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Molarity, Dilution, and pH. Main Idea: Solution concentrations are measured in molarity . Dilution is a useful technique for creating a new solution from a stock solution. pH is a measure of the concentration of hydronium ions in a solution. Properties of Aqueous Solutions.

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Molarity, Dilution, and pH


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    1. Molarity, Dilution, and pH Main Idea: Solution concentrations are measured in molarity. Dilution is a useful technique for creating a new solution from a stock solution. pH is a measure of the concentration of hydronium ions in a solution.

    2. Properties of Aqueous Solutions • Solution-a homogeneous mixture of two or more substances. • Solute-a substance in a solution that is present in the smallest amount. • Solvent-a substance in a solution that is present in the largest amount. • In an aqueous solution, the solute is a liquid or solid and the solvent is always water.

    3. Molarity Review • One of the most common units of solution concentration is molarity. • Molarity (M) is the number of moles of solute per liter of solution. • Molarity is also known as molar concentration, and the unit M is read as “molar.” • A liter of solution containing 1 mol of solute is a 1M solution, which is read as a “one-molar” solution. • A liter of solution containing 0.1 mol of solute is a 0.1 M solution.

    4. Molarity Equation • To calculate a solution’s molarity, you must know the volume of the solution in liters and the amount of dissolved solute in moles. • Molarity (M) = moles of solute liters of solution

    5. Molarity Example A 100.5-mL intravenous (IV) solution contains 5.10 g of glucose (C6H12O6). What is the molarity of the solution? The molar mass of glucose is 180.16 g/mol. SOLUTION: • Calculate the number of moles of C6H12O6 by dividing mass over molar mass = 0.0283 mol C6H12O6 • Convert the volume of H2O to liters by dividing volume by 1000 = 0.1005 L • Solve for molarity by dividing moles by liters = 0.282 M

    6. Preparing Molar Solutions • Now that you know how to calculate the molarity of a solution, how would you prepare one in the laboratory? • STEP 1: Calculate the mass of the solute needed using the molarity definition and accounting for the desired concentration and volume. • STEP 2: The mass of the solute is measured on a balance. • STEP 3: The solute is placed in a volumetric flask of the correct volume. • STEP 4: Distilled water is added to the flask to bring the solution level up to the calibration mark.

    7. http://www.ltcconline.net/stevenson/2008CHM101Fall/CHM101LectureNotes20081022.htmhttp://www.ltcconline.net/stevenson/2008CHM101Fall/CHM101LectureNotes20081022.htm

    8. Properties of Aqueous Solutions • All solutes that dissolve in water fit into one of two categories: electrolyte or non-electrolyte. • Electrolyte-a substance that when dissolved in water conducts electricity • Non-electrolyte-a substance that when dissolved in water does not conduct electricity. • To have an electrolyte, ions must be present in water.

    9. Electrolytic Properties of Aqueous Solutions • NaCl in water. • What happens? • NaCl(s)→ Na+(aq) + Cl–(aq) • Completely dissociates

    10. Strong vs. Weak Electrolytes • How do you know when an electrolyte is strong or weak? • Take a look at how HCl dissociates in water. • HCl(s)→ H+(aq) + Cl–(aq)

    11. Electrolytic Properties of Aqueous Solutions

    12. Electrolytic Properties of Aqueous Solutions

    13. Hydrated Ions

    14. Electrolytic Properties of Aqueous Solutions • What about weak electrolytes? • What makes them weak? • Ionization of acetic acid • CH3COOH(aq)↔ CH3COO–(aq) + H+(aq)

    15. Electrolytic Solutions

    16. Precipitation Reactions • Precipitation Reaction-a reaction that results in the formation of an insoluble product. • These reactions usually involve ionic compounds. • Formation of PbI2: • Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)

    17. Preciptate

    18. Precipitate

    19. Precipitation Reactions • How do you know whether or not a precipitate will form when a compound is added to a solution? • By knowing the solubility of the solute! • Solubility-The maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature. • Three levels of solubility: Soluble, slightly soluble or insoluble.

    20. Precipitation Reactions

    21. Determining Solubility • Determine the solubility for the following: (1) Ag2SO4 (2) CaCO3 (3) Na3PO4

    22. Diluting Molar Solutions • In the laboratory, you might use concentrated solutions of standard molarities, called stock solutions. • For example, concentrated hydrochloric acid (HCl) is 12 M. • You can prepare a less-concentrated solution by diluting the stock solution with additional solvent. • Dilution is used when a specific concentration is needed and the starting material is already in the form of a solution (i.e., acids).

    23. Dilution of Solutions • When you want to dilute a solution, what happens to the number of moles present in the solution? • Do they increase? • Decrease? • Stay the same?

    24. Dilution of Solutions

    25. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? Add water to the 3.0 M solution to lower its concentration to 0.50 M Dilute the solution!

    26. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? But how much water do we add?

    27. moles of NaOH in ORIGINAL solution = moles of NaOH in FINAL solution PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? How much water is added? The important point is that --->

    28. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? Amount of NaOH in original solution = M • V = (3.0 mol/L)(0.050 L) = 0.15 mol NaOH Amount of NaOH in final solution must also = 0.15 mol NaOH Volume of final solution = (0.15 mol NaOH) / (0.50 M) = 0.30 L or 300 mL

    29. PROBLEM: You have 50.0 mL of 3.0 M NaOH and you want 0.50 M NaOH. What do you do? Conclusion: add 250 mL of waterto 50.0 mL of 3.0 M NaOH to make 300 mL of 0.50 M NaOH.

    30. Preparing Solutions by Dilution A shortcut M1 • V1 = M2 • V2 Where M represents molarity and V represents volume. The 1s are for the stock solution and the 2s are for the solution you are trying to create.

    31. How do you get from this…

    32. …to this?

    33. Add an ionic compound!

    34. Colligative Properties • Properties that depend only on the number of solute particles and not on their identity.

    35. Some Colligative Properties are: • Vapor pressure lowering • Boiling point elevation • Freezing Point depression

    36. Vapor Pressure

    37. Vapor Pressure Lowering • The particles of solute are surrounded by and attracted to particles of solvent. • Now the solvent particles have less kinetic energy and tend less to escape into the space above the liquid. • So the vapor pressure is less.

    38. Ionic vs Molecular Solutes • Ionic solutes produce two or more ion particles in solution. • They affect the colligative properties proportionately more than molecular solutes (that do not ionize). • The effect is proportional to the number of particles of the solute in the solution.

    39. How many particles do each of the following give upon solvation? • NaCl • CaCl2 • Glucose

    40. Freezing Point Depression

    41. Example • Salt is added to melt ice by reducing the freezing point of water.

    42. Boiling Point Elevation

    43. Example • Addition of ethylene glycol C2H6O2 (antifreeze) to car radiators.

    44. Freezing Point Depression and Boiling Point Elevation Boiling Point Elevation • ∆Tb =mkb (for water kb=0.51 oC/m) • Freezing Point Depression • ∆Tf=mkf (for water kf=1.86 oC/m) • Note: m is the molality of the particles, so if the solute is ionic, multiply by the #of particles it dissociates to.

    45. Which is more effective for lowering the freezing point of water? • NaCl or CaCl2

    46. Example 1: • Find the new freezing point of 3m NaCl in water.

    47. Example 2: • Find the new boiling point of 3m NaCl in water.

    48. ThepH scaleis a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion.Under 7 = acid 7 = neutralOver 7 = base

    49. Acid-Base Reactions • Acids-generally have a sour taste, change litmus from blue to red, can react with certain metals to produce gas, conduct electricity. • Bases-generally have a bitter taste, change litmus from red to blue, feel slippery, conduct electricity. • BrØnstead Acid-proton donor • BrØnstead Base-proton acceptor

    50. Acid-Base Reactions • Acid or Base? • HCl(aq) + H2O(l)→ H3O+(aq) + Cl–(aq) • NH3(aq) + H2O(l)→ NH4+(aq) + OH–(aq)