• Reading Review Chapter 15In the lecture slides, we stopped after each chapter to cover review questions. Upload the solutions to the problems from chapter 15 here.

CH1000
Fundament
als of
Chemistry
Module 4 – Chapter 13

Liquid State of Matter

• Liquids are an intermediate between gases and solids

• They contain particles close to one another but have fluidity (can
assume the shape of a container)

• Significant attractive forces exist between particles in a liquid.

• Liquid Review:
• Close contact
• Some attractive forces
• Fluid shape

Changes in State

Vaporization
Liquid to Vapor

Molecules in liquid state have
different kinetic energies (KEs)

Those with higher KEs can
overcome attractive forces
between particles and escape to
the gas phase

Sublimation
Solid to Vapor

Phase change from the solid to gas
phase that bypasses the liquid
state

Condensation
Vapor to liquid

Molecules in the gas phase can
strike the surface of a liquid and
return to the liquid phase

In a closed container, an
equilibrium develops between
molecules evaporating and
condensing

Vapor Pressure

• Vapor Pressure is the pressure exerted by a
vapor in equilibrium with its liquid phase.

• Independent of the quantity of liquid or its
surface area

• Increases with increasing temperature

• Depends on the strength of attraction
between molecules in the liquid state

• Volatile liquids have very weak attractive
forces between molecules. Evaporate very
rapidly at ambient temperature. Have high
vapor pressures as a result

Measuring Vapor
Pressure of a Liquid

•Measure using a barometer

•Vapor from the liquid exerts a
force on the Hg and pushes the
column downward

•The difference in height
relative to vacuum provides
the vapor pressure for the
liquid

Surface Tension

•Resistance of a liquid to an increase in surface
area.

•Molecules on a liquid surface are strongly
attracted by molecules within the liquid.

•Surface tension increases with increasing attractive
interactions between molecules.

Capillary
Action

Capillary Action

is the spontaneous rise of a liquid in a
narrow tube

Cohesive forces exist between water molecules in a
liquid

Adhesive forces exist between water molecules and
the walls of the container.

When the cohesive forces between molecules are less
than the adhesive forces between liquid and
container, the liquid will move up the walls of the
container.

Capillary Action in
Action

• Shape of the meniscus reflects the relative strength of cohesive
forces within the liquid and adhesive forces between the liquid
and the tube.

Boiling Point

• Temperature at which the
vapor pressure of a liquid is
equal to the external
pressure above the liquid.

• The normal boiling point is
the boiling temperature
when the vapor pressure is
1 atm.

Freezing/Melting Point

• The freezing/melting point is
the temperature at which the
solid phase of a substance is
in equilibrium with its liquid
phase

• While both solid and liquid
phases are present, the
temperature remains
constant.

• The energy is used to change
the solid to the liquid phase.

Changes of State

•Heat of fusion is the energy
required to change 1 g of a solid
at its melting point to a liquid

• The heat of fusion for water
is 335 J/g.

• Use the heat of fusion as a
conversion factor

•Heat of vaporization is the
energy required to change 1 g of
liquid to vapor at its normal
boiling point.

• The heat of vaporization for
water is 2259 J/g.

• Use the heat of vaporization
as a conversion factor

Intermolecular
Forces

Dipole-Dipole Attractions

• In covalent molecules, due to different atoms having different
electronegativities, molecules are polar

• When polar molecules are put together, they will align to permit
interaction between oppositely polarized portions of the
molecules

Hydrogen Bonding

• A special type of dipole-dipole attraction
• One type of strong intermolecular force/attraction between

molecules
• To form hydrogen bonds, a compound must have covalent

bonds between hydrogen and F, O, or N.

London Dispersion Forces

• Interaction between nonpolar molecules and noble gases
• London forces arise from uneven, instantaneous charge

distributions due to electron movement in nonpolar molecules.

• Attractive forces
between molecules

• These forces allow
for formation of
liquids and solids

• The degree of
intermolecular
forces correlates
with a compound’s
physical properties.

Hydrates

•Hydrates are solids that contain water
molecules as part of their crystalline
structure

•The formula lists the anhydrous formula of
the compound followed by the number of
waters present per structural unit.

•Hydrates are named by placing a prefix
corresponding to the number of water
molecules. Followed by hydrate

•Hydrates will often decompose by losing
water upon heating

•To calculate % water in a hydrate:
• Calculate the molar mass of the

compound

• Calculate the %water of the
compound

Water: A Unique Liquid

• Physical properties of water
• Colorless, odorless, tasteless liquid
• More dense in liquid than solid phase
• High boiling point, high heat of

fusion/vaporization due to hydrogen bonding

• Structure of Water Molecules
• Two OH bonds are formed by the overlap of 1s

orbitals in the H with orbitals on the O

• The molecular geometry of water is bent, due to
the two lone pairs on oxygen

• Water has a permanent dipole due to the
molecules’ shape and the polar O-H bonds.

Osmosis – process by which water
flows through a membrane from a

region of more pure water to a region
of less pure water

• Reverse Osmosis – process by which water flows through a
membrane from a region of less pure water to a region of more
pure water, due to the presence of an external stimulus (typically
pressure)

Reading
Review

What are the three
changes in state?

What is vapor
pressure?

What are the three
types of

intermolecular
forces?

What molecular
shape does water

have?

How do you know
a compound is a
hydrate from its

formula?

• Slide 1

Liquid State of Matter
Changes in State
Vapor Pressure

• Measuring Vapor Pressure of a Liquid

Surface Tension
Capillary Action

• Capillary Action in Action

Boiling Point
Freezing/Melting Point
Changes of State

• Intermolecular Forces

Hydrates
Water: A Unique Liquid

• Slide 15
• Reading Review

CH1000
Fundament
als of
Chemistry
Module 4 – Chapter 15

Arrhenius Acids

Arrhenius Acid: An acid solution contains an excess of H+ ions.

Common Properties of Acids

1. Sour taste

2. Turns litmus paper pink

3. Reacts with:

Metals to produce H2 gas

Bases to yield water and a salt

Carbonates to give carbon dioxide

Arrhenius Bases

Arrhenius Bases: A basic solution contains an excess of OH– ions.

Common Properties of Bases

1. Bitter/caustic taste

2. Turns litmus paper blue

3. Slippery, soapy texture

4. Neutralizes acids

Brønsted-Lowry Acids
and Bases

Lewis Acid-Bases

Summary of
the Acid/Base

Theories

Reactions of Acids

Base Reactions

Bases can be amphoteric (act as either Brönsted acids or bases)

In general:

Zn(OH)2 (aq) + 2 HBr (aq) ZnBr2 (aq) + 2 H2O (l)

As a base:

NaOH and KOH can also react with metals.

2 NaOH (aq) + 2 Al (s) + 6 H2O (l) 2 NaAl(OH)4 (aq) +
3 H2 (g)

base + metal + water salt + hydrogen

Zn(OH)2 (aq) + 2 NaOH (aq) Na2Zn(OH)4 (aq)

As an acid:

Salts

Salts: products from acid-base reactions.

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

Salts are ionic compounds.

Salts contain a cation (a metal or ammonium ion) derived from
the base and an anion (excluding oxide or hydroxide ions)
derived from the acid.

Salts are generally crystalline compounds with high melting
and boiling points.

Electrolytes and
Nonelectrolytes

Electrolytes: compounds
that conduct electricity when
dissolved in water.

Nonelectrolytes:
substances that do not
conduct electricity when
dissolved in water.

Ion movement causes
conduction of electricity in
water.

3 classes of compounds,
acids, bases, and salts are
electrolytes because they
produce ions in water when
they dissolve.

Comparing Solution Conductivity

(Sugar solution) (Salt solution)(Distilled water)

Dissociation of Electrolytes

Salts dissociate into their respective cations and anions when
dissolved in water.

Hydrated sodium (purple) and chloride (green) ions

The negative end of the water dipole is attracted to the
positive Na+ ion.

When NaCl dissolves in water, each ion is surrounded by
several water molecules.

The permanent dipoles in the water molecules cause specific
alignment around the ions.

NaCl (s) Na

+ (aq)

+ Cl- (aq)

Electrolyte Ionization

Ionization: process of ion formation in solution. Ionization
results from the chemical reaction between a compound and
water.

Acids ionize in water, producing the hydronium ion (H3O+) and
a counter anion.

Bases ionize in water, producing the hydroxide ion (OH-) and a
counter cation.

HCl (g) + H2O (l) H3O
+ (aq) + Cl- (aq)

H3PO4 (aq) + H2O (l) H2PO4
– (aq) + H3O

+ (aq)

NH3 (aq) + H2O (l) OH
– (aq) + NH4

+ (aq)

Strong and Weak
Electrolytes

Strong electrolytes:
undergo complete
ionization in water.
Example: HCl (strong
acid)

Weak electrolytes:
undergo incomplete
ionization in water.
Example: CH3COOH
(weak acid)

HCl (left) is 100% ionized.
CH3COOH exists primarily
in the unionized form.

HF (aq) + H2O (l) F
– (aq) + H3O

+ (aq)

Double arrows indicate incomplete
ionization

(typically weak electrolytes).

Salts

Salts can dissociate into more than 2 ions, depending upon the
compound.

A 1 M solution of NaCl produces a total of 2 M of ions.

A 1 M solution of CaCl2 produces a total of 3 M of ions.

NaCl (s) Na+ (aq) + Cl- (aq)

1M 1M 1M

CaCl2 (s) Ca
2+ (aq) + 2 Cl- (aq)

1M 1M 2M

Colligative Properties of
Electrolyte Solutions

Colligative properties: depend only on the number of moles of
dissolved particles present.

This must be taken into consideration when calculating freezing
point depression or boiling point elevation due to the presence
of solute particles.

Example: What is the boiling point elevation of a 1.5 m
aqueous solution of CaCl2? (Kb for water is 0.512 ºC/m).

Because CaCl2 is a strong electrolyte, 3 mol of ions (1 mol
Ca2+ and 2 mol Cl- ions) will be present in the solution.

ΔTb = 1.5 m
CaCl2

= 2.3 ºC

×

3 mol ions
1 mol CaCl2

0.512 ºC
1 m

×

Autoionization of Water

Pure water auto(self) ionizes according to the reaction:

Based on the reaction stoichiometry:

Concentration H3O+ = Concentration OH– = 1 x 10–7 M

[H3O+] x [OH–] = (1 x 10–7)2 = 1 x 10–14

When acid or base is present in water, [H3O+] and [OH-]
change.

In acidic solutions, [H3O+] > [OH–].

In basic solutions, [H3O+] < [OH–].

H2O (l) + H2O (l) H3O
+ (aq) + OH– (aq)

Introduction to

pH

The pH scale

Increasing acidity Increasing basicityHigh H3O

+

Low OH-
Low H3O

+

High OH-

In pure water, [H3O
+] = 1 x 10-7 M, so

pH = – log(1 x 10-7) = 7

pH = – log[H3O
+]

pH
Calculations

pH = – log[H3O
+]

[H3O
+] = 1 x 10-5 M

[H3O
+] = 2 x 10-5 M

If exactly 1

Exponent = pH
pH = 5

If a number between 1 and 10

The pH is between
the exponent and
next lowest whole

number
pH = 4.7

Generalizations

[H3O
+] = 10-pH

pH

Neutralization

General Reaction

Example

Overall Ionic Equation:
H+(aq) + Cl- (aq) + Na+ (aq) + OH- (aq) Na+ (aq) + Cl- (aq) + H2O
(l)

All species are included; soluble compounds shown as ions.

Net Ionic Equation:

Spectator ions (orange) are removed from both sides.

acid + base salt +
water

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

H+ (aq) + OH- (aq)
H2O (l)

Titration

Titration: experiment
where the volume of one
reagent (titrant) required
to react with a measured
amount of another
reagent is measured.

Titrations allow the
amount of an acid or
base present in a sample
to be determined.

Indicators are used to signal the endpoint of a
titration,

the point when enough titrant is added to react
with the acid/base present.

Burets deliver measured amounts of the titrant
into a solution of the unknown reagent.

Endpoin
t

Net Ionic Equations

Rules for Writing Net Ionic Equations

1. Strong electrolytes are written as the corresponding ions.
Example: NaOH (aq) is written as Na+(aq) + OH-(aq)

2. Weak electrolytes and nonelectrolytes are written as
molecules. Example: CH3OH(aq), CH3COOH(aq), etc.

3. Solids and gases are written as their molecular forms.

4. The net ionic equation does not include spectator ions.

5. The net ionic equation must balance atoms and charge.

Acid Rain

1. Emission of nitrogen or sulfur
oxides.

2. Transportation of these chemicals
throughout the atmosphere.

3. Chemical reaction of the oxides
with water.

4. This forms sulfuric and nitric acids.

5. Precipitation carries the acids to
the ground.

General Process for Acid Rain
Formation:

Acid rain: atmospheric
precipitation more acidic

than typical.

Reading Review

1. What ion is present in an acid? What ion
is present in a base?

2. What products are produced in a
reaction between an acid and a metal
oxide?

3. What type of reaction produces salts?

4. What is the pH range for acids?

5. What is titration?

• Slide 1

Arrhenius Acids
Arrhenius Bases

• Brønsted-Lowry Acids and Bases
• Summary of the Acid/Base Theories

Reactions of Acids
Base Reactions
Salts

• Electrolytes and Nonelectrolytes

Dissociation of Electrolytes
Electrolyte Ionization

• Strong and Weak Electrolytes

Salts

• Colligative Properties of Electrolyte Solutions

Autoionization of Water

• Introduction to pH
• pH Calculations

pH
Neutralization
Titration
Net Ionic Equations
Acid Rain
Reading Review