1. NOTES: Chapter 15 – Water and Aqueous Systems

2. Chapter Objectives: Describe hydrogen bonding in water and how it explains water’s unique properties and behaviors; Explain why ice floats; Describe a solution and its components; Explain why some substances dissolve in others; Distinguish among strong electrolyte, weak electrolyte, and nonelectrolyte solutions; Explain how colloids and suspensions differ from solutions.

3. 15.1 - Liquid Water and Its Properties Many of water’s special properties and behaviors can be attributed to its ability to form hydrogen bonds, including:  surface tension,  specific heat capacity, and  heat of vaporization.

5. Water is a POLAR molecule the bonds that hold the hydrogen and oxygen atoms together are POLAR COVALENT BONDS…  this means that the electrons that make up the covalent bonds are shared UNEQUALLY …they spend more time near the oxygen atom than they do near the hydrogen atom  so, a water molecule has a positive end and a negative end RECALL…

6. Water is POLAR… POLAR WATER MOLECULES attract one another, as well as ions and other polar molecules

7. Water Molecules… Water also exhibits HYDROGEN BONDING Each water molecule can form a maximum of 4 hydrogen bonds with neighboring water molecules

8. HYDROGEN BOND: HYDROGEN BOND: A relatively strong intermolecular force in which a hydrogen atom that is covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair on another electronegative atom in the same molecule or one nearby

12. SURFACE TENSION: SURFACE TENSION: the inward force that tends to minimize surface area  Hydrogen bonding increases this inward force; H 2 O is more attracted to itself than the air molecules around it; this leads to a spherical shape (figure 15.4 in the text)  Surfactants, such as soap or detergent, can be used to decrease surface tension in water by interfering with the hydrogen bonding (see figure 15.5)

13. SURFACE TENSION: -water has greater surface tension than most liquids because at the air/water interface the surface water molecules are H-bonded to each other and to the water molecules below (rather than to the air molecules) -causes water to “bead” -creates a “skin” on the surface

15. Related Properties: ► COHESION: molecules are held together by H bonds contributes to upward movement of water in plants from roots to leaves

17. Properties of Water ► ADHESION: water sticks to other surfaces (either by polar attraction or hydrogen bonds) **explains the meniscus in a graduated cylinder!

18. SPECIFIC HEAT CAPACITY: SPECIFIC HEAT = the energy it takes to heat 1 gram of a substance by 1˚C Hydrogen bonds allow water to hold a large amount of energy, leading to a high specific heat capacity EX: the specific heat of: water= 4.18 J / g ˚C = 1.0 cal/g ˚C iron= 0.447 J / g ˚C = 0.11 cal/g ˚C

19. Water’s high specific heat means that it resists temp. changes when it absorbs or releases heat *Heat is absorbed to break H-bonds; and given off when they form SPECIFIC HEAT CAPACITY:

22. Water Vapor and Ice because of hydrogen bonding, water absorbs a large amount of heat as it evaporates (or vaporizes) (it takes a lot of heat energy to break those hydrogen bonds and convert liquid water to water vapor)

24. HEAT OF VAPORIZATION: HEAT OF VAPORIZATION = the energy required to vaporize 1 g of a substance at its boiling point  Hydrogen bonding is a strong intermolecular force that binds groups of molecules together, leading to a higher heat of vaporization required to separate the molecules  heat of vaporization of water = 2260 J/g  compare to liquid methane (which has no hydrogen bonding): H vap = 510 J/g

25. EVAPORATIVE COOLING: after high temp. molecules have evaporated, the remaining liquid is cooler (ex: sweating) -stabilizes temp. in aquatic ecosystems -protects organisms from overheating *(as 1 g of water evaporates from our skin, 539 g of body cools by 1°C)

26. The air in the shower stall is at the same temperature as the air outside, but there's less water vapor outside to condense on the skin.

27. EVAPORATIVE COOLING SPECIFIC HEAT (WHILE WATER IS HEATING) HEAT OF VAPORIZATION (WHILE CHANGING STATE)

28. CONDENSATION: the reverse of vaporization is CONDENSATION (gas condensing into liquid) as water vapor condenses into liquid water, a large amount of heat is released **this is why you can be severely burned if steam condenses on your skin!

29. Water and Climate:  the processes of vaporization and condensation allow water (especially large bodies of water) to moderate the temperature of its surroundings  climates near the ocean tend to fluctuate in temperature less than areas farther inland.

30. Liquid Water vs. Solid Water Solid water is less dense than liquid water

31. Liquid Water vs. Solid Water Hydrogen bonds force water molecules into a rigid framework as kinetic energy decreases (solid water), leading to a structure with a larger volume (and therefore less dense!) than when kinetic energy is higher and water molecules are allowed to slide past one another (liquid water)

32. Benefits of “Floating” Ice: a layer of ice on top of a pond or lake acts as an insulator for the water beneath Oceans and lakes don’t freeze solid!! aquatic life can survive through the winter

34. Melting and Freezing: HEAT OF FUSION = the amount of heat required to melt 1 g of solid to liquid at its melting point Heat of fusion of water = 334 J/g people used to place barrels of water in their root and vegetable cellars to prevent the temperature in there from dropping below freezing in the winter…WHY DID THIS WORK???

35. 15.2: Aqueous Solutions PURE WATER never exists in nature because water dissolves so many substances water samples containing dissolved substances are called AQUEOUS SOLUTIONS

36. WATER IS THE SOLVENT OF LIFE due to its polarity, water is a versatile solvent cell

37. AQUEOUS SOLUTIONS: Solution – mixture consisting of a solute and a solvent Solute – dissolved particles Solvent – dissolving medium Aqueous Solution – solution in which the solvent is water

38. SOLVATION: SOLVATION = Process by which a solute dissolves Solvent has to be attracted to the solute for solvation to occur; if a solute dissolves, it is said to be SOLUBLE If the solute is more attracted to itself rather than the solvent, solvation will not occur and the solute is said to be INSOLUBLE

39. SOLVATION: “Like dissolves like” is a good rule of thumb to tell whether or not a solute will dissolve in a solvent; in other words, solutes and solvents with similar polarities will make a solution  polar with polar (or ionic)  nonpolar with nonpolar

40. SOLVATION: example: NaCl in water δ + H attracts Cl - ions δ - O attracts Na + ions A similar scheme can be used to illustrate the dissolution of a polar molecular substance in water.

41. Compounds that are HYDROPHILIC (“water loving”), are soluble in water: -ionic compounds: charged regions of polar water molecules have an electrical attraction to charged ions -polar compounds: charged regions of polar water molecules are attracted to oppositely charged regions of other polar molecules “Like Dissolves Like”:

42. Compounds that are HYDROPHOBIC (“water fearing”) and are insoluble in water: -Nonpolar compounds: symmetric distribution in charge, or composed of nonpolar bonds “Like Dissolves Like”:

43. ELECTROLYTES: ELECTROLYTES: Compounds that will conduct an electric current in aqueous solution or the molten state Ionic compounds are electrolytes because they have a positive and negative charge that can conduct an electric current Electrolytes can be classified as strong or weak depending on the extent to which the ions dissociate (come apart) in solution

44. NONELECTROLYTES: NONELECTROLYTES: Compounds that will NOT conduct electricity in aqueous solution or the molten state Many molecular compounds are nonelectrolytes (compounds of carbon especially)

45. Water of Hydration / Hydrated Compounds: many substances can exist as a crystal structure that includes water molecules such compounds are called: HYDRATES EXAMPLE: CuSO 4 5H 2 O -shows that for each pair Cu 2+ and SO 4 2- ions, there are 5 water molecules

46. 15.3: Heterogeneous Aqueous Systems Aqueous mixtures can fall into three categories based on particle size: Solutions Suspensions Colloids

47. SOLUTIONS: SOLUTION = a homogeneous mixture may be:  gas in gas:AIR  liquid in gas:water vapor in AIR  gas in liquid:CO 2 in water (soda)  liquid in liquid:HCl in water  solid in liquid:NaCl in water  solid in solid:Cu in Ag (sterling silver)

48. SUSPENSIONS: SUSPENSION = mixtures from which particles settle out upon standing particle size much larger than in solutions (average diameter greater than 100 nm, compared to particles in solutions, 1 nm) EXAMPLES:  clay in water  sand in water

49. COLLOIDS: COLLOIDS = heterogeneous mixtures containing particles intermediate in size (between 1 and 100 nm) the particles are dispersed through the medium, which can be solid, liquid, or gas many have a cloudy appearance when concentrated

50. COLLOID Properties: Tyndall Effect – scattering of visible light in all directions **the same way dust particles scatter a beam of light in a dusty room

51. COLLOID Properties: Emulsions – colloid dispersions of liquids in liquids Emulsions require an EMULSIFYING AGENT EXAMPLE: grease and water don’t mix; but add soap or detergent and they do!

52. Properties of Solutions, Colloids, and Suspensions: PropertySolutionColloidSuspension Particle Type Ions, atoms, small molecules Large molecules or particles Large particles or aggregate Particle Size 0.1 – 1 nm1 – 100 nm> 100 nm Effect of Light No scatteringExhibits Tyndall effect Effect of Gravity Stable no separation Unstable, will separate Filtration No particles on filter Particles on filter Uniformity HomogenousBorderlineheterogeneous Example NaCl (aq) Milk, fogDirt in water