The document discusses liquids and solids from the perspective of the kinetic molecular theory. It explains that in liquids, particles are more closely packed than gases due to intermolecular forces, but are still able to flow freely unlike solids where particle motion is limited to vibration. Key differences between liquids and solids include density, compressibility, diffusion rates, and types of molecular structures (crystalline vs amorphous). The document also covers concepts such as surface tension, intermolecular forces including hydrogen bonding and London dispersion forces, and how force strength influences melting and boiling points.

1. Liquids and SolidsLiquids and Solids
Quick brainstorm
What are some basic differences
between gases and liquids?
between liquids and solids?

2. A look back at the three statesA look back at the three states
Particle arrangement is the key!

3. The Kinetic Molecular TheoryThe Kinetic Molecular Theory
The purpose of the KMT is to help
explain particle behavior in the three
states of matter.
We will re-visit this theory when we
look at gases which is the next unit.
Now: Let’s look at liquids and solids.

4. Liquids and the KMTLiquids and the KMT
• Liquids have a higher density than gases due
to the closer arrangement of the particles.
• Liquids are much less compressible than
gases.
• Particles are not bound to fixed positions,
rather they move constantly.
This motion is fluidity. Fluidity depends on forces
and temperature.
viscosity- resistance to flow

5. Liquids and the KMTLiquids and the KMT
• The speed of motion of liquid particles
depends on the temperature.
Liquids under extremely cold temps. Would be
very viscous- resist flow.
• Liquids diffuse and mix with other liquids.
Speed of diffusion is governed by temperature of
the liquid and the types of forces between the
liquid particles.
• Liquids exhibit surface tension and capillary
action.

6. Surface TensionSurface Tension
The forces felt
by particles on
the surface are
primarly down
and to the sides.
The surface
tension is
directly related
to strength of
forces between
particles.

7. Solids and the KMTSolids and the KMT
• Particles are closely packed, forces play
a strong role.
• Particle motion is limited to vibrations
primarily due to the strength of particle
attractions.
• Most substances are most dense as a
solid.

8. Solids and the KMTSolids and the KMT
• Solids are less compressible than liquids
due to close particle arrangement.
• Diffusion does occur, but the rate is
millions of times slower than in liquids.

9. Two types of solid structuresTwo types of solid structures
crystalline solids- ordered particle
arrangement.
amorphous solids- random
arrangement of particles.

11. Ionic CrystalsIonic Crystals
Attractive
forces between
cations and
anions.
Very strong
attractive
forces.

12. Covalent network crystalsCovalent network crystals
Each atom is
covalently
bonded to a
neighboring
atom.
Examples are
carbon in the
form of
diamond,
graphite,
nanotube and
buckyball
These are all
Allotropes

13. Metallic CrystalsMetallic Crystals
Metal atoms
surrounded
by a sea of
electrons

14. The Forces between particlesThe Forces between particles
We have been focusing on the fact that
particles in a liquid and solid are closer
together, and must experience some
form of attraction (strong or weak) to
neighboring particles.
Now, we will define (describe) those
forces.

15. List 2 ways that people can
cause weathering and erosion
to take place.
Forces
http://en.wikipedia.org/wiki/Sodium_Chloride

16. Intermolecular Forces
• Polar Molecules have
• Dipole-dipole Force:
• + end of one molecule attracts – end of
another (based on electronegativity
difference IN polar molecules)
– PH3

17. PolarityPolarity
How do you determine polar vs. nonpolar?
Draw the molecule:
Perfectly symmetric – nonpolar
= EQUAL sharing of electrons
Asymmetric and has polar bonds – polar
= UNEQUAL sharing of electrons

18. Dipole-DipoleDipole-Dipole
The force of
attraction
between two
polar molecules

19. Hydrogen BondsHydrogen Bonds
(not really bonds)(not really bonds)
STRONG DIPOLE-DIPOLESTRONG DIPOLE-DIPOLE
The force of
attraction between
hydrogen (H) and
the lone pairs of
electrons on a
highly
electronegative
atom (F, O, or N)

20. Non-Polar molecules have
• London Dispersion Forces: nonpolar
molecules have temporary dipoles (due to
electron motion in atoms)
– I2
A second atom or molecule, in turn, can
be distorted by the appearance of the
dipole in the first atom or molecule
(because electrons repel one another)
which leads to an electrostatic attraction
between the two atoms or molecules
http://www.chem.purdue.edu/gchelp/liquids/disperse.html

21. Intermolecular Forces
• Hydrogen bonds: special type of dipole-
dipole force that results from large
electronegativity difference between
hydrogen and nitrogen, oxygen, or fluorine
• H and FON
• Not a true bond

22. London dispersion forcesLondon dispersion forces
The movement of electrons causes an instantaneous
dipole, which can then induce a dipole in a neighboring
atom, generating a temporary force.

23. Intermolecular forces
Force Type of
Compound
Relative
strength
Boiling Points
of
Compounds
with this force
Melting Points
of Compounds
with this force
Ionic
Dipole-
Dipole
Hydrogen
Bonds
London
Dispersion
Generalization: As force strength _____________, energy required to separate
molecules/ ions ___________, therefore MP and BP ______________.

24. Types of intermolecular forcesTypes of intermolecular forces
Is the molecule polar or nonpolar?
Is H bonded to either
F, O, or N?
Hydrogen Bond Dipole – Dipole force
London Dispersion
Force
NO
YES
POLAR NONPOLAR

25. ExamplesExamples
CH4
NH3
H2O
Cl2
HBr

26. Strength of the forcesStrength of the forces
Which of these forces would you
consider the strongest?
How do ionic bonds (aka ionic forces) fit
into the picture?
Hydrogen BondsHydrogen Bonds
Ionic Bonds/Forces are the strongestIonic Bonds/Forces are the strongest
forces between two particles.forces between two particles.

27. Quick Questions:
• Cl2
and Br2
have approximately the same shape
and neither is polar.
– Which intermolecular force affects Cl2 and Br2?
– Upon cooling, both Cl2
and Br2
form solids. Why?
– At 25o
C, chlorine (Cl2
) is a gas whereas bromine (Br2
)
is a liquid. Why?

28. What about energy changes
with physical changes?
-Energy changes during phase
transitions: Heating/ Cooling Curves
- Phase Diagrams: Graphical depictions
of the relationship between the phases
of matter

29. How do we predict how much energy is
required to change a substance’s phase?
• Molecular Mass: as mass increases, the
interaction between molecules increases
(due to more electrons)
• Molecular Geometry: polar or nonpolar
molecules
• Intermolecular Forces
– Forces between two ions or molecules.
– As opposed to intraparticlar forces
• Covalent bonding

30. Thermochemistry
The study of energy
changes during physical
and chemical changes

31. What is Energy?
Definition of Energy: the capacity to do work or produce heat
Equation: E = q (heat) + w (work)
Forms that energy may take:
Kinetic, potential, light, heat, work
Units for energy:
Joules (J) or Kilojoules (kJ)
Calories (cal) = 4.184J = 1 cal

32. Thermochemistry in two big
categories
Energy associated
with
Chemical Changes
Exothermic reactions
Endothermic
reactions
Energy associated
with
Physical Changes
Melting
Freezing
Boiling
Vaporizing
Sublimation
Deposition

33. How is energy involved in phase
changes?
Endothermic Exothermic
Melting Freezing
Boiling
(Evaporating)
Condensing
Sublimating Deposition

34. Physical Change
What is the Kinetic Theory?
- all particles are in constant random motion
http://www.usatoday.com/weather/tg/wevapcon/wevapcon.htm
http://www.nyu.edu/pages/mathmol/modules/water/water_concepts.html

35. Energy Causes Phase Changes
• Adding energy to a substance can cause
two things to occur
– The intermolecular forces within the
substance can weaken due to added heat
causing molecular vibration and increased
temperature.
– The intermolecular forces within the
substance can be reduced to a point at which
the substance changes phase

36. Temperature
Temperature
- When samples of different temperatures
are in contact, ENERGY is transferred.

37. Heat and Temperature
Temperature
Temperature (T) – a
measure of heat flow
An average of the kinetic
energy of all the
atoms/particles of a
substance
- Celsius, Kelvin
- Kelvin = Celsius + 273
Heat
Heat (q) – the total
energy of a substance.
A sum of the kinetic and
potential energies of a
substance
(E=q constant pressure)
The first law of thermodynamics: the law of the conservation of energy.

38. Enthalpy
A substance’s energy can also be measured by enthalpy. (H)
H = q (heat) + w (work)
At constant pressure, no work can be done on or by the
system. So, w = 0
H = q (at constant pressure)
Enthalpy is equivalent to heat if the system is at constant
pressure.

39. What is a heating/ cooling
curve?
Shows the energy changes for a substance
during a phase change Movie
Time
Temperature(°C)

40. What if there is a phase
transition (freezing/ melting)?
Q = mHfus
• Hfus = enthalpy of fusion
– Energy required to change 1 gram of a
substance from a solid to liquid (no
temperature change)
– Units: J/g
• Why is Hfus positive when you’re melting a
substance?

41. What if there is a phase transition
(evaporation/condensing)?
Q = mHvap
• Hvap = enthalpy of vaporization
– Energy required to change 1 gram of a
substance from a liquid to a gas (no
temperature change)
– Units: J/g
• Why is Hvap positive when you’re
evaporating a substance?

42. How are energy changes within
a phase measured?
Q = mCp∆T
(like the calorimetry equation)
There is a temperature change!!!
What is specific heat?

43. What is a Phase Diagram?
http://www.chemistrycoach.com/Phase_diagram.htm
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA2/MAIN/BENZENE/CD2R1.HTM
Pressure(atm)
Temperature (°C)
1.0
0.5
1.5
2.0
-10 57 60
B
C
A
D
E

44. Phase Diagram Terms
• Triple Point: the temperature and pressure
at which the solid, liquid, and gas phases
coexist at equilibrium
• Critical Point: the temperature and
pressure at which the gas and liquid states
of a substance become identical and form
one phase

45. • Normal boiling point: The temperature at
which the liquid and gas phases are in
equilibrium.
• Normal melting point: The temperature at
which the solid and liquid phases are in
equilibrium.

46. What is Vapor Pressure?
• Gas pressure exerted by particles that
have escaped from a liquid
• As temperature increases, more
molecules can escape into the gas phase,
increasing the pressure above the liquid

47. Vapor Pressure (Pvap)Pvap(atm)
Temperature (°C)
Diethyl ether Ethanol Water

48. Physical Change and Phase Diagrams
http://wps.prenhall.com/wps/media/objects/602/616516/Media_Assets/Chapter10/Text_Images/FG10_28.JPG

49. More Phase Diagrams
http://www.chem.neu.edu/Courses/1131Tom/Lecture25/img007.GIF

50. More Phase Diagrams
Temp (°C)
Pressure(atm)
A
B
C
D
E
1
5
15 32

51. Phase Diagrams Think ABOUT
Based on the phase
diagram for water explain
to the person next to you
how you think ice skating
works…
(Hint: think pressure and the solid
liquid phase boundary)

52. Phase diagrams
• What can we determine about density
from a phase diagram?
• Based on the relative density of the liquid
and solid phases of water, why do you
think ice skating works?