Intermolecular Forces, Liquids, and Solids

1. Intermolecular Forces, Liquids, and Solids Chapter 11

2. States of Matter • What are the 3 states of matter? • What physical properties do we know about them? • Liquids and solids are similar compared to gases… • Incompressible • Density doesn’t change with temperature • Because solids and liquids are molecules that are so close together while gases are so far apart. • What holds them so close together?

3. Intermolecular Forces • Physical Properties of solids and liquids is due largely to Intermolecular Forces(IMFs). • IMFs are forces that exist between molecules within a solid or a liquid • There are many different types of IMFs. • By understanding the nature & strength of IMFs we can relate composition & structure to physical properties.

4. Types of IMFs • Strong Intermolecular Forces: • Covalent Bonding • Ionic Bonding • Weak Intermolecular Forces: • Dipole-Dipole • Hydrogen “bonding” • London Dispersion Forces • During phase changes, molecules stay intact… NRG used to overcome forces

5. What is a dipole? Dipole – Dipole Forces: • Attraction between molecules with dipole moments. • Maximizes + to – interactions and minimizes + to + or – to – interactions. • 1% the strength of ionic bonds • Force strength increases as polarity increases • Not important for gases… why? Ion – Dipole Force: • Force between Ion and dipole of a polar molecule. • Weaker than Dipole – Dipole forces.

6. + + - - - + + - + - - + + - - + - + + -

7. Hydrogen Bonding • Special dipole-dipole attraction • Hydrogen covalently bonded to highly electronegative elements • Especially strong with: N, O, F (because they have high electronegativity and they are small enough to get close) • Stronger than other dipole-dipole attractions • Important for bonding of Water and DNA • Effects Boiling Points

8. H2O HF H2Te H2Se NH3 SbH3 H2S HI AsH3 HCl HBr PH3 SnH4 GeH4 SiH4 CH4 100 Boiling Points 0ºC -100 200

9. Water d+ d- d+

10. London Dispersion Forces • Instantaneous dipoles • Electrons move randomly, and creates a momentary nonsymmetrical distribution of charge (even in nonpolar molecules) • Induces short-lived dipole with neighboring molecules • Exist between ALL molecules, but are the weakest forces of attraction

11. d- d d+ d+ d- d+ H H H H H H H H H H H H Example

12. London Dispersion Forces • Weak and short lived • Last longer at Low Temp. • Polarizabliity – ability to distort the electron cloud; “squishiness” of the molecule. • increases with the number of electrons in a molecule… CCl4 experiences greater LDFs than CH4. • Larger molecules = more LDF = higher melting and boiling points

13. IMF Summary • London Dispersion Forces • Dipole-Dipole Forces • Hydrogen Bonding ALL 3 are sometimes referred to as “Van der Waals Forces” • See flow chart on pg 417 to review what types of forces are involved with different types of bonds. Increasing strength

14. The Liquid State • Many of the properties of liquids are due to the internal attraction of atoms: • Surface Tension • Beading • Capillary Action • Viscosity • Stronger IMFs cause each of these to increase.

15. Surface tension • Molecules in the middle are attracted in all directions. • Molecules at the top are only pulled inside. • Minimizes surface area. • High IMF means • Higher surface tension

16. Capillary Action • Liquids spontaneously rise in a narrow tube. • IMFs are cohesive (connect like things) • Adhesive forces form between two separate “like” objects (polar molecules attract to polar bonds in material making up container)

18. Beading • If a polar substance is placed on a non-polar surface. • There are cohesive, • But no adhesive forces. • And Visa Versa H2O Wax Paper

19. Viscosity • Measure of a liquid’s resistance to flow. • Increases with IMFs • Increases with molecular size.

20. Liquids… a structural model • Strong IMFs (like solids) • Considerable molecular motion (like gases)

21. Phase Changes • Phase of a substance determined by… • Temperature • Pressure • Strength of IMFs • Stronger Forces = Bigger Effect • Hydrogen Bonding • Polar Molecule (dipole – dipole) • London Dispersion Forces Decreasing Strength

22. Types of Phase Changes: • Melting / Freezing • Vaporization / Condensation • Sublimation / Deposition We’ll go more into the energy changes, heating curves, and phase diagrams later in this unit…

23. Solids • Two major types. • Amorphous- those with much disorder in their structure. • Crystalline- have a regular arrangement of components in their structure.

24. Crystals • Lattice- a three dimensional grid that describes the locations of the pieces in a crystalline solid. • Unit Cell-The smallest repeating unit in of the lattice. • Three common shapes: • Cubic • Body Centered Cubic • Face Centered Cubic

25. Cubic

26. Body-Centered Cubic

27. Face-Centered Cubic

28. Solids Amorphous Solids: • Components “frozen in place” and lacking orderly arrangement. Ex: Glass and plastic We’ll focus more crystalline solids…

29. Crystalline Solids: • Ionic Solids have ions at lattice points (Salt) • Molecular Solids have molecules (Sugar) • Metallic Solids (Cu, Fe, Al, Pt…) • Covalent Network Solids (Diamonds, Quartz)

30. The book drones on about • Using diffraction patterns to identify crystal structures. • Talks about metals and the closest packing model. • It is interesting, but trivial. • We need to focus on metallic bonding. • Why do metal atoms stay together. • How their bonding effects their properties.

31. Metallic bonding Empty Molecular Orbitals 3p Filled Molecular Orbitals 3s 2p 2s 1s Magnesium Atoms

32. The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around. Empty Molecular Orbitals 3p Filled Molecular Orbitals 3s 2p 2s 1s Magnesium Atoms

33. The 3s and 3p orbitals overlap and form molecular orbitals. Empty Molecular Orbitals 3p Filled Molecular Orbitals 3s 2p 2s 1s Magnesium Atoms

34. Electrons in these energy level can travel freely throughout the crystal. Empty Molecular Orbitals 3p l l Filled Molecular Orbitals 3s 2p 2s 1s Magnesium Atoms

35. This makes metals: -conductors -Malleable because the bonds are flexible. Empty Molecular Orbitals 3p l l Filled Molecular Orbitals 3s 2p 2s 1s Magnesium Atoms

36. Carbon- A Special Atomic Solid • There are three types of solid carbon. • Amorphous- coal uninteresting. • Diamond- hardest natural substance on earth, insulates both heat and electricity. • Graphite- slippery, conducts electricity. • How the atoms in these network solids are connected explains why.

37. Diamond- each Carbon is sp3hybridized, connected to four other carbons. • Carbon atoms are locked into tetrahedral shape. • Strong s bonds give the huge molecule its hardness.

38. Empty MOs Filled MOs Why is it an insulator? The space between orbitals make it impossible for electrons to move around E

39. Graphite is different. • Each carbon is connected to three other carbons and sp2 hybridized. • The molecule is flat with 120º angles in fused 6 member rings. • The p bonds extend above and below the plane.

40. This p bond overlap forms a huge p bonding network. • Electrons are free to move through out these delocalized orbitals. • The layers slide by each other.

41. Molecular solids. • Molecules occupy the corners of the lattices. • Different molecules have different forces between them. • These forces depend on the size of the molecule. • They also depend on the strength and nature of dipole moments.

42. Those without dipoles or H bonding. • Most are gases at 25ºC. • The only forces are London Dispersion Forces. • These depend on size of atom. • Large molecules (such as I2 ) can be solids even without dipoles. • Example: the Halogens

43. The Halogens Boiling points State at 20 oC Gas Gas Liquid Solid F 85 K 239 K 333 K 458 K Cl Br I

44. Those with dipoles. • Dipole-dipole forces are generally stronger than L.D.F. • Hydrogen bonding is stronger than Dipole-dipole forces. • No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds. • Stronger forces lead to higher melting and freezing points.

45. d+ H O H d+ Water is special • Each molecule has two polar O-H bonds. • Each molecule has two lone pair on its oxygen. • Each oxygen can interact with 4 hydrogen atoms.

46. d+ H O d+ H H d+ O d+ H H d+ O H d+ Water is special • This gives water an especially high melting and boiling point.

47. Ionic Solids • The extremes in dipole dipole forces-atoms are actually held together by opposite charges = IONIC BONDS • Huge melting and boiling points. • Atoms are locked in lattice so hard and brittle. • Every electron is accounted for so they are poor conductors-good insulators.

48. Vapor Pressure • Vaporization - change from liquid to gas at boiling point. • Evaporation - change from liquid to gas below boiling point • Heat of Vaporization (DHvap ) - the energy required to vaporize 1 mole at 1 atm.

49. Vaporization is an endothermic process - it requires heat. • Energy is required to overcome intermolecular forces. • Responsible for cool earth. • Why we sweat.

50. Condensation • Change from gas to liquid. • Achieves a dynamic equilibrium with vaporization in a closed system. • What is a closed system? • A closed system means matter can’t go in or out. • What is a “dynamic equilibrium?”