1Chapter 15 Acid-Base Titrations & pH 15.1 Aqueous Solutions & The Concept of pH
2Self-Ionization of Water Autoprotolysis:H2O (l) + H2O (l) → H3O+ (aq) + OH- (aq)Molarity at 25°C1.0 x 10-7 moles H30+ per liter of solution1.0 x 10-7 moles OH- per liter of solution
3Ionization Constant for Water (KW) KW = [H3O+][OH-] = (1.0 x 10-7M)(1.0 x 10-7M) = 1.0 x 10-14M2KW is a constant at ordinary ranges of room temperaturesNeutral: [H3O+] = [OH-]Acidic: [H3O+] > [OH-]Basic: [H3O+] < [OH-]
5Calculating [H3O+] and [OH-] Assume that strong acids and bases are completely ionized in solution:1.0 M H2SO4 = 2.0 M H3O+1.0 M Ba(OH)2 = 2.0 M OH-
6pH Calculations and the Strength of Acids and Bases Weak acids and weak bases cannot be assumed to be 100% ionized[H30+] and [OH-] cannot be determined from acid and base concentrations, and must be determined experimentally
9The pH ScaleDue to the variations in soln’s, there are many possible concentrations of hydronium & hydroxide ions for solutions.Usually this spans M to 1 MIn order to compare substances, the pH scale was developed.Ex:6 M sol’n of HCl has a H3O+ molarity of 6 M6 M sol’n of HC2H3O2 has a H3O+ molarity of 0.01M
10pHpH is a scale in which the concentration of hydronium ions in solution is expressed as a number ranging from 0 to 14.Instead of referring to a scale of 1 to 10-14, the pH scale is much easier to use.pH is the negative of the exponent of the hydronium concentration.
11Calculating pH & pOHpHThe negative of the common logarithm of the hydronium ion concentrationpH = - log [H3O+]pOHThe negative of the common logarithm of the hydroxide ion concentrationpOH = - log [OH-]pH + pOH = 14.0
12pH ExampleA solution with a hydronium concentration of M has a pH of 11.What would be the pH of a solution with a hydronium concentration of 10-6 M?pH = 6
16Interpreting the pH Scale As pH decreases below 7, hydronium ion concentrations increase & hydroxide ion concentrations decrease.** pH values differ in factors of 10.Ex:An acidic sol’n w/ a pH of 3 has 10 times the hydronium concentration as a sol’n w/ a pH of 4.
17Interpreting the pH Scale As pH increases above 7, hydroxide ion concentrations increase & hydronium ion concentrations decrease.Ex: A basic sol’n w/ a pH of 9 has 10 times the hydroxide concentration as a sol’n w/ a pH of 8.A neutral sol’n has equal concentrations of hydronium and hydroxide ions.
19Significant Figures & pH Significant digits when calculations involve logarithms are dependant only on the number of digits to the right of the decimal.Example:[H3O+] = 2.5 x 10-3pH = 2.60This concentration has 2 significant digitsSo the pH will have 2 digits to the right of the decimal pointpH = ?????
20Practice1. Determine the [H3O+] & [OH-] in a 0.01 M solution of HClO4.2. An aqueous solution of Ba(OH)2 has a [H3O+] of 1 x M. What is the [OH-]? What is the molarity of the solution?3. Determine the pH of a 1 x 10-4 M solution of HBr.4. Determine the pH of a 5 x 10-4 M solution of Ca(OH)2.5. What is the pH of a solution whose [H3O+] = 6.2 x 10-9 M?6. Determine the pH of a M solution of NaOH.
21More PracticeWhat are the [H3O+] & [OH-] of a solution if its pH = 9.0?The pH of a solution if What is the concentration of hydroxide ions in the solution? If the solution is Sr(OH)2 (aq), what is its molarity?The pH of a hydrochloric acid solution for cleaning tile is What is the [H3O+] in the solution?A shampoo has a pH of 8.7. What are [H3O+] & [OH-] in the shampoo?
22Practice Answers! [H3O+] = 1 x 10-2 M, [OH-] = 1 x 10-12 M [OH-] = 1 x 10-3 M, [Ba(OH)2] = 5 x 10-4 MpH = 4.0pH = 11.0pH = 8.21pH = 10.87[H3O+] = 1 x 10-9 M, [OH-] = 1 x 10-5 M[OH-] = 1 x 10-4 M, [Sr(OH)2] = 5 x 10-5 M[H3O+] = 0.35 M[H3O+] = 2 x 10-9 M, [OH-] = 5 x 10-6 M
24Indicators and pH Meters Acid-Base IndicatorsCompounds whose colors are sensitive to pHTransition IntervalpH range over which an indicator color change occursIndicators are useful when they change color in a pH range which includes the endpoint of the reaction
26Using Indicators to Measure pH A pH meter is the most accurateway to measure pH.Measures voltage differencebetween two electrodesIt will determine the exact pH of a sol’n.There are also colored dyes that will change in a predictable way according to a standard chart. These are called indicators.
28Acid-Base Titration Titration Controlled addition of the measured amount of a solution of a known concentration required to react completely with a measured amount of sol’n of unknown concentrationEquivalence PointThe point at which the solutions used in a titration are present in chemically equivalent amountsTitration Curves: End pointThe point in a titration at which the rxn is just completed
30Molarity and Titration Standard SolutionA solution that contains the precisely known concentration of a solute, used in titration to find the concentration of the solution of unknown concentrationPrimary StandardA highly purified solid compound used to check the concentration of the known solution in a titration
31Calculations with Molar Titrations Start with the balanced equation for the neutralization reaction and determine the chemically equivalent amounts of the acid and baseDetermine the moles of acid (or base) from the known solution used during the titrationDetermine the moles of solute of the unknown solution used during the titrationDetermine the molarity of the unknown solution
33More Practice!How many moles of HCl are in mL of a M solution?How many moles of NaOH would neutralize 20.0 mL of a 13.9 M solution of H2SO4?How many milliliters of a 2.76 M KOH solution contain mol of KOH?
34More Practice!4. A mL sample of a solution of RbOH is neutralized by mL of a M solution of HBr. What is the molarity of RbOH? 5. If mL of a solution of Ba(OH)2 requires mL of a M solution of HNO3 for complete titration, what is the molarity of the Ba(OH)2 solution? 6. You have a vinegar solution believed to be 0.83 M. You are going to titrate mL of it with a NaOH solution known to be M. At what volume of added NaOH would you expect to see an endpoint?
35Answers! 2.14 x 10-2 mol HCl 5.56 x 10-2 mol NaOH 29.9 mL M RbOHM Ba(OH)232 mL NaOH
36BuffersBuffers have many important biological functions. They keep a solution at a constant pH, when manageable amounts of acid of base are added.Ex: Your blood is a buffer! Its pH is very slightly basic at 7.4. Even though you may eat many different types of foods or medicines, your blood pH stays relatively stable, varying only about 0.1. That means your blood controls its own pH!
37BuffersBuffers contain ions or molecules that react with hydronium or hydroxide if they are added to the solution. That means, even if you add an acid or a base, your pH will stay the same.To make a buffer, you combine a weak acid or a weak base with its corresponding salt.
38Buffers Example: Ammonia is combined with its salt, NH4Cl, in sol’n: If acid is added to this solution, ammonia reacts with the H+ :NH3 (aq) + H+ (aq) → NH4+ (aq)If a base is added to this solution, the NH4+ from the dissolved salt will react with the OH- :NH4+ (aq) + OH- (aq) → NH3 (aq) + H2O (l)
39H2CO3 (aq) + OH- → HCO3- (aq) + H2O (l) BuffersBlood’s pH is regulated by many systems, but dissolved CO2 is a very important method. Carbonic acid, H2CO3, and the hydrogen carbonate ion, HCO3-, are both dissolved in your blood.CO2 (g) + H2O (l) → H2CO3 (aq)If you add OH- :H2CO3 (aq) + OH- → HCO3- (aq) + H2O (l)If you add H+ :HCO3- (aq) + H+ → H2CO3 (aq)
40BuffersYour lungs control the amount of carbon dioxide in your body. If your body takes in too much carbon dioxide, your blood may become too acidic so you may yawn to lower the concentration of carbonic acid by expelling CO2.
41BuffersIf you hyperventilate, too much CO2 is expelled, which causes the concentration of carbonic acid to become too low, and your blood may become too basic. Breathing into a paper bag will increase the concentration of CO2 in your lungs and restore the proper pH.