1. CHAPTER 9
Liquids and Solids 1

2. Description of Liquids & Solids
Solids & liquids are condensed states
atoms, ions, molecules are close to one another
highly incompressible
Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other.
Liquids & gases are fluids
easily flow
Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other.
Intermolecular attractions in liquids & solids are strong 3

3. Description of Liquids & Solids

4. Description of Liquids & Solids
Converting a gas into a liquid or solid requires the molecules to get closer to each other:
cool or compress.
Converting a solid into a liquid or gas requires the molecules to move further apart:
heat or reduce pressure.
The forces holding solids and liquids together are called intermolecular forces.

5. Kinetic-Molecular Description of Liquids & Solids
strengths of interactions among particles &
degree of ordering of particles
Gases< Liquids < Solids 5

6. Intermolecular Attractions
The covalent bond holding a molecule together is an intramolecular forces.
The attraction between molecules is an intermolecular force.
Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl).
When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).
When a substance condenses intermolecular forces are formed.

7. Intermolecular Attractions

8. Intermolecular Attractions
Dipole-Dipole Forces
Dipole-dipole forces exist between neutral polar molecules.
Polar molecules need to be close together.
Weaker than ion-dipole forces:
Q1 and Q2 are partial charges.

9. Intermolecular Attractions
Dipole-Dipole Forces
There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble.
If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

10. Intermolecular Attractions
Dipole-dipole interactions
consider NH3 a very polar molecule 11

11. Intermolecular Attractions
Dispersion Forces
Weakest of all intermolecular forces.
It is possible for two adjacent neutral molecules to affect each other.
The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom).
For an instant, the electron clouds become distorted.
In that instant a dipole is formed (called an instantaneous dipole).

12. Intermolecular Attractions
Dispersion Forces

13. Intermolecular Attractions
Polarizability is the ease with which an electron cloud can be deformed.
The larger the molecule (the greater the number of electrons) the more polarizable.

14. Intermolecular Attractions

15. Intermolecular Attractions
Dispersion Forces
London dispersion forces depend on the shape of the molecule.
The greater the surface area available for contact, the greater the dispersion forces.
London dispersion forces between spherical molecules are lower than between sausage-like molecules.

16. Intermolecular Attractions
Hydrogen bonding
consider H2O 12

17. Intermolecular Attractions
Hydrogen Bonding
Special case of dipole-dipole forces.
By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high.
Intermolecular forces are abnormally strong.

18. Intermolecular Attractions
Hydrogen Bonding
H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N).
Electrons in the H-X (X = electronegative element) lie much closer to X than H.
H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X.
Therefore, H-bonds are strong.

19. Intermolecular Attractions
Hydrogen Bonding

20. Intermolecular Attractions
Hydrogen Bonding
Ice Floating
Solids are usually more closely packed than liquids;
therefore, solids are more dense than liquids.
Ice is ordered with an open structure to optimize H-bonding.
Therefore, ice is less dense than water.
In water the H-O bond length is 1.0 Å.
The O…H hydrogen bond length is 1.8 Å.
Ice has waters arranged in an open, regular hexagon.
Each + H points towards a lone pair on O.
Ice floats, so it forms an insulating layer on top of lakes, rivers, etc. Therefore, aquatic life can survive in winter.

21. Intermolecular Attractions
Hydrogen Bonding
Hydrogen bonds are responsible for:
Protein Structure
Protein folding is a consequence of H-bonding.
DNA Transport of Genetic Information

22. Comparing Intermolecular Attractions

23. Intermolecular Attractions
Coulomb’s law & the attraction energy determine:
melting & boiling points of ionic compounds
the solubility of ionic compounds
Arrange the following ionic compounds in the expected order of increasing melting and boiling points.
NaF, CaO, CaF2 9

24. Intermolecular Attractions and Phase Changes10

25. Intermolecular Attractions and Phase Changes12

26. Evaporation
Process in which molecules escape from the surface of a liquid
T dependent 7

27. Evaporation 8

28. Vapor Pressure
pressure exerted by a liquid’s vapor on its surface at equilibrium
Vap. Press. (torr) for 3 Liquids Norm B.P.
0oC 20oC 30oC
diethyl ether oC
ethanol oC
water oC 9

29. Vapor Pressure
Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid.
These molecules move into the gas phase.
As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid.
After some time the pressure of the gas will be constant at the vapor pressure.

30. Vapor Pressure
Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface.
Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

31. Vapor Pressure
If equilibrium is never established then the liquid evaporates.
Volatile substances evaporate rapidly.
The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

32. Vapor Pressure
Liquids boil when the external pressure equals the vapor pressure.
Temperature of boiling point increases as pressure increases.
Two ways to get a liquid to boil: increase temperature or decrease pressure.
Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required.

33. Boiling Points
Boiling point is temperature at which the liquid’s vapor pressure is equal to applied pressure
normal boiling point is boiling 1 atm 11

34. Distillation
Process in which a mixture or solution is separated into its components on the basis of the differences in boiling points of the components
Distillation is another vapor pressure phenomenon. 12

35. The Liquid State energy associated with changes of state
heat of vaporization
amount of heat required to change 1 g of a liquid substance to a gas at constant T
units of J/g
heat of condensation
reverse of heat of vaporization 3

36. The Liquid State molar heat of vaporization or DHvap
amount of heat required to change 1 mol of a liquid to a gas at constant T
units of J/mol
molar heat of condensation
reverse of molar heat of vaporization 4

37. The Liquid State

38. Phase Changes
Surface molecules are only attracted inwards towards the bulk molecules.
Sublimation: solid  gas.
Vaporization: liquid  gas.
Melting or fusion: solid  liquid.
Deposition: gas  solid.
Condensation: gas  liquid.
Freezing: liquid  solid.

39. Phase Changes
Sublimation: Hsub > 0 (endothermic).
Vaporization: Hvap > 0 (endothermic).
Melting or Fusion: Hfus > 0 (endothermic).
Deposition: Hdep < 0 (exothermic).
Condensation: Hcon < 0 (exothermic).
Freezing: Hfre < 0 (exothermic).
Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization:
it takes more energy to completely separate molecules, than partially separate them.

40. Phase Changes is endothermic. is exothermic.
All phase changes are possible under the right conditions (e.g. water sublimes when snow disappears without forming puddles).
The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
cool gas  condense  cool liquid  freeze 
cool solid
is exothermic.

41. Phase Changes

42. Phase Changes
Plot of temperature change versus heat added is a heating curve.
During a phase change, adding heat causes no temperature change.
These points are used to calculate Hfus and Hvap.
Supercooling: When a liquid is cooled below its melting point and it still remains a liquid.
Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

43. Phase Changes and Heating Curves

44. Critical Temperature and Pressure
Gases liquefied by increasing pressure at some temperature.
Critical temperature: the minimum temperature for liquefaction of a gas using pressure.
Critical pressure: pressure required for liquefaction.

45. Phase Diagrams (P vs T)
convenient way to display all of the different phases of a substance
phase
diagram for
water 8

46. Phase Diagrams (P vs T)
phase diagram for carbon dioxide 9

47. Synthesis Question
Maxwell House Coffee Company decaffeinates its coffee beans using an extractor that is 7.0 feet in diameter and 70.0 feet long. Supercritical carbon dioxide at a pressure of atm and temperature of C is passed through the stainless steel extractor. The extraction vessel contains 100,000 pounds of coffee beans soaked in water until they have a water content of 50%.

48. Synthesis Question
This process removes 90% of the caffeine in a single pass of the beans through the extractor. Carbon dioxide that has passed over the coffee is then directed into a water column that washes the caffeine from the supercritical CO2. How many moles of carbon dioxide are present in the extractor?

49. Synthesis Question

50. Synthesis Question

51. Group Question
How many CO2 molecules are there in 1.0 cm3 of the Maxwell House Coffee Company extractor? How many more CO2 molecules are there in a cm3 of the supercritical fluid in the Maxwell House extractor than in a mole of CO2 at STP?