The Structure of Water

1. The Structure of Water Water is a Polar Molecule. This means that it has a positive and a negative end. This is why water is “sticky.”

2. It also allows us to BEND WATER!!!!

3. Why is water polar? It has to do with something called ELECTRONEGATIVITY. This is an atom’s ability to attract electrons. There is a trend of electronegativity on the periodic table.

4. The Solution Process • Ionic and covalent molecules dissolve differently • Covalent – share e- (from the RS of the PT) • Ionic – give and take e- (from opposite sides of the PT)

6. Dissociation – when the ions separate from the crystal structure • We write dissociation equations like this: AgCl(s)  Ag+(aq) + Cl-(aq) (Pay attention to the states!) • Solvation – the process of surrounding the solute ions with solvent molecules • Hydration – when the solvation process involes water as the solvent. The solute ions are said to be HYDRATED.

7. Heat of Solution • E is required to break attractive forces. • Therefore, the separation of solute particles from one another and the separation of solvent particles from one another are endothermic processes. THEY USE HEAT. • Example:

8. The attraction between solute and solvent particles during solvation is exothermic. THEY RELEASE HEAT. • Example: “Heat of solution” is the next energy change! Energy Absorbed > Energy Released ENDOTHERMIC (cold) Energy Absorbed < Energy Released EXOTHERMIC (hot)

9. Assignment: • Ex 44 1-10 • Draw diagrams of the solution process for • An ionic compound • A covalent molecule This assignment is due tomorrow!!!

10. Factors that Affect the Rate of Solubility 1. Agitation (stirring) 2. Temperature 3. Surface Area

11. Terms you need to know…. Soluble – capable of being dissolved in a particular solvent Insoluble – cannot be dissolved in a particular solvent Miscible – liquids that mix in all proportions and have no max concentration Immiscible – two liquids that form separate layers instead of dissolving

12. Saturated – when a solution contains the max amount of solute it can dissolve at a given temp Unsaturated – when a solution doesn’t contain the max amount of solute it can dissolve at a given temp Supersaturated – more than the max amount of solute dissolved for a specific temp (made by dissolving at high temp and cooling)

13. Saturated solution • The animation represents a saturated solution: • the blue "molecules" escape into solution from the ordered crystal. At the same time, molecules are coming out of solution and depositing on the solid. Since this is a continual process and the concentrations do not change, it is called dynamic equilibrium.

14. Solubility – the concetration of a saturated solution at a specific Temp and Pressure Ex – the solubility of Na2SO4 at 0°C is 4.76 g/100mL H2O

15. Saturated Solubility curve Supersaturated Unsaturated

17. 3 Ways to Saturate a Solution • Add more solute to the saturation point • Evaporate the solvent and……

18. 3. Decrease the temperature to the saturation point

19. SOLUBILITY Depends on: Temperature Pressure Chemical natures of the solute and solvent

20. Temperature In general, as the temp increases, so does solubility (for endothermic reactions) Example: Cold packs - The dissolving reaction is endothermic - requires heat. Heat comes from the surroundings, and the pack feels cold.

22. If the reaction is exothermic, the increase in heat will actually cause the solubility to DECREASE. • This is not very common.

23. Reading graph: at 38 °C the solubility of copper sulphate, CuSO4, is 28g per 100g of water.

24. Reading graph: at 84 °C the solubility of potassium sulphate, K2SO4, is 22g per 100g of water.

25. Ex Q1: How much potassium nitrate will dissolve in 20g of water at 34 °C? • At 34 °C the solubility is 52g per 100g of water, so scaling down, 52 x 20 / 100 = 10.4g will dissolve in 20g of water.

26. Ex Q3: 100 mL of saturated copper sulphate solution was prepared at a temperature of 90 °C. What mass of copper sulphate crystals form if the solution was cooled to 20 °C? • Solubility of copper sulphate at 90 °C is 67g/100g water, and 21g/100g water at 20 °C. Therefore for mass of crystals formed = 67 - 21 = 46g (for 100 cm3 of solution)

27. Exercise 43: 1-9

28. Solubility of Gases • The solubility for gases always decreases when the temperature increases.

29. Solubility of Gases: Temperature • ↑ Temp = ↑ KE Particles will move more, break IMFs and the gas will escape! • Ex – pop that has been sitting open at room temp will be flat as the CO2 has come out of solution.

30. Pressure and Solubility of Gases • With a change in pressure - liquids and solids show no change in solubility • Gases increase in solubility with an increase in pressure. • If the pressure is increased, the gas molecules are "forced" into the solution. The concentration of gas molecules in the solution have increased!

31. Pressure and Solubility of Gases • Carbonated beverages provide the best example of this phenomena. All carbonated beverages are bottled under pressure to increase the carbon dioxide dissolved in solution. • When the bottle is opened, the pressure above the solution decreases. As a result, the solution effervesces and some of the carbon dioxide bubbles off. • Quiz: Champagne continues to ferment in the bottle. The fermentation produces CO2. Why is the cork wired on a bottle of champagne? • Answer: As more CO2is formed , the pressure of the gas increase.The wire is to prevent the cork from blowing off.

32. Quiz: If a diver had the "bends", describe how this can be treated. • Answer: Decompression chambers are used to keep a high pressure and gradually lower the pressure.

33. Quiz: The amount of dissolved oxygen in a mountain lake at10,000 ft and 50oF is __?_ than the amount of dissolved oxygen in a lake near sea level at 50oF. • Answer: Less at higher altitude because less pressure. • A Coke at room temperature will have __?_ carbon dioxide in the gas space above the liquid than an ice cold bottle. • Answer: More gas, because the warm coke can hold less of the gas in solution.

34. BONUS ASSIGNMENT • THERMAL POLLUTION • 3 page research paper on thermal pollution. • Include references • If you plaigerise, you get ZERO.

35. Freezing Point Depression • This happens when the freezing point is lowered by adding a solute to a solvent. • In order for a liquid to freeze, it must achieve a very ordered state that results in the formation of a crystal. • If there are impurities (solute) in the liquid, the liquid is less ordered – the sol’n is now more difficult to freeze and a lower temp is required.

36. Which decreases FP more? • NaCl • CaCl2 NaCl  Na+ + Cl- CaCl2  Ca 2+ + 2Cl- Which has more particles????

37. Boiling Point Elevation Review – VP – the pressure at which a liquid is in equilibrium with its vapour. Boiling – the temp at which the vapour pressure equals the pressure above the liquid

38. If we add solute to a solvent, the vp of the sol’n is lowered. This happens because: • At the surface of the sol’n (where evaporation occurs) there are fewer solvent particles due to the presence of solute particles – lowers vp

39. The solute particles absorb energy and will reduce the energy available to evaporate the solvent – lowers vp • Energy is required to overcome the IMFs between the solute and solvent particles – lowers vp

40. If VP is lowered, the temp must be raised to have the VP = the P above the solution • Example • H2O boils at 100°C • H2O and NaCl boils at > 100°C

41. You need to remember: • Adding a solute to a solvent will: • LOWER THE FREEZING POINT • RAISE THE BOILING POINT • Make sure you’re able to tell me why!!!!

42. LAB TOMORROW!!! • You need to get into a group • You need to bring: • 2 small ziplock bags • 2 large ziplock bags • 1 cup of sugar • 500 mL of milk • toppings

43. Ex 49: 1-8

44. Like Dissolves Like • Remember – “like dissolves like” • This means that polar molecules dissovle polar molecules • Non-polar molecules will dissolve in non-polar molecules

45. g solute g solute x 100 x 100 = g solution g solute + g solvent moles of solute volume in liters of solution Units of Concentrations amount of solute per amount of solvent or solution Percent (by mass) = Molarity (M) = moles = M x VL

47. Examples What is the percent of KCl if 15 g KCl are placed in 75 g water? %KCl = 15g x 100/(15 g + 75 g) = 17% What is the molarity of the KCl if 90 mL of solution are formed? mole KCl = 15 g x (1 mole/74.5 g) = 0.20 mole molarity = 0.20 mole/0.090L = 2.2 M KCl

48. Examples: Example 1:What is the concentration when 5.2 moles of hydrosulfuric acid are dissolved in 500 mL of water? Step one: Convert volume to liters, mass to moles. 500 mL = 0.500 L Step two: Calculate concentration. C = 5.2 mol/0.500 L = 10mol/L

49. Example 2: What is the volume when 9.0 moles are present in 5.6 mol/L hydrochloric acid? • Example 3: How many moles are present in 450 mL of 1.5 mol/L calcium hydroxide? • Example 4: What is the concentration of 5.6 g of magnesium hydroxide dissolved in 550 mL? • Example 5: What is the volume of a 0.100 mol/L solution that contains 5.0 g of sodium chloride?

50. How many Tums tablets, each 500 mg CaCO3, would it take to neutralize a quart of vinegar, 0.83 M acetic acid (CH3COOH)? 2CH3COOH(aq) + CaCO3(s)  Ca(CH3COO)2(aq) + H2O + CO2(g) a quart moles acetic acid = 0.83 moles/L x 0.95 L = 0.79 moles AA the mole ratio mole CaCO3 = 0.79 moles AA x (1 mole CaCO3/2 moles AA) = 0.39 moles CaCO3 molar mass mass CaCO3 = 0.39 moles x 100 g/mole = 39 g CaCO3 number of tablets = 39 g x (1 tablet/0.500g) = 79 tablets