1. Chapter 11 Liquids, solids, and intermolecular forces

2. Intermolecular forces
Intermolecular forces – attractive forces between particles composed of matter
These forces dictate many properties of an atom
Melting point, boiling point, crystal lattice, miscibility, solubility, etc.
Driving factor for intermolecular forces is Coulomb’s law!!!!
Positive and negative charges attract!!!
Intermolecular forces

3. Dispersion forces
Dispersion/London forces – the result of fluctuations in the electron distribution within molecules or atoms
At the point in time when the electrons are localized on one side of the atom, the region has a partial negative charge, while the opposite side of the atom would have a partial positive charge

4. Dispersion force: Atom size
Larger atoms will have greater dispersion forces
More electrons, more space, more volume in which electrons can polarize the atom
Greater dispersion force results in higher melting points and boiling points since it take more energy (higher temperature) to separate the attracted particles

5. Dispersion forces: Molecular shape
In general, higher molar mass results in greater dispersion forces
Molecular shape must also be considered, since it will determine the surface area of interaction

7. Dipole-dipole force
Dipole-dipole force – attractive force between polar molecules
If a molecule has a net dipole it is polar and the positive region of one molecule can attract to the negative region in another molecule

8. Dipole-dipole vs dispersion
Dipole-dipole interactions are a stronger force than dispersion forces since the dipoles are permanent instead of temporary
Every molecule has dispersion forces, but only polar molecules have dipole-dipole forces

9. Dipole-dipole force: Dipole moment
The greater the magnitude of the positive and negative charge in the dipole, the greater the force
Increasing dipole moment = stronger dipole-dipole force

10. Dipole-dipole force: Miscibility
Miscibility – the ability to mix without separating into two states
Dipole-dipole forces are present between one polar molecule and another, so they will be miscible
Dipole-dipole forces are not present between polar molecules and nonpolar molecules
Water is a polar molecule
Oil is a non polar molecule
Water and oil are not miscible, they don’t mix

11. Hydrogen bonding
Hydrogen bonding – attractive force that occurs between the H that is bonded to a very electronegative atom (O, N, F) on one molecule and a very electronegative atom on another molecule
Since there is a large electronegativity difference, the magnitude of the dipole moment is larger, resulting in a very strong intermolecular force, larger than regular dipole-dipole forces

13. Ion-dipole force
Ion-dipole force – attractive force between ions and a polar compound
Ionic compound composed of cation and anion
Cation is attracted to negative end of dipole
Anion is attracted to positive end of dipole

14. Summary of intermolecular forces

15. Identify the intermolecular forces at play in a sample of the following molecules (assume solid or liquid phase)
Cl2
NH3 KI
Identify the major interaction expected between the following molecules
Water and ammonia (NH3)
Formaldehyde (CH2O) and methane (CH4)
Practice

16. Problem
C2H6O has two isomers (molecules with the same formula but with different structures) ethanol and dimethyl ether. Draw the two possible Lewis dot structures for C2H6O and then determine which should have a higher boiling point.

17. Surface tension
Intermolecular forces lower the potential energy of molecules
More interactions equal lower energy
Less interactions equal higher energy
Surface tension – the energy required to increase the surface area by a unit amount
It takes energy to increase the surface area

18. Surface tension continued
Placing an object on the surface will increase the surface area of the liquid
Even if the object is more dense, it needs to overcome the surface tension of the liquid before it can sink, it needs to overcome the energy to move interior molecules to the surface
Surface tension continued

19. Viscosity
Intermolecular forces attract molecules to each other, making them less likely to flow
Viscosity – the resistance of a liquid to flow
Lower viscosity, faster flow
Water
Higher viscosity, slower flow
Maple syrup
All factors that enhance intermolecular forces (force type, shape and size) will affect viscosity
Temperature also affects viscosity
Higher temperature, lower viscosity
Lower temperature, higher viscosity

20. Viscosity trends

21. Capillary action – the ability of a liquid to flow against gravity up a narrow tube
Adhesive force – the attraction of the molecules with the surface of the tube
Cohesive force – the attraction of the molecules in the liquid to each other
Adhesive forces pull the liquid up a tube until gravity balances out the capillary action
Capillary action

22. Tube width and meniscus shape
Liquid will travel higher up more narrower tubes
Narrower tube increases adhesive forces per volume of liquid
Stronger adhesive forces cause liquid to creep up the sides to produce a concave meniscus shape
If cohesive forces are stronger then the meniscus will have a convex shape

23. Capillary action video

24. Vaporization Vaporization – the transition from liquid to gas
The thermal energy of molecules in a sample is represented by a distribution curve that has a dependence on temperature
Molecules on the higher end of the distribution curve can have enough energy to transition into the gas phase
Vaporization

25. Heat of vaporization ΔHvap
It requires energy to go from the liquid to the gas phase (to vaporize), this process is endothermic
Heat of vaporization (ΔHvap) – the amount of heat required to vaporize one mole of a liquid to a gas
At 100° C (boiling point of water)
ΔHvap = kJ/mol H2O(l)  H2O(g)

26. Condensation
Condensation – the transition from the gas phase to the liquid phase
This is the opposite of vaporization and is an exothermic process
For the reverse process the heat is the same magnitude, but with opposite sign
At 100° C (boiling point of water)
ΔHvap = kJ/mol H2O(g)  H2O(l)

27. How much heat does it take to boil 1 mol of water at its boiling point?
How much heat does it take to boil 45.8g of water at its boiling point?
Calculate the mass of water (in grams) that can be vaporized at its boiling point with 155 kJ of heat.
Practice

28. Vapor pressure
In a sealed container liquid molecules can enter the gas phase, but they can also condense back down into the liquid phase
The rate of vaporization stays constant, while the rate of condensation increases until both rates are equal. At this point, a dynamic equilibrium has been reached
The pressure of a gas in a dynamic equilibrium with its liquid is its vapor pressure

30. Temperature dependence
Temperature affects the vapor pressure of a liquid
Increasing temperature increases the vapor pressure increase

31. Boiling point – the temperature at which the liquid’s vapor pressure equals the external pressure
Surface molecules and interior molecules now have enough energy to enter the gas phase
Normal boiling point – the temperature at which the vapor pressure of a liquid is equal to 1 atm
Boiling point

32. Pressure dependence Lower pressure will lower the boiling point
Lower pressure makes it easier for liquid molecules to enter the gas phase, lowering the temperature at which the liquid “boils”
Pressure dependence

33. Heating a liquid in a sealed container will increase the vapor pressure, producing more gas molecules, increasing the pressure of the container
Since the container is sealed and more gas molecules are being forced into the same space, the density of the gas increases
At the same time the density of the liquid is decreasing with the temperature increase
Eventually this will cause the liquid and gas phases to combine to make a single phase, at this point it is considered a supercritical fluid
The temperature at this point is the critical temperature (Tc)
The pressure at this point is the critical pressure (Pc)
Critical point

35. Heating a liquid

36. Practice
Will cooking macaroni in a pot of water occur faster in Norco, California (elevation = 195 m) or in Denver, Colorado (elevation = 1609 m)?
Which element should have a higher boiling point, Ne or Xe?

37. Solid phase changes
Sublimation – transition from the solid to gas phase
Deposition – transition from gas to solid phase
Melting/fusion – transition of a solid to liquid phase
Freezing – the transition of a liquid to a solid
Melting/freezing point – the temperature at which a liquid freezes, or a solid melts

38. Heat of fusion and sublimation
Heat of fusion (ΔHfus ) – the amount of heat required to melt 1 mol of solid
ΔHfus = kJ/mol
H2O(s)  H2O(l) endothermic process
Heat of sublimation (ΔHsub ) – the amount of heat required to sublime 1 mol of solid
ΔHsub = kJ/mol
H2O(s)  H2O(g) endothermic process
Solid to liquid to gas, overall process is sublimation
ΔHsub = ΔHfus + ΔHvap
Heat of fusion and sublimation

39. ΔHfus vs ΔHvap

40. Problems
How much energy is needed to convert 25.6 g of H2O(s) to H2O(l) at 0 °C?
For H2O, ΔHfus = kJ/mol

41. Build a phase diagram activity

42. Phase diagrams
Temperature and pressure affect the phase of a substance
Phase diagrams show at a given temperature and pressure what a phase will be

43. Navigating a phase diagram

44. Phase diagrams
Intermolecular forces for different substances change the properties of that substance (melting point, boiling point, critical point, triple point)
Phase diagrams are different for each substance

45. Chapter 11 done s
Intermolecular