Chapter 13 Properties of Solutions

1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 13Properties of Solutions John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

2. December 10 • The solution process • Why a solution forms? • Chapter HW • 1,3,4,6 • Solution process13, 15 • Saturated solutions – Factors affecting solubility19, 21, 25, 27, 29, 31

3. Section 13.1The Solution process • Energy Changes and Solution Formation • Solution Formation, Spontaneity and Disorder. • Solution Formation and Chemical Reactions

4. Solutions • Solutions are homogeneous mixtures of two or more pure substances. • In a solution, the solute (present in smaller amount) is dispersed uniformly throughout the solvent (present in largest amount).

5. Solutions The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

6. How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

7. How Does a Solution Form If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

8. Energy Changes and SolutionFormation Three processes affect the energetic of the process: • Separation of solute particles D H1 • Separation of solvent particles D H2 • New interactions between solute and solvent D H3

9. Energy Changes and Solution Formation • We define the enthalpy change in the solution process as • Hsoln = H1 + H2 + H3. • Hsoln can either be positive or negative depending on the intermolecular forces.

10. Breaking attractive intermolecular forces is always endothermic. • Forming attractive intermolecular forces is always exothermic.

11. Energy Changes in Solution The enthalpy change of the overall process depends on H for each of these steps.

12. Why Do Endothermic Processes Occur? Things do not tend to occur spontaneously (i.e., without outside intervention) unless the energy of the system is lowered.

13. To determine whether Hsoln is positive or negative, we consider the strengths of all solute-solute and solute-solvent interactions: • H1 and H2 are both positive. • H3 is always negative. • It is possible to have either H3 > (H1 + H2) or H3 < (H1 + H2).

14. Examples: • NaOH added to water has Hsoln = -44.48 kJ/mol. • NH4NO3 added to water has Hsoln = + 26.4 kJ/mol. • “Rule”: LIKE DISSOLVES LIKE!!! • polar solvents dissolve polar solutes. Non-polar solvents dissolve non-polar solutes. Why? • If Hsoln is too endothermic a solution will not form. • NaCl in gasoline: the ion-dipole forces are weak because gasoline is non-polar. Therefore, the ion-dipole forces do not compensate for the separation of ions. • Water in octane: water has strong H-bonds. There are no attractive forces between water and octane to compensate for the H-bonds.

15. Why Do Endothermic Processes Occur? Yet we know that in some processes, like the dissolution of NH4NO3 in water, heat is absorbed, not released.

16. Enthalpy Is Only Part of the Picture The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system.

17. Entropy- Disorder So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered.

18. Solution formation -Spontaneity • Spontaneous change tend to occur if the process results in • a) lower energy for the whole system • So the changes with D H < 0 are favored. • b) an increase in the total disorder, randomnes, degree of dispersal or entropy D S > 0

19. A solution will form except in the cases that the attraction between solute-solute/solvent-solvent are too strong compared with the solute-solvent attractions. • Solution formation always increase the Entropy of the system.

20. Solution formation and chemical reaction • Dissolution is a physical change—you can get back the original solute by evaporating the solvent. • If you can’t, the substance didn’t dissolve, it reacted.

21. Student, Beware! Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.

22. Solution Formation and Chemical Reactions • Consider: • Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g). • Note the chemical form of the substance being dissolved has changed (Ni  NiCl2). • 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. • NiCl2·6H2O is a hydrate that contains Ni2+ ions.

23. Example: • 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.

24. Section 13.3 • Factors Affecting Solubility

25. Types of Solutions • Saturated • Solvent holds as much solute as is possible at that temperature. • Dissolved solute is in dynamic equilibrium with solid solute particles.

26. Types of Solutions • Unsaturated • Less than the maximum amount of solute for that temperature is dissolved in the solvent.

27. Types of Solutions • Supersaturated • Solvent holds more solute than is normally possible at that temperature. • These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask. • NaC2H3O2 Sodium acetate usually forms supersaturated solutions.

28. Factors Affecting Solubility • *Solute-Solvent interactions • *Pressure effects (only for gases) • *Temperature effects

29. Factors Affecting Solubility • Chemists use the axiom “like dissolves like”: • Polar substances tend to dissolve in polar solvents. • Nonpolar substances tend to dissolve in nonpolar solvents.

30. Factors Affecting Solubility The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

31. Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not.

32. Factors Affecting Solubility • Solute-Solvent Interaction • Polar substances tend to dissolve in polar solvents. • Miscible liquids: mix in any proportions. • Immiscible liquids: do not mix. • Intermolecular forces are important: water and ethanol are miscible because the broken hydrogen bonds in both pure liquids are re-established in the mixture. • The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water.

33. Solute-Solvent Interaction • The number of -OH groups within a molecule increases solubility in water. • Generalization: “like dissolves like”. • The more polar bonds in the molecule, the better it dissolves in a polar solvent. • The less polar the molecule the less it dissolves in a polar solvent and the better is dissolves in a non-polar solvent.

34. 5 1 Example – Place the following substances in order of increasing solubility in water: 3 2 4

35. 5 1 Example – Place the following substances in order of increasing solubility in hexane (C6H14): 4 3 2

36. Solute-Solvent Interaction

37. Solute-Solvent Interaction

38. Solute-Solvent Interaction • Network solids do not dissolve because the strong intermolecular forces in the solid are not re-established in any solution. • Pressure Effects • Solubility of a gas in a liquid is a function of the pressure of the gas.

39. Gases in Solution • In general, the solubility of gases in water increases with increasing mass. • Larger molecules have stronger dispersion forces.

40. Gases in Solution • The solubility of liquids and solids does not change appreciably with pressure. • The solubility of a gas in a liquid is directly proportional to its pressure.

41. Henry’s Law Sg = kPg where • Sg is the solubility of the gas; • k is the Henry’s law constant for that gas in that solvent; • Pg is the partial pressure of the gas above the liquid.

42. Henry’s Law Constant • Is different for each solute-gas pair. • Varies with temperature

43. The higher the pressure, the more molecules of gas are close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution. • Therefore, the higher the pressure, the greater the solubility. • The lower the pressure, the fewer molecules of gas are close to the solvent and the lower the solubility.

44. Example – What is the concentration of CO2 in water in a soda bottled under a pressure of 20.0 atm of CO2? (k = 3.1 x 10-2 mol L-1 atm-1) 0.62 mol L-1 or 0.62 M

45. Pressure Effects • Carbonated beverages are bottled with a partial pressure of CO2 > 1 atm. • As the bottle is opened, the partial pressure of CO2 decreases and the solubility of CO2 decreases. • Therefore, bubbles of CO2 escape from solution.

46. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature. Note the exception Ce2(SO4)3

47. Temperature • The opposite is true of gases: • Carbonated soft drinks are more “bubbly” if stored in the refrigerator. • Warm lakes have less O2 dissolved in them than cool lakes.

48. Ways of Expressing Concentrations of Solutions

49. Homework • 33, 35, 37, 41, 45, 47, 49

50. Concentrations * Qualitative Expressions Dilute – Concentrate *Quantitative Expressions • Molarity-Molality • %by mass, ppm,ppb • Mole Fraction