Intermolecular forces

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. 431 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 • 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 ForcesLondon 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 ForcesLondon 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 ForcesHydrogen 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 • 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 forcesbind molecules to each other. • Adhesive forcesbind molecules to a surface • Meniscusis 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 • CsClStructure • Cs+ has a coordination number of 8. • Cationto anion ratio is 1:1. • Zinc BlendeStructure (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.