1. Unit 13: Solutions

2.  Solution - homogeneous mixture Solvent – substance that dissolves the solute Solute - substance being dissolved

3. Solute - KMnO 4 KMnO 4 Solution Solvent - H 2 O

4. Solvation  Solvation – the process of dissolving solute particles are separated and pulled into solution solute particles are surrounded by solvent particles

5. Solvation Strong Electrolyte Non- Electrolyte solute exists as ions only - + salt - + sugar solute exists as molecules only - + acetic acid Weak Electrolyte solute exists as ions and molecules View animation online.animation

6. Solvation  Dissociation  separation of an ionic solid into aqueous ions NaCl(s)  Na + (aq) + Cl – (aq)

7. Solvation  Molecular Solvation  molecules stay intact C 6 H 12 O 6 (s)  C 6 H 12 O 6 (aq)

8. Solvation NONPOLAR POLAR “Like Dissolves Like”

9. Solubility  Unsaturated  When the solvent holds less solute than it normally can at a given temperature.

10. Solubility  Saturated  When the solvent holds as much of a solute as it normally can at a given temperature.

11. Solubility  Supersaturated  When the solvent holds more dissolved solute than it normally can at that temperature.  Formed when a saturated solution is formed at high temperatures and then cooled slowly.

12. Solubility SATURATED SOLUTION no more solute dissolves UNSATURATED SOLUTION more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form concentration

13. Solubility Curves  Solubility Curves show:  maximum grams of solute that will dissolve in 100 g of solvent at a given temperature  varies with temperature  based on a saturated solution

14. Solubility  Solubility Curve  shows how temperature affects the solubility of a substance.

15. Solid Solubility  Solids are more soluble at...  high temperatures

16. Gas Solubility  Gases are more soluble at...  low temperatures &  high pressures (Henry’s Law).  EX: nitrogen narcosis, the “bends,” soda

17. Solubility Curve

18.  Things to remember  The graph is set in 100 mL or 100 grams of water.  The line represents the saturation point at that temperature.  Anything above the line (at that temp) is supersaturated  Anything below the line (at that temp) is unsaturated.

19. Example What amount of NaCl would make a saturated solution in 100 mL of water at 80 0 C? Just look it up on the graph. Answer : 40 grams GRAPH

20. BACK

22.  What if it is not 100 mL or 100 grams of water?  Set up a proportion: solid solid = liquidliquid liquidliquid From graph From problem Always 100 mL This will be x

23. Example  What amount of KNO 3 would make a saturated solution in 177mL of water at 50 0 C? 80 gx = 100mL 177mL = 141.6 g GRAPH

24.  What if there is no amount of water given?  Set up a proportion: solid solid = liquidliquid liquidliquid From graph From problem Always 100 mL This will be x

25. Example  What amount of water at 20 0 C would make a saturated solution with 63g of KNO 3 ? 30 g63 g = 100mL X 100mL X = 210 mL GRAPH

26. Precipitation Reactions: What are they?  Type of double replacement reaction.  Two solutions of ionic compounds are mixed  One of the products of the reaction is an insoluble salt called a precipitate (chunky milk)  Use solubility rules to determine the precipitate.

27. Solubility Rules Solubility (in water) CompoundsExceptions Soluble Alkali metal (Group 1) salts none Ammonium salts (NH 4 + ) none Acetates (C 2 H 3 O 2 - ) none Nitrates (NO 3 - ) none Chlorides (Cl - ), Bromides (Br - ), Iodides (I - ) Compounds of Ag, Hg, and Pb Sulfates (SO 4 2- ) Compounds of Sr, Ba, Hg, and Pb **************************************************************** Insoluble Carbonates (CO 3 2- ), Phosphates (PO 4 3- ), Sulfites (SO 3 2- ), Chromates (CrO 4 2- ) Compounds of Alkali metal (Group 1)and NH 4 + Compounds of Alkali metal (Group 1) and NH 4 + Sulfides (S 2- ), Hydroxides (OH - ) Compounds of Alkali metal (Group 1)and NH 4 +, Ca, Sr, Ba Compounds of Alkali metal (Group 1) and NH 4 +, Ca, Sr, Ba

28. Example using chart  Pb(OH) 2  NH 4 Cl  K 2 SO 4  Ba 3 (PO 4 ) 2 Insoluble Soluble Insoluble RULES

29. How to use it.  Write the products for the double replacement reactions and balance.  Then, using the solubility rules  identify the precipitate by a subscript (s) for solid (insoluble in water).  Identify the soluble compounds by marking them with a subscript (aq) for aqueous (soluble in water).

30. Example  Ag(NO 3 ) 2 + LiCl  Ag(NO 3 ) 2 + LiCl  Ag(NO 3 ) 2 + LiCl  Ag(NO 3 ) 2 + 2 LiCl  AgCl 2 + 2 LiNO 3 Ag(NO 3 ) 2 + 2 LiCl  AgCl 2 + 2 LiNO 3 AgCl 2 + LiNO 3 +2 +1 s aq RULES

31. Concentration  The amount of solute in a solution.  Describing Concentration  % by mass - medicated creams  % by volume - rubbing alcohol  Molarity and Molality - used by chemists

32. Concentration (parts per million and parts per billion) SAWS Water Quality Report - June 2000

33. B. Molarity solvent only 1000mL = 1 L

34. Molarity  Find the Molarity of a solution containing.5 mol of glucose in 2 L of water. =.25 M GIVEN: M = ? Mol =.5 mol V = 2 L WORK: M= mol solute/ liters solvent M=.5 mol 2 L

35. Dilution  Preparation of a desired solution by adding water to a concentrated solution.  Moles of solute remain the same.

36. Dilution  What volume of 15.8M HNO 3 is required to make 250 mL of a 6.0M solution? GIVEN: M 1 = 15.8M V 1 = ? M 2 = 6.0M V 2 = 250 mL WORK: M 1 V 1 = M 2 V 2 (15.8M) V 1 = (6.0M)(250mL) V 1 = 95 mL of 15.8M HNO 3