1. Chapter 11 Intermolecular Forces, Liquids, and Solids
Chemistry, The Central Science, 11th edition
Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten
Chapter 11 Intermolecular Forces, Liquids, and Solids
John D. Bookstaver
St. Charles Community College
Cottleville, MO
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2. Aim: What are intermolecular forces, and how can we assign them to molecules?
Do Now: Which is stronger?
Intermolecular forces or Intramolecular forces?
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3. States of Matter
The fundamental difference between states of matter is the distance between particles.
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4. States of Matter
Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.
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5. The States of Matter
The state a substance is in at a particular temperature and pressure depends on two antagonistic entities:
the kinetic energy of the particles;
the strength of the attractions between the particles.
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6. Intermolecular Forces
The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.
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7. 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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8. Intermolecular Forces
These intermolecular forces as a group are referred to as van der Waals forces.
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9. van der Waals Forces Dipole-dipole interactions Hydrogen bonding
London dispersion forces
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10. Ion-Dipole Interactions
Ion-dipole interactions (a fourth type of force), are important in solutions of ions.
The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.
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11. Dipole-Dipole Interactions
Molecules that have permanent dipoles are attracted to each other.
The positive end of one is attracted to the negative end of the other and vice-versa.
These forces are only important when the molecules are close to each other.
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12. Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling point.
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13. Critical Thinking
What do you notice is different in solid CH3CN compared to liquid CH3CN?
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14. London Dispersion Forces
While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.
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15. London Dispersion Forces
At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.
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16. London Dispersion Forces
Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.
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17. London Dispersion Forces
London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.
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18. 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.
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19. Factors Affecting London Forces
The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).
This is due to the increased surface area in n-pentane.
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20. Factors Affecting London Forces
The strength of dispersion forces tends to increase with increased molecular weight.
Larger atoms have larger electron clouds which are easier to polarize.
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21. Aim: How do intermolecular forces affect physical properties of liquids.
Do Now: What types of van der Waals forces do you expect to see between molecules of HF?
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22. 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 be the dominating force.
If one molecule is much larger than another, dispersion forces will likely determine its physical properties.
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23. How Do We Explain This?
The nonpolar series (SnH4 to CH4) follow the expected trend.
The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.
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24. Hydrogen Bonding
The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
We call these interactions hydrogen bonds.
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25. Hydrogen Bonding
Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.
Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.
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26. Sample Exercise 11.1 (Pg. 450)
In which of these substances is hydrogen bonding likely to play an important role in determining physical properties? Methane (CH4) Hydrazine (H2NNH2) Methyl fluoride (CH3F) Hydrogen sulfide (H2S)
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27. Summarizing Intermolecular Forces
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28. Go Figure (Pg. 453)
At what point in the flow chart, would a distinction be made between SiH4 and SiH2Br2?
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29. Exercise 11.2 (Pg. 454)
List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling point.
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30. Aim: How do intermolecular forces affect properties of liquids
Aim: How do intermolecular forces affect properties of liquids? (continued)
Do Now: Order the following molecules from lowest to highest in terms of boiling points.
AgNO3, CH4, SiF4, H2O, SO3
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31. Exercise 11.3
(a) Identify the intermolecular attractions present in the following substances, and (b) select the substance with the highest boiling point: CH3CH3 , CH3OH, and CH3CH2OH.
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32. 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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33. 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. Increases with increasing molecular weight
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34. Surface Tension
Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.
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35. Review: How do viscosity and surface tension change?
As temperature increases
As intermolecular forces of attraction become stronger
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36. Aim: What are the energy changes associated with phase changes, and what is vapor pressure?
Do Now: In terms of molecules, what happens to viscosity as temperature increases?
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37. Phase Changes
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38. 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.
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39. Energy Changes Associated with Changes of State
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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40. 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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41. Sample Exercise 11.3 (Pg.460)
Calculate the enthalpy change upon converting 1.00 mol of ice at -25∘C to steam at 125∘C under a constant pressure of 1 atm. The specific heats of ice, liquid water, and steam are 2.03, 4.18, and 1.84 J/g-K, respectively. For H2O, ∆Hfus= 6.01 kJ/mol and ∆Hvap =40.67 kJ/mol.
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42. Practice Exercise 1 (Pg. 460)
What information about water is needed to calculate the enthalpy change for converting 1 mol H2O(g) at 100∘C to H2O(l) at 80∘C?
Heat of fusion
Heat of vaporization
Heat of vaporization and specific heat of H2O(g)
Heat of vaporization and specific heat of H2O(l)
Heat of fusion and specific heat of H2O(l)
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43. Vapor Pressure
At any temperature some molecules in a liquid have enough energy to escape.
As the temperature rises, the fraction of molecules that have enough energy to escape increases.
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44. Vapor Pressure
As more molecules escape the liquid, the pressure they exert increases.
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45. Vapor Pressure
The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.
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46. Vapor Pressure
The boiling point of a liquid is the temperature at which it’s vapor pressure equals atmospheric pressure.
The normal boiling point is the temperature at which its vapor pressure is 760 torr.
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47. Practice Exercise 1 (Pg. 464)
In the mountains, water in an open container will boil when:
Its critical temperature exceeds room temperature
Its vapor pressure equals atmospheric pressure
Its temperature is 100∘C
Enough energy is supplied to break covalent bonds
None of these is correct
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48. 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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49. Aim: What are the types of solids?
Do Now: Which solids will have the highest boiling point?
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50. Solids We can think of solids as falling into two groups:
crystalline, in which particles are in highly ordered arrangement.
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51. 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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52. 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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53. Crystalline Solids
Because of the ordered in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.
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54. Crystalline Solids
There are several types of basic arrangements in crystals, like the ones depicted above.
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55. Types of Bonding in Crystalline Solids
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56. Covalent-Network and Molecular Solids
Diamonds are an example of a covalent-network solid, in which atoms are covalently bonded to each other.
They tend to be hard and have high melting points.
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57. Covalent-Network and Molecular Solids
Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces.
They tend to be softer and have lower melting points.
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58. Metallic Solids
Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.
In metals valence electrons are delocalized throughout the solid.
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