Titrimetric (volumetric) analysis is a quantitative analytical method where the volume of a solution of known concentration that reacts completely with the analyte is used to determine the concentration of the analyte. It involves a titrant of known concentration, an indicator to signal the endpoint, and standardization. Common titration reactions include acid-base neutralization, precipitation, complexation, and redox. Indicators are used to detect the endpoint visually through a color change. Primary standards have a known high purity and are used directly, while secondary standards are standardized against a primary standard. Burets and volumetric flasks are used to deliver and contain solutions accurately for titrations.

1. PRINCIPLES OF TITRIMETRIC (VOLUMETRIC)
ANALYSIS

2. Titrimetric methods
 as an analytical method in which the volume of a
solution of known concentration consumed during an
analysis is taken as a measure of the amount of active
constituent in a sample analyzed

3. Analyte – or the active constituent in the sample
Titrant – solution of known concentration
Indicator – chemical which changes color at a point
equivalent quantities of analyte and titrant have
reacted

4. Stoichiometric point or Equivalence point – the
theoretical point at which equivalent amounts of each
have reacted
Endpoint – a sudden change apparent by use of
indicators

5. Standardization – determination of the exact
concentration of the solution
Standardized using 2 types of standards:
Primary standard – subs of known and high
degree of purity; used in direct standardization
Secondary standard – std. soln used in indirect
standardization
- std solution of known conc usually
standardized by primary std.

6. Requirements for Primary Standard:
1. High purity = 99.9%
2. Definite and known composition
3. Not affected by drying
4. Soluble in water

7. Ex.
Stdzn. Of NaOH
4 primary stds:
1. Potassium acid phthalate (KHP)/ C6H4COOK .
COOH
2. Sulfamic acid (HSO3)
3. Benzoic acid
4. Potassium acid iodate

8. Example: standardization of HCL
 Standardize titrimetrically using sodium carbonate of
known purity as a primary standard or
 using standard NaOH as a secondary standard

9. Standardization using Primary
Standard

10. Std. solutions Primary std Indicator
NaOH KHP/sulfamic acid php
HCl Na2CO3 php
AgNO3 NaCl K2CrO4
EDTA CaCO3 hydroxynaphthol
KMnO4 Na2C2O4
Iodine As2O3 starch
Sodium Methoxide Benzoic acid Thymol blue
Ceric sulfate As2O3 orthophenanthr
oline
Karl Fischer reagent Sodium tartrate

11.  A solution of known concentration used to standardize
another solution

12. Secondary
Standard
Indicator
HCl NaOH php
NH4SCN AgNO3 Ferric alum
Na2S2O3 Iodine Starch

14. ANALYTE TYPE TITRANT INDICA
TOR
Acetic acid Alkalimetric-direct NaOH php
Milk of Magnesia Acidimetric- residual NaOH mr
Aspirin Alkalimetric-residual HCl php
NaCl Volhard-precipitation. NH4SCN FAS
Povidone-Iodine NH4SCN FAS
Calcium lactate Complexation EDTA hydroxynap
hthol
H2O2 Redox-direct KMnO4
NaNO2 Redox- residual KMnO4
Ammonium
Potassium tartrate
Iodimetry Iodine starch
CuSO4 Iodometry Na2S2O3 starch

15. TYPE
Menadione (vit. K) REDOX
Selenium sulfide IODOMETRY
Phenol IODOMETRY- residual
Malic acid INDIRECT
PERMANGANATE
OXIDATION
Sulfonamide DIAZOTIZATION
NaHCO3 DIRECT ACIDIMETRY
ZnO ACIDIMETRY-
RESIDUAL/
EDTA-DIRECT
CaCO3 EDTA-DIRECT

16. Types of Titration:
DIRECT TITRATION – one titrant used, one volumetric
solution
RESIDUAL TITRATION – two titrants, two volumetric
solutions
1st VS – is added in excess
2nd VS – used to titrate the excess

17. Residual Titration - is used whenever the direct
titration is not practicable
- for compounds which react too slowly
with titrant
- poor solubility
- volatile substances are involved
Blank Determination – process of repeating the
procedure but omitting the sample

18. Computation:
DIRECT TITRATION
% Purity = N x V X meq. wt./f x 1000 x100
wt of sample

19. RESIDUAL TITRATION
% Purity = N1V1 - N 2V2 X meq. wt. x100
wt of sample

20. Acids :
- # of replaceable Hydrogen
Ex: HCl = 1
H2 SO4 = 2
H3 PO 4 = 3
CH3 COOH = 1
Bases:
- # of replaceable OH
Ex: NaOH = 1
Mg(OH)2 = 2
Al(OH)3 = 3
Salts:
- total positive or negative charges/ cation or anion
Ex: NaCl = 1
MgO = 2
Ca3 (PO4)2 = 6

21.  Results in drug assays are usually expressed in terms of
%w/w, %w/v and %v/v

22. In the assay of conc. HCL:
 samples are taken by weight
 Difficult to measure small vol. of conc. acid
 Results are express in % w/ w

23. In the assay dil. HCl:
 Express in % w/ v

24. Standard Solution
 is a solution of known normality or molarity
NORMALITY
 eq/L or meq/ml
MOLARITY
 moles/L or mmoles/mL

25. N = wt.
V x meq. wt.
= g
mL x g/mole
meq/mole

26. M = moles
L
= mmoles
mL

27. 1. Find the molarity of HCl soln. which contains the
volume of 2400 mL and also it contains 230 g of
HCl. MW of HCl=36.46
2. Find the normality of 20 g NaOH diluted to a volume
of 1L. MW of NaOH= 40

28. 3. Compute for normality of HCl that make use of a
primary standard (Na2CO3) that weighs 1.5 g. It
consumed 30 mL of HCl after titration. MW of
Na2CO3 = 106
4. Compute for normality of HCl using secondary
standard (NaOH) having a concentration of 2 N. It
consumed 35 mL of NaOH. The volume sample used
was 25 mL.

29. 5. A soln. with a final volume of 500 mL was prepared by
dissolving 25mL of methanol (density= 0.7914) in
chloroform. Calculate the molarity of methanol in the
soln. MW = 32
6. Calculate the normality of a NaCl soln. prepared by
dissolving 2.5 g of NaCl in water and then tapping it
off with more water to a total volume of 500 mL. MW
of NaCl= 58.44 g/mole

30. 7. What mass of oxalic acid, H2C2O4, is required to
make 250 mL of 0.05M soln.(MW=90)
8. Sulfamic acid is a primary standard that can be used
to standardize NaOH. What is the molarity if 34.26 mL
reacts with 0.3337 g of sulfamic acid. MW= 97

31. 9. What is the normality of H2SO4 containing 73.5 g/
500 mL of solution? MW of H2SO4= 98 g/mole

32. 10. The molecular wt. of NaOH is 40. How many grams
of NaOH pellets are needed to make 500 mL of 1.5 N
soln?

33. 11. Convert the following concentration as required:
0.5 M KOH to N
0.025 M H2SO4 to N
2 N NaCl to M

34. Express in decimal:
12. One twentieth N of HCl
13. One tenth M of H2SO4

35. Standardization
 determination of normality or molarity of solution
 accomplished by the use of another standard solution
known as a SECONDARY STANDARD or by the use of
known purity substance as PRIMARY STANDARD

36.  The wt. of substance chemically equivalent to 1 mL of
std. soln.
 Express in mg/ mL
 To get titer value:
 N x meq wt.
 1 meq X 60 g x 1000 mg
mL 1000 meq 1 g
= 60 mg/mL

37. Ex.
1. What would the titer value be of the soln. in terms of
CaCl2 of 0.05 N soln.? MW = 111
2. Compute for percent purity of 0.1 g sodium ascorbate
if the titer value is 9.905 mg/mL. It consumed 10 mL
thru direct titration with iodine.

38. Volumetric Apparatus
2 types:
 to deliver: burets, pipets
 to contain: volumetric flask, graduated cylinders

39. Burets
 Graduated glass tubes of uniform bore throughout the
whole length
 Closed at the bottom by glass or stopcock
 Volumes read at lower meniscus except highly colored
liquid

40. BURETS
 Calibration by means of:
1. Ostwald pipet
2. Kiehl buret

41. CHEMICAL REACTIONS USED IN TITRIMETRY
 1. Neutralization (acid-base)
 2. precipitation
 3. complexation
 4. redox

42. NEUTRALIZATION REACTION
 a chemical process in which an acid (proton donor)
reacts with a base (proton acceptor)
 the products are : water and salt

43. Strong acid and Strong Base
HCL + NaOH → NaCl + H2O

44. INDICATORS
- are complex organic compounds used:
 to determine end point
 to determine ion concentration
 indicators may be acids or bases

45. Mixed or combined indicators
 are used to give sharp color change

46. Commonly used pH indicators
Color Change
Indicator pH Range Acid Base
 Malachite green 0.0 – 2.0 yellow green
 Methyl Yellow 2.9 – 4.0 red yellow
 Bromophenol blue 3.0 - 4.6 yellow blue
 Methyl Orange 3.2 – 4.4pink yellow
 Bromocresol green 4.0 -5.4 yellow blue

47.  Methyl red 4.2 – 6.2 red yellow
 Bromocresol purple 5.2 – 6.8 yellow purple
 Bromothymol blue 6.0 – 7.6 yellow blue
 Phenol red 6.8 – 8.2 yellow red
 Cresol red 7.2 – 8.8 yellow red
 Thymol blue 8.0 – 9.2 yellow blue
 Phenolphthalein 8.0 – 10.0 colorless red
 Thymolphthalein 9.3 – 10.5 colorless blue

48. Indicators:
Aqueous :
 Phenolphthalein
 Methyl orange
 Methyl red
 Strong Acid + Strong Base : php, MO, MR
 Weak Acid + Strong base : php
 Weak base + strong acid : MR

49. Strong Acids:
 HI
 HBr
 HCl
 H2SO4
 HNO3
Weak Acid :
 Acetic acid
 Oxalic Acid
 HF
 HCN

50. Strong Bases:
 NaOH
 KOH
Weak Bases:
 Alanine, C3H5O2NH2
Ammonia, NH3
Methylamine, CH3NH2
Pyridine, C5H5N

51. Standard Acid Solutions used in Acidimetry and
Alkalimetry
1. HCl – more preferable to sulfuric in the titration of
compounds that yield precipitate
2, Sulfuric Acid – preferable in hot titrations since HCl
will volatilize

52. Standard Alkali Solutions are:
 1. NaOH
 2. KOH
 3. Ba(OH)2

53. All alkali solutions should be
restandardized frequently