The document provides an overview of analytical chemistry principles, focusing on quantitative chemical analysis, including methods for both qualitative and quantitative analysis. It explains various methods of chemical analysis, including physical and chemical methods, and delves into standard solutions, titrations (direct and back titration), and titrimetry. Additionally, it covers acid-base titrations, detection of end points, and features of titration curves for strong and weak acids.

1. Analytical chemistry II (CHEM 321)
Contact Hr-2
CHAPTER ONE
Review of principles of quantitative chemical
analysis and concentration unit

2. Analytical chemistry
Deals with separating, identifying and determining the relative
amount of components in a sample of material.
Deals with elucidating the structural configration of susbstance
Chemical analyses can be
 Qualitative - identification of the chemical components
in a sample
- answers what is the sample type question
 Quantitative - determination of the amount of a certain
component in the sample.
- answer how much type question

3. STEPS IN CHEMICAL ANALYSIS
 Selecting a methjods of analysis based on
accuracy, reliability, economy
 Sampling
 Preparing laboratory sample
 Defining replicate sample
 Preparing solution of the sample
 Eliminating interference using precipitaion,
extraction, chromatography, distilation etc
 Completing the analysis and calibration
 calculating result and report

4. Method Of Chemical Analysis
A. Physical method (physicochemical methods)
 uses physical features of substance being analyzed
 physical state, color, luster, ability to melt, sublime,
color of flame, hardness
 potential
 Conductance
 Current strength
B. chemical method
by adding one or a series of chemical reagents to the analyte and
observing chemical reactions and their products
example Gravimetric and titrimetric methods

5. Solution – homogeneous mixture containing two or more
than two components
Molarity
Normality
Equvalent weight
Equivalent weight of acid base reaction
H2SO4 + NaOH Na2SO4 + H2O
H2SO4 + NaOH Na2SO4 + H2O
KMnO4 + KI + H2SO4 K2SO4 + MnSO4 + I2 + H2O redox reaction

6. Stadard solution
Is a solution containing a precisely known concentration
of an element or a substance
Is a solution of accurately known concentartion prepared
from primary standard
For example, with NaCl, its formula mass is 58.43 g/mol.
When preparing this standardized solution, we would need
to add 23.4 g of NaCl to a container that has 400 ml.

7. Desirable properties of the standard solution
• Be sufficiently stable
• React rapidly with the analyte
• React more or less completely
• Under go selective reaction with the analyte
Types of standard solution
A. primary standard solution
 is a solution with a high purity and less
reactivity.

8. Requirment for primary standard
High purity
Stability in air
Must not be hygroscopic (not absorb water)
Moderate cost and easly available
Large formula weight solid (to minimize weight error)
Reasonable solubility

9. commonly used primary standards
Potassium hydrogen phthalate (KHP, KHC8H4O4)
 Used to standardize base using phenolphthalein indicator
Potasium hydrogen iodate (KH(IO3)2)
 Strong acids to standardize bases
 Any indicator with end point b/n pH 5-9 may be
used
Sodium carbonate (Na2CO3)
 bases used to standardized strong acids to an end point
of pH 4-5
Benzoic acid (C6H5COOH), sodium oxalate
(Na2C2O4), borax (Na2B4O710H2O) mercuric oxide
(HgO) are some other primary standard.

10. Questions
1. Exactly 0.8032 g of the primary standard KHP
(molecular weight = 204.23g) was weighted into an
erlemnyer flask and dissolved in about 30 ml of water.
About two drops of an indicator was added and the
solution was treated with NaOH solution until color of
the dolution is changed from colorless to light pink.
Calculate the concentration of NaOH solution
corresponding to an end point volume of 35.20ml
KHP + NaOH H2O + NaKP
2. A 25.00ml acidic sample of FeCl2 was treated to an end
point with 15.32ml of 0.09158M KMnO4. calculate the
concentration and mass of FeCl2.
5Fe2+
+MnO4
-
+8H+
5Fe3+
+Mn2+
+4H2O

11. B. secondary standard
is not pure and is chemically reactive than primary
standards.
is a solution that is made specifically for a
certain analysis.
is a substance whose active agent contents have
been found by comparison against a primary
standard.
 Example NaOH. Commercially available NaOH
contains impurities of NaCl, Na2CO3, and Na2SO4,
and readily absorbs H2O from the atmosphere. To
determine the concentration of NaOH in a solution,
we titrate it against a primary standard weak acid,
such as potassium hydrogen phthalate, KHC8H4O4.

12. question
1. A highly pure solid KHP was dissolved in water
untill the volume is 100ml and this solution is
used to titrate a small portion of a 50ml of NaOH
solution. Then un known HCl solution is titrated
with NaOH. Mention the primary standard,
secondary standard and the anlyte from the above
volumetriic analysis.

13. THE DIFFERENCE BETWEEN PRIMARY AND SECONDARY STANDARDS

14. UNIT TWO
TITRIMETRIC METHODS OF ANALYSIS
TITRIMETRIC METHODS
 Analytical procedures in which the amount of analyte
(volume or mass) is determined from the amount of
standard reagent required to react completley with analyte.
Volumetric titrimetry involves measuring the volume of a
solution of known concentration that is needed to react
essencially and completly with the analyte.
gravimetric titrimetry involves measuring the mass of the
precipitate formed in the titration.

15. • Titration:-
• is a technique in which a solution of known concentration is used to determine the
concentration of an unknown solution.
• Is the process of adding a measured and controlled volume of standard solution from a
burrete to a solution of known volume containing a substance to be determined
(analyte)
• The process in wich standard reagent is added to a solution of analyte untill the reaction
b/n the analye and the reagent is judged complete.
• Titrant:- is solution of known concentration.
- Added from a burette to a known quantity of the analyte which is made to react.
• Titrand/analyte:- Is the substance being determined during titration
- is solution of unknown concentration.
- solution kept in the Erlenmeyer flask

16. Titration set up

17. Classes of volumetric titrimetric method
Acid base or neutralization titration
•Based upon chemical reaction b/n acid and base

18. Precipitation titration
•Based upon the reaction that yield ionic cpds which are
slightly soluble
Complex formation titration
•Based on the reaction occurring b/n the metal ion (lewis
acid) and ligand (lewis base)
•Redox titration
•Is one in which the substance to be determined is either
oxidized or reduced

19. Titrimetric method based on mode of operation
A. Direct titration
Known volume of analyte is titrated with standard
solution of a reagent.
 examples Acid-base titrations titration
Question in titration of sulfuric acid against sodium
hydroxide, 32.2ml of 0.25M NaOH is required to neutralize
26.6ml of H2SO4. calculate the molarity of sulfuric acid.
H2SO4(aq)+2NaOH(aq)→Na2SO4(aq)+2H2O(l)
ans..0.151MH2SO4

20. B. Back/indirect titration
•Is the process in which excess of the standard
solution used to react with the analyte is determined
by titration with with a second standard solution.
For example the amount of phosphate in a sample can be
determined by adding a measured excess of the standard
silver nitrate to a solution of the sample, which leads to the
formation of insoluble silver phosphate.
The unreacted excess of the silver nitrate is then treated with
a standard solution of potasium thiocyanate.
3Ag+(excess )+ PO4
-3
Ag3PO4(s) + Ag+
(unreacted)
Ag+
(unreacted) + SCN- AgSCN(s)

21. Question
•A student was asked to determine the concentration of
ammonia, First the student pipetted 25.00 mL of the cloudy
ammonia solution into a 250.0 mL conical flask. 50.00 mL
of 0.100 mol L-1
HCl(aq) was immediately added to the
conical flask which reacted with the ammonia in solution.
The excess (unreacted) HCl was then titrated with 0.050 mol
L-1
Na2CO3(aq). 21.50 mL of Na2CO3(aq) was required.
Calculate the concentration of the ammonia in the cloudy
ammonia solution.

22. Purpose of back titration
We use back titration if
The analyte may be in solid form like chalk (not
soluble in water)
 The analyte may contain impurities which may
interfere with direct titration.
 when the end point is more easly identified than
in forward reaction.

23. Difference between direct and back titration
Direct titration Back titration
Reactions one chemical reaction occurs. two chemical reactions occur
Titration done b/n a known compound and
an unknown compound.
titration is done between two
known compounds
Titrand is the unknown compound. is the remaining amount of the
reagent added in excess.
Application used when the endpoint of the
titration can easily be obtained.
used to determine the exact
endpoint when there are sharp
color changes.

24. End point and equivalence poin during titration
End point
•A point at which a physical change (color change)
associated with the equivalence point appears
Equivalence point
•Point in a titration where the reagent added is chemically
equivalent to the analyte.
•It is the point at which the reaction is theoretically complete.
Note:- under ideal condition the equivalence point and the
end point are identical.
- In most practical analysis the end point is near, equal
but not identical with the equivalence point.

25. Detection of end point during titration
Visual indicators
Change of color of the added indicators
Formation of soluble product
Appearance of a precipitate
Formation of soluble colored product

26. UNIT THREE ACID-BASE TITRATION
•It is a process of determining the concentration of
an acid or base solution.
•For example if the solution of unkown concentarion
is acidic, a standard base solution is added to acidic
solution untill it is neutralized. On the other hand if
the solution of unkown concentration is basic, a
standard solution of acid is added to the basic
solution untill it is neutralized.

27. DETECTION OF END POINTS
•The point of chemical equivalence can be is indicaed by
chemical indicators or instrumental methods.
•Due to simplicity of the technique and availlability of many
chemical indicators, equivalence points in acid base titration
are determined most often by acid-base indicators.
Acid base indicators
Are weak organic acids or weak organic bases that indicate
whether a solution is acidic, basic or neutral.

28. Acid type indicators
Represented by idealized formula HIn which ionizes to give
its conjugate base In-
HIn(aq) + H2O(l) H3O+(aq) + In-(aq)
Example phenolphthalien (colorless for acid and pink for
base)
Base type indicators
Represented by idealized formula In- which ionizes to give
its conjugate acid Hin
HIn(aq) + H2O(l) H3O+(aq) + In-(aq)
Example methyl orange(red for acid and light yellow color
for base)

29. STRON ACID –STRONG BASE TITRATION
• Due to complete ionization of strong acid and strong base the net
ionic reaction is the reaction of hydronium ion and hydroxide ion to
give water.
• Features of the titration curve for the titration of a strong acid with a
strong base.
 The pH is low at the beginning of the titration.
 The pH change slowly until just before the equivalence point.
 Just before the equivalence point, the pH rises sharply.
 At the equivalence point, the pH is 7.00.
 Just past the equivalence point, the pH continues its sharp
rise.
 Further beyond the equivalence point, the pH continues to increase,
but much less slowly.
Question what will be the features of titration of strong base with
strong acid?

30. Titration curves of stron acid with strong base

31. •To drive titration curves of strong acid with strong base
three different regions must be considered
Before equivalence point
At this region pH can be calculated before and after
addtion of titrants (i.e strong base).
Before addition of titrant the concentration of hydronium
ion is the concentration of strong acid.
After addition of titrant when strong base is added some of
hydronium ion is consumed, so the remaining hydronium
can be calculated by
H3O+=

32. Equivalence point
At equivalence point equivalent amount of acid and base are
brought, no more concentration of acid and base is left. To
calculate the pH of the solution dissociation of water must
be considered.
Kw=[H3O+][OH-]
[H3O+]=
After equivalence point
At this stage no more concentarion of acid is left, the
concentration of hydroxide ion can be calculated by
OH-=
[H3O+]=

33. Example Calculate the pH for titration of 20 mL of 0.5 M
HCl (a strong acidwith 0.5 M NaOH (strong base)
a. Before additon ofNaOH
b. After addition of 5ml of NaOH
c. After addtion of 10ml of NaOH
d. After addition of 15ml of NaOH
e. After addion of 20 ml of NaOH
f. After addtion of 25ml of NaOH
g. After addtion of 30 ml of NaOH

34. Titration of weak acid with strong base
• The conjugate base of a weak acid will undergo hydrolysis, which
will affect the pH of the solution.
• Example consider weak acid HA and its conjugate base A-
A-+H2O HA + OH-
Features of titration of a weak acid with a strong base
 The initial pH is higher because the weak acid is only partially
ionized.
 At the half-neutralization, pH = pKa.
The pH is greater than 7 at the equivalence point because the anion
of the weak acid hydrolyzes.

35. Titration curves of weak acid with strong base

36. •pH calculation for titration of weak acid with strong base
Befor addition of any strong base
• pH can be calculated by considering dissociation of weak
acid
HA +H2O H3O++ A-
ka=
[H3O+] = ka, [H3O+] =
[H3O+]=

37. Before equivalence point
At this stage we have mixture of unreacted weak acid and conjugated
base formed from reaction of weak acid and strong base and the pH
of the solution can be calculated using Hendersen-Hasselbalch
equation.
pH=pKa +log
Concentration unreacted acid =
Concentration of conjugate base =
At Equivalence Point
All the acid and base is consumed but the hydrolysis of conjugate
base is considered.
A-+H2O HA + OH-
Kb = = [H3O+] =

38. After equivalence point
•At this point all acids are consumed and excess of base is
remained.
•Even if the conjugate base is able to under go hydrolysis,
the strong base prevents its hydrolysis.
[H3O+] =

39. Example
•consider the titration of 20 mL of 0.5 M acetic acid, CH3COOH,
with 0.5 M NaOH. calculate pH at this different region
a. Before additon ofNaOH
b. After addition of 5ml of NaOH
c. After addtion of 10ml of NaOH
d. After addition of 15ml of NaOH
e. After addion of 20 ml of NaOH
f. After addtion of 25ml of NaOH
g. After addtion of 30 ml of NaOH