1. Chapter 15

2. Section 1

3.  The Ionization Constant of Water (K W )  Water self-ionizes to a small extent to form H 3 O + and OH -  The concentrations of these ions, represented by [H 3 O + ] and [OH - ], equal 1.0 x 10 -7 M.  K W = [H 3 O + ][OH - ] = 1.0 x 10 -14 M.

4.  Acids – INCREASE [H 3 O + ] and DECREASE [OH - ]  Bases – INCREASE [OH - ] and DECREASE [H 3 O + ] **During the ionization process here K W remains constant at 1.0 x 10 -14 M. THUS, knowing the concentration of either allows you to calculate the concentration of the OTHER

5.  Assume that strong acids and bases are completely ionized in solution: 1.0 M H 2 SO 4 = 2.0 M H 3 O + 1.0 M Ba(OH) 2 = 2.0 M OH -

7.  Determine the hydronium and hydroxide ion concentrations in a solution that is 3.0 x 10 -2 M NaOH.

8.  Determine the hydronium and hydroxide ion concentrations in a solution that is 1.0 x 10 -4 M Ca(OH) 2.

9.  pH  pH 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.

11. pH = -log[H 3 O + ] pOH = -log[OH - ] pH + pOH = 14 (-logK W )

12.  Neutral Solution – pH, pOH BOTH = 7  Acidic Solution – pH 7  Basic Solution – pH > 7pOH < 7  **pH values differ in factors of 10.**  ** pH values differ in factors of 10.**  An acidic solution w/ a pH of 3 has 10 times the hydronium concentration as a solution w/ a pH of

13.  The number of SFs in the [H 3 O + ] OR [OH - ] determines the number of SFs to the RIGHT of the decimal place in pH and pOH.  IF [H 3 O + ] = 1 x 10 -7 M, THEN pH = 7.o  IF [H 3 O + ] = 1.00 x 10 -7 M, THEN pH = 7.000

14.  Determine the pH of the following solutions:  1 x 10 -3 M HCl  1.0 x 10 -5 M HNO 3  1 x 10 -4 M NaOH  1.0 x 10 -2 M KOH

15.  Determine the pH of the following solutions:  [H 3 O + ] = 6.7 x 10 -4 M  [H 3 O + ] = 2.5 x 10 -2 M  2.5 x 10 -6 M HNO 3  2.0 x 10 -2 M Sr(OH) 2

16.  Determine [H 3 O + ] for the following solutions:  pH = 5.0  pH = 12.0  pH = 1.50  pH =3.67  [OH - ]?

17.  Unlike STRONG Acids and Bases, these DO NOT ionize completely.  Thus pH must be determined EXPERIMENTALLY.

18. 1. Determine the [H 3 O + ] & [OH - ] in a 0.01 M solution of HClO 4. 2. An aqueous solution of Ba(OH) 2 has a [H 3 O + ] of 1 x 10 - 11 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.

19. 1. What is the pH of a solution whose [H 3 O + ] = 6.2 x 10 -9 M? Determine the pH of a 0.00074 M solution of NaOH. hat are the [H 3 O + ] & [OH - ] of a solution if its pH = 9.0? The pH of a solution if 10.0. What is the concentration of hydroxide ions in the solution? If the solution is Sr(OH) 2 (aq), what is its molarity? 1. The pH of a hydrochloric acid solution for cleaning tile is 0.45. What is the [H 3 O + ] in the solution?

20. Section 2

21.  Compounds whose colors are sensitive to pH  Usually weak acids or bases in which the color of the ion (In - ) changes along with pH  The free anion is ONE color, the bonded anion is a DIFFERENT color

22.  Methyl Red (4.4-6.2)  Bromthymol Blue (6.2-7.6)  Methyl Orange (3.1-4.4)  Bromphenol Blue (3.0-4.6)  Phenolphthalein (8.0-10.0)  Phenol Red (6.4-8.0)

24.  Transition Interval – the pH range over which an indicator changes color.  pH Meter – determines the pH of a solution by measuring the voltage between two electrodes that are placed IN the solution.  pH Strips – contain the universal indicator a solution containing several different indicators

25.  The controlled addition and measurement of the amount of a solution of known concentration required to completely react with a measured amount of a solution of unknown concentration. **The point is to determine when the acid and the base are present in chemically equivalent amounts**

26.  Equivalence Point – the point at which the two solutions used in a titration are present in chemically equivalent amounts  End Point – the point in a titration at which the indicator changes color

28.  Strong Acid/Strong Base – E.P. at pH = 7  Strong Acid/Weak Base – E.P. at pH < 7  Weak Acid/Strong Base -- E.P. at pH > 7  Weak Acid/Weak Base -- E.P. at pH ??? **CHOOSE YOUR INDICATOR ACCORDINGLY**

29.  Standard Solution A solution that contains the precisely known concentration of a solute, used in titration to find the concentration of the solution of unknown concentration  Primary Standard A highly purified solid compound used to check the concentration of the known solution in a titration

30.  Set up the Titration apparatus.  Rinse each buret THREE times with acid/base to be used

31.  Fill the acids buret to a point ABOVE the calibration mark.

32.  Release enough of the acid to bring the volume down below the calibration mark.  RECORD this volume.

33.  Fill the flask with the APPROXIMATE amount of acid required.  RECORD the new buret volume.

34.  Add three drops of the appropriate indicator to the flask.

35.  Fill the OTHER buret with the KNOWN base to a point ABOVE the calibration point.

36.  Lower the level of known base to a point BELOW the calibration point.  RECORD this volume.

37.  Place the flask with the unknown acid under the buret with the buret extending BELOW the flask’s mouth.

38.  SLOWLY release base from the buret, swirling CONSTANTLY. Continue this until the color takes a noticeably LONGER time to disappear (a few seconds)

39.  At this point, only add A DROP AT A TIME of known base, swirl until color fades.

40.  STOP adding known base when the color takes 30 seconds to disappear (or DOESN’T disappear).  RECORD the new volume of known base

41.  DO THE MATH

42.  Obtain the balanced neutralization equation.  Determine the number of moles of KNOWN acid or base used to reach the equivalence point  Determine the number of moles of UNKNOWN acid or base present.  Determine the molarity of the UNKNOWN acid or base solution.

44.  A 15.5mL sample of 0.215M KOH solution required 21.2mL of aqueous acetic acid solution in a titration experiment. Calculate the molarity of the acetic acid solution.

45.  By titration, 17.6mL of aqueous H 2 SO 4 neutralized 27.4mL of 0.0165M LiOH solution. What was the molarity of the aqueous acid solution?

46.  How many moles of HCl are in 31.15 mL of a 0.688 M solution?  How many moles of NaOH would neutralize 20.0 mL of a 13.9 M solution of H 2 SO 4 ?  How many milliliters of a 2.76 M KOH solution contain 0.0825 mol of KOH?  4. A 25.00 mL sample of a solution of RbOH is neutralized by 19.22 mL of a 1.017 M solution of HBr. What is the molarity of RbOH?

47.  If 29.96 mL of a solution of Ba(OH) 2 requires 16.08 mL of a 2.303 M solution of HNO 3 for complete titration, what is the molarity of the Ba(OH) 2 solution?  You have a vinegar solution believed to be 0.83 M. You are going to titrate 20.00 mL of it with a NaOH solution known to be 0.519 M. At what volume of added NaOH would you expect to see an endpoint?

48.  Buffers 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!

49.  Buffers 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.

50.  Example: Ammonia is combined with its salt, NH 4 Cl, in sol’n:  If acid is added to this solution, ammonia reacts with the H + : NH 3 (aq) + H + (aq) → NH 4 + (aq)  If a base is added to this solution, the NH 4 + from the dissolved salt will react with the OH - : NH 4 + (aq) + OH - (aq) → NH 3 (aq) + H 2 O (l)

51.  Blood’s pH is regulated by many systems, but dissolved CO 2 is a very important method. Carbonic acid, H 2 CO 3, and the hydrogen carbonate ion, HCO 3 -, are both dissolved in your blood. CO 2 (g) + H 2 O (l) → H 2 CO 3 (aq)  If you add OH - : H 2 CO 3 (aq) + OH - → HCO 3 - (aq) + H 2 O (l)  If you add H + : HCO 3 - (aq) + H + → H 2 CO 3 (aq)

52.  Your 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 CO 2.

53.  If you hyperventilate, too much CO 2 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 CO 2 in your lungs and restore the proper pH.