Chapter 12-13: Mixtures and Aqueous Solutions

1. Chapter 12-13: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.

3. In a solution, the dissolved particles cannot be easily seen or separated from the solution. • Alloys are solutions of metals!

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

5. 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.

6. 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.

7. Colloids • Mayonnaise and hair gel are good examples of colloids. • There are 7 types of colloids, found on page 404…

8. 7 Types of Colloids Page 404 Two groups of colloids: Heterogeneous colloids – two phases are clearly seen Homogeneous colloids – appears to be one phase

9. 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…

10. Results of Tyndall Effect • 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

11. The Tyndall Effect

12. Electrolytes • Electrolytes • Solutions that conduct electricity. • Ionic solutions are electrolytes. • Covalent solutions are nonelectrolytes.

13. What do you think? Is saltwater (NaCl in water) an electrolyte? Is sugar water (C6H12O6 in water) an electrolyte? Conductivity tester (meter) • can tell us if a solution is an electrolyte, and sometimes, how strong an electrolyte is.

14. Warm Up- Would the following solutes be electrolytes or nonelectrolytes CaCl2 CCl4 SrBr2 How do you know??

15. Warm Up Be sure you hand in any HW that you may not have had on test day… Article Analysis Review 12.2 Enthalpy WS 10.4 and 16.1 Notes

16. 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

17. Solubility • Increasing temperature increases the solubility of solids in liquids. • Increasing temperature decreases the solubility of gases in liquids! …

18. Reading Solubility Curves

19. Solid-Liquid solubility with temperature

20. Gas-Liquid solubility with temperature

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

22. Like Dissolves Like! • Some solvents are polar, 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.

23. Like Dissolves Like Water is very polar. Does it dissolve polar substances or non polar substance?

24. Saturation • Saturated Solution • solution has as much solute in it as it will allow (equal to solubility) • Unsaturated Solution • more solute can still 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 …

25. Solid-Liquid solubility with temperature

26. Solubility Graph Practice Work on the front side— 15 minutes Work on the back side- 15 minutes

27. KNO3 Lab

28. Warm Up

29. Concentration - Molarity • The “Stoichiometry” of Solutions • Concentration • the quantitative amount of solute present in a solution • Molarity (M) – moles/liter • number of moles solute in liters of solution

30. 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

31. 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

32. Solution Preparation By dilution of a standard solution: • Use the relationship M1V1=M2V2 2. Calculate volume of the “standard solution” to use to get desired volume of desired molarity solution.

33. Solution Formation • The nature of the solvent and solute affects whether a substance will dissolve • Other factors determine how fast a soluble substance dissolves • Agitation (shaking) • Temperature • Solute particle size

34. Will it dissolve? (Solubility Rules) • Not all ionic solids (salts) will dissolve. • We use solubility rules to decide if the substance will dissolve. • Salts containing… • Alkali metal cations (+) are soluble. • NH4+, NO3-, SO42- are soluble. • Pb+, Ag+, Hg2+ are insoluble. • CO3-, PO43-, S2- are insoluble. • Which of the following salts are soluble? • NaCl, HgCO3, Ca(NO3)2, AgF, PbI2, FeSO4 BaSO4, SrSO4, and PbSO4 are insoluble

35. Dissociation and Ions Present • Dissociation = a salt dissolving into its ions: • How many moles of ions are in a solution of 1 mole of NaCl? • How many moles of ions are in solutions of 1 mole of each of the following?:

36. Net Ionic Equations • When we write a balanced chemical equation, we show all species present (all reactants and all products): • In a net ionic equation, we show only precipitates formed, and the reactants that form them: • The chemicals that stay ions are called spectator ions, And are left out (Na+, NO3-) Remember to Balance

37. Net Ionic Equation Practice • Write the net ionic equations for the following:

38. 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.