Intermolecular Forces, Liquids, and Solids

1. Intermolecular Forces, Liquids, and Solids Seneca Valley SHS AP Chemistry Chapter 10

2. Intermolecular Forces • The attractive forces holding solids and liquids together are called intermolecular forces. The covalent bond holding a molecule together is an intramolecular force • Intermolecular forces are much weaker than intramolecular forces (chemical bonds) (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl). • When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). • When a substance condenses intermolecular forces are formed.

3. Intermolecular Forces

4. Intermolecular Forces Dipole-Dipole Forces • Molecules orient themselves to maximize the + --- - interactions and minimize the + --- + and - --- -forces. • There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. • If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

5. Intermolecular Forces London Dispersion Forces • Weakest of all intermolecular forces. • Exist primarily between noble gases and nonpolar molecules • The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). • For an instant, the electron clouds become distorted. • In that instant a dipole is formed (called an instantaneous dipole). • Relatively weak and short-lived

6. Intermolecular Forces London Dispersion Forces

7. Intermolecular Forces London Dispersion Forces • One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom). • Polarizability is the ease with which an electron cloud can be deformed. • London dispersion forces exist between all molecules. • The greater the surface area available for contact, the greater the dispersion forces. • London dispersion forces between spherical molecules are lower than between sausage-like molecules.

8. Intermolecular Forces London Dispersion Forces The larger the molecule (the more electrons it contains), the stronger the dispersion force becomes.

9. Intermolecular Forces Hydrogen Bonding • Special case of dipole-dipole forces. • By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. • Intermolecular forces are abnormally strong. • H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N). • Electrons in the H-X (X = electronegative element) lie much closer to X than H. • H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X. • Therefore, H-bonds are strong.

10. Intermolecular Forces Hydrogen Bonding

11. A Molecular Comparison of Liquids and Solids

12. Intermolecular Forces Hydrogen Bonding • Hydrogen bonds are responsible for: • Ice Floating • Solids are usually more closely packed than liquids; • therefore, solids are more dense than liquids. • Ice is ordered with an open structure to optimize H-bonding. • Therefore, ice is less dense than water. • In water the H-O bond length is 1.0 Å. • The O…H hydrogen bond length is 1.8 Å. • Ice has waters arranged in an open, regular hexagon. • Each + H points towards a lone pair on O. • Ice floats, so it forms an insulating layer on top of lakes, rivers, etc. Therefore, aquatic life can survive in winter.

13. Intermolecular Forces Hydrogen Bonding Hydrogen bonds are responsible for: • Protein Structure • Protein folding is a consequence of H-bonding. • DNA Transport of Genetic Information

14. Intermolecular Forces Comparing Intermolecular Forces

15. Some Properties of Liquids Viscosity • 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 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.

16. Some Properties of Liquids Surface Tension

17. Bonding in Solids • There are four types of solid: • 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.

18. Bonding in Solids

19. Bonding in Solids Molecular Solids • Intermolecular forces: dipole-dipole, London dispersion and H-bonds. • Weak intermolecular forces give rise to low melting points. • Room temperature gases and liquids usually form molecular solids at low temperatures.

20. Bonding in Solids Covalent Network Solids • Intermolecular forces: dipole-dipole, London dispersion and H-bonds. • Atoms held together in large networks. • Examples: diamond, graphite, quartz (SiO2), silicon carbide (SiC), and boron nitride (BN).

21. Bonding in Solids Covalent Network Solids

22. Bonding in Solids Covalent Network Solids • In diamond: • each C atom is tetrahedral; • there is a three-dimensional array of atoms. • Diamond is hard, and has a high melting point (3550 C). • In graphite: • each C atom is arranged in a planar hexagonal ring; • layers of interconnected rings are placed on top of each other; • the distance between layers is large (3.41 Å);

23. Bonding in Solids Ionic Solids

24. Bonding in Solids Ionic Solids • NaCl Structure • Face-centered cubic lattice. • Cation to anion ratio is 1:1. • Examples: LiF, KCl, AgCl and CaO • CsCl Structure • Different from the NaCl structure (Cs+ is larger than Na+). • Cation to anion ratio is 1:1.

25. Bonding in Solids Metallic Solids

26. Phase Changes • Surface molecules are only attracted inwards towards the bulk molecules. • Sublimation: solid  gas. • Vaporization: liquid  gas. • Melting or fusion: solid  liq. • Deposition: gas  solid. • Condensation: gas  liquid. • Freezing: liquid  solid. Energy Changes Accompanying Phase Changes • Sublimation: Hsub > 0 (endo) • Vaporization: Hvap > 0 (endo) • Melting or Fusion: Hfus > 0 (endo) • Deposition: Hdep < 0 (exo) • Condensation: Hcon < 0 (exo) • Freezing: Hfre < 0 (exo)

27. Phase Changes Energy Changes Accompanying Phase Changes • All phase changes are possible under the right conditions (e.g. water sublimes when snow disappears without forming puddles). • The sequence heat solid  melt  heat liquid  boil  heat gas is endothermic. • The sequence cool gas  condense  cool liquid  freeze  cool solid is exothermic.

28. Phase Changes Energy Changes Accompanying Phase Changes

29. Phase Changes 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.

30. Vapor Pressure Explaining Vapor Pressure on the Molecular Level • Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid. • These molecules move into the gas phase. • As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. • After some time the pressure of the gas will be constant at the vapor pressure.

31. Vapor Pressure Explaining Vapor Pressure on the Molecular Level • 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.

32. Vapor Pressure 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.

33. Vapor Pressure Vapor Pressure and Boiling Point • Liquids boil when the external pressure equals the vapor pressure. • Temperature of boiling point increases as pressure increases. • Two ways to get a liquid to boil: increase temperature or decrease pressure. • Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked. • Normal boiling point is the boiling point at 760 mmHg (1 atm).

34. Phase Diagrams • Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases. • Given a temperature and pressure, phase diagrams tell us which phase will exist. • Features of a phase diagram: • Triple point: temperature and pressure at which all three phases are in equilibrium. • Vapor-pressure curve: generally as pressure increases, temperature increases. • Critical point: critical temperature and pressure for the gas. • Normal melting point: melting point at 1 atm.

35. Phase Diagrams • Any temperature and pressure combination not on a curve represents a single phase.

36. Phase Diagrams The Phase Diagrams of H2O and CO2 • Triple point at 0.0098C and 4.58 mmHg. • Normal melting (freezing) point is 0C. • Normal boiling point is 100C. • Critical point is 374C and 218 atm • Triple point at -56.4C and 5.11 atm. • Normal sublimation point is -78.5C. • (At 1 atm CO2 sublimes it does not melt.) • Critical point occurs at 31.1C and 73 atm.

37. Intermolecular Forces, Liquids, and Solids End of Chapter 10