Chem. 31 – 10/23 Lecture

Announcements Statistical Calculations Lab Resubmissions due 10/25 AA Lab – Scheduled due date is 10/30 Today’s Lecture Chapter 6 Polyprotic acids Chapter 7 – Titrations Overview and definitions Detection of endpoints Back titrations Precipitations (covered qualitatively this semester)

Polyprotic Acids Release more than 1 H+ per molecule Examples: H2SO4, H3PO4, H2C2O4 H3PO4 has 3 Ka values (Ka1, Ka2, Ka3) for 3 reactions losing H+: H3PO4 ↔ H2PO4- + H+ Ka1 H2PO4- ↔ HPO42- + H+ Ka2 HPO42- ↔ PO43- + H+ Ka3

Chapter 7 - Titrations Introduction – Overview Chapter 7 covers general titrations (quantitation, practical aspects, types of titrations, shape of precipitation titration curve – not covering calculations due to time) Chapter 11 covers titration curves for acid-base titrations - covered later Other Chapters (12, 16) cover other types of titrations – not covered

Titrations Definitions Titrant: Reagent solution added out of buret (concentration usually known) Analyte solution: Solution containing analyte Equivalence Point: point where ratio of moles of titrant to moles of analyte is equal to the stoichiometric ratio titrant analyte solution for: Al3+ + 3C2O42- → Al(C2O4)33- n(Al3+)/n(C2O42-) = 1/3 at equivalence pt.

Titrations Practical Requirements The equilibrium constant must be large Size of K value depends on desired precision and concentration of analyte Typically K ~ 106 is marginal, K > 1010 is better The reaction must be fast It must be possible to “observe” the equivalence point observed equivalence point = end point

Titrations Detection of Endpoints An endpoint is defined as the point in the titration when the equivalence point is observed Ways to detect endpoints: Use of colored reactants example: MnO4- + H2C2O4 (aq) → Mn2+ + CO2 (g) Use of indicators An indicator changes color in response to the change in a reactant’s concentration Use of simple instruments Must respond quickly, but typical equipment is low cost PINK Clear Clear

Titrations Detection of Endpoints Simple instruments electrodes (typically respond to log of ion concentrations) spectroscopic measurements (measurement of absorption of light) Can improve titration precision vs. using indicators Titration Error = Difference between end point and equivalence point = systematic error Note: It is possible to have large errors or uncertainties in detection of reagent conc. by various methods without having great titration errors to meter

Titrations Other Definitions Standardization vs. Analyte Titrations To accurately determine an analyte’s concentration, the titrant concentration must be well known This can be done by preparing a primary standard (high purity standard) Alternatively, the titrant concentration can be determined in a standardization titration (e.g. vs. a known standard) Rationale: many solutions can not be prepared accurately from available standards Example: determination of [H2O2] by titration with MnO4- neither compound is very stable so no primary standard instead, [MnO4-] determined by titration with H2C2O4 in standardization titration then, H2O2 titrated using standardized MnO4-

Titrations What Makes a Titration Sharp? SHARP TITRATION A sharp titration has a large slope (absolute value) Slope at endpoint seen in plot of log[analyte] vs. V(titrant) With a sharp titration, errors or uncertainties in V(equivalence point) are small uncertainties in log[analyte] [reactant] at eq. point Log[analyte] V(eq. pt.) V(titrant) small uncertainty in V results NON-SHARP TITRATION Log[analyte] V(titrant) larger unc. in V

Titrations Other Definitions Direct vs. Back Titration In a direct titration, the titrant added slowly to the analyte until reaching an end point In a back titration, a reagent is added to the analyte in excess, and then that reagent is titrated to an end point Often done to get sharper Endpoint

Titrations Back Titration Example Titration to determine moles of Na2CO3 in a sample: First, direct titration: Na2CO3 + 2HCl → H2CO3 + NaCl (we will do as next to last lab) HCl not that sharp Na2CO3

Titrations Back Titration Example NaOH HCl Titration to determine moles of Na2CO3 in a sample: Now, via back titration: excess HCl added to sample Na2CO3 + HCl → NaCl + H2CO3 +heat → NaCl + H2O + CO2(g) After heating only NaCl and excess HCl left Excess HCl titrated with NaOH to NaCl + H2O Na2CO3 Very Sharp Excess HCl

Titrations Some Questions List two requirements for a titration to be functional. In a back titration, what is actually being titrated? (a) analyte b) reagent added c) excess reagent d) secondary reagent) Why might one want to standardize a prepared solution of 0.1 M NaOH rather than prepare it to exactly 0.100 M? NaOH is a hygroscopic solid that also absorbs CO2.

Titrations Back Titration Example The Mass percent of carbonate is determined in a soil sample by a back titration. A 1.00 g soil sample is placed in a flask and then 10.00 mL of 1.00 M HCl is added. The sample is heated to drive out CO2, and the excess HCl requires 38.11 mL of 0.0825 M NaOH. What is the percent carbonate (CO32-) in the soil sample? 15

Precipitation Titrations - covering qualitatively Example: Titration of Hg22+ by CrO42- Hg22+ + CrO42- → Hg2CrO4 (s) Ksp(Hg2CrO4) = 2.0 x 10-9 K = 1/Ksp = 5 x 108 = large (reaction near full to product) Titration has 3 regimes: Before equivalence point (excess Hg22+ in flask) – [Hg22+] is high At equivalence point (nHg2^2+/nCrO4^2- = 1/1) [Hg22+] is rapidly decreasing After equivalence point (excess CrO42- in flask) [Hg22+] is low CrO42- Hg22+ This is different than text example

Titrations Shapes of Titration Curves – Precipitation Example However, moles are not readily measured. Concentration or log[Hg22+] more readily measured. Log[Hg22+] or pHg22+ ( = -log[Hg22+]) is plotted on y-axis Plot of moles in flask vs. V(titrant) Easier to understand At equivalence point both Hg22+ and CrO42- are present in low amounts moles analyte moles titrant [Hg22+] = Ksp1/2 V(titrant) V(eq. pt.)