1. Chapter 13
Liquids and solids

2. 13.1 Intermolecular forces
The physical properties of a substance generally depend on:
the spacing between particles (atoms, molecules, ions) that make up the substance
the forces of attraction among them.

3. Intermolecular forces determine a substances state at room temperature as well as melting point and boiling point

4. Reviewing what we know gases Solids Low density Highly compressible
Expand to fill container
High density
Only slightly compressible
Rigid (keeps its shape)

5. Intermolecular vs. intramolecular forces
Inter- means between or among
Intermolecular forces can hold together identical particles or two different types of particles
Weaker than intramolecular forces (bonds)

6. Forces in water:

7. London Dispersion forces
Occur in all substances
Weak and short lived
Only intermolecular force nonpolar molecules experience
Atoms develop a temporary dipolar arrangement of charge as electrons move around the nucleus
This instantaneous dipole can then induce a similar dipole in a neighboring atom.

9. Become stronger as the size of atoms or molecules increase

11. Dipole-Dipole Attraction
Attraction between oppositely charged regions of polar molecules
Polar molecule =
Neighboring polar molecules orient themselves so that oppositely charged regions align

13. Hydrogen bonding
Dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a fluorine, oxygen, or nitrogen atom

14. What types of intermolecular forces would be present in the liquid state of each of the following?
CH3OH
co2
h2o Kr
NH3

15. Rank the following from lowest to highest boiling point
CH3OH
co2
h2o Kr
NH3

16. 13.2 Phase Changes Phase changes are physical changes
no chemical bonds are broken/made

17. Energy requirements of phase changes
Molar heat of fusion – energy required to melt 1 mole of a substance (water = 6.01 kJ/mol)
Molar heat of vaporization – energy required to vaporize 1 mole of a substance (water = 40.7 kj/mol)

19. Liquid  steam: 40.7 kJ/mol
Csteam = 1.70 J/goC
CWater = 4.18 J/goC
solid  liquid: 6.01 kJ/mol
Cice= 2.10 J/goC

20. How much heat must be absorbed to melt 150.0 g of ice?

21. How much heat is released when 50.0 g of steam cools to 40oC?

22. How much energy is required to completely change 25 g of liquid water at 25 degrees Celsius to steam?

23. 13.3 Liquids
Kinetic molecular theory also applies to liquids and solids
Must take intermolecular forces into account to apply it

24. Much denser than gasses
Due to intermolecular forces holding particles together
Fluidity – both gases and liquids are classified as fluids because they can flow and diffuse
Liquids diffuse more slowly because intermolecular attractions interfere with the flow

25. Viscosity - measure of the resistance of a liquid to flow
Attractive forces – stronger intermolecular forces = higher viscosity
Particle size – larger molecules = higher viscosity
Temperature – lower temperature = higher viscosity

26. Surface tension – the energy required to increase the surface area of a liquid by a given amount
Caused by intermolecular forces pulling down on the particles on the surface of a liquid which stretches it tight like a drum

27. Cohesion – force of attraction between identical molecules
Adhesion – force of attraction between molecules that are different

28. Evaporation
To escape into the vapor phase a molecule must have enough energy to overcome the intermolecular forces of the liquid
Rate increases as temperature increases
Why does sweating cool you down?

29. Vapor pressure - the pressure exerted by a vapor over a liquid
In a closed container will eventually reach equilibrium

30. https://www.youtube.com/watch?v=JsoawKguU 6A

31. Boiling point – temperature at which vapor pressure = atmospheric pressure

33. Predict which substance in each of the following pairs would have the largest vapor pressure
H2O vs. CH3OH
Ch3OH vs. Ch3Ch2Ch2oh

34. 13.4 Solids
Solid particles have as much kinetic energy as liquids or gasses but much stronger attractive forces between particles
Limit the motion of particles to vibrations

35. Density of solids – almost always greater than density of liquids
Exception = water

36. Crystalline solids – solid whose atoms, ions, or molecules are arranged in an orderly, geometric structure

38. Classify the type of solid:

39. Ionic solids Made of cation + anion
Each ion is surrounded by ions of the opposite charge
High melting point
Brittle
Held together by strong forces between ions

40. Molecular solids Fundamental particle is a molecule (relatively small)
Melt at relatively low temperatures
Held together by weak intermolecular forces

41. Atomic solid Fundamental particle is the atom Properties vary greatly
Group low melting points
Diamond - very high melting point

42. Metallic Bonding
Metals are held together by nondirectional covalent bonds (called the electron sea model) among the closely packed atoms

43. Amorphous solids
Particles are not arranged in a regular, repeating pattern
Does not contain crystals
Forms when molten material cools too quickly for crystals to form
Glass
Rubber
Some plastics

44. Name the type of crystalline solid formed by each of the following:
Ammonia
Iron
Cesium fluoride
Argon
sulfur

45. Which of the following is most likely to be solid at room temperature?
Bao Hf O2