Transcript Chapter 14 LIQUIDS AND SOLIDS

Chapter 14 “LIQUIDS AND SOLIDS”
What are the properties of the
‘condensed states’ of matter?
T H Witherup 02/06
Honors (rev 07)
Ch. 14 OBJECTIVES

Show how the Kinetic-Molecular (K-M) Theory accounts
for the physical properties of liquids & solids.
 Describe different types of intermolecular forces, and
how they affect properties of liquids & solids.
 Learn about viscosity & surface tension, and explain
their relationship to intermolecular forces.
 Compare crystalline & amorphous substances.
 Relate the structure & bonding in the four types of
crystalline solids to the properties they exhibit.
 Describe the changes of state (vaporization,
condensation, boiling, sublimation, deposition, melting,
freezing).
 Learn how to interpret “Heating/Cooling Curves” and
“Phase Diagrams.”
14-1 Condensed States of Matter:
Liquids and Solids

Condensed matter has much higher density
(mass/volume) than gases.
 Unlike gases, particles of condensed matter
experience different amounts and types of
attractive forces.
 Kinetic-Molecular Theory can help explain the
properties of condensed matter.

The state of a substance at room temperature
depends on the attractive forces between its particles.
Comparing the States of Matter

Gas




Liquid




Total disorder
Particles free to move past each other
Particles far apart
Disorder
Particles free to move past each other
Particles close together
Solid



Ordered arrangement
Particles vibrate, but remain in a fixed position
Particles close together
-
Comparing Properties of Gases, Liquids
& Solids
PROPERTY
COMPRESSIBILITY
DENSITY
VOLUME
SHAPE
DIFFUSION
EXPANSION
(with Heat)
GAS
LIQUID
SOLID
Comparing Properties of Gases, Liquids & Solids
PROPERTY
GAS
LIQUID
SOLID
COMPRESSIBILITY
DENSITY
High
Slight
Low
VOLUME
SHAPE
DIFFUSION
EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids
PROPERTY
GAS
LIQUID
SOLID
COMPRESSIBILITY
DENSITY
High
Slight
Low
Low
High
High
VOLUME
SHAPE
DIFFUSION
EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids
PROPERTY
GAS
LIQUID
SOLID
COMPRESSIBILITY
DENSITY
High
Slight
Low
Low
High
High
Fills
container
Definite
Rigid
VOLUME
SHAPE
DIFFUSION
EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids
PROPERTY
GAS
LIQUID
SOLID
COMPRESSIBILITY
DENSITY
High
Slight
Low
Low
High
High
Fills
container
Of container
Definite
Rigid
Of container
Fixed
VOLUME
SHAPE
DIFFUSION
EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids
PROPERTY
GAS
LIQUID
SOLID
COMPRESSIBILITY
DENSITY
High
Slight
Low
Low
High
High
Fills
container
Of container
Definite
Rigid
Of container
Fixed
Rapid
Slow
Very slow (@
surface)
VOLUME
SHAPE
DIFFUSION
EXPANSION
(with Heat)
Comparing Properties of Gases, Liquids & Solids
PROPERTY
GAS
LIQUID
SOLID
COMPRESSIBILITY
DENSITY
High
Slight
Low
Low
High
High
Fills
container
Of container
Definite
Rigid
Of container
Fixed
DIFFUSION
Rapid
Slow
EXPANSION
(with Heat)
High
Low
Very slow (@
surface)
Low
VOLUME
SHAPE
Review of Chemical Bonding

Ionic: transfer of electrons from a metal to a
nonmetal.


Metallic: sharing of valence electrons of the
metal atoms.


Resulting ions have opposite charge and are
attracted to each other, forming an ionic bond.
Results in a network of positive ions in a “sea of
electrons.”
Covalent: strong intramolecular forces from the
sharing of valence electrons between atoms.


Results in individual molecules with specific shapes.
Intermolecular forces exist between molecules.
Intermolecular Forces have…
 …a
wide range of strengths.
 They are much weaker than ionic,
covalent and metallic bonds.
 Basic types of Intermolecular Forces:

Dispersion Forces = attraction between
temporary ‘induced dipoles.’
• Consider noble gas boiling points and Fig. 14-8.
Dispersion Forces: Boiling Points of Noble
Gases (°C)
Boiling Point (deg C)
0
-50
0
20
40
100
-107 Xe
-150
-250
80
-67 Rn
-100
-200
60
-153 Kr
-186 Ar
-246 Ne
-269 He
-300
Atomic Number
What causes the boiling point of helium to be so low and that of radon to be
so high? Consider the effects shown in Figure 14-8 of text.
(Data from Figure 14-7, page 462 of your textbook.)
Intermolecular Forces have…
 …a
wide range of strengths.
 They are much weaker than ionic,
covalent and metallic bonds.
 Basic types of Intermolecular Forces:

Dispersion Forces = attraction between
temporary ‘induced dipoles.’
• Consider noble gas boiling points and Fig. 14-8.

Dipole-Dipole = attraction between polar
molecules (dipoles) having permanent charge
separation.
• Consider HCl, HBr, CO2 (linear) and SO2 (bent).
Dipole-Dipole Forces: Polar Molecules
Consider HCl (a dipole or polar molecule):
δ+
δ-
H
Cl
δ+
:::::::::: H
::::::::::
δCl
Compare CO2 and SO2 (covalent molecules with polar bonds):
Polar or Nonpolar? Why?
+
+
+ +
B.P. = -78°C
O
C
O
B.P. = -10°C
S
O
O
Intermolecular Forces have…
…a wide range of strengths.
 They are much weaker than ionic, covalent and
metallic bonds.
 Basic types of Intermolecular Forces:


Dispersion Forces = attraction between temporary
‘induced dipoles.’
• Consider noble gas boiling points and Fig. 14-8.

Dipole-Dipole = attraction between polar molecules
(dipoles) having permanent charge separation.
• Consider HCl, HBr, CO2 (linear) and SO2 (bent).

Hydrogen Bonds = strong attraction between H atom
of a molecule and a very electronegative atom (F,O,N)
of another molecule.
• See data on the next slides.
Consider Boiling Points

As mass increases, we expect Boiling Point to rise.
 CH4, SiH4, GeH4, SnH4


NH3, PH3, AsH3, SbH3


Observed: HCl < HBr < HI << HF (anomaly)
H2O, H2S, H2Se, H2Te


Observed: PH3 < AsH3 < SbH3 ~= NH3 (anomaly)
HF, HCl, HBr, HI


Observed: CH4 < SiH4 < GeH4 < SnH4 (follow trend)
Observed: H2S < H2Se < H2Te << H2O (anomaly)
See Table on next slide.
Consider Boiling Points (kelvins)
Formula
Period↓
Period 2
Period 3
Period 4
Period 5
XH4
XH3
XH2
XH
CH4
(93)
SiH4
(163)
NH3
(243)
PH3
(183)
H2O
(373)
H2S
(213)
HF
(303)
HCl
(193)
GeH4
(193)
SnH4
(223)
AsH3
(203)
SbH3
(243)
H2Se
(243)
H2Te
(273)
HBr
(213)
HI
(253)
What causes these anomalies (shown in yellow)? HYDROGEN BONDING!
Data is plotted on page 466 of text.
Hydrogen Bonding Forces
Involves F-H, N-H or O-H bonds.
How different are the electronegativity values of these atom pairs?
What does this do to the bond polarity?
Consider water. As you will soon see, many of water’s unusual
Properties result from “hydrogen bonding”!
*******
*******
Comparison of Intramolecular & Intermolecular Forces
FORCE
ATTRACTION
ENERGY,
kJ/mol
EXAMPLE
Ionic
Anion-cation
400 – 4000
NaCl
Covalent
Shared Electrons
150 - 1100
Cl-Cl
Metallic
Cations in “Sea of
Electrons”
75 - 1000
Cu
Ion-Dipole
Ion with Dipole
40 – 600
Cl-…H2O
Dipole-Dipole
Dipole charges
5 – 25
Br-Cl…Br-Cl
H-Bond
H to N, O, F
10 - 40
H2O to H2O
Ion-induced
dipole
Ion & e- cloud of
neighbor
3 – 15
Fe2+…O2
Dipole-induced Dipole charge & edipole
cloud of neighbor
2 – 10
H-Br…Br2
Dispersion
(London)
0.05 – 40
Cl-Cl…Cl-Cl
Electron clouds of
neighbors
Summary: Intermolecular Forces have…
…a wide range of strengths.
 They are much weaker than ionic, covalent and
metallic bonds.
 Basic types of Intermolecular Forces:




Dispersion Forces = attraction between temporary
‘induced dipoles.’
Dipole-Dipole = attraction between polar molecules
(dipoles) having permanent charge separation.
Hydrogen Bonds = strong attraction between H atom
of a molecule and a very electronegative atom
(F,O,N) of another molecule.
14-2 Properties of Liquids


Intermolecular forces generally determine the physical
properties of liquids, such as...
Viscosity = resistance to motion between molecules of a
liquid as they move past each other.


Surface Tension = unbalanced attractive forces at the
surface of a liquid that causes the surface to act like a
film.


Water bugs, and the paperclip experiment!
Capillarity = the tendency of a liquid to flow through a
small opening or tube.


Syrup, motor oil, hot fudge!
“Wicking” of fabric when it gets wet.
Density = mass/volume.

Wide range, from <1g/mL (butane) to >13.6g/mL (mercury).
Video Clip (Water)
Comparing Small Molecules
Compound
Formula
Mass (u)
State @ 25C MP*
BP*
Methane
CH4
~16
Gas
90
112
Ammonia
NH3
~17
Gas
195
240
Water
H2O
~18
Liquid
273
373
Nitrogen
N2
~28
Gas
63
77
Oxygen
O2
~32
Gas
55
90
*MP = Melting Point; BP = Boiling Point (kelvins).
Water: A very special compound
The most abundant substance on Earth’s surface.
(The “Blue Planet.”)
 Critical to life forms as we know them.
 Some unusual properties of water are:









Unexpectedly high boiling point for its size.
Unusually high specific heat.
Solid form (ice) has lower density than liquid.
High surface tension.
High of vaporization.
Excellent solvent (“universal solvent”).
Why? HYDROGEN BONDING!
What are some consequences of these properties?
14-3 The Nature of Solids

Crystalline Solids






Highly ordered, repeating arrangement of particles.
Ionic (NaCl), covalent (sugar) or metallic (Fe).
Characterized by specific ‘unit cells.’ (Fig. 14-20)
Generally have sharp melting points.
Fracture occurs along definite planes when stressed.
Amorphous Solids (“supercooled liquids”)





Highly disordered, random arrangement of particles.
Wax, plastics (PET, PE, PS, Nylon®).
Characterized by lack of organized structure.
Generally soften over a wide temperature range.
Fracture occurs randomly with stress.
Bonding in Solids
 Metallic

Solids
All metallic elements.
 Molecular

Most ‘organic’ compounds (contain carbon) &
many ‘inorganic’ compounds (CO2, H2O, SO2).
 Ionic

Solids
Solids
Typical salts (NaCl, KBr, CaCl2)
 Covalent-Network

Solids
Diamond, graphite, silicon.
See the link: http://undergrad-ed.chemistry.ohio-state.edu/chemapplets/
Crystals/ClosestPackedStructures.html
Properties of Crystalline Solids (Fig 14-23)
Type
Particles
Forces Between
Particles
Properties
Metallic
Atoms
Metallic bond
Soft to hard; variable
melting points; good
conductivity; malleable;
ductile.
Molecular Atoms or H-bond, dipolemolecules dipole, dispersion
Soft; variable melting
points; poor conductors.
Ionic
Electrostatic
Hard; brittle; high melting
points; poor conductors.
Covalent bonds
Very hard; very high
melting points; poor
conductivity
Ions
Covalent- Atoms
network
14-4 Changes of State (Phase Changes)

Phase Change – conversion of a substance
from one physical state of matter to another.

Six types of phase changes to consider:






Melting (solid to liquid)
Freezing (liquid to solid)
Vaporization (evaporation, boiling) (liquid to gas)
Condensation (gas to liquid)
Sublimation (solid to gas)
Deposition (gas to solid)
Relative Potential Energy (P.E.) Of Matter
During Phase Changes (from Fig. 14-30)
DEPOSITION
SUBLIMATION
CONDENSATION
VAPORIZATION
FREEZING
MELTING
POTENTIAL ENERGY
GAS
LIQUID
SOLID
Energy Changes & Phase Changes

To change a substance from one state of matter
to another always involves a gain or loss of
energy.



Melting a solid requires energy input, but freezing a
liquid removes energy from the liquid.
Boiling a liquid requires energy input, but condensing
a gas removes energy from the gas.
Subliming a solid requires energy input, but
deposition of a gas removes energy from the gas.
Vaporization & Condensation

According to K-M Theory, temperature is a measure of the
average kinetic energy of the particles of a substance.

Fast molecules near the surface of a liquid may escape
the liquid (evaporation).


This results in some vapor in the space above a liquid.
Volatile liquids have a tendency to easily evaporate due to the poor
attraction among their molecules, resulting in high concentrations
of vapor around the liquid.
• Gasoline, for example, poses high hazards because of this property.

What happens as fast molecules leave the liquid?



The remaining molecules have lower average kinetic energy,
which is observed by the lower temperature of the liquid.
Evaporative cooling, sweating, wind chill are examples.
Liquid-Vapor Equilibrium occurs in closed vessels.


Rates of evaporation and condensation are the same.
Dynamic equilibrium (rates of opposing processes are equal).
Kinetic Energy Distribution
http://undergrad-ed.chemistry.ohio-state.edu/chemapplets/
KineticMolecularTheory/Maxwell.htmlKinetic Energy Distribution
In Chapter 13 we looked at two simulations at this website:
“Gases – Kinetic Molecular Theory.” (Maxwell distribution.)
“Maxwell distribution of a gas (or liquid) at different temperatures.”
If you need to refresh your memory, revisit the website to look
at the Maxwell distribution shapes.
The general shapes of the curves apply to liquids as well.
VAPOR PRESSURE

Not all molecules in a sample of a substance have
the same amount of kinetic energy.





Some molecules are moving faster than others, so there
is a distribution of kinetic energy.
The average kinetic energy of the substance is
proportional to its temperature.
Molecules that move very fast may achieve ‘escape
velocity’ and leave the liquid entirely. (They evaporate.)
This tendency of a liquid to become a gas at a given
temperature is its ‘vapor pressure,’ and all liquids exert a
vapor pressure against the container (or atmosphere).
Vapor pressure increases with temperature. (Why?)
•
•
•
•
Consider what happens to a puddle of water.
Does the water evaporate?
What if the water is in a closed jar? What would you see?
This is an example of “liquid-vapor equilibrium.”
HEAT OF VAPORIZATION is…

…the amount of heat needed to vaporize a
given amount of liquid at its boiling point.


Energy is added to the liquid, but the
temperature does not change.



Molar heat of vaporization of water = 40.7 kJ/mol or
540. cal/mol.
The added energy is expended to overcome the
intermolecular attractions of molecules of the liquid,
converting the substance from liquid to gas.
Liquids with strong intermolecular attractions (such as
water) have high heats of vaporization.
Condensation is the opposite of evaporations.


Heat of Condensation = -(Heat of Vaporization)
Why is a burn from steam more damaging to tissue
than that from boiling water?
“IN CLASS” ACTIVITY


Explore Heating & Cooling Curves
Explore Phase Diagrams

To do these, get one of the laptops from the cart,
and go to the websites shown on the following
slides.





Read the information.
Perform the simulations as directed.
Take notes as appropriate.
Record your observations.
Answer the questions.
Heating/Cooling Curves
“Phase Changes” Animation and Activity:
http://www.chm.davidson.edu/chemistryapplets/Phase
Changes/HeatingCurve.html
In class, go to this website, read the information,
and perform the Heating Curve exercise.
Answer all questions.
Also,
Doc Brown’s website has a nice summary of the states
of matter and “heating/cooling” curves. Check it out!
http://www.wpbschoolhouse.btinternet.co.uk/page03/3_52states.htm
HEATING/COOLING CURVE SUMMARY

Plots time (x-axis) vs. Temperature (y-axis).
 Identical in shape, but opposite in direction.
 Ramps indicate temperature changes.
 Plateaus indicate no change in temperature, but show
absorption/release of energy.



Two phases co-exist where there is a plateau.
Length of plateau is proportional to the energy change that occurs
during a phase change.
During a phase change, all of the energy is used to induce
the change, so the temperature of the mixture of phases
stays constant.


For example, water on a stove will continue to boil at the same
temperature until all of the water has been converted to a gas.
Similarly, a mixture of ice and water will remain at zero C until all of
the ice melts, after which the liquid begins to heat.
Phase Diagrams
In class, go to this animation to observe what happens
to phases as temperature and pressure change:
http://www.chm.davidson.edu/ChemistryApplets/PhaseChanges/
PhaseDiagram.html
At this website (in class) you are to explore all of the following applets:
Phase Diagram
Phase Diagram Part 1, Part 2, Part 3, Part 4, Part 5.
Read the information, perform the simulations and answer all questions.
To learn more about phase diagrams, check out this website:
http://www.chemistrycoach.com/Phase_diagram.htm
PHASE DIAGRAM SUMMARY

A plot of pressure and temperature is
determined experimentally to show the changes
in phase that occur under various conditions.
 Phases in equilibrium may be identified from the
Phase Diagram.
 Triple Point and Critical Point may be identified.



Triple Point – The temperature and pressure at
which all three phases exist in equilibrium.
Critical Point – The temperature and pressure
beyond which the substance can only exist as a gas.
http://www.activotec.com/graphics/products/sc
fexpl.gif
Ch. 14 OBJECTIVES

Show how the Kinetic-Molecular (K-M) Theory accounts
for the physical properties of liquids & solids.
 Describe different types of intermolecular forces, and
how they affect properties of liquids & solids.
 Learn about viscosity & surface tension, and explain
their relationship to intermolecular forces.
 Compare crystalline & amorphous substances.
 Relate the structure & bonding in the four types of
crystalline solids to the properties they exhibit.
 Describe the changes of state (vaporization,
condensation, boiling, sublimation, deposition, melting,
freezing).
 Learn how to interpret “Heating/Cooling Curves” and
“Phase Diagrams.”