Today, Let’s Review An Acid-Base Titration

In acid-base titrations, a buret is used to deliver measured volumes of an acid or base solution of known titration (the titrant) to a flask that contains a solution of a base or an acid, respectively, of unknown concentration (the unknown). If the concentration of the titrant is known, then the concentration of the unknown can be determined. Plotting the pH changes that occur during an acid-base titration against the amount of acid or base added produces a titration curve; the shape of the curve provides important information about what is occurring in solution during the titration. An Acid-Base Titration

Let’s WatchAn Acid-Base Titration —A Titration of a Strong Acid with a Strong Base DID YOU NOTE:When a strong base is added to a strong acid, At 1st- pH is low noting excess H+; as base is added, OH- neutralizes the H+ and pH increases. The equivalence point should be at a pH of ______ Addition of a strong base after the equivalence causes an excess of OH– and produces a rapid increase in pH.

Titrations of Strong Acids and Bases A pH titration curve shows a sharp increase in pH in the region near the equivalence point and produces an S-shaped curve; the shape depends only on the concentration of the acid and base, not on their identity.

For the AP Exam, you have to be able to.. Calculate the pH at Each region of Titration Curve

Calculations of Titration of a Strong Acid by a Strong Base What is the pH at each of the following points in the titration of 25 ml of 0.1 M HNO3 by 0.1 M NaOH? Before addition of ANY NaOH(Initial pH) After 24 ml of 0.1 M NaOH is added(before the equivalence point) After 25 ml of 0.1 M NaOH is added(at the equivalence point) After 26 ml of 0.1 M NaOH is added(after the equivalence point)

1st in all our calculations, Volume X Molarity = MOLES We will utilize: Volume X Molarity = MOLES 2nd – The Equivalence Pt. can always be found using VaMa = VbMb since this is where # moles of acid = # moles of base

Before addition of ANY NaOH (Initial pH) 25 ml of 0.1 M HNO3 is added to the flask HNO3 100% dissociates --> H + + NO3 - Initial H+ = 0.1M = 10-1 M pH = 1

After 24 ml 0.1 M NaOH added no. moles H+ =(0.025 L)(0.1 M HNO3 ) = .0025 mmoles H+ no. moles OH - = (0.024 L)(0.1 M NaOH) = .0024 mmoles OH -

H3O+ + OH - ----> 2H2O Initial 0.0025 0.0024 Reacting -0.0024 -0.0024 After 0.0001mol [H+] = 0.0001 mmol/.025 +.024L= [H+] =2.04 x 10-3 M pH = 2.69 Note: TOTAL VOLUME USED

After 25 ml 0.1 M NaOH added H3O+ + OH - ----> 2H2O Initial 0.0025 0.0025 Reacting -0.0025 -0.0025 After 0.0 0.0 NEUTRALIZATION After 1 x 10-7 M (ONLY H3O+ from H2O) pH = 7.0

AFTER 26 ML OF 0.1 M NAOH H3O+ + OH - ----> 2H2O Initial 0.0025 0.0026 Reacting-0.0025 -0.0025 After 0 0.0001 [OH -] = .0001 mmol/.025 + .026 L= [OH -] =1.96 x 10-3 M pOH = 2.71 , pH = 11.29

Summary: For strong acid strong base titration 1. The pH has a low value at the beginning 2. The pH changes slowly just before the equivalence point 3. At the equivalence point, the pH rises sharply with addition of only 0.l ml base 4. Any indicator whose color changes between pH 4-10 is suitable for this titration.

I’m going back to the movie of Acid-Base Titration to see if we can calculate the pH after each addition of NaOH. We know the concentration of acid is 0.1M Fill in your chart as I fill in mine! Movie

Practice Problem: We’ll do a second on paper! Find the pH of the solution if 0.1M HCl is titrated with .1 M NaOH. Starting with 10 ml HCl after 1 ml 0.1 M NaOH Added after 5 ml 0.1 M NaOH Added after 10 ml 0.1 M NaOH Added after 15 ml 0.1 M NaOH Added after 25 ml 0.1 M NaOH Added

Titrations of Weak Acids and Bases The shape of the titration curve for a weak acid or a weak base depends dramatically on the identity of the acid or base and the corresponding value of Ka or Kb.

Titrations of Weak Acids and Bases The pH changes much more gradually around the equivalence point in the titration of a weak acid or a weak base. [H+] of a solution of a weak acid (HA) is not equal to the concentration of the acid but depends on both its pKa and its concentration. Only a fraction of a weak acid dissociates, so [H+] is less than [HA]; therefore, the pH of a solution of a weak acid is higher than the pH of a solution of a strong acid of the same concentration.

Calculating the pH of a solution of a WEAK ACID OR BASE – If the initial concentration of a weak acid or base are known, one can calculate the volume at the equivalence point Utilizing VaMa = VbMb

Calculating Titration Curve: Weak acid by a Strong Base (same 4 points) What is the pH at each of the following points in the titration of 25 ml of 0.1 M HC2H3O2 by 0.1 M NaOH --Before ANY NaOH added(initial pH) --After 10 ml 0.1 M NaOH added(before) --After 25 ml 0.1 M NaOH added(neutralization) --After 26ml 0.1 M NaOH added(after)

Initial pH H C2H3O2 <-----> H+ + C2H3O2 - 0.1M-x x x Ka = 1.8 x 10-5 = x2/ .1-x As2 x = 1.34 x 10-3 M = [H+ ] pH = 2.88

After 10 ml of 0.1M NaOH VNaOH MNaOH = .010 L(0.1 M) = = .001 mol OH - H C2H3O2 + OH -<-----> H2O + C2H3O2- init .0025 mol .001 mol react - .001 mol - .001 mol + .001 mol After .0015 mol 0 .001 mol

Calculation- H C2H3O2 <-----> H+ + C2H3O2 - .0015 mmol -x X .001 mmol +x (1st calc. MH C2H3O2 : .0015 mol/.035 L= 0.0429M 2ndcalc. MNaOH: .001 mol/.035 L= 0.0286 M) Ka = 1.8 x 10-5 = x (0.0286 M)/ 0.0429 M [H+]=x=2.7 x 10-5 pH = 4.58

Want to try Henderson -Hasselbach? pH = pKa + Log [C2H3O2 -] / [H C2H3O2 ] pH = 4.76 + Log (0.0286 M / 0.0429 M) pH = 4.76 + (-.17) pH = 4.58

After 25 ml NaOH added(Neutralization) VNaOH MNaOH =.025 L(0.1 M) = = .0025 mol OH - H C2H3O2 + OH -<-----> H2O + C2H3O2- init .0025 mol .0025 mol react -.0025 mol -.0025 mol +.0025 mol after 0 mmol 0 mmol .0025 mol

Acetate ion hydrolyzes: ONLY C2H3O2- ION LEFT!!! Acetate ion hydrolyzes: C2H3O2- + H2O <-----> H C2H3O2 + OH – .0025 mol -x x x Kh or b = Kw/Ka= 5.7 x 10-10 5.7 x 10-10 = x2 / .0025 mol / .050 L (0.05 M) [OH - ] = X = 5.7 x 10-6 pOH = 5.28 pH = 8.75

After 26 ml 0.1 M NaOH VNaOH MNaOH = .026 L(0.1 M) = = .0026 mol OH - H C2H3O2 + OH -<-----> H2O + C2H3O2- init .0025 mol .0026 mol react -.0025mmol -.0025mol +0025mol After 0 mmol 0.0001mol .0025mmol

That .0001 mol of Strong base controls the pH Calculation That .0001 mol of Strong base controls the pH [OH - ] = .0001 mol / .051 L [OH - ] = 1.96 x 10-3 M pOH = 2.71 pH = 11.29 WHEW !!

Principal Features: Weak Acid:Strong Base Titration Curve 1. Initial pH higher than in strong acid titration. 2. Sharp increase at start of titration 3. Long gradual change in pH before equivalence point. 4. At 1/2 neutralization point, pH = pKa 5. pH at equivalence point > 7.

6. Beyond equivalence point, titration curve identical to strong acid:strong base titration curve. 7. Steep portion at equivalence point- short.

Practice Problem: WA/WB Starting with 15 ml of 0.1 M H C2H3O2 What is pH initially? What is the pH After 5 ml 0.1 M NaOH added? After 10 ml 0.1 M NaOH added? After 15 ml 0.1 M NaOH added? After 25 ml 0.1 M NaOH added?

Important FINAL Note! BEST Titration curve video The midpoint of a titration is defined as the point at which exactly enough acid (or base) has been added to neutralize one-half of the acid (or base) originally present and occurs halfway to the equivalence point. At the midpoint of the titration of an acid, [HA] = [A–]. The pH at the midpoint of the titration of a weak acid is equal to the pKa of the weak acid.

Titrations of Polyprotic Acids or Bases When a strong base is added to a solution of a polyprotic acid, the neutralization reaction occurs in stages. 1. The most acidic group is titrated first, followed by the next most acidic, and so forth 2. If the pKa values are separated by at least three pKa units, then the overall titration curve shows well-resolved “steps” corresponding to the titration of each proton

Titrations of Polyprotic Acids or Bases

Indicators Most acid-base titrations are NOT monitored by recording the pH as a function of the amount of the strong acid or base solution used as a titrant Instead, an acid-base indicator is used, and they are compounds that change color at a particular pH and if carefully selected, undergo a dramatic color change at the pH corresponding to the equivalence point of the titration Acid-base indicators are typically weak acids or bases whose changes in color correspond to deprotonation or protonation of the indicator itself

Indicators

HLt <-> H+ + Lt- Ex Litmu Indicators The chemistry of indicators are described by the general equation Hn(aq) ⇋ H+ (aq) + n–(aq), where the protonated form is designated by Hn and the conjugate base by n– HLt <-> H+ + Lt- Ex Litmu Red Blue The ionization constant for the deprotonation of indicator Hn is Kin = [H+] [n–] / [Hn] The value of pKin determines the pH at which the indicator changes color

Indicators • A good indicator must have the following properties: 1. Color change must be easily detected 2. Color change must be rapid 3. Indicator molecule must not react with the substance being titrated 4. The indicator should have a pKin that is within one pH unit of the expected pH at the equivalence point of the titration

• Choosing the correct indicator for an acid-base titration Indicators • Choosing the correct indicator for an acid-base titration 1. For titrations of strong acids and strong bases (and vice versa), any indicator with a pKin between 4 and 10 will do 2. For the titration of a weak acid, the pH at the equivalence point is greater than 7, and an indicator such as phenolphthalein or thymol blue, with pKin > 7, should be used 3. For the titration of a weak base, where the pH at the equivalence point is less than 7, an indicator such as methyl red or bromcresol blue, with pKin < 7, should be used

Indicators

Indicators • Paper or plastic strips that contain combinations of indicators estimate the pH of a solution by simply dipping a piece of pH paper into it and comparing the resulting color with standards printed on the container