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Chapter 10

Chapter 10. Liquids and Solids. Solids vs. Liquids vs. Gases. Intermolecular Forces. Occur between (not within) molecules Covalent/Ionic Bonding Dipole-Dipole London Dispersion Forces. Important to Note. Changing states is a PHYSICAL (not chemical) change

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Chapter 10

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  1. Chapter 10 Liquids and Solids

  2. Solids vs. Liquids vs. Gases

  3. Intermolecular Forces • Occur between (not within) molecules • Covalent/Ionic Bonding • Dipole-Dipole • London Dispersion Forces

  4. Important to Note • Changing states is a PHYSICAL (not chemical) change • The molecules do NOT break apart in a physical (phase) change!

  5. Dipole-Dipole Forces • Two molecules with dipoles (polarity) can have an attraction to one another called a dipole-dipole attraction • This will make the most +, - attractions and -, - and +, + repulsions • Hydrogen bonding is a common (stronger) example Dipole-Dipole < Covalent < Ionic

  6. Hydrogen Bonding • Dipole-dipole attractions that are unusually strong • Hydrogen is small and can arrange itself closely to other atoms/molecules • Bond polarities also allow hydrogen to be attracted to the slightly negative end of a polar molecule • Causes higher boiling points • Important in organic molecules (especially large ones)

  7. London Dispersion Forces • Dipole-dipole force between noble gas atoms and nonpolar molecules • Instantaneous dipole (NO permanent dipole) which causes a similar dipole in an adjacent atom • Weak, short-lived attraction causes the freezing points of LARGER molecules to be lower • Larger molecules have greater LDF’s

  8. Surface Tension • Molecules at a liquids’ surface are attracted to other molecules near the surface (next to) and directly below them • To increase the surface area, energy is required to move molecules from the inside out (combatting surface tension) • Liquids with large IMF will have higher surface tensions

  9. Capillary Action • Spontaneous rising of a liquid in a narrow tube • Typically exhibited by polar liquids • Cohesive forces: intermolecular forces between liquid molecules • Adhesive forces: forces between liquid molecules and a container

  10. How Interesting • When a liquid is polar, it forms a concave meniscus. When a liquid is nonpolar, it forms a convex meniscus.

  11. Viscosity • Measures a liquids resistance to flow • Higher IMF = higher viscosity • More complex molecules = higher viscosity

  12. Modeling • Solids and gases are easy to model • Liquids are not… • Spectroscopy is used

  13. SOLIDS • Classified into crystalline and amorphous • Crystalline: highly regular arrangement • Amorphous: disorderly structure • Crystalline represented by various lattice structures (3D positioning) called crystal lattice structures

  14. Types of Solids • Ionic Solids: have ions at lattice points • Molecular Solids: have covalently bonded molecules at lattice points • Atomic Solids: have atoms at lattice points (examples: metallic solids, network solids, and Group 8A solids) • Metallic: delocalized nondirectional covalent bonding • Network: atoms bonded with strong directional covalent bonds • Group 8A: noble gas elements attracted with LDF

  15. Solid Classification

  16. Packing In Solids • Metals have characteristics based on their packing/bonding • “Closest Packing” spherical atoms are packed together in an arrangement that most efficiently uses the available space

  17. Types of Packing • Hexagonal Closest Packed (hcp) = hexagonal unit cell…abababab… • Cubic Closest Packed (ccp) = face-centered cubic unit cell…abcabcabc…

  18. Molecular Solids • Contain strong covalent bonding within the molecules but relatively weak forces between the molecules • The type of bond that forms between molecules will depend on the atoms present (LDF, hydrogen bonding, etc.)

  19. Ionic Solids • Stable, high-melting substances held together by strong electrostatic forces • Packing occurs so that the electrostatic forces among oppositely charged ions are maximized and repulsions are minimized • Larger ions (anions) are packed in hcp or ccp • Smaller ions (cations) are packed into holes made by the larger ions.

  20. Trigonal holes formed by 3 spheres in one layer (smallest) • Tetrahedral holes formed by one sphere sitting in a dimple formed by three other spheres (medium) • Octahedral holes formed between two sets of three spheres (largest)

  21. Vaporization • Endothermic (requires energy) • Heat of vaporization (∆Hvap) = amount of energy required to vaporize 1 mole of a liquid at 1 atm • Liquids with high vaporization - VOLATILE • Vapor pressure is partly determined by intermolecular forces (large IMF = lower VP) • Vapor pressure increases with temperature

  22. Dynamic Equilibrium • Rate of evaporation = rate of condensation

  23. Clausius-Clapeyron Equation • Can be used to determine the (∆Hvap) by measuring Pvap at different temperatures and evaluating a line of best fit… ln(Pvap,T1/Pvap,T2) = (∆Hvap/R)(1/T2-1/T1) ∆Hvap = must be in J (X 1000) R = constant (8.3145 J/Kmol) T = Kelvin ***to solve ln(x), do e^ for both sides

  24. Example • The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kJ/mol. Calculate the vapor pressure of water at 50°C. • Answer: 93.7 torr

  25. Sublimation • Solid directly to gas without passing a liquid state

  26. Heat/Enthalpy of Fusion(∆Hfus) • Occurs when solids melt • Energy is used to disrupt the bonds present between solid molecules • Temperature remains constant

  27. Heating Curve

  28. Normal Melting/Boiling Point • Temperature at which melting/boiling occurs at STANDARD pressure (1atm)

  29. Fancy Vocab • Supercooling: Liquid remains a liquid at temperatures below 0C at 1atm due to lack of solid arrangement necessary • When orderly arrangement necessary occurs, liquid will rapidly solidify • Superheating: liquid temperature raises above boiling point. Bubbles make it so the molecules with the most KE/highest temp collect in one area • Boiling chips are used to avoid these bubbles

  30. Phase Diagram • Shows different conditions of temperature and pressure and the phases that occur for different substances • Only works with a closed system • Triple Point: all three phases exist in equilibrium

  31. Critical Point: temp/press at which neither a true liquid/gas is present

  32. Chapter 11 Properties of Solutions

  33. Important Vocabulary • Homogeneous means there is only one phase (compositions do not vary) • Ex: Kool Aid, air, steel • Solute: Gets dissolved • Solvent: Does the dissolving • Solution: Homogeneous mixture consisting of a solute and solvent

  34. Dilute vs. Concentrated • Can’t be used in calculations • Molarity, mass percent, and mole fraction can be used to show solution concentrations

  35. Molarity • Moles of solute/liters of solution • Represented by M • Example: A solution was prepared by adding 5.84 g of formaldehyde, H2CO, to 100.0 g of water. The final volume of the solution was 104.0 mL. Calculate the molarity. • Answer: 1.87 M H2CO

  36. Mass Percent • Percent by mass of the solute in the solution • Mass Percent = (mass of solute/mass of solution) X 100% • Example: A solution was prepared by adding 5.84 g of formaldehyde, H2CO, to 100.0 g of water. The final volume of the solution was 104.0 mL. Calculate the mass percent. • Answer: 5.52 % H2CO, 94.48% H2O

  37. Mole Fraction • Represented by X • Moles of part/moles of solution X 100% Mole Frac. A = XA = nA/(nA+nB) • Example: A solution was prepared by adding 5.84 g of formaldehyde, H2CO, to 100.0 g of water. The final volume of the solution was 104.0 mL. Calculate the mole fraction. • Answer: XH2CO = 0.0338, XH2O = 0.9662

  38. Molality • Represented by m • Moles of solute per kilogram of solvent Molality = moles of solute/kilogram of solvent • Example: A solution was prepared by adding 5.84 g of formaldehyde, H2CO, to 100.0 g of water. The final volume of the solution was 104.0 mL. Calculate the molality. • Answer: 1.94 m H2CO

  39. NEW: Normality (N) • Number of “equivalents” per liter of solution • Equivalents - depends on reaction: • Acid-base reactions…mass of acid/base that can use/accept ONE mole of protons • Oxidation-reduction…quantity of oxidizing/reducing agent that will react with ONE mole of electrons • NOT ON AP EXAM!!

  40. Normality Example • Given the following reaction: H3PO4 + 3NaOH --> PO43- + 3H2O + 3Na+ If we have 28.42 g H3PO4 in 800 mL of water, what is the normality of the solution? Answer: 1.09 N H3PO4

  41. Solubility • Shows what will dissolve in what • “Like dissolves like” = polar solvents will dissolve polar/ionic solutes and nonpolar solvents will dissolve nonpolar solutes

  42. Factors Affecting Solubility 1. Structure 2. Pressure 3. Temperature

  43. 1. Structure Effects • Polarity of solute/solvent (like dissolves like) • Example: vitamins are fat-soluble and water-soluble • Fat-soluble = nonpolar, hydrophobic (water-fearing), build up/stored in fatty tissue, too much = hypervitaminosis • Water-soluble = polar, hydrophilic (water-loving), extra are excreted by the body

  44. 2. Pressure Effects • Doesn’t affect liquids/solids, but has a large affect on gases • Gas solubility increases as the partial pressure of the gas above the solution increases

  45. Henry’s Law • Shows relationship between gas pressure and concentration of dissolved gas: C = kP • C = concentration of dissolved gas • K = constant for particular solution • P = partial pressure of gas above solution • Works best with gases that don’t dissociate in/react with solvent

  46. Henry’s Law Example • The solubility of O2 is 2.2 X 10-4 M at 0C and 0.10 atm. Calculate the solubility of O2 at 0C and 0.35 atm. • Answer: 7.7 X 10-4 M O2

  47. 3. Temperature Effects • For most solids, solubility increases as temperature increases • For most gases, solubility decreases as temperature increases • Thermal Pollution in lakes: increase in temp. lowers dissolved oxygen concentrations

  48. Vapor Pressure of Solutions • If a solution contains a nonvolitile (not easily vaporized) solute, its vapor pressure is LOWER than the pure solvent. • Shells of water solvation make it so it’s harder for the solvent to vaporize

  49. Molecules that do not dissociate (break up) in water (solvent) have higher vapor pressures than ionic compounds that do dissociate • The decrease in a solution’s vapor pressure is proportional to the number of particles the solute makes in solution.

  50. Answer This… • Which compound affects the vapor pressure of a solution the least: glucose, sodium chloride, or calcium chloride? • Solutions with covalent compounds > Solutions with ionic compounds

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