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

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

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

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  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. 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

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

  11. 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

  12. London Forces • Instantaneous dipole-induced dipole attractions • London Forces • Dispersion forces • Operate between all molecules • Neutral or net charged • Nonpolar or polar

  13. London Dispersion Forces • Ease with which dipole moments can be induced and thus London Forces depend on • Polarizability of electron cloud • Points of attraction • Number atoms • Molecular shape (compact or elongated)

  14. 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

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

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

  17. How Intermolecular Forces Determine Physical Properties • Hexane, C6H14 • BP 68.7 °C • More sites (marked with *) along its chain where attraction to other molecules can occur Propane, C3H8 BP –42.1 °C

  18. 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

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

  20. 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  +  + +  + 

  21. 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

  22. 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

  23. 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

  24. Examples of Hydrogen Bonding

  25. 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

  26. 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

  27. 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

  28. 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

  29. Ex. Ion-Dipole Attractions: AlCl3·6H2O • Attractions between ion and polar molecules • 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

  30. 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

  31. 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

  32. 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

  33. Using Intermolecular Forces • Often can predict physical properties (like BP, MP and many others) by comparing strengths of intermolecular attractions • Ion-Dipole • Hydrogen Bonding • Dipole-Dipole • London Forces • Larger, longer, and therefore heavier molecules often have stronger intermolecular forces • Smaller, more compact, lighter molecules have generally weaker intermolecular forces Strongest Weakest

  34. 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

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

  36. 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

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

  38. 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 Diffusion

  39. Surface Tension Why does H2O bead up on a freshly waxed car instead of forming a layer? • 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

  40. 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

  41. 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

  42. 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

  43. 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

  44. 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

  45. Viscosity • Resistance to flow • Measure of fluid’s resistance to flow or changing form • Related to intermolecular attractive forces • Also called internal friction • Depends on intermolecular attractions

  46. Viscosity • Viscosity decreases when temperature increases • Most people associate liquids with viscosity • Syrup more viscous than water • Gases have viscosity • Respond almost instantly to form-changing forces • Solids, such as rocks and glass have viscosity • Normally respond very slowly to forces acting to change their shape

  47. Acetone Polar molecule Dipole-dipole and London forces Ethylene glycol Polar molecule Hydrogen-bonding Dipole-dipole and London forces Effect of Intermolecular Forces on Viscosity Which is more viscous?

  48. Your Turn! For each pair given, which is has more viscosity? CH3CH2CH2CH2OH, CH3CH2CH2CHO C6H14, C12H26 NH3(l ), PH3(l) A. CH3CH2CH2CH2OH C6H14NH3(l ) B. CH3CH2CH2CH2OH C12H26NH3(l ) C. CH3CH2CH2CHO C6H14 PH3(l ) D. CH3CH2CH2CHO C12H26NH3(l ) E. CH3CH2CH2CH2OH C12H26PH3(l )

  49. Solubility • “Like dissolves like” • To dissolve polar substance, use polar solvent • To dissolve nonpolar substance, use nonpolar solvent • Compare relative polarity • Similar polarity means greater ability to dissolve in each other • Differing polarity means that they don’t dissolve, they are insoluble • Surfactants • Both polar and non-polar characteristics • Used to increase solubility

  50. Your Turn! Which of the following are not expected to be soluble in water? • HF • CH4 • CH3OH • All are soluble