1. Chem. 31 – 10/23 Lecture

2. 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)

3. 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

4. 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

5. 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.

6. 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

7. 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

8. 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

9. 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-

10. 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

11. 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

12. 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

13. 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

14. 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 M? NaOH is a hygroscopic solid that also absorbs CO2.

15. 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 mL of 1.00 M HCl is added. The sample is heated to drive out CO2, and the excess HCl requires mL of M NaOH. What is the percent carbonate (CO32-) in the soil sample? 15

16. 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

17. 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.)