2. Heterogeneous Mixtures Heterogeneous Mixtures: Not evenly blended Suspensions: a mixture containing particles that settle out if left undisturbed Colloids: a mixture containing particles that are small enough not to settle out. Held up by Brownian motion (erratic motion of particles bumping into one another Tyndall Effect- light scattered due to Brownian motion

4. Homogenous Mixtures Solutions= homogeneous mixtures containing two or more substances Atoms, molecules, or ions evenly mixed (same composition/dissolved). Solutions may exist as gases, liquids, or solids

5. Solutions Solute: The substance being dissolved in a solution Generally lesser in quantity Solvent: The dissolving medium in a solution Generally greater in quantity Normally WATER!!!!

6. Check For Understanding In the following examples, identify what is the SOLUTE and what is the SOLVENT #1 Making Kool-Aid with water #2 Putting sugar in your coffee #3 Dissolving an Alka-Seltzer Tablet #4 Mixing iodine into ethanol

7. Solvation Solvation is the process of surrounding solute particles with solvent particles to form a solution.

8. Factors that Affect Rate of Solvation Agitation- shaking or stirring Moves dissolved solute particles away from the contact surfaces more quickly allowing new collisions between solute and solvent particles. Surface Area- breaking the solute into small pieces Allows more collisions to occur between solute and solvent particles. Temperature- solubility of solids increase as temperature increases. As temperature increases, particles move faster allowing more collisions between solute and solvent particles

9. Two Types of Solutes Electrolytes: Solutes that gives solution ability to conduct electricity Normally ionic compounds Ex: NaCl Nonelectrolytes: Solution does NOT conduct electricity Normally covalent compounds Ex: Sugar.

10. Solubility and Miscibility Soluble= a substance that can be dissolved in a particular solvent (e.g. salt in water) Insoluble= a substance that cannot be dissolved in a particular solvent (e.g. sand in water) Miscible= two liquids that can be mixed (e.g. vinegar and water) Immisible= two liquids that cannot mix (e.g. oil and water)

11. Solubility and Miscibility General Rule: Like Dissolves Like Polar solvents dissolve polar solutes (and ionic solutes) Nonpolar solvents dissolve nonpolar solutes Two polar liquids are miscible Two nonpolar liquids are miscible A polar and a nonpolar liquid are immiscible

12. Solubility of a Gas Gases can dissolve into liquids as well (e.g. soda) Pressure increases solubility The more pressure of the gas on the liquid, the more that gas will dissolve. Temperature decreases solubility Solubility of a gas decreases as temperature increases.

14. Saturation Saturated solutions contain the maximum amount of dissolved solute for a given solvent at a specific temperature Unsaturated solutions contain less dissolved solute for a given temperature than possible Supersaturated solutions contain more dissolved solute than a saturated solution Dissolved at high temperatures, then cooled slowly

15. Concentration Concentration is a measure of how much solute is in a solution Molarity Molality Mole Fraction Percent by Mass Percent by Volume

16. What is Molarity? Molarity: Tells us how CONCENTRATED a solution is based on moles of solute and volume of solution. Formula Molarity = moles of a solute liters of a solution Symbol for Molarity (M).

17. Example 1 What would be the molarity of a solution that had 1.5 moles of NaCl dissolved in a 0.76 liter solution? Molarity = 1.5 moles = 2.0 M 0.76 liters

18. Example 2 What is the molarity of a solution that contained 56 grams of KCl in 1.45 liters of solution? 56 grams KCl x 1 mol KCl = 0.75 mol 74.6 g KCl Molarity = 0.75 mol = 0.52 M 1.45 liters

19. Example 3 How many grams of MgO would I need to dissolve if I wanted to create 6.8 L of a solution that was 0.23 M? mol solute =.23 M = 1.6 mol MgO 6.8 liters 1.6 mol MgO x 40.3 g MgO = 64 g MgO 1 mol MgO

20. There are 2.3 moles of NaCl in 0.45 L solution. What is the molarity? Example 4

21. 120.1 g of CaCO 3 is put into 1.22 L of solution. What is the molarity? Example 5

22. You need to make 2.00 L of a 6.00 M solution. How many grams of HCl do you need? Example 6

23. You need to make 1.5 L of a 2.0 M solution. How many grams of NaOH do you need? Example 7

24. Molality Molality states CONCENTRATION in terms of moles solute and mass of solvent. Formula: Molality=mol of SOLUTE kg of SOLVENT Molality Symbol (m)

25. What is the molality of a solution that contains.500 mol HC 2 H 3 O 2 in 0.125 kg H 2 O? Example 1

26. Example 2 What is the molality of a solution in which 27.0 g of S is dissolved in 450. g of alcohol? 27.0 g S x 1 mole S = 0.841 mol S 32.1 g S 450. g x 1 kg = 0.450 kg alcohol 1000 g 0.841 mol S = 1.87 m 0.450 kg alcohol

27. What is the molality of a solution that contains 58.0 g HNO 3 in 0.500 kg H 2 O? Example 3

28. What mass of HC 2 H 3 O 2 must be dissolved in 800. g H 2 O to produce a 6.25 m solution? Example 4

29. What mass of water is required to dissolve 100. g NaCl to prepare a 1.50 m solution? Example 5

30. Mole Fraction Shows CONCENTRATION by relating moles of solute and moles of solvent to moles of the whole solution Formula

31. What is the mole fraction of sulfur dioxide in gas containing 128 g sulfur dioxide dissolved in every 1500 g of carbon dioxide? Example 1

33. A gas mixture contains 50.4g of dinitrogen monoxide gas and 65.2g of oxygen gas. What is the mole fraction of dinitrogen monoxide? Example 2

35. Example 3 You have a solution that contains 39.5 grams sodium chloride in 58.7 grams of water. What is the mole fraction for EACH SUBSTANCE? Change grams of NaCl to moles Change grams of water to moles So total moles = 3.94 mol X NaCl = X water = 0.675 mol NaCl 3.26 mol water 0.675 mol 3.94 mol = 0.171 3.26 mol 3.94 mol = 0.827

36. A gas mixture contains following mole fractions: N 2 0.450, O 2 0.334, CO 2 0.023, SO 2 0.017, and N 2 O 4 0.120 The mixture also contains argon. What is the mole fraction for argon? Total mole fraction =1 0.450 + 0.334 + 0.023 + 0.017 + 0.120 = 0.944 1 - 0.944 = 0.056 which means 0.056 = X Ar Example 4

37. Percent by Mass Tells what percentage of a solution is the solute based on MASS

38. Example 1 34 grams of Kool-Aid powder is mixed with 154 grams of water. What would be the percent by mass of the Kool- Aid powder? 34 grams solute 188 grams solution x 100= 18.09 % = 18 %

39. Example 2 You add 13.5 grams of salt to a pot full of 284 grams of water. What is the percent by mass for the salt? 13.5 grams solute 297.5 grams solution x 100= 4.54 %

40. Example 3 If you have 350 g of solution with a % by mass of 5.2% Ca(OH)2, how many grams of Ca(OH)2 are in the solution?

41. Percent by Volume Tells what percentage of a solution is the solute based on VOLUME

42. Example 1 0.32 L of bleach is mixed with water to make a solution of 3.77 L. What is the percent by volume? 0.32 L solute 3.77 L solution x 100= 8.49 % = 8.5 %

43. Example 2 You put 450 mL of detergent in your washer that holds 80.0 L of solution. What is the percent by volume of detergent in your washer? 0.45 L solute 80.0 L solution x 100= 0.56 %

44. Example 3 If you have 25.0 mL of solution with a 32.0% by volume of ethanol, how many mL of ethanol are in the solution?

46. Colligative Properties Physical properties of solutions that are affected by the number of particles but not by the identity of dissolved solute particles are called colligative properties. Colligative means depending on the collection Colligative Properties include Vapor pressure lowering Boiling point elevation Freezing point depression

47. Vapor Pressure Lowering Vapor pressure is the pressure exerted in a closed container by liquid particles that have escaped the liquid’s surface and entered the gaseous state. Reaches a state of equilibrium If a solute is added to a liquid, a mixture of solvent AND solute will occupy the surface area, causing there to be less liquid particles near the surface to escape the solution. If less particles escape the liquid phase, there is less vapor pressure. Thus, vapor pressure lowering is due to the number of solute particles in the solution.

48. Vapor Pressure Lowering

49. Boiling Point Elevation Boiling Point is temperature at which the vapor pressure is equal the atmospheric pressure. When a solute lowers the vapor pressure of a solvent, the boiling point is also affected. More heat is needed to supply additional kinetic energy to raise the vapor pressure to atmospheric pressure. The temperature difference between a solution’s boiling point and a pure solvent's boiling point is called the boiling point elevation.

50. Freezing Point Depression At a solvent's freezing point temperature, particles no longer have sufficient kinetic energy to overcome intermolecular attractive forces. The freezing point of a solution is always lower than that of the pure solvent. Solute particles interfere with the attractive forces among solvent particles. A solution's freezing point depression is the difference in temperature between its freezing point and the freezing point of the pure solvent.