1. Intermolecular Forces!!! AKA the forces that hold stuff together

2. HOW IS THIS POSSIBLE?!?!?! BELLRINGER
Get a “Unit 5 Guide” from the front table
On the front lab table, there is a paper clip floating on top of a petri dish filled with water.
HOW IS THIS POSSIBLE?!?!?!
Discuss with a neighbor.

3. A Molecular Comparison of Liquids and Solids

4. A Molecular Comparison of Liquids and 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. Intermolecular Forces
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).

6. Intermolecular Forces

7. Intermolecular Forces
Ionic Bonds
Remember that ionic bonds are TECHNICALLY intermolecular forces, not intramolecular
Strongest of all intermolecular forces.

8. Intermolecular Forces
Ion-Dipole Forces
Interaction between an ion and a dipole (e.g. water).
Next strongest intermolecular force.

9. Van-der-Waals Forces
"Intermolecular forces" is a general term referring to any of the attractions between molecules or atoms
“Van-der-Waals forces” are more specific, and occur between molecules or atoms of the same type. Ionic bonds are not Van der Waals forces because they occur between DIFFERENT atoms
In other words, all Van der Waals forces are intermolecular forces, but not all intermolecular forces are Van der Waals forces.

10. Intermolecular Forces
Dipole-Dipole Forces
Dipole-dipole forces exist between neutral polar molecules.
Polar molecules need to be close together.
Weaker than ion-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.

11. Intermolecular Forces
Dipole-Dipole Forces

12. Intermolecular Forces
Dipole-Dipole Forces

13. Intermolecular Forces
London 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).

14. Intermolecular Forces
London Dispersion Forces
One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom).
The forces between instantaneous dipoles are called London dispersion forces.

15. Intermolecular Forces
London Dispersion Forces
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.
London dispersion forces increase as molecular weight increases.
London dispersion forces exist between all molecules.
London dispersion forces depend on the shape of the molecule.

16. Intermolecular Forces
London Dispersion Forces
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.

17. Intermolecular Forces
London Dispersion Forces

18. Intermolecular Forces
London Dispersion Forces

19. Intermolecular Forces
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.

20. Intermolecular Forces
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.

21. Hydrogen Bonding

22. Hydrogen Bonding

23. Intermolecular Forces
Hydrogen Bonding
Hydrogen bonds are responsible for:
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.

24. Intermolecular Forces
Hydrogen Bonding

25. Intermolecular Forces

26. Some Properties of Liquids
Viscosity
Viscosity is the resistance of a liquid to flow.
A liquid flows by sliding molecules over each other.
The stronger the intermolecular forces, the higher the viscosity.
Surface Tension
Bulk molecules (those in the liquid) are equally attracted to their neighbors.

27. Some Properties of Liquids
Viscosity

28. Surface Tension

29. Some Properties of Liquids
Surface Tension
Surface molecules are only attracted inwards towards the bulk molecules.
Therefore, surface molecules are packed more closely than bulk molecules.
Surface tension is the amount of energy required to increase the surface area of a liquid.
Cohesive forces bind molecules to each other.
Adhesive forces bind molecules to a surface.

30. Some Properties of Liquids
Surface Tension
Meniscus is the shape of the liquid surface.
If adhesive forces are greater than cohesive forces, the liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e.g. water in glass).
If cohesive forces are greater than adhesive forces, the meniscus is curved downwards.
Capillary Action: When a narrow glass tube is placed in water, the meniscus pulls the water up the tube.

31. 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.

32. Phase Changes

33. Phase Changes Energy Changes Accompanying 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).

34. Phase Changes Energy Changes Accompanying Phase Changes
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.

35. Phase Changes

36. Phase Changes Energy Changes Accompanying Phase Changes
All phase changes are possible under the right conditions.
The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
cool gas  condense  cool liquid  freeze  cool solid
is exothermic.

37. Phase Changes Heating Curves
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.

39. Phase Changes 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.

40. Phase Changes
Critical Temperature and Pressure

41. Vapor Pressure Explaining Vapor Pressure on the Molecular Level
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.

42. Vapor Pressure
Explaining Vapor Pressure on the Molecular Level

43. Vapor Pressure Explaining Vapor Pressure on the Molecular Level
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.
Volatility, Vapor Pressure, and Temperature
If equilibrium is never established then the liquid evaporates.
Volatile substances evaporate rapidly.

44. Vapor Pressure Volatility, Vapor Pressure, and Temperature
The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

45. Vapor Pressure
Volatility, Vapor Pressure, and Temperature

47. Vapor Pressure Vapor Pressure and Boiling Point
Liquids boil when the external pressure equals the vapor pressure.
Temperature of boiling point increases as pressure increases.

48. Vapor Pressure Vapor Pressure and Boiling Point
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.
Normal boiling point is the boiling point at 760 mmHg (1 atm).

49. Phase Diagrams
Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases.
Given a temperature and pressure, phase diagrams tell us which phase will exist.
Any temperature and pressure combination not on a curve represents a single phase.

50. Phase Diagrams Features of a phase diagram:
Triple point: temperature and pressure at which all three phases are in equilibrium.
Vapor-pressure curve: generally as pressure increases, temperature increases.
Critical point: critical temperature and pressure for the gas.
Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid.
Normal melting point: melting point at 1 atm.

51. Phase Diagrams

52. Phase Diagrams
The Phase Diagrams of H2O and CO2

53. Phase Diagrams The Phase Diagrams of H2O and CO2 Water:
The melting point curve slopes to the left because ice is less dense than water.
Triple point occurs at C and 4.58 mmHg.
Normal melting (freezing) point is 0C.
Normal boiling point is 100C.
Critical point is 374C and 218 atm.

54. Phase Diagrams The Phase Diagrams of H2O and CO2 Carbon Dioxide:
Triple point occurs at -56.4C and 5.11 atm.
Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.)
Critical point occurs at 31.1C and 73 atm.