1. Solution Homogeneous mixture in a single phase.

2. Classification of Matter Matter Pure Substances ElementsCompounds Mixtures Homogeneous Mixtures Heterogeneous Mixtures Also called solutions Also called suspensions

3. Solvent Substance you have the most of. Substance that retains its phase. Dispersing medium.

4. Solute Substance you have the least of. Substance that dissolves.

5. Dilute Solution Relatively small amount of solute.

6. Concentrated Solution Relatively large amount of solute.

7. Aqueous Solution Water is the solvent.

8. NaCl(aq) Solution of NaCl dissolved in water!

9. Soluble Capable of being dissolved.

10. Solubility Maximum amount of 1 substance that will dissolve in a given amount of another substance.

11. Factors that influence solubility Temperature for all systems & pressure for systems involving gases.

12. Factors that influence the rate of dissolving. TemperatureTemperature Stirring or agitationStirring or agitation Surface Area of soluteSurface Area of solute Amount of solute already presentAmount of solute already present

13. What kind of change is dissolving? Physical change

14. How does a chemist represent the dissolving of a covalently bonded substance? C 6 H 12 O 6 (s)  C 6 H 12 O 6 (aq)

15. How does a chemist represent the dissolving of an ionically bonded substance? NaCl(s)  Na +1 (aq) + Cl -1 (aq)

16. What happens to an ionic substance when it dissolves? It separates into ions.

17. What is the equation for the dissolving of CaBr 2 ? CaBr 2 (s)  Ca +2 (aq) + 2Br -1 (aq)

18. Molecule-Ion Attraction Interaction between water molecules and ions in solution.

19. What kind of molecule is H 2 O? Polar – the O end is a bit negative & the H’s are a bit positive.

20. Is the orange ion positive or negative & how do you know?

22. Units of Solubility? Grams of solute per 100 grams of solvent

23. What happens to the solubility of all gases as the temperature increases? The solubility of gases  as the temperature 

24. What happens to the solubility of most solids as the temperature increases? The solubility of most solids  as the temperature .

25. What do you need to conduct electricity? Mobile charged particles!

26. Electrolyte? A substance that dissolves in water to produce a solution that conducts an electric current!

27. Nonelectrolyte? A substance that dissolves in water to produce a solution that does not conduct an electric current!

28. Saturated solution Contains the maximum amount of dissolved solute at that temperature.

29. Supersaturated solution Contains more than the maximum amount of dissolved solute at that temperature.

30. Unsaturated solution Contains less than the maximum amount of dissolved solute at that temperature.

31. Precipitation The opposite of dissolving. A solid comes out of solution.

32. Dynamic Equilibrium Term used to describe a saturated solution. Precipitation & dissolving are ocurring at the same rates. No net change.

33. How do you test a solution for saturation? Throw a crystal of the solute into the solution & observe what happens.

34. What are 3 possible outcomes of the saturation test? Crystal dissolves – Soln was unsaturated.Crystal dissolves – Soln was unsaturated. No change in crystal – Soln was saturated.No change in crystal – Soln was saturated. Crystal gets larger – Soln was supersaturated.Crystal gets larger – Soln was supersaturated.

35. What are 3 regions of a solubility curve? On the trace – saturated solution.On the trace – saturated solution. Above the trace – supersaturated solution.Above the trace – supersaturated solution. Below the trace – unsaturated solution.Below the trace – unsaturated solution.

36. Concentration A number that describes how much solute compared to how much solution or how much solvent.A number that describes how much solute compared to how much solution or how much solvent.

37. Percent Part over Whole X 100%

38. Molarity (M) Molarity = # moles solute Liters of solution

39. No. of Particles No. of Moles No. of Grams No. of Liters X formula mass X 22.4 L/mol X 6.02 X 10 23  by F.M.  by 6.02 X 10 23  by 22.4 Mole Map

40. Parts per Million (PPM) PPM = grams solute X 1000000 grams solution

41. How much KCl will dissolve in 300 grams of water at 50  C? Use the graph to set up a proportion. Problem: 42 g KCl = X g KCl 100 g H 2 O300 g H 2 O X = 126 g KCl

42. 88 g KNO 3 in 100 g H2O at 50  C. 20 g KNO 3 in 100 g H 2 O at 10  C. 88 g – 20 g = 68 g KNO 3 precipitates. 50  to 10  - How much KNO 3 precipitates?

43. source Measuring Heat of Solution If Temperature of H 2 O , dissolving was exothermic. If temperature of H 2 O , dissolving was endothermic.

44. Factors affecting Solubility TemperatureTemperature PressurePressure Nature of the Solvent & SoluteNature of the Solvent & Solute

45. Nature? IonicIonic Polar CovalentPolar Covalent Nonpolar CovalentNonpolar Covalent “Like Dissolves Like”

46. Nature of Solute Nonpolar Solvent Polar Solvent NonpolarSolubleInsoluble PolarInsolubleSoluble IonicInsolubleSoluble

47. Molarity C X Volume C = Molarity D X Volume D M C X V C = M D X V D Dilution Problems

48. Colligative Property Depends on the concentration of the solute, NOT on the nature of the solute. Does not matter if ions or neutral particles are in solution.

49. What are 2 colligative properties? Freezing Point Depression & Boiling Point Elevation

50. Which solution has the most dissolved particles? a)1 mole of C 6 H 12 O 6 in 1 Liter H 2 O b)1 mole of NaCl in 1 Liter H 2 O c)1 mole of CaBr 2 in 1 Liter H 2 O