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Hlanganani Tutu, C403 School of Chemistry Email: hlanganani.tutu@wits.ac.za PowerPoint Presentation
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Hlanganani Tutu, C403 School of Chemistry Email: hlanganani.tutu@wits.ac.za

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  1. Hlanganani Tutu, C403 School of Chemistry Email: hlanganani.tutu@wits.ac.za Chapter 13

  2. The Solution Process • Solution - homogeneous mixture of solute and solvent. • In solutions, intermolecular forces become rearranged. Chapter 13

  3. Examples of solutions • gas in gas – e.g. air • gas in liquid -- e.g. soda • gas in solid -- e.g. gas on solid, catalyst • liquid in liquid • liquid in solid -- e.g. mercury amalgam • solid in liquid • solid in solid -- e.g. 14-karat gold, brass Chapter 13

  4. Consider NaCl (solute) dissolving in water (solvent): • Interruption of water H-bonds, • NaCl→Na+ + Cl-, • ion-dipole forces form: Na+…-OH2 and Cl- …+H2O. • If water is the solvent, we say the ions are hydrated. Chapter 13

  5. Energy Changes and Solution Formation • 3 energy steps in forming a solution: • separation of solute molecules (H1), • separation of solvent molecules (H2), and • formation of solute-solvent interactions (H3). Chapter 13

  6. Hsoln = H1 + H2 + H3. • Hsoln can be +ve or -ve depending on the intermolecular forces. Chapter 13

  7. Chapter 13

  8. “Rule”: polar solvents dissolve……………………….? Non-polar solvents dissolve..............................................? Chapter 13

  9. Exercise: Why doesn’t gasoline dissolve NaCl? Exercise: Why doesn’t water and octane mix well (immiscible)? Remember: the resultant solution’s interactions must be stronger than the interactions in the original substance Chapter 13

  10. Solution Formation, Spontaneity, and Disorder • When the energy of the system decreases (e.g. dropping a book and allowing it to fall to a lower potential energy), the process is spontaneous. Chapter 13

  11. Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids. Therefore, they spontaneously mix Chapter 13

  12. Chapter 13

  13. There are solutions that form by physical processes and those by chemical processes. Chapter 13

  14. Consider: • Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g). • When all the water is removed from the solution, no Ni is found only NiCl2·6H2O. Therefore, Ni dissolution in HCl is a chemical process. Chapter 13

  15. Consider: • NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq). • When the water is removed from the solution, NaCl is found. Therefore, NaCl dissolution is a physical process. Chapter 13

  16. Saturated Solutions and Solubility • Dissolve: solute + solvent  solution. • Crystallization: solution  solute + solvent. • Saturation: crystallization and dissolution are in equilibrium. Chapter 13

  17. Solubility: amount of solute required to form a saturated solution. • Supersaturation: reached when more solute is dissolved than in a saturated solution. Chapter 13

  18. Chapter 13

  19. Factors Affecting Solubility • Solute-Solvent Interaction • Miscible liquids: mix in any proportions. • Immiscible liquids: do not mix. • Intermolecular forces are important • The more C atoms, the less the solubility in water. Chapter 13

  20. The -OH groups in a molecule increase solubility in water….“like dissolves like” Chapter 13

  21. Chapter 13

  22. Which of these two would be more soluble in water? Chapter 13

  23. Network solids do not dissolve. Why? Chapter 13

  24. Pressure Effects • Solubility of a gas in a liquid is a function of the pressure of the gas. Chapter 13

  25. Pressure Effects Chapter 13

  26. The higher the pressure, the more molecules of gas are close to the solvent Chapter 13

  27. Henry’s Law gives: where: Sg - solubility of a gas, k is a constant, and Pg is the partial pressure of a gas Chapter 13

  28. Example 27g of acetylene, C2H2, dissolves in 1L of acetone at 1.0 atm pressure. If the partial pressure of acetylene is increased to 12 atm, what is the solubility in acetone? Solution: S1 = kP1…………(1) S2 = kP2…………(2) Ans: 3.2 x 102 g Chapter 13

  29. Carbonated beverages are bottled with a partial pressure of CO2 > 1 atm. • What happens when a bottle is opened? Chapter 13

  30. Temperature Effects • As temperature increases, solubility of solids generally increases, e.g. sugar in warm water • Sometimes, solubility decreases as temperature increases (e.g. Ce2(SO4)3). Chapter 13

  31. Temperature Effects • Gases - less soluble at high temperature • Thermal pollution in dams and rivers – loss of O2 Chapter 13

  32. Ways of Expressing Concentration • Mass Percentage, ppm, and ppb • Definitions: Chapter 13

  33. Example: How would you prepare 425 g of an aqueous solution containing 2.40% by mass of sodium acetate, NaC2H3O3? Ans: Mass of NaC2H3O3 = 10.2 g Mass of H2O = mass of solution - mass of NaC2H3O3 = 415 g Chapter 13

  34. Chapter 13

  35. Chapter 13

  36. Mole Fraction, Molarity, and Molality Chapter 13

  37. Converting between molarity (M) and molality (m) requires density. Exercise: 0.2 mol of ethylene glycol is dissolved in 2000 g of water. Calculate the molality Chapter 13

  38. Example: What is the molality of a solution containing 5.67 g of glucose, C6H12O6 (Mr = 180.2 g), dissolved in 25.2 g of water? (Calc. the mole fractions of the components as well). • Solution: • Think about the solute!................glucose (express in moles) • Think about the solvent!...............water (express in kilograms) Ans: 1.25 m Chapter 13

  39. Example: Converting molarity to molality An aqueous solution is 0.907M Pb(NO3)2. What is the molality of lead nitrate, Pb(NO3)2, in this solution? The density of the solution is 1.252 g/mL. (Molar mass of Pb(NO3)2 = 331.2 g) Ans: 0.953 m Pb(NO3)2 Chapter 13