1. Intermolecular forces
Liquids and Solids
Chapter 11

2. Liquids vs. Solids
Physical properties are due to intermolecular forces
Understood in terms of kinetic-molecular theory
Gases are highly compressible and assume the shape and volume of their container
Liquids are almost incompressible, assume the shape but not the volume of the container
Solids are incompressible and have a definite shape and volume

3. Liquids vs. Solids Solids and liquids are condensed phases
Converting a gas into a liquid or solid requires the molecules to get closer to each other
Forces holding solids and liquids together are called intermolecular forces

4. Intermolecular forces

5. Intermolecular Forces
Attraction between molecules
Weaker than ionic or covalent bonds (16 kJ/mol vs kJ/mol for HCl)
Melting or boiling breaks intermolecular forces
Condensing forms intermolecular forces
Melting points / Boiling points reflect strength of intermolecular forces
High melting/boiling points indicates strong attractive forces

6. Intermolecular Forces
Van der Waals forces exist between neutral molecules
Includes London-dispersion forces, dipole-dipole forces, and hydrogen-bonding forces
Ion-dipole interactions are important in solutions
ALL are WEAK electrostatic interactions
(~15% as strong as a covalent or ionic bond)

7. Van der Waals Forces Ion-Dipole Dipole-Dipole
Interaction between an ion and the partial charge on the end of a polar molecule (dipole)
Important in formation of solution between ionic substances in polar liquids (ex. NaCl in water)
Dipole-Dipole
Exist between neutral polar molecules
Polar molecules attract each other
Need to be close together to form strong attractions
Weaker than ion-dipole forces

8. Van der Waals Forces London Dispersion Forces
Weakest of all intermolecular forces
Possible for two adjacent neutral molecules to affect each other
Nucleus of one molecule (atom) attracts the electrons in an adjacent molecule (atom)
Electron “clouds” become distorted – temporary
Temporary distortion creates an instantaneous dipole
One instantaneous dipole can create an instantaneous dipole in a nearby molecule (atom)
Temporary dipoles attract each other

9. Van der Waals Forces London Dispersion Forces
Molecules must be very close together for these attractive forces to occur
Polarizability is the ease with which an electron cloud can be deformed
The larger the molecule- the more polarizable it is
Forces increase as molecular weight increases
Forces depend on the shape of the molecule

10. Van der Waals Forces Hydrogen Bonds
Boiling points of compounds with hydrogen bonded to an electronegative atom are abnormally high
Special case of dipole-dipole interactions
Requires:
H bonded to a small electronegative element
An unshared pair of electrons on a nearby small electronegative atom/ion
Hydrogen only has one electron, so in an electronegative bond it is “electron bare”

11. Properties in liquids

12. Properties in Liquids Viscosity Surface Tension
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.
Surface molecules are only attracted inwards towards the bulk molecules

13. Surface Tension

14. Surface Tension
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
Meniscus is the shape of the liquid surface.
Adhesive > Cohesive : U-shaped meniscus (water)
Capillary Action: When a narrow glass tube is placed in water, the meniscus pulls the water up the tube

15. Phase changes

16. Phase Changes

17. Enthalpy of 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).

18. 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.

20. 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.

21. Vapor Pressure on a Molecular Level

22. 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.
Volatility, Vapor Pressure, and Temperature
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.

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

24. Phase diagrams

25. Phase Diagrams

26. Phase Diagrams Water vs. Carbon Dioxide

27. Solids

28. Unit Cells
Crystalline solid: well-ordered, definite arrangements of molecules, atoms or ions.
Crystals have an ordered, repeated structure.
The smallest repeating unit in a crystal is a unit cell.
Unit cell is the smallest unit with all the symmetry of the entire crystal.
Three-dimensional stacking of unit cells is the crystal lattice.

29. Unit Cell vs. Lattice

30. Three common types of unit cell.
Primitive cubic, atoms at the corners of a simple cube,
each atom shared by 8 unit cells;
Body-centered cubic (bcc), atoms at the corners of a cube plus one in the center of the body of the cube,
corner atoms shared by 8 unit cells, center atom completely enclosed in one unit cell;
Face-centered cubic (fcc), atoms at the corners of a cube plus one atom in the center of each face of the cube,
corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells.

31. Unit Cells

32. Solids: Four Types
Molecular (formed from molecules) - usually soft with low melting points and poor conductivity.
Covalent network (formed from atoms) - very hard with very high melting points and poor conductivity.
Ions (formed from ions) - hard, brittle, high melting points and poor conductivity.
Metallic (formed from metal atoms) - soft or hard, high melting points, good conductivity, malleable and ductile.

33. Covalent Network Solid

34. Ionic Lattice CsCl Structure Cs+ has a coordination number of 8.
Cation to anion ratio is 1:1.
Zinc Blende Structure (ZnS).
S2- ions adopt a fcc arrangement.
Zn2+ ions have a coordination number of 4.
The S2- ions are placed in a tetrahedron around the Zn2+ ions.
Fluorite Structure (CaF2).
Ca2+ ions in a fcc arrangement.
There are twice as many F- per Ca2+ ions in each unit cell.