Entropy
If bricks fall off of a truck, what would the
pile look like?
• More entropy, +∆S, is favorable.
• Nature has a tendency to favor disorder.
• Does your room get messier over time or does
your room get cleaner over time?
– You have to do work to organize/clean your room
because it tends to get messy over time!
Entropy
• Food coloring demonstration.
• Particles also like chaos/spreading out!
– Ex: diffusion occurs because particles have a
natural tendency to spread out and become more
disordered. It would take work to reverse this
process and force the particles to come back
together.
• A gas in a container that is opened will float
out into the air. Why?
– More space in the air to spread out = more
favorable disorder!
Entropy- What to Look For
• In state changes look at the states on both sides of the
equation. If the substance is changing into a state with
more motion, +ΔS. If not, -ΔS.
– Order of states (least to most disorder): s < l < aq < g
– Ex: H2O (g)  H2O (l)
-ΔS
• In chemical reactions, look at states and the number of
moles of different substances on each side.
– Ex: 2Na (s) + 2H2O (l)  2NaOH (aq) + H2 (g) +ΔS
What Can ∆H & ∆S Tell Us?
• Based on the signs of ∆H and ∆S, we can more
easily determine whether a process will occur (is
spontaneous) or not.
• Lower energy is favorable, -∆H, and higher entropy
is favorable, +∆S.
• If a process releases energy and at the same time
increases its entropy, what can you conclude about
the likelihood of that process occurring?
– Very likely because it is favorable in terms of both
enthalpy and entropy!
What Can Both ∆H & ∆S Tell Us?
• What can you also say about the opposite: if a
process absorbs energy and at the same time
decreases its entropy, will the process occur?
– Think about the signs for ∆H and ∆S!
• +∆H and -∆S.
– Very unlikely because it is unfavorable in terms of
both enthalpy and entropy!
• But what about processes that are favorable in
terms of one and not the other? In other words
when both signs of ∆H and ∆S are the same?
– T is the last component that determines if the
change will occur or not!
Spontaneity & Gibbs Free Energy
• There must be another quantity that determines if
a change will occur besides just ∆H and ∆S.
• ∆G = ∆H -T∆S  Notice T is considered!
• G stands for Gibbs free energy: energy in a system
that’s available for work.
– If ∆G is negative (the system gives off free energy), the
process is spontaneous- it will occur.
– If ∆G is positive (the system takes in free energy), the
process is nonspontaneous- it will not occur.
– If ∆G equals zero, the process is at equilibrium (forward
and reverse reaction rates are the same).
Spontaneity & Gibbs Free Energy
• ∆G = ∆H -T∆S
• This formula allows you to make conclusions
about spontaneity even if one term is favorable
and the other is not.
– Remember, it depends on T!
• When ΔH and ΔS are positive, at high
temperatures the reaction will be spontaneous.
• When ΔH and ΔS are negative, at low
temperatures the reaction will be spontaneous.
Summary
Calculations
• ∆G = ∆H -T∆S
• ∆Hfusion is used for fusion (melting) or freezing;
∆Hvaporization is used for vaporization or
condensation.
– Number is the same, sign changes (+ if E taken in, if E released).
• Same is true for ∆S.
– + if disorder of new phase increases, - if disorder of
new phase decreases.
• T is temperature in Kelvin, K.
• We can calculate the value for any of these.
Calculating Gibbs Free Energy
• The sign of ∆G indicates whether the state will
change or not at a given temperature.
• Ex: Will ice melt at 273.00K?
• ∆H fusion: 6,009 J/mol
• ∆S fusion: 22.00 J/(mol·K)
Just under 0°C.
• Signs of ∆H and ∆S?
– Melting requires energy be taken in (+∆H = +6,0009 J/mol).
– Changing from solid to liquid means more entropy (+∆S =
+22.00 J/mol·K).
• ∆G = 6,009 J/mol – (273.00K) x (22.00 J/mol·K)
• ∆G = 6,009 J/mol – 6,006 J/mol
• ∆G = + 3 J/mol  Ice will not melt!
∆G During State Changes
• During state changes, states are in equilibrium with
each other (both phases are present).
• Therefore, ∆G = 0.
• Can rearrange formula to find melting and boiling
points:
∆G = ∆H -T∆S
0 = ∆H -T∆S
T∆S = ∆H
• T = ∆H/∆S
Tmp= ∆Hfus/∆Sfus & Tbp = ∆Hvap/ ∆Svap
A Note About Pressure
• So far we have only looked at the effect of T
on phase changes (besides G, H, and S).
• P also has an effect, but only when gases are
involved.
– Gases are P dependent because they are
compressible.
– Liquids and solids are not compressible, so not
affected by P during phase changes.
• This is examined more closely in AP chemistry.