2. What is Analytical Chemistry ?
Analytical chemistry deals with separating, identifying, and quantifying the relative amounts of the components of an analyte.
Analyte = the thing to analyzed; the component(s) of a sample that are to be determined.

3. Analytical Chemistry analyze: "what is it? (qualitative analysis)
"how much is there?“ (quantitative analysis)

4. The role of analytical chemistry: central science
The relationship between analytical chemistry and the other sciences
Chemistry : Biological, Inorganic, Organic, Physical
Physics : Astrophysics, Astronomy, Biophysics
Biology : Botany, Genetics, Microbiology, Molecular biology, Zoology
Geology : Geophysics, Geochemistry, Paleontology, Paleobiology
Environmental science : Ecology, Meteorology, Oceanography
Medicine : Clinical, Medicinal, Pharmacy, Toxicology
Material science : Metallurgy, Polymers, Solid state
Engineering : Civil, Chemical, Electronical, Mechanical
Agriculture : Agronomy, Animal, Crop, Food, Horticulture, Soil
Social Science : Archeology, Anthropology, Forensics
Analytical
chemistry

5. Several different areas of analytical chemistry:
Clinical analysis - blood, urine, feces, cellular fluids, etc., for use in diagnosis.
Pharmaceutical analysis - establish the physical properties, toxicity, metabolites, quality control, etc.
3. Environmental analysis - pollutants, soil and water analysis, pesticides.
4. Forensic analysis - analysis related to criminology; DNA finger printing, finger print detection; blood analysis.
5. Industrial quality control - required by most companies to control product quality.
6. Bioanalytical chemistry and analysis - detection and/or analysis of biological components (i.e., proteins, DNA, RNA, carbohydrates, metabolites, etc.).
This often overlaps many areas.
Develop new tools for basic and clinical research.

6. History of Analytical Methods
Classical methods: early years (separation of analytes) via precipitation, extraction or distillation
Qualitative: recognized by color, boiling point, solubility, taste
Quantitative: gravimetric or titrimetric measurements
Instrumental Methods: newer, faster, more efficient
Physical properties of analytes: conductivity, electrode potential, light emission absorption, mass to charge ratio and fluorescence, many more…

7. Types of Analysis Gravimetric Methods Titrimetric (Volumetric) Methods
measure the mass of an analyte (or something chemically equivalent to the analyte)
Titrimetric (Volumetric) Methods
measure the quantity of a reagent needed to completely react the analyte
Electroanalytical Methods
measure the change in the electrical potential, current, resistance or charge produced by an analyte
Spectroscopic Methods
measure the interaction between electromagnetic radiation (light, UV, IR, etc.) and the analyte
Chemical Separations
separate and measure the analyte of interest by chemical means (chromatography)
Other Methods

8. 1.Define the information you need 2.Select an analysis method
Process of Analysis
1.Define the information you need
2.Select an analysis method
3.Obtain a sample & 'clean' it up
4.Prepare the sample, solutions and standards
5.Do the analysis!
6.Account for interferences
7.Calculate results and estimate reliability
8.Convert results to information

9. Expressing Analysis Results
percent composition (% composition) - X's 100
%W/W %W/V %V/V
part per thousand (ppt) - X's 1000
parts per million (ppm) - X's 106
parts per billion (ppb) - X's 109
e.g.
22 ppm (w/v) lead
124 ppb (w/w) atrazine in soil

10. Titrations Introduction 1.) Buret Evolution Primary tool for titration
Gay-Lussac (1824)
Blow out liquid
Mohr (1855)
Compression clip
Used for 100 years
Descroizilles (1806)
Pour out liquid
Henry (1846)
Copper stopcock
Mohr (1855)
Glass stopcock

11. Principles of Volumetric Analysis
titration
titrant
analyte
indicator
equivalence point vs. end point
titration error
blank titration

12. Principles of Volumetric Analysis
standardization
standard solution
Methods for establishing concentration
direct method
secondary standard solution

13. Titrations Introduction Controlled Chemical Reaction
2.) Volumetric analysis
Procedures in which we measure the volume of reagent needed to react with an analyte
3.) Titration
Increments of reagent solution (titrant) are added to analyte until reaction is complete.
- Usually using a buret
Calculate quantity of analyte from the amount of titrant added.
Requires large equilibrium constant
Requires rapid reaction
- Titrant is rapidly consumed by analyte
Controlled Chemical Reaction

14. Titrations Introduction 4.) Equivalence point
Quantity of added titrant is the exact amount necessary for stoichiometric reaction with the analyte
- Ideal theoretical result
Analyte
Oxalic acid
(colorless)
Titrant
(purple)
Equivalence point occurs when 2 moles of MnO4- is added to 5 moles of Oxalic acid

15. CuCl Titration with NaOH Before any addition of NaOH
Titrations
Introduction
5.) End point
What we actually measure
- Marked by a sudden change in the physical property of the solution
- Change in color, pH, voltage, current, absorbance of light, presence/absence ppt.
CuCl Titration with NaOH
After the addition of 8 drops of NaOH
End Point
Before any addition of NaOH

16. Titrations Introduction 5.) End point
Occurs from the addition of a slight excess of titrant
- Endpoint does not equal equivalence point
Analyte
Oxalic acid
(colorless)
Titrant
(purple)
After equivalence point occurs, excess MnO4- turns solution purple  Endpoint

17. Titrations Introduction 5.) End point Titration Error Primary Standard
- Difference between endpoint and equivalence point
- Corrected by a blank titration
i. repeat procedure without analyte
ii. Determine amount of titrant needed to observe change
iii. subtract blank volume from titration
Primary Standard
- Accuracy of titration requires knowing precisely the quantity of titrant added.
- 99.9% pure or better  accurately measure concentration
Analyte
Oxalic acid
(colorless)
Titrant
(purple)

18. Titrations Introduction 6.) Standardization Titration Standardization
Required when a non-primary titrant is used
- Prepare titrant with approximately the desired concentration
- Use it to titrate a primary standard
- Determine the concentration of the titrant
- Reverse of the normal titration process!!!
Titration
Standardization
titrant known concentration
titrant unknown concentration
analyte unknown concentration
analyte known concentration

19. Principles of Volumetric Analysis
primary standard
1. High purity
2. Stability toward air
3. Absence of hydrate water
4. Available at moderate cost
5. Soluble
6. Large F.W.
secondary standard solution

20. Titrations Introduction 7.) Back Titration
Add excess of one standard reagent (known concentration)
- Completely react all the analyte
- Add enough MnO4- so all oxalic acid is converted to product
Titrate excess standard reagent to determine how much is left
- Titrate Fe2+ to determine the amount of MnO4- that did not react with oxalic acid
- Differences is related to amount of analyte
- Useful if better/easier to detect endpoint
Analyte
Oxalic acid
(colorless)
Titrant
(purple)
(colorless)
(colorless)

21. Titrations Titration Calculations
1.) Key – relate moles of titrant to moles of analyte
2.) Standardization of Titrant Followed by Analysis of Unknown
Calculation of ascorbic acid in Vitamin C tablet:
Starch is used as an indicator: starch + I3-  starch-I3- complex
(clear) (deep blue)
(ii) Titrate ascorbic acid with I3-:
1 mole ascorbic acid  1 mole I3-

22. Titrations Titration Calculations
2.) Standardization of Titrant Followed by Analysis of Unknown
Standardization: Suppose mL of I3- solution is required to react with g of pure ascorbic acid, what is the molarity of the I3- solution?

23. Titrations Titration Calculations
2.) Standardization of Titrant Followed by Analysis of Unknown
Analysis of Unknown: A vitamin C tablet containing ascorbic acid plus an inert binder was ground to a powder, and g was titrated by mL of I3-. Find the weight percent of ascorbic acid in the tablet.

24. Titrations Spectrophotometric Titrations
1.) Use Absorbance of Light to Follow Progress of Titration
Example:
- Titrate a protein with Fe3+ where product (complex) has red color
- Product has an absorbance maximum at 465 nm
- Absorbance is proportional to the concentration of iron bound to protein
Analyte
(colorless)
titrant
(colorless)
(red)
As Fe3+ binds protein
solution turns red

25. Titrations Spectrophotometric Titrations
1.) Use Absorbance of Light to Follow Progress of Titration
Example:
- As more Fe3+ is added, red color and absorbance increases,
- When the protein is saturated with iron, no further color can form
- End point – intersection of two lines (titrant has some absorbance at 465nm)
When all the protein is bound to Fe3+,
no further increase in absorbance.
As Fe3+ continues to bind protein
red color and absorbance increases.

26. Titrations Precipitation Titration Curve
1.) Graph showing how the concentration of one of the reactants varies as titrant is added.
Understand the chemistry that occurs during titration
Learn how experimental control can be exerted to influence the quality of an analytical titration
- No end point at wrong pH
- Concentration of analyte and titrant and size of Ksp influence end point
- Help choose indicator for acid/base and oxidation/reduction titrations
Sharpness determined by titration condition
Monitor pH, voltage, current, color, absorbance, ppt.

27. Volumetric Procedures and Calculations
relate the moles of titrant to the moles of analyte
# moles titrant = # moles analyte
#molestitrant=(V*M)titrant =
#molesanalyte=(V*M)analyte

28. EXAMPLE: Describe the preparation of 2. 000 L of 0
EXAMPLE: Describe the preparation of L of M KHP (potassium hydrogen phthalate) solution (f.w ) from primary standard solid.

29. EXAMPLE: Describe the preparation of 2. 000 L of 0
EXAMPLE: Describe the preparation of L of M KHP (potassium hydrogen phthalate) solution (f.w ) from primary standard solid.
(2.000 Lsoln)( molKHP)( gKHP)
#g KHP =
(1 L soln) (1 mol KHP)

30. EXAMPLE: Describe the preparation of 2. 000 L of 0
EXAMPLE: Describe the preparation of L of M KHP (potassium hydrogen phthalate) solution (f.w ) from primary standard solid.
(2.000 Lsoln)( molKHP)( gKHP)
#g KHP =
(1 L soln) (1 mol KHP)
= g KHP

31. EXAMPLE: Describe the preparation of 2. 000 L of 0
EXAMPLE: Describe the preparation of L of M KHP (potassium hydrogen phthalate) solution (f.w ) from primary standard solid.
40.85 g of primary standard grade KHP is weighed on a balance and transferred to a 2-L volumetric flask. Carbonate free water is added to the flask until it reaches the bottom of the neck of the flask. The solution is then mixed. More carbonate free water is now added o bring the volume up to the line on the neck of the flask. That line represents a volume of L.

32. EXAMPLE: Describe the preparation of 1 L of 0. 1 M NaOH solution (f. w
EXAMPLE: Describe the preparation of 1 L of 0.1 M NaOH solution (f.w ) from reagent grade solid.

33. EXAMPLE: Describe the preparation of 1 L of 0. 1 M NaOH solution (f. w
EXAMPLE: Describe the preparation of 1 L of 0.1 M NaOH solution (f.w ) from reagent grade solid.
(1 Lsoln)(0.1 molNaOH)(40.00 gNaOH)
# g NaOH = (1 L soln) (1 mol NaOH)

34. EXAMPLE: Describe the preparation of 1 L of 0. 1 M NaOH solution (f. w
EXAMPLE: Describe the preparation of 1 L of 0.1 M NaOH solution (f.w ) from reagent grade solid.
(1 Lsoln)(0.1 molNaOH)(40.00 gNaOH)
# g NaOH = (1 L soln) (1 mol NaOH)
= 4 g NaOH)

35. EXAMPLE: Describe the preparation of 1 L of 0. 1 M NaOH solution (f. w
EXAMPLE: Describe the preparation of 1 L of 0.1 M NaOH solution (f.w ) from reagent grade solid.
(1 Lsoln)(0.1 molNaOH)(40.00 gNaOH)
# g NaOH = (1 L soln) (1 mol NaOH)
= 4 g NaOH)
4 g of NaOH are weighed and added to 1 L of carbonate free water.

36. EXAMPLE: What is the molarity of a NaOH solution when 37
EXAMPLE: What is the molarity of a NaOH solution when mL of the NaOH solution are required to react with g KHP primary standard solid?

37. EXAMPLE: What is the molarity of a NaOH solution when 37
EXAMPLE: What is the molarity of a NaOH solution when mL of the NaOH solution are required to react with g KHP primary standard solid?
( gKHP)(1 molKHP)(1 molNaOH)(1000 mLsoln)
M NaOH =
(37.85 mLsoln)( gKHP)(1 molKHP)(1 Lsoln)

38. EXAMPLE: What is the molarity of a NaOH solution when 37
EXAMPLE: What is the molarity of a NaOH solution when mL of the NaOH solution are required to react with g KHP primary standard solid?
( gKHP)(1 molKHP)(1 molNaOH)(1000 mLsoln)
M NaOH =
(37.85 mLsoln)( gKHP)(1 molKHP)(1 Lsoln)
= molar

39. EXAMPLE: What is the % KHP in an unknown if 42
EXAMPLE: What is the % KHP in an unknown if mL of the above NaOH soln are required to titrate g of unknown?

40. EXAMPLE: What is the % KHP in an unknown if 42
EXAMPLE: What is the % KHP in an unknown if mL of the above NaOH soln are required to titrate g of unknown?
(42.06mLsoln)( molNaOH)(1 mol KHP)
% KHP =
(1.545 g sample) (1 L soln) (1 mol NaOH)
( gKHP) (1 L soln)
X 100
(1 mol KHP)(1000 mL soln)

41. EXAMPLE: What is the % KHP in an unknown if 42
EXAMPLE: What is the % KHP in an unknown if mL of the above NaOH soln are required to titrate g of unknown?
(42.06mLsoln)( molNaOH)(1 mol KHP)
% KHP =
(1.545 g sample) (1 L soln) (1 mol NaOH)
( gKHP) (1 L soln)
X 100
(1 mol KHP)(1000 mL soln)
= % KHP

42. EXAMPLE: What is the molarity of an HCl solution if it took 39
EXAMPLE: What is the molarity of an HCl solution if it took mL of the above NaOH solution to titrate mL HCl solution?

43. Acid-Base Indicators

44. Precipitation Titration Curve
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
titration curve => pAg vs. vol. AgNO3 added

45. Precipitation Titration Curve
p-function pX = - log10[X]
precipitation titration curve
four types of calculations
initial point
before equivalence point
equivalence point
after equivalence point

46. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
titration curve => pAg vs. vol. AgNO3 added
initial point
after 0.0 mL of AgNO3 added
at the initial point of a titration of any type, only analyte is present, no titrant is present, therefore pAg can not be calculated.

47. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
before equivalence point
pAg can be accurately calculated only after some AgBr has started to form. This may take a few mL of titrant

48. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
before equivalence point
after 5.0 mL of AgNO3 added
VNaBr*MNaBr - VAgNO3*MAgNO3
MNaBr unreacted =
VNaBr + VAgNO3

49. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
before equivalence point
after 5.0 mL of AgNO3 added
(50.00mL* M) -(5.00mL* M)
MNaBr =
( )mL
MNaBr unreacted = 3.64 X 10-3M
Ksp = [Ag+][Br-] = 5.2 X 10-13M2

50. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
equivalence point
at mL of AgNO3 added
becomes when [Ag+] = [Br-]
[Ag+]2 = 5.2 X 10-13M2
[Ag+] = 7.21 X 10-7M
pAg = 6.14

51. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
after equivalence point
After the equivalence point there is very little change in the amount of precipitate present (except very close to the equivalence point)

52. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
after equivalence point
thus, at mL of AgNO3 added
VAgNO3*MAgNO3 - VNaBr*MNaBr
MAgNO3 unreacted =
VAgNO3 + VNaBr

53. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
after equivalence point
thus, at mL of AgNO3 added
VAgNO3*MAgNO3 - VNaBr*MNaBr
MAgNO3 unreacted =
VAgNO3 + VNaBr
(25.10 mL * ) - (50.00 mL * )
MAgNO3 =
( )mL

54. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
after equivalence point
thus, at mL of AgNO3 added
(25.10 mL * M) - (50.00 mL * M)
MAgNO3 =
unreacted ( )mL
MAgNO3 unreacted = 1.33 X 10-5M

55. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
after equivalence point
thus, at mL of AgNO3 added
MAgNO3 unreacted = 1.33 X 10-5M
[Ag+]total = [Ag+]AgNO3 unreacted + [Ag+]dissolved AgBr

56. EXAMPLE: Derive a curve for the titration of 50. 00 mL of 0
EXAMPLE: Derive a curve for the titration of mL of M NaBr with M AgNO3.
after equivalence point
thus, at mL of AgNO3 added
[Ag+]total = [Ag+]AgNO3 unreacted + [Ag+]dissolved AgBr
[Ag+]total = 1.33 X 10-5M + [Ag+]dissolved AgBr
[Ag+]total ~ 1.33 X 10-5M
pAg = 4.88