Chapter 13-14: Mixtures and Aqueous Solutions

1. Chapter 13-14:Mixtures and Aqueous Solutions We use solutions all the time What are they? Where do we find them? How do we describe them?

2. Soluble versus insoluble • Some solids are soluble in water, ie: table salt, NaCl. Soluble means: able to be dissolved. • Soluble ionic solids (made of cation and anion) dissociate into their ions in water. • Soluble covalent solids (like sugar) dissolve because they are relatively polar. • In a solution, the dissolved particles cannot be easily seen or separated from the solution. • Alloys are solutions of metals!

3. Parts of a solution • The dissolving medium is the solvent (what does the dissolving…the dissolver) • The dissolved substance is the solute (what gets dissolved…the dissolvey) • The solute andsolventtogether form the solution. • Solvents and solutes can be any phase. solution

4. Special types of mixtures - Suspensions • Suspensions • mixtures where the solutes particles are very large, so they don’t completely dissolve into their solvent. • Solute particles will settle out of the solution if left undisturbed. – this creates two phases. • Muddy water and Italian salad dressing are good examples of suspensions.

5. Special types of mixtures - Colloids • Colloids • mixtures where the solute particle is smaller than particles in a suspension, but not small enough to dissolve. • Colloids have two phases: • Dispersed phase – the solute • Dispersing medium – the solvent. • Mayonnaise and hair gel are good examples of colloids. • There are 7 types of colloids, found on page 398…

6. 7 Types of Colloids Page 398 Two groups of colloids: Heterogeneous colloids – two phases are clearly seen Homogeneous colloids – appears to be one phase

7. The Tyndall Effect John Tyndall, Brittish, c1860 • The Tyndall effect allows us to distinguish between solutions, colloids, and suspensions. • It works by shining a beam of light into the mixture. If… • Light doesn’t pass through • the mixture is a suspension or a heterogeneous colloid. • Light passes through unobstructed • the mixture is a solution. • Light passes, but the beam can be seen in the mixture • the mixture is a homogeneous colloid

8. Electrolytes • Electrolytes • Solutions that conduct electricity. • Ionic solutions are electrolytes. • Covalent solutions are nonelectrolytes. • Is saltwater (NaCl in water) an electrolyte? • Is sugar water (C6H12O6 in water) an electrolyte? • Conductivity tester • can tell us if a solution is an electrolyte, and sometimes, how strong an electrolyte is.

9. Solubility • Solubility • The extent to which a solute will dissolve in a solvent. (how much solute will dissolve) • High solubility • large amounts of solute will dissolve in a solvent • Low solubility • only small amounts of solute will dissolve • Increasing temperature increases the solubility of solids in liquids • Increasing temperature decreases the solubility of gasses in liquids! …

10. Solid-Liquid and Gas-Liquid solubility with temperature

11. Gasses in liquids • In addition to cold temperatures, high pressures increase solubility of gasses in liquids. • Henry’s Law: • solubility of a gas in a liquid increases with increasing pressure of that gas above the liquid.

12. Like Dissolves Like! • Some solvents are polar, like magnets, having partialnegative and partial positive ends. (H2O) • Other solvents are nonpolar, having no “+” “-” poles • Polar solutes tend to dissolve well in polar solvents… • Nonpolar solutes tend to dissolve well into nonpolar solvents. • “Like dissolves like” • Water is very polar. Does it dissolve polar substances or non polar substance?

13. Saturation • Saturated Solution • solution has as much solute as it will allow (equal to solubility) • Unsaturated Solution • more solute can dissolve into solution (less than solubility) • Supersaturated Solution • too much solute in solution-some will fall out (more than solubility) • We express the quantitative amount of solute in a solution with concentration …

14. Concentration - Molarity • Concentration • the quantitative amount of solute present in a solution • Molarity (M) – moles/liter • number of moles solute in liters of solution • We can use the T-chart method to find moles of solute present.

15. Try these Molarity questions • What is the concentration [in Molarity] when 3 moles of NaCl are dissolved in 2 Liters of water? • How much (in liters) of a 0.1 M solution do you need to get 2 moles of solute? • How many moles of NaOH are present in 300mL of a 1M solution? • How many grams of HCl are found in 100mL of a 2M solution? 1.5 M “molar” 20 L .3 moles 7.2 grams

16. Concentration - Molality • Molality (m) – moles/kilogram – number of moles solute in kilogram of solvent. • molality is used less often, but is important when we discuss colligative properties • Remember, where Molarity is “per liter solution”, molality is “per kilogram of solvent”

17. Try these molality questions • What is the concentration in molality when 2 moles of NaCl are dissolved in 4kg of water? • How many moles of solute are present in 1 kg of a 3 m solution? • What mass of water do you need to add to 4 moles of NaCl to make a 2 m solution? • What is the molality of a solution created by dissolving 3.5 moles methanol in 340g of CCl4. .5 m “molal” 3 moles 2 kilograms 10.3 m

18. Solution Preparation • By solid dissolving: • 1. calculate how many grams are needed to create our volume of our desired molarity solution • 2. weigh out that mass, and add it to a volumetric flask • 3. add some water and allow to dissolve • 4. add water to the desired volume • By dilution of a standard solution: • 1. use the relationship M1V1=M2V2 • 2. calculate volume of standard molarity solution to use to get desired volume of desired molarity solution. End of chapter 13

19. Begin C14: Colligative Properties • “Colligative” means depends on amount. • A colligative property of a solution depends on the amount of solute dissolved in solution. • Physical Properties of a solution change because solute particles act like impurities, getting in the way of solvent particles. • We add impurities to lower freezing points, increase boiling points, or reduce vapor pressure. The more impurities, the greater the change Ethyl glycol is added to water in your car’s radiator to increase water’s boiling point.

20. Osmosis and Osmotic Pressure • Osmosis is the travel of a solvent from an area of low concentration (high purity) to high concentration (low purity). • Examples of osmosis: • A Cucumber placed in a conc. NaCl solution (brine) loses water, shrivels up, and becomes a pickle. • Limp carrots and celery, placed in water, become firm because water enters via osmosis. From pure to impure

21. Strong/Weak Electrolytes • Recall that a solid compound made up of a cation and anion is called a salt. • Salts that dissolve completely into their ions when put in water dissociate completely. • Salts that dissociate completely form strong electrolytes – solutions that conduct electricity well. • Some salts only partially dissociate, forming weak electrolytes – solutions that conduct electricity, but do so poorly.

22. H+ / OH-Ions – (Acids and Bases) • When a H+ ion is released into solution, a H3O+ ion is produced, called Hydronium ion. • When a OH- ion is produced, we call this a Hydroxide ion. • Hydronium (H3O+) and Hydroxide (OH-) are the fundamental ions involved in acid/base chem. • Acids that dissociate completely, releasing H+ ions form strong electrolytes. • Bases that dissociate completely releasing OH- ions form strong electrolytes. End of chapter 14 – Problem set on next slide

23. Chapter 14 Quiz Review Problems • In your text, on page 447, do the following (for a grade) as a quiz review: • #1-2 : Solubility of salts • #13-15 : Dissociation of aqueous salts • #16-18 : Precipitation reactions, net ionic equations, and spectator ions