Chapter 17 “Water and Aqueous Systems”

1. Chapter 17“Water and Aqueous Systems”

2. Section 17.1Liquid Water and it’s Properties • OBJECTIVES: • Describe the hydrogen bonding that occurs in water.

3. Section 17.1Liquid Water and it’s Properties • OBJECTIVES: • Explain the high surface tension and low vapor pressure of water in terms of hydrogen bonding.

4. The Water Molecule • Water is a simple triatomic molecule. • Each O-H bond is highly polar, because of the high electronegativity of the oxygen • bond angle = 105 o • due to the bent shape, the O-H bond polarities do not cancel. This means water as a whole is polar. • Fig. 17.2, p.475

5. The Water Molecule • Water’s bent shape and ability tohydrogen bond gives water many special properties! • Water molecules are attracted to one another. • This gives water: high surface tension, low vapor pressure, high specific heat, high heat of vaporization, and high boiling point

6. High Surface Tension • liquid water acts like it has a skin • glass of water bulges over the top • Water forms round drops • spray water on greasy surface • All because water hydrogen bonds. • Fig. 17.4, p.476

7. O H H O H H Surface Tension d- • One water molecule hydrogen bonds to another. • Also, hydrogen bonding occurs to other molecules all around. d+ d+ d- d+ d+

8. Surface Tension • A water molecule in the middle of solution is pulled in all directions.

9. Surface Tension • Not true at the surface. • Only pulled down and to each side. • Holds the molecules together. • Causes surface tension.

10. Surface Tension • Water drops are round, because all molecules on the edge are pulled to the middle- not to the air!

11. Surface Tension • Glass has polar molecules. • Glass can hydrogen bond. • Attracts the water molecules. • Some of the pull is up a cylinder.

12. Meniscus • Water curves up along the side of glass. • This makes the meniscus, as in a graduated cylinder • Plastics are non-wetting; no attraction

13. Meniscus In Plastic In Glass

14. Surface tension • All liquids have surface tension • water is higher than most others • How to decrease surface tension? • Use a surfactant - surface active agent • a wetting agent, like detergent or soap • interferes with hydrogen bonding

15. Low vapor pressure • Fig. 17.6, p.477 • Hydrogen bonding also explains water’s unusually low vapor pressure. • Holds water molecules together, so they do not escape • good thing- lakes and oceans would evaporate very quickly!

16. Specific Heat Capacity • Water has a high heat capacity (also called specific heat). • It absorbs 4.18 J/gºC, while iron absorbs only 0.447 J/gºC. • Remember: SH = heat Mass x DT • If we calculate the heat need to raise the temperature of both iron and water by 75ºC - water is almost 10 x more!

17. Section 17.2Water Vapor and Ice • OBJECTIVES: • Account for the high heat of vaporization and the high boiling point of water, in terms of hydrogen bonding.

18. Section 17.2Water Vapor and Ice • OBJECTIVES: • Explain why ice floats in water.

19. Evaporation and Condensation • Because of the strong hydrogen bonds, it takes a large amount of energy to change water from a liquid to a vapor. • 2,260 J/g is the heat of vaporization. • This much energy to boil 1 gram water • You get this much energy back when it condenses. • Steam burns, but heats things well.

20. Ice • Most liquids contract (get smaller) as they are cooled. • They get more dense. • When they change to solid, they are more dense than the liquid. • Solid metals sink in liquid metal. • But, ice floats in water. • Why?

21. Ice • Water becomes more dense as it cools until it reaches 4ºC. • Then it becomes less dense. • As the molecules slow down, they arrange themselves into honeycomb shaped crystals. • These are held together by hydrogen bonds. (Fig. 17.9, p.481)

22. H H H O O H H H H H O O H O H O H H H H O H O O H H H H H H O O O H H H Liquid Solid

23. Ice • 10% greater volume than water. • Water freezes from the top down. • The layer of ice on a pond acts as an insulator for water below • It takes a great deal of energy to turn solid water to liquid water. • Heat of fusion is: 334 J/g.

24. Section 17.3Aqueous Solutions • OBJECTIVES: • Explain the significance of the statement “like dissolves like”.

25. Section 17.3Aqueous Solutions • OBJECTIVES: • Distinguish among strong electrolytes, weak electrolytes, and nonelectrolytes, giving examples of each.

26. Solvents and Solutes • Solution - a homogenous mixture, that is mixed molecule by molecule. • Solvent- the dissolving medium • Solute -the dissolved particles • Aqueous solution- a solution with water as the solvent. • Particle size about 1 nm; cannot be separated by filtration!

27. Aqueous Solutions • Water dissolves ionic compounds and polar covalent molecules best. • The rule is: “like dissolves like” • Polar dissolves polar. • Nonpolar dissolves nonpolar. • Oil is nonpolar. • Oil and water don’t mix. • Salt is ionic- makes salt water.

28. How Ionic solids dissolve • Called solvation. • Water breaks the + and - charged pieces apart and surrounds them. • Fig. 17.12, p. 483 • In some ionic compounds, the attraction between ions is greater than the attraction exerted by water • Barium sulfate and calcium carbonate

29. H H H H O O O H H H H O O H H O O H H H H H H O H O H How Ionic solids dissolve

30. Solids will dissolve if the attractive force of the water molecules is stronger than the attractive force of the crystal. • If not, the solids are insoluble. • Water doesn’t dissolve nonpolar molecules because the water molecules can’t hold onto them. • The water molecules hold onto each other, and separate from the nonpolar molecules. • Nonpolars? No repulsion between them

31. Electrolytes and Nonelectrolytes • Electrolytes- compounds that conduct an electric current in aqueous solution, or in the molten state • all ionic compounds are electrolytes (they are also salts) • barium sulfate- will conduct when molten, but is insoluble in water!

32. Electrolytes and Nonelectrolytes • Do not conduct? Nonelectrolytes. • Many molecular materials, because they do not have ions • Not all electrolytes conduct to the same degree • there are weak electrolytes, and strong electrolytes • depends on: degree of ionization

33. Electrolytes and Nonelectrolytes • Table 17.3, p.485 lists some common electrolytes and nonelectrolytes • How do you know if it is strong or weak? Rules on handout sheet.

34. Electrolyte Summary • Substances that conduct electricity when dissolved in water, or molten. • Must have charged particles that can move. • Ionic compounds break into charged ions: NaCl ® Na1+and Cl1- • These ions can conduct electricity.

35. Nonelectrolytes do not conduct electricity when dissolved in water or molten • Polar covalent molecules such as methanol (CH3OH) don’t fall apart into ions when they dissolve. • Weak electrolytes don’t fall completely apart into ions. • Strong electrolytes do ionize completely.

36. + heat - heat Water of Hydration(or Water of Crystallization) • Water molecules chemically bonded to solid salt molecules (not in solution) • These compounds have fixed amounts of water. • The water can be driven off by heating: • CuSO4.5H2O CuSO4 + 5H2O • Called copper(II)sulfate pentahydrate.

37. Hydrates • Table 17.4, p.486 list some familiar hydrates • Since heat can drive off the water, the forces holding it are weak • If a hydrate has a vapor pressure higher than that of water vapor in air, the hydrate will effloresce by losing the water of hydration

38. Hydrates • Some hydrates that have a low vapor pressure remove water from the air to form higher hydrates- called hygroscopic • used as drying agents, or dessicants • packaged with products to absorb moisture

39. Hydrates • Some compounds are so hygroscopic, they become wet when exposed to normally moist air- called deliquescent • remove sufficient water to dissolve completely and form solutions • Fig. 17.17, p.487 • Sample Problem 17-1, p.488 for percent composition

40. Section 17.4Heterogeneous Aqueous Systems • OBJECTIVES: • Explain how colloids and suspensions differ from solutions.

41. Section 17.4Heterogeneous Aqueous Systems • OBJECTIVES: • Describe the Tyndall effect.

42. Mixtures that are NOT Solutions • Suspensions: mixtures that slowly settle upon standing. • Particles of a suspension are greater in diameter than 100 nm. • Can be separated by filtering (p.490) • Colloids: heterogeneous mixtures with particles between size of suspensions and true solutions (1-100 nm)

43. Mixtures that are NOT Solutions • The small particles are the dispersedphase, and are spread throughout the dispersion medium • The first colloids were glues. Others include mixtures such as gelatin, paint, aerosol sprays, and smoke • Table 17.5, p.491 list some common colloidal systems and examples

44. Mixtures that are NOT Solutions • Many colloids are cloudy or milky in appearance when concentrated, but almost clear when dilute • do not settle out • cannot be filtered out • Colloids exhibit the Tyndall effect- the scattering of visible light in all directions. • suspensions also show Tyndall effect

45. Mixtures that are NOT Solutions • Flashes of light are seen when colloids are studied under a microscope- light is reflecting- called Brownian motiontodescribe the chaotic movement of the particles • Table 17.6, p.492 summarizes the properties of solutions, colloids, and suspensions

46. Mixtures that are NOT Solutions • Emulsions- colloids dispersions of liquids in liquids • an emulsifying agent is essential for maintaining stability • oil and water not soluble; but with soap or detergent, they will be. • Oil and vinegar dressing? • Mayonnaise? Margarine?