Liquids and Solids

1. Liquids and Solids • Gas • low density • high compressibility • completely fills its container • Solid • high density • only slightly compressible • rigid • maintains its shape

2. Liquids and Solids • Liquids • properties lie between those of solids and gases • H2O(s) --> H2O(l) DHofus = 6.02 kJ/mol • H2O(l) --> H2O(g) DHovap = 40.7 kJ/mol • large value of DHvap suggests greater changes in structure in going from a liquid to a gas than from a solid to liquid • suggests attractive forces between the molecules in a liquid, though not as strong as between the molecules of a solid

3. Liquids and Solids • Densities of the three states of water • H2O(g) D = 3.26 x 10-4g/cm3 (400oC) • H2O(l) D = 0.9971 g/cm3 (25oC) • H2O(s) D = 0.9168 g/cm3 (OoC) • Similarities in the densities of the liquid and solid state indicate similarities in the structure of liquids and solids

4. Intermolecular Forces • Bonds are formed between atoms to form molecules • intramolecular bonding (within the molecule)

5. Intermolecular Forces • The properties of liquids and solids are determined by the forces that hold the components of the liquid or solid together • may be covalent bonds • may be ionic bonds • may weaker intermolecular forces between molecules

6. Intermolecular Forces • During a phase change for a substance like water • the components of the liquid or solid remain intact • the change of state is due to the changes in the forces between the components • e.g., H2O(s) --> H2O (l) …the molecules are still unchanged during the phase change

7. Dipole-Dipole Forces • Polar molecules • line up in an electric field • positive end of molecule will line up with the negative pole of the electric field while the negative end of the molecule will line up with the positive pole • can attract each other • positive end of one molecule will attract the negative end of another molecule

8. Dipole-Dipole Forces • Dipole-dipole forces • about 1% as strong as covalent or ionic bonds • become weaker with distance • unimportant in the gas phase

9. Hydrogen Bonding • A particularly strong dipole-dipole force • When hydrogen is covalently bonded to a very electronegative atom such as N, O, or F • Very strong due to • great polarity of the bond between H and the N, O or F • close approach of the dipoles due to H’s small size

10. Hydrogen Bonding • H-bonding has a very important effect on physical properties • For example, boiling points are greater when H-bonding is present

11. London Dispersion Forces • aka Van der Waals forces • Nonpolar molecules must exert some kind of force or they would never solidify

12. London Dispersion Forces • London dispersion forces (LDF) • due to an instantaneous dipole moment • created when electrons move about the nucleus • a temporary nonsymmetrical electron distribution can develop (I.e., all the electrons will shift to one side of the molecule)

13. London Dispersion Forces • The instantaneous dipole moment can induce an instantaneous dipole moment in a neighboring molecule, which could induce another instantaneous dipole moment in a neighboring molecule, etc. (like a “wave” in the stands of a football game)

14. London Dispersion Forces • The LDF is very weak and short-lived • To form a solid when only LDF exists requires very low temperatures • the molecules or atoms must be moving slowly enough for the LDF to hold the molecules or atoms together in a “solid” unit

15. London Dispersion Forces • Element Freezing Point (oC) Helium -269.7 Neon -248.6 Argon -189.4 Krypton -157.3 Xenon -111.9

16. London Dispersion Forces • Notice that as the MM of the noble gas increases, the freezing point increases • This implies that the LDF between the atoms is stronger as the MM increases • Large atoms with many electrons have an increased polarizability (the instantaneous dipole would be larger), resulting in a larger London Dispersion Force between the atoms than between smaller atoms

17. The Liquid State • Properties of liquids • low compressibility • lack of rigidity • high density (compared to gases)

18. The Liquid State • Surface Tension • results in droplets when a liquid is poured onto a surface • depends on IMF’s

19. The Liquid State • Molecules at the surface experience an uneven pull, only from the sides and below. Molecules in the interior are surrounded by IMF’s • Uneven pull results in liquids assuming a shape with minimum surface area • Surface tension is a liquids resistance to an increase in surface area. • Liquids with high IMF’s have high surface tensions

20. The Liquid State • Capillary Action • Exhibited by polar molecules • The spontaneous rising of a liquid in a narrow tube • due to two different forces involving the liquid

21. The Liquid State • Cohesive forces - IMF between the liquid molecules • Adhesive forces - forces between the liquid molecules and the polar (glass) container • adhesive forces tend to increase the surface area • cohesive forces counteract this • Concave meniscus (water) - indicates adhesive forces of water towards the glass is greater than the cohesive forces between the water molecules. • Convex meniscus (nonpolar substances such as mercury) shows cohesive forces is greater than adhesive forces.

22. The Liquid State • Viscosity • Measure of a liquid’s resistance to flow • Depends on strength of IMF’s between liquid molecules • molecules with large IMF’s are very viscous • Large molecules that can get tangled up with each other lead to high viscosity

23. The Liquid State • So what does a liquid “look like?” • A liquid contains many regions where the arrangements of the components are similar to those of a solid • There is more disorder in a liquid than in a solid • There is a smaller number of regions in a liquid where there are holes present

24. Types of Solids • Ways to classify solids • Crystalline vs. Amorphous Solid • Crystalline solids • regular arrangement of components • positions of components represented by a lattice • unit cell - smallest repeating unit of the lattice

25. Types of Solids • three common unit cells exist • simple cubic • body centered cubic • face centered cubic

26. Types of Solids • Amorphous Solids • noncrystalline • glass is an example • disorder abounds

27. Types of Solids • X-ray diffraction • used to determine the structures of crystalline solids • diffraction occurs when beams of light are scattered from a regular array of points • obtain a diffraction pattern • Bragg equation: nl = 2d sinq

28. Types of Solids • Where n is an integer • l is the wavelength of the x-rays • d is the distance between the atoms • q is the angle of incidence and reflection • Use x-ray diffraction to determine bond lengths, bond angles, determine complex structures, test predictions of molecular geometry

29. Types of Solids • Example: • x-rays of wavelength 1.54 A were used to analyze an aluminum crystal. A reflection was produced at q = 19.3 degrees. Assuming n = 1, calculate the distance d between the planes of atoms producing the reflection. • (D = 2.33 A)

30. Types of Solids • Types of Crystalline Solids • Ionic Solids (e.g. NaCl) • Molecular Solids (e.g. C6H12O6) • Atomic Solids which include: • Metallic Solids • Covalent Network Solids

31. Types of Solids • Classify solids according to what type of component is found at the lattice point (of a unit cell) • Atomic Solids have atoms at the lattice points • Molecular Solids have discrete, relatively small molecules at the lattice points • Ionic solids have ions at the lattice points

32. Types of Solids • Different bonding present in these solids results in dramatically different properties • Element (atomic solid) M.P. (oC) Argon -189 C(diamond) 3500 Cu 1083

33. Structure and Bonding inMetals • Properties of Metals • high thermal conductivity • high electrical conductivity • malleability (metals can be pounded thin) • ductility (metals can be drawn into a fine wire) • durable • high melting points

34. Structure and Bonding inMetals • Properties are due to the nondirectional covalent bonding found in metallic crystals • Metallic crystal • contains spherical atoms packed together • atoms are bonded to each other equally in all directions

35. Structure and Bonding inMetals • Closest Packing • most efficient arrangement of these uniform spheres • Two possible closest packing arrangements • Hexagonal Closest Packed Structure • Cubic Closest Packed Structure

36. Structure and Bonding inMetals • Hexagonal Closest Packed Structure (hcp) • aba arrangement • First Layer • each sphere is surrounded by six other spheres

37. Structure and Bonding inMetals • Second Layer • the spheres do not lie directly over the spheres in the first layer • the spheres lie in the indentations formed by three spheres • Third Layer • the spheres lie directly over the spheres in the first layer

38. Structure and Bonding inMetals • Cubic Closest Packed Structure (ccp) • abc arrangement • First and Second Layers are the same as in hexagonal closest packed structure • Third Layer • the spheres occupy positions such that none of the spheres in the third layer lie over a sphere in the first layer

39. Structure and Bonding inMetals • Finding the net number of spheres in a unit cell • important for many applications involving solids (when I figure it out, I’ll let you know…or when it shows up on the ACS or AP test…then I’ll figure it out!)

40. Structure and Bonding inMetals • Examples of metals that are ccp • aluminum, iron, copper, cobalt, nickel • Examples of metals that are hcp • zinc, magnesium • Calcium and some other metals can go either way

41. Structure and Bonding inMetals • Some metals, like the alkali metals are not closest packed at all • may be found in a body centered cubic (bcc) unit cell where there are only 8 nearest neighbors instead of the 12 in the closest packed structures

42. Bonding Models for Metals • The model must account for the typical physical properties of metals • malleability • ductility • efficient and uniform conduction of heat and electricity in all directions • durability of metals • high melting points

43. Bonding Models for Metals • To account for these physical properties, the bonding in metals must be • strong • nondirectional • It must be difficult to separate atoms, but easy to move them (as long as the atoms stay in contact with each other

44. Bonding Models for Metals • Electron Sea Model (simplest picture) • Positive Metal ions (Metal cations) are surrounded by a sea of valence electrons • the valence electrons are mobile and loosely held • these electrons can conduct heat and electricity • meanwhile, the metal ions can move around easily

45. Bonding Models for Metals • Band Model or Molecular Orbital (MO) model • related to the electron sea model • more detailed view of the electron energies and motions

46. Bonding Models for Metals • MO model • electrons travel around the metal crystal in molecular orbitals formed from the atomic orbitals of the metal atoms • In atoms like Li2 or O2, the space between the energies of the molecular orbitals is relatively wide (big energy difference between the orbitals)

47. Bonding Models for Metals • However, when many metal atoms interact, the molecular orbital energy levels are very close together • Instead of separate, discrete molecular orbitals with different energies, the molecular orbitals are so close together in energies, that they form a continuum of levels, called bands

48. Bonding Models for Metals • Core electrons of metals are localized • the core electrons “belong” to a particular metal ion • The valence electrons of metals are delocalized • the valence electrons occupy partially filled, closely spaced molecular orbitals

49. Bonding Models for Metals • Thermal and Electrical conductivity • metals conduct heat and electricity because of highly mobile electrons • electrons in filled molecular orbitals get excited (from added heat or electricity) • these electrons move into higher energy, empty molecular orbitals

50. Bonding Models for Metals • Conduction electrons • the electrons that can be excited to empty MO’s • Conduction bands • the empty MO’s that can accept the conducting electrons