Chapter 12: Intermolecular Attractions and the Properties of Liquids and Solids

1. Chapter 12: Intermolecular Attractions and the Properties of Liquids and Solids Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

2. Intermolecular Forces • Important differences between gases, solids, and liquids: • Gases • Expand to fill their container • Liquids • Retain volume, but not shape • Solids • Retain volume and shape

3. Intermolecular Forces • Physical state of molecule depends on • Average kinetic energy of particles • Recall KE  Tave • Intermolecular Forces • Energy of Inter-particle attraction • Physical properties of gases, liquids and solids determined by • How tightly molecules are packed together • Strength of attractions between molecules

4. Intermolecular Attractions • Converting gas liquid or solid • Molecules must get closer together • Cool or compress • Converting liquid or solid gas • Requires molecules to move farther apart • Heat or reduce pressure • As T decreases, kinetic energy of molecules decreases • At certain T, molecules don’t have enough energy to break away from one another’s attraction

5. Inter vs. Intra-Molecular Forces • Intramolecular forces • Covalent bonds within molecule • Strong • Hbond(HCl) = 431 kJ/mol • Intermolecular forces • Attraction forces between molecules • Weak • Hvaporization(HCl) = 16 kJ/mol Intermolecular attraction (weak) Covalent Bond (strong)

6. Electronegativity Review Electronegativity: Measure of attractive force that one atom in a covalent bond has for electrons of the bond

7. Bond Dipoles • Two atoms with different electronegativity values share electrons unequally • Electron density is uneven • Higher charge concentration around more electronegative atom • Bond dipoles • Indicated with delta (δ) notation • Indicates partial charge has arisen

8. Net Dipoles • Symmetrical molecules • Even if they have polar bonds • Are non-polarbecause bond dipoles cancel • Asymmetrical molecules • Are polarbecause bond dipoles do not cancel • These molecules have permanent, net dipoles • Molecular dipoles • Cause molecules to interact • Decreased distance between molecules increases amount of interaction

9. POLAR COVALENT BOND COVALENT BOND IONIC BOND CHCl3 TiO2 F2 CaBr2

10. Group Problem Identify the overall dipole moment for CHCl3 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

11. Group Problem Identify the overall dipole moment for these molecules: Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

12. Solubility LIKE DISSOLVES LIKE polar molecules dissolve in polar solvents nonpolar molecules dissolve in nonpolar solvents Polar Solvents Water: H2O Methanol: CH3OH Ethanol: CH3CH2OH Acetone: (CH3)2CO Acetic Acid: CH3CO2H Ammonia: NH3 Acetonitrile: CH3CN Nonpolar Solvents Pentane: C5H12 Hexane: C6H14 Cyclohexane: C6H12 Benzene: C6H6 Toluene: CH3C6H5 Chloroform: CHCl3 Diethylether: (CH3CH2)2O

13. Intermolecular Forces • When substance melts or boils • Intermolecular forces are broken • Not covalent bonds • Responsible for non-ideal behavior of gases • Responsible for existence of condensed states of matter • Responsible for bulk properties of matter • Boiling points and melting points • Reflect strength of intermolecular forces

14. Three Important Types of Intermolecular Forces • London dispersion forces • Dipole-dipole forces • Hydrogen bonds • Ion-dipole forces • Ion-induced dipole forces

15. London Forces • When atoms near one another, their valence electrons interact • Repulsion causes electron clouds in each to distort and polarize • Instantaneous dipoles result from this distortion • Effect enhanced with increased volume of electron cloud size • Effect diminished by increaseddistance between particles and compact arrangement of atoms

16. London Forces • Affect All molecules, both polar and nonpolar • Boiling point (BP) is an indication of relative intermolecular force strength • Ease with which dipole moments can be induced and thus London Forces depend on • Distance between particles • Polarizability of electron cloud • Points of attraction • Number of atoms • Molecular shape (compact or elongated)

17. Polarizability • Ease with which the electron cloud can be distorted • Larger molecules often more polarizable • Larger number of less tightly held electrons • Magnitude of resulting partial charge is larger • Larger electron cloud

18. Group Problem Arrange the following atoms in order of increasing polarizability: Ar, He, Kr, Ne and Xe.

19. Table 12.1 Boiling Points of Halogens and Noble Gases Larger molecules have stronger London forces and thus higher boiling points.

20. Number of Atoms in Molecule • London forces depend on number atoms in molecule • Boiling point of hydrocarbons demonstrates this trend

21. Group Problem Which of the following molecules will have the highest boiling point? More sites (marked with *) along its chain where attraction to other molecules can occur

22. Molecular Shape • Increased surface area available for contact = increased London forces • London dispersion forces between spherical molecules are lower than chain-like molecules • More compact molecules • Hydrogen atoms not as free to interact with hydrogen atoms on other molecules • Less compact molecules • Hydrogen atoms have more chance to interact with hydrogen atoms on other molecules

23. Physical Origin of Shape Effect • Small area for interaction • Larger area for interaction More compact – lower BP Less compact – higher BP

24. Dipole-Dipole Attractions +  +  • Occur only between polar molecules • Possess dipole moments • Molecules need to be close together • Polar molecules tend to align their partial charges • Positive to negative • As dipole moment increases, intermolecular force increases  +  + +  + 

25. Dipole-Dipole Attractions • Tumbling molecules • Mixture of attractive and repulsive dipole-dipole forces • Attractions (- -) are maintained longer than repulsions(- -) • Get net attraction • ~1–4% of covalent bond

26. Dipole-Dipole Attractions • Interactions between net dipoles in polar molecules • About 1–4% as strong as a covalent bond • Decrease as molecular distance increases • Dipole-dipole forces increase with increasing polarity

27. Hydrogen Bonds • Special type of dipole-dipole Interaction • Very strong dipole-dipole attraction • ~10% of a covalent bond • Occurs between H and highly electronegative atom (O, N, or F) • H—F, H—O, and H—N bonds very polar • Electrons are drawn away from H, so high partial charges • H only has one electron, so +H presents almost bare proton • –X almostfull –1 charge • Element’s small size, means high charge density • Positive end of one can get very close to negative end of another

28. Examples of Hydrogen Bonding

29. Hydrogen Bonding in Water • Responsible for expansion of water as it freezes • Hydrogen bonding produces strong attractions in liquid • Hydrogen bonding (dotted lines) betweenwater molecules in ice form tetrahedral configuration

30. Hydrogen Bonding in Water 1.97 Å 0.957 Å

31. Your Turn List all intermolecular forces for CH3CH2OH. A. Hydrogen-bonds B. Hydrogen-bonds, dipole-dipole attractions, London dispersion forces C. Dipole-dipole attractions D. London dispersion forces E. London dispersion forces, dipole-dipole attractions

32. Your Turn In the liquid state, which species has the strongest intermolecular forces, CH4, Cl2, O2 or HF? A. CH4 B. Cl2 C. O2 D. HF

33. Ion-Dipole Attractions • Attractions between ion and charged end of polar molecules • Attractions can be quite strong as ions have full charges (a) Negative ends of water dipoles surround cation (b) Positive ends of water dipoles surround anion

34. Ex. Ion-Dipole Attractions: AlCl3·6H2O • Positive charge of Al3+ ion attracts partial negative charges – on O of water molecules • Ion-dipole attractions hold water molecules to metal ion in hydrate • Water molecules are found at vertices of octahedron around aluminum ion • Attractions between ion and polar molecules

35. Ion-Induced Dipole Attractions • Attractions between ion and dipole it induces on neighboring molecules • Depends on • Ion charge and • Polarizability of its neighbor • Attractions can be quite strong as ion charge is constant, unlike instantaneous dipoles of ordinary London forces • E.g., I– and Benzene

36. Group Problem List the intermolecular forces and rank in order of strength.

37. Summary of Intermolecular Attractions Dipole-dipole • Occur between neutral molecules with permanent dipoles • About 1–4% of covalent bond • Mid range in terms of intermolecular forces Hydrogen bonding • Special type of dipole-dipole interaction • Occur when molecules contain N—H, H—F and O—H bonds • About 10% of a covalent bond

38. Summary of Intermolecular Attractions London dispersion • Present in all substances • Weakest intermolecular force • Weak, but can add up to large net attractions Ion-dipole • Occur when ions interact with polar molecules • Strongest intermolecular attraction Ion-induced dipole • Occur when ion induces dipole on neighboring particle • Depend on ion charge and polarizability of its neighbor

39. Melting & Boiling Point Often can predict physical properties by comparing strengths of intermolecular attractions: Boiling Point increases when intermolecular forces increase Melting Point increases when intermolecular forces increase

40. Physical Properties that Depend on How Tightly Molecules Pack • Compressibility • Measure of ability of substance to be forced into smaller volume • Determined by strength of intermolecular forces • Gases highly compressible • Molecules far apart • Weak intermolecular forces • Solids and liquids nearly incompressible • Molecules very close together • Stronger intermolecular forces

41. Intermolecular Forces Determine Strength of Many Physical Properties • Retention of volume and shape • Gases, expand to fill their containers • Weakest intermolecular attractions • Molecules farthest apart • Liquids retainvolume, but not shape • Attractions intermediate • Solids retain both volume and shape • Strongest intermolecular attractions • Molecules closest

42. Intermolecular Forces and Temperature • Decrease with increasing temperature • Increasing kinetic energy overcomes attractive forces • If allowed to expand, increasing temperature increases distance between gas particles and decreases attractive forces

43. Diffusion • Movement that spreads one gas though another gas to occupy space uniformly • Spontaneous intermingling of molecules of one gas with molecules of another gas • Occurs more rapidly in gases than in liquids • Hardly at all in solids

44. Diffusion • In Gases • Molecules travel long distances between collisions • Diffusion rapid • In Liquids • Molecules closer • Encounter more collisions • Takes a long time to move from place to place • In Solids • Diffusion close to zero at room temperature • Will increase at high temperature

45. Surface Tension • Inside body of liquid • Intermolecular forces are the same in all directions • Molecules at surface • Potential energy increases when removing neighbors • Molecules move together to reduce surface area and potential energy Why does H2O bead up on a freshly waxed car instead of forming a layer?

46. Surface Tension • Causes a liquid to take the shape (a sphere) that minimizes its surface area • Molecules at surface have higher potential energy than those in bulk of liquid and move to reduce the potential energy • Wax = nonpolar • H2O = polar • Water beads in order to reduce potential energy by reducing surface area

47. Surface Tension • Liquids containing molecules with strong intermolecular forces have high surface tension • Allows us to fill glass above rim • Gives surface rounded appearance • Surface acts as “skin” that lets water pile up • Surface resists expansion and pushes back • Surface tension increases as intermolecular forces increase • Surface tension decreases as temperature increases

48. Wetting • Ability of liquid to spread across surface to form thin film • Greater similarity in attractive forces between liquid and surface, yields greater wetting effect • Occurs only if intermolecular attractive force between surface and liquid about as strong as within liquid itself

49. Wetting Ex. H2O wets clean glass surface as it forms H–bonds to SiO2 surface • Does not wet greasy glass, because grease is nonpolar and water is very polar • Only London forces • Forms beads instead Surfactants • Added to detergents to lower surface tension of H2O • Now water can spread out on greasy glass

50. Surfactants (Detergents) • Substances that have both polar and non-polar characteristics • Long chain hydrocarbons with polar tail • Nonpolar end dissolves in nonpolar grease • Polar end dissolves in polar H2O • Thus increasing solubility of grease in water