The Kinetic Molecular Theory of Liquids & Solids – Chapter 9

1. The Kinetic Molecular Theory of Liquids & Solids– Chapter 9 • Distance between gas molecules are so great at ordinary temperatures and pressures (25 *C and 1atm) that there is no real interaction between gas molecules. • Liquids – the molecules are so close together that there is little empty space. Allowing for a definite volume but taking the shape of it’s container. • Solids – molecules are held rigidly in a position with virtually no freedom of motion. So that they have a definite volume and shape.

3. Intermolecular Forces • These are attractive forces between molecules: • Dispersion (London) Forces • Dipole-Dipole Forces • Hydrogen Bonding • These forces are partly responsible for the non-ideal gas law behavior discussed earlier. And these forces are why liquids and solids do not present “ideal behavior”. • Keep in mind that intramolecular forces are forces within a molecule, and intermolecular forces are between molecules.

4. Intermolecular Forces…………(cont.) • Boiling points and melting points often reflect the strength of these intermolecular forces. • Molecular substances tend to have the following characteristics: • Non-conductors of electricity when pure, for example pure water and ethyl alcohol are 2 examples of molecular substances that, when pure, will not conduct electricity. • Generally, are insoluble in water but soluble in non-polar solvents. • They tend to have low melting and boiling points. The stronger the intermolecular force the higher the boiling or melting point.

5. 1. Dispersion (London) Forces: • The attractive forces that arise as a result of temporary dipoles induced in atoms or molecules. • This happens when something (cation or polar compound) “throws off” the distribution of electrons in an atom or molecule. • These forces can only occur in non-polar molecules. • As molar mass (MM) increases the dispersion forces become stronger, consequently the boiling point of non-polar molecules tend to increase with MM. • Reference: Table 9.2 – page 234.

7. 2. Dipole-Dipole Forces: • These are attractive forces between polar molecules. • The larger the dipole moment then the greater the dipole force. • Polar molecules tend to have slightly higher boiling and melting points than non-polar substances of similar MM. • Reference: Table 9.3 – page 235.

9. 3. Hydrogen Bonding: • This a special, strong type of dipole-dipole interaction between hydrogen and polar bond in elements nitrogen, oxygen, and fluorine. Only between N—H, O—H, and F—H bonds have hydrogen bonding. • These types of bonds tend to not follow the relationship between MM and boiling points. • Reference: Table 9.4 – page 237. • *It is important to remember that the 3 intermolecular forces mentioned above are relatively weak in comparison to the covalent bond within a molecule.

10. Polarizability • The ease with which the electron cloud of a particle can be distorted is called its polarizability. Smaller atoms (or ions) are less polarizable than larger ones because their electrons are closer to the nucleus and therefore are held more tightly. Thus, we observe several trends: • Polarizability increases down a group because atomic size increases, so the larger electron clouds are farther from the nucleus and, thus, easier to distort. • Polarizability decreases from left to right across a period because the increasing Zeff shrinks atomic size and holds the electrons more tightly. • Cations are less polarizable than their parent atoms because they are smaller; anions are more polarizable because they are larger.

13. Example • What types of intermolecular forces exist between the following pairs: a. Hydrogen bromide and hydrogen sulfide b. Diatomic chlorine and carbon tetrabromide c. Diatomic iodine and the nitrate ion d. ammonia and benzene

14. Example 9.5 – page 237 • Would you expect to find hydrogen bonds in: • Acetic acid • Diethyl ether • Hydrazine

15. Example 9.6 – page 238 • What types of intermolecular forces are present in: • nitrogen • chloroform • carbon dioxide • ammonia

16. Properties of Liquids • Intermolecular forces give rise to a number of structural features and properties of liquids. Notable, two features: surface tension and viscosity. • Surface Tension: the amount of energy required to stretch or increase the surface of a liquid be a unit area (for example 1 cm2). • Liquids that have strong intermolecular forces have high surface tensions. So because of hydrogen bonding, water has a greater surface tension than most other liquids. • Viscosity: a measure of a fluids resistance to flow (unit = N * s/m2). • Liquids that have strong intermolecular forces have higher viscosities than those with weaker intermolecular forces. • Examples: Water = 1.01 x 10-3 Acetone = 4 x 10-3 Blood = 3.16 x 10-4 Glycerol= 1.49

17. Network Covalent Solids • Most covalent solids are molecular however some form a network of repeating patterns, these are called network covalent solids. • Plastics, and allotropes of carbon (diamond and graphite) do this. • They tend to have high melting points compared to molecular covalent.

18. Ionic Solids • Consist of anions and cations that are held in a regular repeating arrangement by strong ionic bonds, or electrostatic interactions. • Ionic solids tend to have high boiling and melting points. • Ionic solids do not conduct electricity, because the ions are fixed. • Many are soluble in water. Would solutions of ionic solids be good or poor conductors of electricity. • The strength of the ionic bond depends on 2 things: • The charge of the ion, CaO has a stronger ionic bond than NaCl • The size of the ions: d = rcation + ranion

19. Metallic Solids • There is no real “bond” between metals. Basically the electrons of metals arrange around each other’s nuclei. This forms a “sea” of electrons, so that the electrons move about in the crystal. • They are good conductors of electricity, have high thermal conductivity, tend to have luster (reflect light), are ductile and malleable, and are generally insoluble in water.

20. Example 9.7- page 243 • Reference: Table 9.5- page 244 • For each species in a column A, choose the description that best applies in column B.

21. Crystal Structures • Ionic solids tend to crystallize in definite geometric forms. These geometric forms are made up of unit cells (the smallest structural unit in the 3-D repeating pattern). • There are 14 possible crystalline structures but we will focus on the main 3.

22. 3 Main Crystal Structures • Simple Cubic (SC) • Face Centered Cubic (FCC) • Body Centered Cubic (BCC)

23. How they differ… • Number of atoms per unit cell SC FCC BCC • The relation between side of cell, s (in nm), and the atomic radius, r. This relation is expressed in the following equation and offers and experimental way to determine the atomic radius of a metal: SC: 2r = s BCC: 4r = s1/3 FCC: 4r = s21/2 • The percentage of empty space – the greater the amount of empty space the more unstable the structure. • Reference table 9.6 – page 245

25. Example 9.8 – page 246 • Silver crystallizes with a face centered cubic unit cell 0.407 nm on an edge. Calculate the atomic radius of silver.

26. A Vapor • A gas is a substance that is normally in the gaseous state at ordinary temperatures and pressures. • A vapor is the gaseous form of a substance that is a liquid or a solid at normal temperature and pressure – generally 25*C and 1 atm.

27. Vapor Pressure • The pressure exerted by the vapor over the liquid remains constant (in a sealed container the rate of condensation becomes equal to the rate of evaporation). • Diagram: • Thus there is a state of dynamic equilibrium between the liquid and the vapor. • So as long as the pressure exerted by the liquid is less than that of the atmosphere than the liquid will not boil. • Diagram: • Once the vapor pressure exerted by the liquid is equal to that of the atmosphere the liquid boils (this is considered the normal boiling point) • Diagram:

29. Vapor Pressure………………..(cont.) • Vapor pressure always increases with temperature. The higher the average kinetic energy then the more collisions and the stronger the force exerted by the liquid (the vapor pressure). • What pressure would be required to boil water at a temperature of 70*C if water exerts a pressure of 24mmHg? • The stronger the intermolecular forces in the liquid the lower the vapor pressure of that liquid. Examples: Water – 24 mmHg Ether – 537 mmHg • Why?

30. The Clausius-Clapeyron Equation: • Refer to page 227 • R = 8.31 J/mol K • T = temperature in K • P = pressure in atm or mmHg

31. Example 9.2 – page 228 • Benzene has a vapor pressure of 183 mmHg at 40 *C. Taking its heat of vaporization to be 30.8 kJ/mol. Calculate its vapor pressure at 25 *C.

32. Important Terms • Critical temperature – the temperature above which the liquid phase of a pure substance cannot exist. Basically, above this temperature the substance is a gas and below it the substance is a liquid. • Reference: Table 9.1 – page 230 • Critical pressure – conversely, this is the pressure that must be applied at the critical temperature to cause condensation. • Supercritical Fluid – just means that the substance is above the critical temperature (refers to a gas). • Melting point – pressure has little effect on melting point. Most solids are denser than their liquid form. There is a very important exception to the aforementioned, guess what it is?

33. Heating Curve

34. Triple Point Diagrams