Unit 5: Solids, Liquids, and Solutions

1. Unit 5: Solids, Liquids, and Solutions Adapted from: http://www.sciencegeek.net/APchemistry/Powerpoints.shtml and http://www.mrayton.com/ap-chemistry.html

2. Intermolecular vsIntramolecular Forces • Thoughts? • Analogy: intermolecular (interstate or international) vs. intramolecular (intrastate commerce)

3. POGIL Activity • Complete with groups!

4. Refresher • Rank (and explain) the following substances from strongest to weakest intermolecular forces. NaCl MgCl2 AlCl3MgSNaBr • Solution: Based on products of charges and then distance between nuclei (Coulomb’s Law) • MgS; AlCl3; MgCl2; NaCl; NaBr

5. Intermolecular Forces • Types of bonding forces vary in their strength as measured by average bond energy. (direct relationship to melting and boiling points!) Strongest Ionic bonds Covalent bonds Hydrogen bonding Dipole-dipole interactions London Dispersion Forces Weakest

6. Hydrogen Bonding • Bonding between hydrogen and more electronegative neighboring atoms such as oxygen, nitrogen, and fluorine.

7. Dipole Related Forces • Stronger than dipole-related forces: Ionic, metallic, and covalent bonding • Weaker than dipole-related forces: London dispersion forces

9. London Dispersion Forces • The temporary separations of charge that lead to the London force attractions are what attract one NONPOLAR molecule to its neighbors. • London forces increase with the size of the molecules. • Also Known As (aka): London forces; dispersion forces; and dispersion-interaction forces

11. Check for Understanding • Rank and provide an explanation of your rankings for the following substances from strongest to weakest intermolecular forces. He NH3 NF3NaCl Solution: Ion-ion>hydrogen>dipole-dipole>dispersion forces

12. Why can you never trust atoms? • They make up everything! • If the Silver Surfer and Iron Man team up, they’d be alloys.

13. Check for Understanding • Rank (and explain) the following substances from strongest to weakest intermolecular forces. HF F2 FCl • Solution: Hydrogen bonding> dipole-dipole>dispersion

14. Check for Understanding • Are ammonia (NH3) and water (H2O) miscible, why or why not? • Solution: Yes, both exhibit hydrogen bonding. Interact with each other through dipole-dipole forces. • Would you expect NaBrto be a solid, liquid, or gas at room temperature? Why?

15. Individual Practice IMFs • Page 475 • #29, 30, 31, 33, and 35 • Read pages 425-471

16. Liquids • Surface Tension • The resistance of a liquid to an increase in its surface area. • Capillary Action • Spontaneous rising of a liquid in a narrow tube. • Polar vs Nonpolar capillary action • Cohesive Force vs Adhesive Force and relationship to the Force of Gravity (concave vs convex meniscus)

17. Viscosity (Liquids are fluid) • A measure of a liquid’s resistance to flow. • High viscosity is an indication of strong intermolecular forces!

18. Check for Understanding • The narrower the capillary tube, the farther the water travels. Why?

19. Properties of Solids • Types • Crystalline Solids • Highly regular arrangement of their components • Amorphous Solids • Considerable disorder in their structures (glass, plastic)

20. Components in a Crystalline Solid • Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance. • Unit Cell: the smallest repeating unit of the lattice.

21. Video or Video: Start at 5:08

22. Metal Alloys • Substitutional Alloy: • Some metal atoms replaced by others of similar size • Example- brass (Cu and Zn) • Interstitial Alloy: • Interstices (holes) in closet packed metal structure are occupied by small atoms. • Example – steel (Fe and C)

23. Network Atomic Solids (Allotropes) • Some covalently bonded substances DO NOT form discrete molecules. • Diamond • Network of covalently bonded carbon atoms (tetrahedral arrangement) • Graphite • Network of covalently bonded carbon atoms (layers arranged in fused six-membered rings)

24. Graphene • One-atom thick layer of graphite. • Characteristics • Strong • Light • Transparent • Good heat and electricity conductor

25. Semiconductors • Conductivity increases at higher temperature • Conductivity of silicon can be increased by combining with other elements. • N-type : more valence electrons (i.e. phosphorus) • P-type : less valence electrons (i.e. aluminum) • Elements with more valence electrons than the “host” crystal

26. Solids • Molecular • Strong covalent forces within molecules • Weak covalent forces between molecules • Ionic • At room conditions generally crystal lattice of alternating cations and anions (typically cubic crystals) • Stable, high-melting substances • Held together by the strong electrostatic forces that exist between oppositely charged ions. (packed to maximize attractions and minimize repulsions)

27. Changes of State • Vaporization • The process in which a the molecules of a liquid can escape the liquid’s surface and form a gas. • Endothermic (requires energy to overcome intermolecular forces in the liquid) • Heat of vaporization, ΔHvap, (energy required to vaporize 1 mole of a liquid at 1 atm)

28. Vapor Pressure and Changes of State • Condensation • Process by which vapor molecules re-form a liquid. • Equilibrium- • The point at which no further net change occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other. (dynamic) • Pressure at which equilibrium is reached is called vapor pressure. • Liquids with high vapor pressures are said to be volatile. • Low vapor pressures are associated with strong intermolecular forces

29. Phase Changes • Sublimation • Process in which a solid goes directly to the gaseous state. • Melting • Heat of fusion, ΔHfus, when a solid melts. • Ionic solids (i.e. NaCl and NaF) have very high melting points and enthalpies of fusion because of the strong ionic forces in these solids. • Molecular solids (i.e. O2) contain nonpolar molecules with weak intermolecular forces and thus have low melting points.

31. Key Components of Phase Diagrams • Critical temperature • Temperature above which the vapor cannot be liquefied. • Critical pressure • Pressure required to liquefy at the critical temperature. • Critical point • Critical temperature and pressure

33. Individual Practice • Page 475 • Questions 12, 15, 17, 18, 20, 22, 23 and 91

34. Solutions

35. Warm-up: • How many grams of sodium chloride are needed to prepare 1.50 liters of 0.500 M NaCl solution? • Solution: 43.8 g NaCl

36. Refresher: Solution Concentration • Mole Fraction: the ratio of moles of solute to total moles of solution • Molarity is the ratio of moles of solute to liters of solution

37. Mass percent • Another way to describe solution composition • Aka: weight percent • Percent by mass of the solute in the solution

38. Molality • The number of moles of solute per kilogram of solvent • Another way to describe solution concentration • Very dilute aqueous solutions, the magnitude of the molality and the molarity are almost the same. • Since molarity depends on the volume of the solution, it changes slightly with temperature. Molality is independent of temperature because it depends only on mass.

39. Concentration Units Model • 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 104.0mL. Calculate the molarity, molality, mass percent, and mole fraction of the formaldehyde in the solution. • Solution: 1.87 M H2CO; 1.94 m H2CO; 5.52% H2CO; 0.0338

40. Concentration Calculations • The density of a 10.0% (by mass) solution of NaOH is 1.109 g/cm3. Calculate the concentration of this solution in molarity, molality, and mole fraction. • Solution: 2.77M; 2.78m; 0.0476

41. Concentration Calculations • A hydrochloric acid solution was made by adding 59.26 g HCl to 100.g H2O. The density of the solution was 1.19 g/cm3. Calculate the concentration of HCl in molarity, molality, mass percent, and mole fraction. • Solution:12.1M; 16.3m; 37.2%; 0.226

42. Individual Practice • Page 519 • 9, 11, 25, 26, 28, 29, and 32

43. Heat of Solution…remember when… • The amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent.

44. Flash back … “Like dissolves Like” • Polar solvents dissolve polar solutes. • Nonpolar solvents dissolve nonpolar solutes. • Application/Connections: • What are some examples of this concept that we have discussed and/or observed in the class?

45. Steps in Solution Formation • Step 1ΔH1: Separating the solute into its individual components (expanding the solute). • Step 2 ΔH2: Overcoming intermolecular forces in the solvent to make room for the solute (expanding the solvent). • Step 3 ΔH3: Allowing the solute and solvent to interact to form the solution.

47. Predicting Solution Formation

48. Factors Affecting Solubility • Separate the behavior of solids and liquids from gases. • Solubility of gases is relatively independent of structure. Whereas, solid and liquids are highly dependent on structure. • Pressure has little effect on the solubilities of liquids and solids. • For gases, the solubility is impacted by the pressure (Henry’s Law)

49. Henry’s Law • Relationship between the partial pressure of a gas above a solution to the concentration of the gas dissolved in the solution. **Only holds when there is NO chemical reaction between the solute and solvent. (not HCl and HI) • Mathematical Relationship

50. Henry’s Law • What will happen to the pressure above the solution when more gas is added to the sealed container? • What will happen to the pressure above the solution when the container is compressed? • Solution: The pressure increases as more gas is added the thus the frequency of collisions of the gas and the solution surface increase. This results in more gas molecules dissolving in the solution, thus the concentration of the gas in solution increases. (direct relationship)