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1. 1 Intermolecular Forces and Liquids and Solids Chang, Chapter 11
2. 2 Phase Changes Many important properties of liquids and solids relate to the ease with which they change from one state to another.
Water evaporates from a glass.
An ice cube melts.
Solid CO2 (Dry Ice) sublimes.
These transformations are called phase changes or changes in state.
3. 3 Liquid-vapor equilibrium Evaporation is the process of transforming a liquid to a gas.
Evaporation depends on temperature.
The higher the temperature, the greater the KE of the molecules.
4. 4 Vapor Pressure Suppose we conduct an experiment in which we place a quantity of ethanol (C2H5OH) in an evacuated, closed container (a).
The ethanol will quickly begin to evaporate.
The pressure exerted by the vapor will begin to increase.
After a short time the pressure will attain a constant value, called the vapor pressure of ethanol (b).
5. 5 Vapor Pressure What happens at the molecular level during evaporation?
One-way traffic at first—liquid molecules move to the empty space and establish a vapor phase.
As the concentration of gaseous molecules increases, some condense (return to the liquid phase).
Dynamic equilibrium established.
6. 6 Vapor Pressure The molecules of a liquid move at various speeds.
At any instant some of the molecules on the surface have sufficient energy to escape.
The weaker the attractive forces, the larger the number of molecules that are able to escape and the higher the vapor pressure.
The movement of molecules from the liquid to the gas phase and vice versa goes on continuously in a dynamic equilibrium.
7. 7 Vapor Pressure
8. 8 Vapor Pressure A measure of the strength of intermolecular forces in a liquid is the molar heat of vaporation (DHvap).
Defined as the energy required to vaporize one mole of liquid.
9. 9 Vapor Pressure The relationship between vapor pressure (P) and temperature (T) can be described by the Clausius-Clapeyron equation:
P = pressure
T = absolute temperature
R = the gas constant (8.314 J/mol-K)
DHvap = the enthalpy of vaporization per mole
C = a constant
This equation predicts that a graph of ln P versus 1/T should be a straight line with slope of -DHvap/R.
10. 10 Vapor Pressure
11. 11 Boiling A liquid boils when its vp equals the external pressure acting on its surface.
At this point, bubbles of vapor are able to form within the interior of the liquid.
Normal boiling point is the boiling point of a liquid at 1 atm pressure.
The bp is related the the DHvap.
The greater the DHvap, the higher the bp.
Ultimately, both are determined by the strength of intermolecular forces.
12. 12 Vapor Pressure
13. 13 Critical temperature and pressure Above the critical temperature (Tc), a gas cannot be made to liquefy, no matter what its pressure.
This is the highest temp at which a substance can exist as a liquid.
KE is so high that molecules break away from their attractive forces.
Critical pressure (Pc) is the minimum pressure that must be applied to bring about liquefaction at the Tc.
14. 14 Critical temperature and pressure Tc and Pc values are determined by the strength of inter-molecular forces.
15. 15 Freezing and Melting A liquid solidifies (freezes) when its molecules have such low energy that they are unable to move past their neighbors.
The normal freezing point of a substance is the temperature at which it freezes at 1 atm. Liquid and solid are in equilibrium.
A solid melts when heat is applied, which allows molecules to overcome the attractive forces in the solid.
16. 16 Freezing and Melting
17. 17 Freezing and Melting Supercooling occurs when a liquid is temporarily cooled below its freezing point.
Occurs when heat is removed so rapidly that the molecules don’t have time to assume an ordered structure.
Stirring or addition of a seed crystal causes rapid solidification.
18. 18 Freezing and melting A measure of the strength of intermolecular forces in a liquid is the molar heat of fusion (DHfus).
Defined as the energy required to melt one mole of liquid.
19. 19 Solid-Vapor Equilibrium Sublimation is a process in which molecules go directly from the solid into the vapor phase (deposition is the reverse process).
Mothballs (naphthalene), iodine and carbon dioxide sublime at room temperature.
20. 20 Phase Diagrams A phase diagram is a graphical way to summarize the conditions under which equilibria exist between different states of matter.
Contains three important lines or phase boundaries:
A to B – vapor pressure curve
A to C – vapor pressure of the solid as it sublimes
A to D – change in melting point of the solid with increasing pressure
Contains two important points:
A – triple point
B – critical point
21. 21 Phase Diagrams
22. 22 Phase Diagrams Triple point
The point on a phase diagram representing the temperature and pressure at which the three phases of a substance coexist in equilibrium.
Critical point
The point at which the temperature and pressure have their critical values (Tc and Pc).
23. 23 Phase Diagrams
24. 24 Phase Diagrams
25. 25 Phase Diagrams
26. 26 You Should Know How To… Explain how DHvap and bp of a compound are related to the strength of its intermolecular forces.
Explain how dispersion forces arise and how they vary with the polarizability of an atom and the size and shape of a molecule.
Describe hydrogen bond and explain why they are stronger than other kinds of intermolecular forces.
Describe how viscosity and surface tension vary with temperature and the strength of intermolecular forces.
Distinguish metals, ionic solids, covalent network solids, and molecular solids by their structures and by their properties.
Distinguish cubic and hexagonal close-packing.
27. 27 You Should Know How To… Distinguish simple cubic, face-centered cubic and body-centered cubic structures.
Define dynamic equilibrium and explain the relationship between the rate of evaporation and vapor pressure of a substance.
Interpret the major features of a phase diagram.
Describe why liquid water is denser than ice.
Predict the relative boiling points, vapor pressures, and enthalpies of vaporization of two substances from the strengths of their intermolecular forces.
Identify molecules that can experience hydrogen bonding.
28. 28 You Should Know How To… Give the coordination number of an atom of ion in a given crystal lattice.
Find the number of atoms or ions in a given unit cell.
Identify the stable phase or phases of a substance at a given temperature and pressure from its phase diagram.
Predict the relative densities of the liquid and solid phases of a substance from its phase diagram.
List the properties of water that are anomalous as a result of hydrogen bonding.
Describe the arrangement of cubic, hexagonal close-packed, primitive, and body-centered cubic structures.
