Chapter 5. LIQUIDS AND SOLIDS

1. Chapter 5. LIQUIDS AND SOLIDS INTERMOLECULAR FORCES 5.1 The Origin of Intermolecular Forces 5.2 Ion-Dipole Forces 5.3 Dipole-Dipole Forces 5.4 London Forces 5.5 Hydrogen Bonding 5.6 Repulsions LIQUID STRUCTURE 5.7 Order in Liquids 5.8 Viscosity and Surface Tension 2012 General Chemistry I

2. INTERMOLECULAR FORCES (Sections 5.1-5.6) 5.1 The Origin of Intermolecular Forces • Phase: uniform in both chemical composition and • physical state - Condensed phase: simply a solid or liquid phase - Condensed phases form when attractive intermolecular forcesbetween molecules pull them together;repulsions dominate at even shorter separations.

3. - Intermolecular forces are weak compared with bonding forces, but boiling points and sublimation points depend on their strength. - All interionic and almost all intermolecular forces can be traced to the coulombic interaction between charges.

4. - Distance dependence of potential energy of interaction 1/r : between ions (ionic bonding) 1/r2 : between ions and dipoles 1/r3 : between stationary dipoles 1/r6 : between rotating dipoles

5. TABLE 5.1 Interionic and Intermolecular Interactions

6. 5.2 Ion-Dipole Forces • This is the attractive force between ions and polar molecules in liquid or solid phase. • Hydration: attachment of water molecules to ionic solute particles is an example of ion-dipole interaction. H2O

7. The potential energy of ion-dipole interactions (~15 kJ mol-1) z = the charge number of the ion m = the electric dipole moment of the polar molecule

8. Water of crystallization: smaller and highly charged cations strongly attract polar water molecules in the solid phase. - Note hydrated salts of Li and Na vs. anhydrous salts of K, Rb, Cs, and NH4+ E.g. Na2CO3.10H2O compared with K2CO3 (Na+; 102 pm, K+; 138 pm) - Note BaCl2 · 2H2O vs. anhydrous KCl (Ba2+; 136 pm, K+; 138 pm)

9. 5.3 Dipole-Dipole Forces • This is the attractive force between polar molecules. • Between stationary polar molecules in the liquid phase (~2 kJ mol-1) • Between rotating polar molecules in the gas phase (~0.3 kJ mol-1)

11. 5.4 London Forces • London force (induce dipole-induced dipole force, dispersion force) ~2 kJmol-1: it exists between all molecules but is the only interaction between nonpolar molecules. - Attractive interactions due to instantaneous fleeting dipole moments - Fluctuation of the electron distribution in one molecule→ temporary dipole → second temporary dipole in the other molecule → ··· Time

12. - Potential energy (strength) of the London interaction is given by • Polarizability a is the ease with which molecular electron clouds can be distorted: a increases with number of electrons. - A large linear molecule is more likely to have stronger London interactions (and hence a higher boiling point) than a smaller or nonlinear one. Examples Alkanes C5H12; mobile liquid C15H32; viscous liquid C18H38; waxy solid

13. - Halogens: gases (F2, and Cl2); liquid (Br2); solid (I2) - Rod-shaped (pentane; Tb = 36 oC) vs. spherical (2,2-dimethylpropane; Tb = 10 oC)

14. TABLE 5.2 Melting and Boiling Points of Substances

15. Allied intermolecular interactions • Dipole-induced dipole interaction between a polar molecule and a nonpolar molecule (~2 kJ mol-1) – Van der Waals interactions is the collective name for dipole-dipole forces between rotating polar molecules, London forces, and dipole-induced dipole forces

16. EXAMPLE 5.2 Explain the trend in the boiling points of the hydrogen halides: HCl, -85 oC; HBr, -67 oC; HI, -35 oC. - Electronegativity differences: HCl > HBr > HI - Number of electrons and London forces: HCl < HBr < HI → not by dipole-dipole forces, but by London forces

17. 5.5 Hydrogen Bonding Some compounds are characterized by exceptionally high Tb due to hydrogen bonding: examples include NH3, H2O, HF – see Fig. 5.9. Hydrogen bonding London forces

18. Hydrogen bonding: an attraction in which a hydrogen atom bonded to a small, strongly electronegative atom, specifically N, O, or F, is attracted to a lone pair of electrons on another N, O, or F atom. Intermolecular and intramolecule types exist. - strong electrostatic interaction ~20 kJ mol-1 O…….H-O linear but asymmetric (101 pm vs 175 pm)

19. - Hydrogen fluoride, (HF)n - Acetic acid dimer (vapor) - DNA double helix - Protein folding

21. 179s Example5.1 Identify the kinds of intermolecular forces that might arise between molecules of each of the following substances: (a) NH2OH; (b) CBr4; (c) H2SeO4; (d) SO2 Solution

22. 179s Example5.5 Suggest, giving reasons, which substance in each of the following pairs is likely to have the higher normal melting point (Lewis structures may help your arguments): • HCl or NaCl; (b) C2H5OC2H5 (diethyl ether) or C4H9OH (butanol); • (c) CHI3 or CHF3; (d) C2H4 or CH3OH. Solution

23. 5.6 Repulsions • Intermolecular repulsions arise from the overlap of orbitals on neighboring molecules and the requirements of the Pauli exclusion principle. • - They are important only at very short distances:

24. LIQUID STRUCTURE (Sections 5.7-5.8) 5.7 Order in Liquids - The liquid phase lies between the extremes of the gas and solid phases. gas phase: moving with almost complete freedom minimal intermolecular forces solid phase: locked in place by intermolecular forces oscillate around an average location

25. - In the liquid phase, molecules have short-range order but not long-range order. - Water loses only 10% of hydrogen bonds upon melting and the rest are continuously broken and reformed.

26. 5.8 Viscosity and Surface Tension • Viscosity: resistance to flow, indication of the intermolecular force strength - Water and glycerol: very viscous due to hydrogen bonding - Hydrocarbon oils and grease: viscous due to tangling long chains - Viscosity usually decreases with temperature due to higher energy of molecules.

27. Viscosities of common liquids Linear alkane chains in Heavy hydrocarbon oil

28. Surface tension: the net inward pull, an indication of the intermolecular force strength - Water: three times larger than many other liquids, due to hydrogen bonds - Mercury: more than six times that of water, partially covalent

29. Wetting: strong interactions of water with the materials’ surface. Water maximizes its contact with the materials by hydrogen bonding. • Capillary action: adhesive forces between a liquid and surface vs. cohesive forces within the liquid H2O Hg – Meniscus: indication of the relative strength of adhesion and cohesion

30. 181s Example5.21 Predict how each of the following properties of a liquid varies as the strength of intermolecular forces increases and explain your reasoning: (a) boiling point; (b) viscosity; (c) surface tension. Solution

31. 181s Example5.23 Predict which liquid in each of the following pairs has the greater surface tension: (a) cis-dichloroethene or trans-dichloroethene; (b) benzene at 20 oC or benzene at 60 oC. Solution

32. 181s Example5.25 Rank the following molecules in order of increasing viscosity at50 oC: C6H5SH, C6H5OH, C6H6. Solution

33. Chapter 5. LIQUIDS AND SOLIDS SOLID STRUCTURES 5.9 Classification of Solids 5.10 Molecular Solids 5.11 Network Solids 5.12 Metallic Solids 5.13 Unit Cells 5.14 Ionic Structures THE IMPACT ON MATERIALS 5.15 Liquid Crystals 5.16 Ionic Liquids 2012 General Chemistry I

34. SOLID STRUCTURES (Sections 5.9-5.14) 5.9 Classification of Solids • Crystalline solid: a solid in which the atoms, ions, or molecules lie in an orderly array with crystal faces • Amorphous solid: one in which the atoms, ions, or molecules lie in a random jumble amorphous silica quartz

35. Classification of Crystalline Solids According to the bonds that hold their atoms, ions, or molecules in place: Metallic: consisting of cations held together bya sea of electrons Ionic: built from the mutual attractions of cations and anions Molecular: assemblies of discrete molecules held in place by intermolecular forces Network: consisting of atoms covalently bonded to their neighbors throughout the extent of the solid

37. 182s Example5.35 Classify each of the following solids as ionic, network, metallic, or molecular: (a) quartz, SiO2; (b) limestone, CaCO3; (c) dry ice, CO2; (d) sucrose, C12H22O11; (e) polyethylene, a polymer with molecules consisting of chains of thousands of repeating –CH2CH2- units. Solution

38. 5.10 Molecular Solids Molecular solids consist of molecules held together by intermolecular forces; physical properties depend on the strengths of those forces. Amorphous molecular solids: as soft as paraffin wax Crystalline molecular solids: - sucrose: numerous hydrogen bonds between OH groups account for high melting point at 184 oC - ultrahigh-density polyethylene: smooth yet tough

39. Molecular Solids and Liquids: Melting and Freezing Most substances increase in density on freezing: water is an important exception. Ice at 0 oC is less dense than water at 0 oC due to a more open hydrogen-bonded structure. - ice water benzene

40. 184s Example5.33 Glucose, benzophenone (C6H5COC6H5), and methane are examples of compounds that form molecular solids. The structures of glucose and benzophenone are given here. (a) What types of forces hold these molecules in a molecular solid? (b) Place the solids in order of increasing melting point.

41. 184s Solution

42. 5.11 Network Solids Network solids are characterized by a strong covalent bond network throughout the crystal: they are very hard and rigid, with high Tm and Tb E.g. Two common allotropes (forms of an element that differ in the way in which the atoms are linked) of carbon have very different network structures. Diamond, with an sp3 hybrid s-bonding framework, is one of the hardest substances

43. Graphite has flat sheets of sp2 hybrid s-bonds with weak bonding between sheets. It conducts electricity parallel to the sheet, and is soft and slippery - Ceramic materials: noncrystalline inorganic oxides, great strength

44. 5.12 Metallic Solids In metallic solids, the cations are bound together by their interaction with the sea of the electrons that they have lost. • Close-packed structure: the spheres stack together with the least waste of space - Hexagonal close-packed structure (hcp): packed with the sequence of ABABAB··· Coordination number = 12 (3 plane below + 6 own plane + 3 plane above): this is the maximum. • Coordination number: the number of • nearest neighbors of each atom

45. - Cubic close-packed structure (ccp): packed with the sequence of ABCABC··· Coordination number of ccp = 12 Occupied space in a ccp:

46. Holes: the gaps (interstices) between the atoms in a crystal - Octahedral hole: a dip in a layer coincides with a dip in the next layer - Tetrahedral hole: a dip between three atoms is directly covered by another atom

47. 5.13 Unit Cells • Unit cell: the smallest unit that, when stacked together repeatedly without any gaps and without rotations, can reproduce the entire crystal. - Face centered cubic (fcc, cubic F) - Body centered cubic (bcc, cubic I) - Primitive cubic (cubic P) Cubic P Cubic I Cubic F

48. Bravais lattices: 14 basic patterns of unit cell in 3D crystalline systems; P = primitive; I = body-centered; F = face-centered; C = with lattice point on two opposite faces; R = rhombohedral

49. Unit cells are characterized by lengths a, b, c and angles a, b, g - Cubic unit cells - Primitive cubic (cubic P) - Body centered cubic (bcc, cubic I) - Face centered cubic (fcc, cubic F)