1. Liquids and Solids
Write your thoughts about the following question in your notebook.
Look around the room. There are solids, liquids, and gases. But, all of the room is 22 ○C. Why isn’t everything in the same state?
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2. Compare evaporation rates.
Need:Wax paper, Water (H2O), ammonia (NH3), isopropanol (C3H7OH)
Procedure:
Use pipettes to drop each substance separately on the wax paper.
Use the same number of drops for each substance.
Record the order in which they evaporate
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3. Intermolecular Forces, Liquids, and Solids
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4. Review NH3 CaO Na3N CO2 C8H8 CO2 C8H8 CaO Na3N NH3 Ionic Covalent
Learning goal from Bonding Unit Compare and Contrast ionic and covalent compounds in terms of elements present.
NH3
CaO Na3N
CO2 C8H8
CO2 C8H8
CaO Na3N
Ionic
NH3
Covalent
Usable definition
Ionic compounds: metal and nonmetal
Covalent compounds: two or more nonmetals
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5. Intramolecular Bonds
The forces that hold one molecule together are intramolecular.
Intra means within.
Why do some molecules remain as individual units (gases), while some are generally found as liquids or solids?
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6. States of Matter
The fundamental difference between states of matter is the distance between particles.
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7. The States of Matter
The state of substances, at any specific temperature and pressure, depends on two characteristics:
the kinetic energy of the particles;
the strength of the attractions between the particles.
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8. Intermolecular forces
The attractions between molecules are called intermolecular forces.
Think highways:
“interstate” is between two different states.
“intrastate is within just one state.
Intramolecular forces are what hold one molecule together.
Intermolecular forces are what hold many molecules together.
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9. Learning Check Intramolecular or Intermolecular?
HCl – hydrochloric acid
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10. Which is stronger intermolecular or intramolecular forces?
The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.
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11. Intermolecular Forces
They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.
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12. Intermolecular Forces
These intermolecular forces as a group are referred to as van der Waals forces.
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13. van der Waals Forces Dipole-dipole interactions Hydrogen bonding
London dispersion forces
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14. Intermolecular or Intramolecular.
A student heated a solid solution. The solution melted into a liquid. Tests on both the solid and liquid showed it to be the same substance. What type of bonds were broken –
intermolecular
intramolecular
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15. Intermolecular or Intramolecular.B
intermolecular
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16. Review of Polar Molecules
More electronegative elements ( Red in table below) draw electrons in a bond towards them.
electronegativity→↑
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17. Review of Polar Molecules
The more electronegative
element will pull e-
towards it, creating
a negative pole. The area of the molecule with a
deficit of e- will be the positive pole.There must be a 0.5 or higher difference in electronegativity.
H2CO
Electronegativity
Hydrogen
2.1
Carbon
2.5
Nitrogen
3.0
Oxygen
3.5
NH3
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18. Polar Molecules
This polarity is permanent. In ammonia, N will always have a – charge and H will always have a + charge
Ammonia molecule - +
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19. Which molecule has polar bonds?
Fluoride trihydride
Phosphorus trihydride
Methane CH4
Electronegativity
Boron
2.0
Hydrogen
2.1
Phosphorus
Carbon
2.5
Nitrogen
3.0
Oxygen
3.5
Fluoride
4.0
Only #1
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20. Dipole-Dipole Interactions
Molecules that have permanent dipoles (polar molecules)are attracted to each other.
The positive end of one is attracted to the negative end
of the other and vice-versa.
Dipole-Dipole
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21. Dipole-Dipole Interactions
Boiling Point
Isopropanol
82.3˚C
Household ammonia
~ 97˚C
Water
100˚C
The more polar the molecule, the higher its boiling point. Why?
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22. Structure and Polarity
Isopropanol
Water
Ammonia _ _ _ + + + + + +
What is so different about the structure of isopropanol that its boiling point is so much lower than water or ammonia?
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23. Which molecule can form dipole dipole bonds?
Electronegativity
Boron
2.0
Hydrogen
2.1
Phosphorus
Carbon
2.5
Nitrogen
3.0
Oxygen
3.5
Fluoride
4.0
Fluoride trihydride
Phosphorus trihydride
Methane CH4
Only #1
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24. London Dispersion Forces+ _ +
Electrons generally repel each other (and,
therefore, tend to stay far away from each other).
Occasionally, randomly they are more concentrated
on one side of the atom.
This creates a temporary dipole.
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25. London Dispersion Forces
The atom or molecule is temporarily polar, with an excess of electrons on the left side and a shortage on the right side. _ _ + +
This induces charges in neighboring atoms or molecules. Dispersion bonding results.
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26. London Dispersion Forces
These forces are present in all molecules, whether they are polar or nonpolar.
The tendency of an electron cloud to distort in this way is called polarizability.
The strength of dispersion forces tends to increase with increased molecular weight. More electrons to form dispersion forces.
Larger atoms have larger electron clouds which are easier to polarize
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27. Which Have a Greater Effect
Which Have a Greater Effect? Dipole-Dipole Interactions or Dispersion Forces
If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force.
If one molecule is much larger than another, dispersion forces will likely determine its physical properties.
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28. Hydrogen Bonding
The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong, because these atoms are very electronegative.
We call these interactions hydrogen bonds. They are a strong type of dipole-dipole.
Water molecule
Hydrogen Bond
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29. Learning Check Which type of intermolecular bonding is each molecule
Electronegativity
Boron
2.0
Hydrogen
2.1
Phosphorus
Carbon
2.5
Sulfur
2.6
Nitrogen
3.0
Oxygen
3.5
Fluoride
4.0
Which type of intermolecular
bonding is each molecule
capable of?
C3H8
C3H7OH
CS2
CH3Cl
All molecules are capable of London Dispersion bonding
C3H7OH (hydrogen bonding)
CH3Cl dipole dipole bonding
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30. Learning Check
Which molecule would have stronger London Dispersion forces?
C3H8 C10H22
C10H22
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31. Summarizing Intermolecular Forces
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32. Oobleck – A Non Newtonian Fluid
Measure 3 – 4 spoons of cornstarch and put into a bowl.
Then gradually add approximately a small amount of water to the cornstarch.
Stir well (this will take some time).
Add small amounts of more water or cornstarch until you get a mixture which is almost solid – almost liquid.
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33. Intermolecular Forces Affect Many Physical Properties
The strength of the attractions between particles can greatly affect the properties of a substance or solution.
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34. Viscosity Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other.
Viscosity increases with stronger intermolecular forces and decreases with higher temperature.
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35. Surface Tension
Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.
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36. Hydrogen bonding, the strongest IMF, gives water a high surface tension.
Observe the pictures at bottom left. What can you say about the intermolecular bonding of water compared to that of mercury?
water
mercury
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37. Surfactants
Detergents and soaps belong to a group called surfactants.
Surfactants work by decreasing the surface tension of a liquid, usually water.
Usually long carbon hydrogen chain
Usually some small salt structure
Oil
water
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38. © 2009, Prentice-Hall, Inc.

39. Surfactants in your Lungs??
Surfactant reduces the surface tension of fluid in the lungs and helps make the small air sacs in the lungs (alveoli) more stable.
This keeps them from collapsing when an individual exhales.
In preparation for breathing air, fetuses begin making surfactant while still in the womb.
Babies that are born very
prematurely often lack adequate
surfactant and must receive
surfactant replacement therapy
immediately after birth in order
to breathe.
Pulmonary surfactant
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40. Capillary Action
When the intermolecular attractions between the molecules of the liquid and those of the tube is greater than the combined effects of gravity and the attractive forces within the liquid, the liquid rises in the tube until equilibrium is restored.
Capillary action is very important in nature, particularly in the transport of water in plants and through the soil.
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41. Anti-fog agents (wetting agents) were initially developed by NASA during the Project Gemini, for use on helmet visors.
Wetting agents reduce surface tension of water and prevent drop formation.
Dive masks
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42. Phase Changes
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43. Energy Changes Associated with Changes of State
The heat of fusion is the energy required to change a solid at its melting point to a liquid.
The heat of vaporization is defined as the energy
required to change a liquid at its boiling point to
a gas.
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44. Energy Changes Associated with Changes of State
The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.
The temperature of the substance does not rise during a phase change.
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45. Kinetic Molecular Theory
A theory used to explain the behaviors and characteristics of gases.
and sometimes solids and liquids

46. KMT Assumptions
A gas is made up of many small particles that move constantly and randomly in straight lines.
The molecules in a gas occupy no volume.
When gas molecules collide, they don’t lose energy due to friction or gain energy either.
Gas molecules are not attracted to each other at all.
The kinetic energy of gas molecules depends only on the temperature of the gas.

47. KMT Assumptions
Ideal gases would be exactly like the description on the previous slide. It is useful to use them as a model. However, they do not actually exist.
Real gases :
really are small, constantly moving particles
but, the molecules do have some volume
and, they do lose energy due to friction in collisions
and, they are slightly attracted to each other
their energy is really only dependent on temperature

48. If a Real gas is at a high temperature and low pressure, it behaves very much like an Ideal gas.
Why?
At high temperatures, the molecules have a lot of energy – hard to notice really small losses and they can escape any attraction to another molecule.
Ideal Gas
Real Gas
At low pressures, the molecules are not forced close to each other – so volume doesn’t matter and they are not close enough to be attracted very much to each other.

49. Phase Diagrams
Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.
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50. Phase Diagrams The circled line is the liquid-vapor interface.
It starts at the triple point (T), the point at which all three states are in equilibrium.
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51. Phase Diagrams
It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.
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52. Phase Diagrams
Each point along this line is the boiling point of the substance at that pressure.
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53. Phase Diagrams
The circled line in the diagram below is the interface between liquid and solid.
The melting point at each pressure can be found along this line.
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54. Phase Diagrams
Below the triple point the substance cannot exist in the liquid state.
Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.
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55. Phase Diagram of Water
Note the high critical temperature and critical pressure.
These are due to the strong van der Waals forces between water molecules.
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56. Phase Diagram of Water The slope of the solid-liquid line is negative.
This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.
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57. Solids We can think of solids as falling into two groups:
crystalline, in which particles are in highly ordered arrangement.
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58. Solids We can think of solids as falling into two groups:
amorphous, in which there is no particular order in the arrangement of particles.
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59. Attractions in Ionic Crystals
In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.
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