Introductory Chemistry: Concepts & Connections 4th Edition

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Transcript Introductory Chemistry: Concepts & Connections 4th Edition

Introductory Chemistry:
Concepts & Connections
4th Edition by Charles H. Corwin
Chapter 13
and Solids
Christopher G. Hamaker, Illinois State University, Normal IL
© 2005, Prentice Hall
Properties of Liquids
• Unlike gases, liquids do not respond dramatically
to temperature and pressure changes.
• We can study the liquid state and make 5 general
1. Liquids have a variable shape, but a fixed
• Liquids take the shape of their container
2. Liquids usually flow readily.
• However, not all liquids flow at the same rate.
Chapter 13
Properties of Liquids Continued
3. Liquids do not compress or expand significantly
• The volume of a liquid varies very little as the
temperature and pressure change.
4. Liquids have a high density compared to gases.
• Liquids are about 1000 times more dense than gases.
5. Liquids that are soluble mix homogeneously.
• Liquids diffuse more slowly than gases but
eventually will form a homogeneous mixture.
Chapter 13
Intermolecular Bond Concept
• An intermolecular bond is an attraction between
molecules, whereas intra molecular bonds are
between atoms in a molecule.
• Some properties of liquids, such as vapor pressure,
viscosity, and surface tension, are determined by
the strength of attraction between molecules.
• Intermolecular bonds are much weaker than
intramolecular bonds.
Chapter 13
Intermolecular Bonds
• Recall, that a polar molecule has positive and
negative charges concentrated in different regions
due to unequal sharing of electrons in bonds.
• This uneven distribution of electrons in a
molecule is called a dipole.
• Intermolecular attractions result from temporary
or permanent dipoles in molecules.
• There are three intermolecular forces: dispersion
forces, dipole forces, and hydrogen bonds.
Chapter 13
Dispersion Forces
• Dispersion forces, or London forces, are the result
of a temporary dipole.
• Electrons are constantly shifting and a region may
become temporarily electron poor and slightly
positive while another region becomes slightly
• This creates a temporary dipole
and two molecules with
temporary dipoles are attracted
to each other.
Chapter 13
Dispersion Forces Continued
• Dispersion forces are the weakest intermolecular
• Dispersion forces are present in all molecules.
• The strength of the dispersion forces in a molecule
is related to the number of electrons in the
– The more electrons in a molecule, the stronger the
dispersion forces.
Chapter 13
Dipole Forces
• Polar molecules have a permanent dipole.
• The oppositely charged ends of polar molecules
are attracted to each other, this is the dipole force.
• The strength of a dipole force is typically 10% of
a covalent bond’s strength.
• Dipole forces are stronger
than dispersion forces.
Chapter 13
Hydrogen Bonds
• Hydrogen bonds are a special type of dipole
• Hydrogen bonds are present when a molecule has
an N-H, O-H, or F-H bond.
• Hydrogen bonds are especially important in water
and living organisms.
Chapter 13
Physical Properties of Liquids
• There are 4 physical properties of liquids that we
can relate to the intermolecular attractions present
in the molecules:
– Vapor Pressure
– Boiling Point
– Viscosity
– Surface tension
Chapter 13
Vapor Pressure
• At the surface of a liquid, some molecules gain
enough energy to escape the intermolecular
attractions of neighboring molecules and enter the
vapor state. This is evaporation.
• The reverse process is condensation.
• When the rates of evaporation and condensation
are equal, the pressure exerted by the gas
molecules above a liquid is called the vapor
Chapter 13
Vapor Pressure Continued
• The stronger the intermolecular forces between
the molecules in the liquid, the less molecules
escape into the gas phase.
• As the attractive force between molecules
increases, vapor pressure decreases.
Chapter 13
Vapor Pressure Comparison
• Lets compare water and ether.
– Water has strong intermolecular attractions and ether
has only weak intermolecular attractions.
• At 0C, neither has a significant vapor pressure.
• At 35C, ether has a significant vapor
pressure and water does not.
Chapter 13
Vapor Pressure vs. Temperature
• As the temperature
increases, the vapor
pressure of a liquid
• Again, the stronger
the intermolecular
attractions, the
lower the vapor
pressure at a given
Chapter 13
Boiling Point
• The normal boiling point of a substance is the
temperature where the vapor pressure is equal to
the standard atmospheric pressure.
• As we saw in the previous graph, the stronger the
intermolecular attractions, the higher the boiling
point of the liquid.
• A liquid with a high boiling point has a low vapor
Chapter 13
• The viscosity of a liquid is a liquid’s resistance to
• Viscosity is the result of an attraction between
• The stronger the intermolecular forces, the higher
the viscosity.
Chapter 13
Surface Tension
• The attraction between molecules at the surface of
a liquid is called surface tension.
• For an object to sink in a liquid, it must first break
through the surface.
• The stronger the intermolecular attractions, the
stronger the surface tension of a liquid.
Chapter 13
Properties of Solids
• Unlike gases, solids do not respond dramatically
to temperature and pressure changes.
• We can study the solid state and make 5 general
1. Solids have a fixed shape and volume.
• Unlike liquids, solids are rigid.
2. Solids are either crystalline or noncrystalline.
• Crystalline solids contain particles in a regular,
repeating pattern.
Chapter 13
Properties of Solids Continued
3. Solids do not compress or expand to any degree
• Assuming no change in physical state, temperature
and pressure have a negligible effect on the volume
of a solid.
4. Solids have a slightly higher density than their
corresponding liquid
• One important exception is water; ice is less dense
than liquid water.
5. Solids do not mix by diffusion
• The particles are not free to diffuse in a solid
heterogeneous mixture.
Chapter 13
Crystalline Solids
• There are three types of crystalline solids
examples of which are shown below:
– Ionic solids like NaCl, (a)
– Molecular solids like H2O, (b)
– Metallic solids like Cu, (c)
Chapter 13
Ionic Solids
• A crystalline ionic solid is
composed of positive and
negative ions arranged in a
regular, repeating pattern.
• In table salt, NaCl, sodium
ions and chloride ions are
arranged in a regular threedimensional structure
referred to as a crystal lattice.
• Other ionic compounds will have different crystal
Chapter 13
Molecular Solids
• A crystalline molecular
solid has molecules
arranged in a particular
• In water, H2O, the
molecules are arranged
in a regular threedimensional structure.
• Other examples of crystalline molecular solids are
table sugar, C12H22O11, and sulfur, S8.
Chapter 13
Metallic Solids
• A crystalline metallic solid is
composed of metal atoms
arranged in a definite pattern.
• A metallic crystal is made up
of positive metal ions
surrounded by valance electrons.
• Metals are good conductors of
electricity because electrons are free
to move about the crystal.
• This is referred to as the “electron sea” model.
Chapter 13
General Properties of Solids
Chapter 13
Changes of Physical State
• Heat is necessary to raise the temperature and
change the physical state of a substance.
• Specific heat is the amount of heat required to
raise 1.00 g of a substance 1C.
• Water is the reference and its specific heat is
1.00 cal/(g×C).
• The specific heats of ice and steam are about half
that of liquid water.
Chapter 13
Solid/Liquid Phase Changes
• As a solid melts, the temperature is constant until
all of the solid is changed to liquid.
• The amount of heat required to melt 1.00 g of
substance is called the heat of fusion (Hfusion).
For water, the heat of fusion is 80.0 cal/g.
• When a liquid changes to a solid, the heat change
is the heat of solidification (Hsolid).
• The value of Hfusion is the same as the value of
Chapter 13
Liquid/Gas Phase Changes
• As a liquid vaporizes, the temperature is constant
until all of the liquid is changed to gas.
• The amount of heat required to vaporize 1.00 g of
substance is called the heat of vaporization
(Hvapor). For water, the value is 540 cal/g.
• When a gas changes to a liquid, the heat change is
the heat of condensation (Hcond).
• The value of Hvapor is the same as the value of
Chapter 13
Temperature/Energy Graphs
• We can graph the amount of energy required to
change the temperature and physical state of a
• The heating curve
for water is shown
• As energy is added
the temperature
increases and
changes the
physical states.
Chapter 13
Energy from Heating Curves
• We can use the heating curve and heat values for
water to calculate how much energy is required to
change the temperature of a sample of water.
• The amount of energy required to raise the
temperature of a substance is calculated using the
following formula:
heat = (specific heat) × (change in temperature) × (mass of sample)
• The amount of energy required to change the state
of a substance is calculated as follows:
heat = (Hxxx) × (mass)
Chapter 13
Energy Calculation
• How much energy is required to raise 25.5 g of
ice at –5.0C to steam at 100.0C?
• Looking at the heating curve for water, there are
4 regions:
1) Heating of solid ice from –5.0C to 0.0C
2) Melting of ice at 0.0C
3) Heating of liquid water from 0.0C to 100.0C
4) Vaporization of water at 100.0C
Chapter 13
Energy Calculation
• The total energy is the sum of the energy in steps
1) through 4). Calculate the energy for each step.
1) (25.5 g) × [0.0 – (–5.0)]C × (0.50 cal/g×C) = 64 cal
2) (25.5 g) × (80.0 cal/g) = 2040 cal
3) (25.5 g) × [100.0 – 0.0)]C × (1.00 cal/g×C) = 2550 cal
4) (25.5 g) × (540 cal/g) = 13,800 cal
• The total energy is:
64 cal + 2040 cal + 2550 cal + 13800 cal =
18,500 cal
Chapter 13
Structure of Water
• Lets start with the electron dot formula for water:
• Water has a bent molecular shape and the bond
angle is 104.5.
• Water is a polar
molecule that exhibits
strong hydrogen
Properties of Ice
• Water is a one of the few substances that is less
dense as a solid than as a liquid.
• As water freezes, the hydrogen bonds organize the
water molecules into a three-dimensional structure
where the molecules are farther apart in the liquid.
• Liquid water has a
density of 1.00 g/mL
while ice has a density
of 0.917 g/mL.
Chapter 13
Physical Properties of Water
• Water has unusually melting and boiling points,
especially compared to the other hydrogen
compounds of Group IVA/16.
• This is due to
hydrogen bonding
which is present in
water but not present
in H2S, H2Se, or
Chapter 13
Chemical Properties of Water
• Water can undergo an electrolysis reaction to
produce hydrogen and oxygen:
2 H2O(l)
2 H2(g) + O2(g)
• Water reacts with active metals to produce
hydrogen and a metal hydroxide:
2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g)
• Water reacts with metal oxides to produce a base:
CaO(s) + H2O(l) → Ca(OH)2(aq)
• Water reacts with nonmetal oxides to produce an
CO2(g) + H2O(l) → H2CO3(aq)
• A hydrate is a crystalline ionic compound that
contains water:
• The dot indicates that water molecules are bonded
directly to each unit of hydrate.
• Heating a hydrate drives off the water and
produces an anhydrous compound (without
CuSO45H2O(s) → CuSO4(s) + 5 H2O(l)
Chapter 13
• There are three types of intermolecular bonds:
– Dispersion forces
– Dipole forces
– Hydrogen bonds
• Dispersion forces are the weakest and hydrogen
bonds are the strongest.
• These intermolecular attractions affect the
physical properties of substances.
Chapter 13
Conclusions Continued
• There are 5 properties of liquids that are affected
by intermolecular bonds:
– Vapor pressure decreases as intermolecular forces
– Boiling point increases as intermolecular forces
– Viscosity increases as intermolecular forces increase
– Surface tension increases as intermolecular forces
Chapter 13
Conclusions Continued
• There are 3 types of crystalline solids:
– Ionic solids
– Molecular solids
– Metallic solids
• Energy is required to raise the temperature of a
• Water displays many unique properties due to the
presence of strong hydrogen bonds.
Chapter 13