1. Solutions!

2. What is a solution? A homogeneous mixture! Made up of a solute and solvent.

3. 2 Parts of a Solution Solvent – Does the dissolving. Usually present in the larger amount. Solute – Is dissolved. Usually present in the smaller amount.

4. Aqueous Solutions Water is a good solvent because the molecules are polar. The oxygen atoms have a partial negative charge. The hydrogen atoms have a partial positive charge.

5. Hydration The process of breaking the ions of salts apart. Ions have charges and attract the opposite charges on the water molecules.

6. Hydration H H O H H O H H O H H O H H O H H O H H O H H O H H O

7. NaCl  Na + + Cl - Al(NO 3 ) 3  Al 3+ + 3NO 3 - MgCl 2  Mg 2+ + 2Cl -

8. Soluble Vs. Insoluble Soluble: when attraction between ions and water molecules stronger than the attraction between ions Insoluble: when attraction between the ions is stronger than the attraction between the ions and water molecules

9. Electrolytes Conducts an electric current when placed in aqueous solution or in molten state. (Soluble) Non-electrolytes do not conduct an electric current in either state. (Insoluble)

10. Properties of Solutions Solubility – how much dissovles Determined by a few factors. - Nature of solvent and solute. - Agitation. - Surface area of solute. - Temperature.

11. Nature of Solvent and Solute “Like dissolves like.” Water is polar. It has small charges. It dissolves ionic and polar compounds. Non-polar will not dissolve.

12. Agitation Shaking increases the interaction between water and the solute. What do you do when you make Kool- Aid?

13. Surface Area More exposed surface area means the water molecules can interact more with the solute. A powder dissolves better than a solid.

14. Solubility of Solids Increases with temp

15. Solubility of Gases - Decreases with temp inc.

16. Concentration A quantitative measure of the amount of solute dissolved in a given quantity of solution Can use a few terms to define concentration. Solubility Molarity Molality

17. Solubility The amount of a substance that dissolves in a given quantity of a solvent at a given temperature. (g solute/100g solvent) Saturated vs. Unsaturated vs. Supersaturated

18. Solubility Saturated Solvent can’t dissolve more solute Unsaturated Solvent can dissolve more solute Supersaturated Solvent dissolved more than expected

20. Molarity M = moles solute/liters of solution Uses the symbol M Example: 2 moles of glucose are dissolved to make 2 liters of solution. M = 2 moles/2L = 1M

21. Sample Problem #1 A saline solution contains 0.90g of NaCl in exactly 100mL of solution. What is the molarity of the solution?

22. Sample Problem #2 How many moles are present in 250mL of 0.24M calcium chloride?

23. Molality Another term used to describe the concentration. Uses the symbol m. m = moles solute/kg solvent

24. Sample Problem #1 What is the resulting molality when 30g of NaCl is dissolved in 100g of water?

25. Dilutions Use a concentrated stock solution to make a more dilute solution. M 1 V 1 = M 2 V 2

26. Dilutions We want 4.5L of a 1M HCl solution. How much 12M HCl should we use?

27. Colligative Properties Depend on the number of particles dissolved in a given mass of solvent. Freezing point depression Boiling point elevation Vapor pressure lowering

28. Colligative Properties Freezing point depression and boiling point elevation both use a similar equation. ∆T = K(m)(i) K is a constant based on the solvent m is the molality of the solution i is the van’t Hoff factor, the number of particles a solute creates in solution

29. Van’t Hoff Factor NaCl Al 2 (CO 3 ) 3 MgCl 2

30. Colligative Properties Which of the following will lower the freezing point of water the most? a. NaCl b. Glucose c. CaI 2 d. Al 2 (CO 3 ) 3

31. Types of Reactions  Precipitation reactions When aqueous solutions of ionic compounds are poured together a solid forms. A solid that forms from mixed solutions is a precipitate If you’re not a part of the solution, you’re part of the precipitate

32. Precipitation reactions NaOH(aq) + FeCl 3 (aq)   NaCl(aq) + Fe(OH) 3 (s) is really Na + (aq)+OH - (aq) + Fe +3 (aq) + Cl - (aq)   Na + (aq) + Cl - (aq) +Fe(OH) 3 (s) So all that really happens is OH - (aq) + Fe +3 (aq)  Fe(OH) 3 (s) Double replacement reaction

33. Precipitations Reactions Only happen if one of the products is insoluble Otherwise all the ions stay in solution- nothing has happened. Need to memorize the rules for solubility

34. Solubility Rules  All nitrates are soluble  Alkali metals ions and NH 4 + ions are soluble  Halides are soluble except Ag +, Pb +2, and Hg 2 +2  Most sulfates are soluble, except Pb +2, Ba +2, Hg +2,and Ca +2

35. Solubility Rules  Most hydroxides are insoluble except NaOH and KOH  Sulfides, carbonates, chromates, and phosphates are insoluble  Lower number rules supersede so Na 2 S is soluble

36. Precipitation reaction We can predict the products Can only be certain by experimenting AgNO 3 (aq) + KCl(aq)  Zn(NO 3 ) 2 (aq) + BaCr 2 O 7 (aq) 

37. Three Types of Equations Molecular Equation- written as whole formulas, not the ions. K 2 CrO 4 (aq) + Ba(NO 3 ) 2 (aq)  Complete Ionic equation show dissolved electrolytes as the ions. 2K + + CrO 4 -2 + Ba +2 + 2 NO 3 -  BaCrO 4 (s) + 2K + + 2 NO 3 - Spectator ions are those that don’t react.

38. Three Type of Equations Net Ionic equations show only those ions that react, not the spectator ions Ba +2 + CrO 4 -2  BaCrO 4 (s) Write the three types of equations for the reactions when these solutions are mixed. iron (III) sulfate and potassium sulfide