1. Unit 7: Solutions, Acids, & Bases

2. I. Definitions and Types of Solutions A. What exactly IS a Solution?

3. B. Types: [Vocabulary terms…] 1] Unsaturated 2] Saturated 3] Supersaturated 4] Suspensions 5] Colloids 6] Alloys

4. Question… HOW do we create solutions?

5. C. Forming Solutions Polar solvents mix best with __________solutes! How do IMF’s relate to this concept? Nonpolar solvents mix best with ________________ solutes! Examples…

6. Like Dissolves Like… See Table 7.1 on pg. 119 of Review book! Solute Nonpolar Solvent Polar Solvent Nonpolar Polar Ionic

7. II. Collision Theory Molecules need the proper orientation and proper amount of kinetic energy to create a solution Remember the formation of Salt Water demonstrations? How did the orientation of the water impact the solvation of the Na+ and Cl- ions?

8. Combining the Solute and Solvent Two major criteria need to be met: 1: “Like Dissolves Like” = need compatible bond types and IMF’s 2: Effective Collisions = For a substance to dissolve into a solvent, there must be an “effective collision” between solute and solvent Need sufficient KE and proper orientation!

9. Factors Affecting the Formation of a Solution Temperature 1] Temperature Increase in Temp = Increase in KE, & increase in # of total collisions Polarity 3] Polarity Similar bond types and IMF’s form solutions faster Surface Area 2] Surface Area Increase in S.A. means more particles contact the solution, & increases the total # of collisions Pressure (gases only) 4] Pressure (gases only) Increases in Pressure cause decreases in volume, and make gases more soluble

10. Other Solutions Terms… Precipitate out = this refers to the solid precipitate that ‘falls out’ of a solution when an irreversible double replacement reaction occurs; Solid may be collected! Miscible = substances that completely and evenly mix together at any concentrations Immiscible = do NOT mix to form a solution

11. Forming Solutions Temperature Changes: Increases = most solids have higher solubilities EXCEPTION = gases! Lower solubility at higher temperatures Pressure Changes: Increases = most gases have higher solubilities Liquids and solids = Exceptions! Little to no solubility differences with changes in pressure

12. III. Collision Theory Collision theory = scientific description of how solutions are formed May have several factors that help collisions occur faster Need _____________ collisions with particles that have sufficient __________ for a solution to occur! Do you have any ideas on how to speed up this process…?

13. II. Solubility Curves and Guidelines A. Table G: Solubility Guidelines for Aqueous Solutions ABOVE Supersaturated = ABOVE ON Saturated = ON the line BELOW Unsaturated = BELOW

14. B. Table F: Solubility Guidelines for Aqueous Solutions These charts tell if an anion or cation is generally soluble in water Exceptions occur for half of those listed General rules of solubility here; not a total and complete list!

15. Determine solubility…

16. C. Net Ionic Equations http://www.youtube.com/watch?v=RjBjwQF276A Net ionic equations are generated when ions react within a solution to form a precipitate [solid] The ions involved in creating the solid will be eliminated from the solution, eventually The ions that are eliminated produce the “net” reaction, or equation! [example on board] http://www.mhhe.com/physsci/chemistry/animations/chang_7e_es p/crm3s2_3.swf

17. IV. Measuring Concentrations A. Percent Mass (w/w) B. Percent Volume (v/v) C. Parts per Million (ppm) D. Parts per Billion (ppb)

18. Let’s practice… What is the concentration of a solution, in parts per million, if 0.02 gram of NaCl is dissolved in 1000. grams of solution?

19. Regents Practice Problems An aqueous solution has 0.0070 gram of oxygen dissolved in 1000. grams of water. ------------------------------------------------------------ Calculate the dissolved oxygen concentration of this solution in parts per million. [Your response must include both a correct numerical setup and the calculated result.]

20. D. Molarity moles solute Molarity (M) = Liter solution Solute MUST be in units of ‘moles’ Solvent MUST be in units of ‘Liters’ May use Molarity to find grams, moles, or molar mass

21. Molarity Examples 1] Calculate the molarity of a solution containing 35.0g NaCl dissolved into 500mL of water. 2] How many grams of NaCl are needed to make 650mL of a 0.010M solution?

22. E. Dilutions Dilutions use Molarity and Volume to distribute the moles of a substance M c V c = M d V d (Molarity) x (volume) in Liters = moles! Extremely useful equation!!!

23. Dilution Examples 1] Calculate the new molarity of 1500.mL of solution made from 25.mL of 12M. 2] How much 18M acid is needed to make 2500mL of 0.10M solution?

24. V. Colligative Properties Colligative properties… properties of solutions that depend on the number of molecules in a given volume of solvent and not on the properties/identity (e.g. size or mass) of the molecules

25. Colligative properties These include: lowering of vapor pressure elevation of boiling point depression of freezing point They are based on the number of particles or electrolytes produced in a solution by a solute

26. A. Electrolytes vs. Particles Electrolytes are…. They conduct electricity! Number of electrolytes formed per molecule depends on the # of ions it breaks into! Ex.] Particles are… They DO NOT conduct electricity! Produce only ONE particle per molecule dissolved! Ex.]

27. B. Factors Affected Freezing Point Depression 1] Freezing Point Depression Adding particles/electrolytes causes the freezing point of the solution to be less than that of the pure substance Magnitude of change depends on # of particles/ions in solution Boiling Point Elevation 2] Boiling Point Elevation Adding particles/electrolytes will cause the boiling point to increase for the solution Can calculate the new boiling point using an equation

28. Web demos http://group.chem.iastate.edu/Greenbowe/sections/proje ctfolder/flashfiles/propOfSoln/colligative.html

29. 2. Vaporization and Boiling Vapor pressure = the pressure exerted by a layer of the gas phase on the surface of a liquid or solid High Vapor pressure = weaker IMF’s, weaker bonds, changes to gas phase easily Low vapor pressure = STRONGER IMF’s, stronger bonds, and prefers to remain a liquid phase

30. Table H Vapor Pressure of Four Liquids

31. 3. Boiling Point vs. Normal Boiling Point Boiling Point = temperature at which vapor pressure of the liquid equals the atmospheric pressure of the surroundings NORMAL boiling point = temperature when the vapor pressure equals standard pressure [1atm]

32. What changes the boiling point? Changes in atmospheric pressure cause changes in the boiling point Low pressure at higher altitudes cause the boiling point to be lower than normal [pressure is lower than normal…] High pressure occurring below sea level cause boiling points to increase […?]

33. Vapor Pressure… REVIEW… INDIRECTLY Vapor pressure changes INDIRECTLY with variations is atmospheric pressure Explain…… Vapor pressure is a function of IMF’s within the liquid phase Stronger IMF’s = lower vapor pressure Weaker IMF’s = higher vapor pressure

34. VI. Acids and Bases Acids = substances that react with a base; often has a low pH May be strong or weak Actual definition depends upon the type of acid/base Bases = substances reacting with an acid; often has a high pH May be strong or weak 3 major types of acids and bases here

35. Properties of Acids and Bases Acids: Taste sour Conduct current in solution React with bases to form water and salt React with some metals to make H2(g) Low pH Bases: Taste bitter Slippery/soapy feeling Conduct current in solution React with acids to form salt water Have low pOH or high pH

36. 1. Arrhenius Acid/Base Acid = substances that release hydrogen ions in aqueous solutions proton donors Ex.] HCl  H + + Cl - Base = substances that release hydroxide ions in aqueous solutions Hydroxide donors Ex.] NaOH  Na + = OH -

37. 2. Bronsted-Lowry Acid/Base Acid = any substance that can donate hydrogen ions in solution Proton donors Ex.] HCl  H + + Cl - Base = any substance that can accept hydrogen ions in solution Proton acceptors Ex.] NH3 + H +  NH4 +

38. Lewis Acids and Bases [non-Regents] Acid = any species that can accept an electron pair from another species in solution Ex.] H + + NH3  Base = any species that can donate an electron pair to another species in solution NH4 +

39. Weak vs Strong acids and bases STRONG acids/bases will completely dissociate [ionize] in solution Ex.] HCl, HNO 3, H 2 SO 4, NaOH, KOH, etc. WEAK acids/bases produce very few ions per molecule in solution Ex.] vinegar, H 3 PO 4, formic acid, citric acid, etc.

40. B. Conjugates Acid/Base Conjugates = species that are formed in solution as a result of the dissociation of an acid or base Ex.] HCl + NaOH  Na + + Cl - + HOH Cl - is the conjugate base Na + is the conjugate acid

41. Acid-Base Reactions Acid + Base = Salt + water Acids and bases neutralize each other H+ + OH-  H2O

42. C. pH and pOH pH = the scale, 0-14, that defines the acidity or basicity of a solution Based on the concentration of hydrogen ions 0-6 = acid 7= neutral 8-14 = base

43. pOH pOH = the measure of the concentration of OH- ions in solution Opposite of pH Scale 0-14 0-6 = strong base 7= neutral 8-14 = acidic

44. Relating pH and pOH pH + pOH = 14 always! pH = -log[H + ] pOH = -log[OH - ]

45. D. pH and pOH Calculations Ex. 1] What is the pH of a solution containing 9.15 x 10 -6 M H + ? Ex. 2] What is the pH of a solution containing [OH - ] = 8.11 x 10 -5 M in 350.mL of solution?

46. E. Indicators Indicators = compounds that change color when the pH changes Color changes indicate the pH of the solution! Use several indicators to pinpoint the final pH!

47. Indicators: Table M Find the pH using the indicator and its color in the solution… Tutorial on indicators: http://www.kentchemis try.com/links/Acids Bases/Indicators.ht m

48. Expanded Indicators Chart

49. F. Titrations Titration = method used to determine the pH of an unknown solution Process to find the concentration of an unknown acid/base by neutralizing it with a base/acid of known concentration An indicator signals the equivalence point and tells that the neutralization is complete

50. Titrations, [continued] Uses burets, a standard solution of known pH, an indicator, and a fixed volume of a solution with an unknown pH Titration Formula: M a V a = M b V b