phc-109-lec.pdf

Analytical
Chemistry Methods
James Verano Senido, RPh

May be categorized based on:
• Size of the sample
• Extent of Determination
• Nature of Analytical Methods
Quantitative Analysis
Pharmaceutical Analysis I

Description Amount (g)
Macro > 1 gram
Semimicro (Meso) 0.1 g to 1 g
Micro 0.01 g to 0.1 g
Submicro 0.001 to 0.01 g
Ultramicro <0.001 gram
Analysis based on SAMPLE SIZE or
AMOUNT

PROXIMATE ASSAYS
• Determination of the TOTAL of a CLASS/GROUP of active
principles in a given sample.
Examples: Assay of Carbohydrates, Assay of Alkaloids
ULTIMATE ASSAYS
• Determination of a SINGLE CHEMICAL SPECIES from a
sample.
Examples: Assay of Sucrose, Assay of Caffeine
Analysis based on EXTENT OF
DETERMINATION

Analysis based on NATURE OF
ANALYTICAL METHOD

• Also known as Wet/Stoichiometric Method
• Analyte is made to react with another substance according
to a well-defined chemical equation.
• The amount is calculated from the amount of reagent used.
Classical Methods
Classical Methods

• The separation by extraction, precipitation, or other means, of
the constituent to be determined either in the natural state or in
the form of a definite compound of known composition and then
weighing of the resulting product.
• The mass of the analyte or some compound is to be
determined.
Gravimetric Method

1
• Preparation of a solution containing a known
weight of the sample
2
• Separation of the desired constituent
3
• Weighing the isolated constituent
4
• Computation of the amount of the particular
constituent in the sample from the observed
weight of the isolated substance.
Gravimetric Method

• Determination of the volume of solution of
known concentration (titrant) required to
complete a chemical reaction with a
substance being analyzed (analyte).
• Also known as Volumetric analysis
Analyte + Titrant → Product
HCl + NaOH → NaCl + H2O
Titrimetric Analysis

1. The titration reaction must be well defined without side
reactions.
2. The reaction must be virtually complete (approximately 100%
of analyte must be converted to product).
3. Other substances in the sample must not react with the
reagent.
4. The reaction rate should be high.
5. There must be a method to detect when the equivalence point
is reached.
6. The exact concentration of reagent must be known.
The following requirements should be
fulfilled to make a titration feasible:

• Analyte/Titrand- chemical substance
being analyzed.
• Titrant- solution of known
concentration added to react with
analyte
• End point- point where titration is
stopped, characterized by a sudden
change in the property of a reaction
mixture.
• Indicator- a chemical which changes
color at or very near the point in the
titration where equivalent quantities of
the titrant and analyte have reacted.

Stoichiometric point
• aka Equivalence point
• Theoretical point at which equivalent
amounts of each titrant and analyte
have reacted.

Gram-Equivalent Weight (GEW)
• Weight in grams that is chemically equivalent to 1 gram atom of
hydrogen.
• Weight of a substance in grams which contains, furnishes, reacts
with directly or indirectly, or replaces 1 gram-atom or ion of
hydrogen (1.0079).
Gram-milliequivalent weight (GMEW)
• GEW/1000

Standard Solution
• A solution of known molarity or normality.
Standardization
• The determination of the exact concentration (i.e. N or M) of a
solution
Primary Standard
• A carefully weighed sample of a substance of known purity
Secondary Standard
• Another standard solution aside from primary standard used for
standardization purposes.

TITER
• Weight of a substance chemically equivalent to 1 ml of standard solution.
Titer (grams/ml) = N x ml x
𝑀𝑊
𝑓 𝑥 1000
Example: Calculate the titer value (in mg/ml) of Calcium Hydroxide (Ca(OH)2) for
0.1 N HCl standard solution (Ca = 40.08; O = 16, H= 1)
N x ml x
𝑀𝑊
𝑓 𝑥 1000
= 0.1 N x 1 ml x
(40.08 𝑥 1) +(16 𝑥 2) +(1 𝑥 2)
2 𝑥 1000
= 0.003704 grams = 3.704 mg/ml

Calculate the following titer values, in grams per ml, for 1 N
Sulfuric Acid:
• Potassium Bicarbonate (KHCO3)
• Potassium Carbonate (K2CO3)
• Calcium Carbonate (CaCO3)
(Ca = 40.08; Carbon = 12; Hydrogen = 1; Oxygen = 16,
Potassium = 39)
Checkpoint:

Potassium Bicarbonate, KHCO3
1 N x 1 ml x
39 𝑥 1 + 1 𝑥 1 + 12 𝑥 1 + 16 𝑥 3
1 𝑥 1000
= 0.1 gram/ml
Potassium Carbonate, K2CO3
1 N x 1 ml x
39 𝑥 2 + 12 𝑥 1 + 16 𝑥 3
2 𝑥 1000
= 0.069 gram/ml
Calcium Carbonate, CaCO3
1 N x 1 ml x
40.08𝑥 1 + 12 𝑥 1 + 16 𝑥 3
2 𝑥 1000
= 0.05 gram/ml
Solution

1. Neutralization (Acid-Base) Reactions
2. Oxidation-Reduction (Redox)
3. Precipitation
4. Complexation
Chemical Reactions used in Titrimetry

Acid-Base Titration

• Also known as Neutralization Reaction
• A technique to determine the concentration of an acid or base
analytes by neutralizing the unknown concentration of an acidic or
basic analyte with the known concentration of an acid or base.
• A neutralization reaction is determined by using an indicator where
the indicator will change its color at the end point of the titration.
This method allows the quantitative analysis for the unknown
acid/alkali concentration.
Acid-Base Titration

Arrhenius definition (from Svante Arrhenius)
• Acids: substances that increases hydrogen
(H+) ions when added to water; form hydronium
ions (H3O+) when combined in water.
HCl + H2O → H3O + Cl-
• Bases: substances that produce hydroxide
ions (OH-) when in water.
NaOH → Na + OH-
Definition of Acids and Bases

Bronsted-Lowry definition (Johannes
Nicolaus Brønsted and Thomas Martin
Lowry)
• Acids: proton (hydrogen ion) donor
HCl + H2O → H3O + Cl-
• Bases: proton (hydrogen ion)
acceptor
NH3 + H2O → NH4 + OH-
Definition of Acids and Bases

• A strong acid dissociates (or ionizes) completely in aqueous
solution to form more hydronium ions (H3O)
• A weak acid does not dissociate completely in aqueous solution
Strong and Weak Acids/Bases

A strong base dissociates completely in aqueous solution to form
hydroxide ions (OH-)
A weak base does not dissociate completely in aqueous solution to
form hydroxide ions (OH-)
Strong and Weak Acids/Bases

Product: Salt and Water
HCl (aq) + NaOH (aq) → NaCl + H2O
Acid-Base (Neutralization) Titration

• Equivalence point: point in titration at which the amount of titrant
added is just enough to completely neutralize the analyte
solution. At the equivalence point in an acid-base titration, moles of
base = moles of acid and the solution only contains salt and water.
Acid-Base (Neutralization) Titration

• are usually weak organic acids or bases in which the undissociated
molecule has one color, and the anion and cation produced by
dissociation has another color.
They are used to:
• To determine the end-point in neutralization process.
• To determine H+ concentration or pH
• To indicate that a desired change in pH has happened
Acid-Base Indicators

• A titration curve is a graph of the pH
versus the amount of the reagent
progressively added to the original
sample.
A titration curve can be used to
determine:
• The equivalence point of an acid-
base reaction
• The pH of the solution at equivalence
point is dependent on the strength of
the acid and strength of the base
used in the titration.
Titration curve

Point 1: No NaOH added yet, so the
pH of the analyte is low
Point 2: This is the pH recorded at a
time point just before complete
neutralization takes place.
Point 3: This is the equivalence point
(halfway up the steep curve).
Point 4: Addition of NaOH continues,
pH starts becoming basic because
HCl has been completely neutralized
Titration of a strong acid with a strong base
For strong acid-strong base titration, pH = 7 at
equivalence point

Titration of a weak acid with a strong base
For weak acid-strong base titration,
pH > 7 at equivalence point

Titration of a strong acid with a weak base
For strong acid-weak base titration, pH < 7 at
equivalence point

Titration of a weak alkali with a weak acid

• Never titrate weak acid/base with weak solutions = no sharp end point
• Appearance of color is more observable than disappearance
• Mixed indicators may be used for no sharp end-point
“Use 3 drops unless otherwise directed”
Titrant Analyte Indicator
Strong Alkali Strong Acid Methyl orange, Methyl
red, or Phenolphthalein
Strong Alkali Weak Acid Phenolphthalein
Strong Acid Weak Alkali Methyl Red
Rules for the Use of Indicators

• Standardization is the process of determining the exact
concentration (normality or molarity) of a solution. Titration is one
type of analytical procedure often used in standardization.
Commonly used standard solutions:
ACIDS: Hydrochloric Acid, Sulfuric Acid
ALKALIS: Sodium Hydroxide, Potassium Hydroxide, Barium
Hydroxide
Standardization

Acidimetric Analysis
• The direct and residual titrimetric analysis of bases using an
accurately measured volume of acid.
Alkalimetric Analysis
• Analyzing the concentration of acids using a known
concentration of alkaline solution.
Types of Acid-Base Titration

Direct Titration
• A basic titration method that involves the reaction between the
unknown compound and the compound of known concentration.
Burette is filled with titrant solution
Indicator is added to the solution to be titrated
Titrant is added to flask with swirling until the end point
Read and record the volume of titrant delivered

Residual Titration (Back Titration)
• A process in which the excess of a standard solution used to
consume an analyte is determined by titration with a second
standard solution.
• Used when the end point of a direct titration deviates appreciably
from the stoichiometric points for some reason.
Possible reasons: The sample is insoluble, the rate of its reaction
with the standard acid is relatively slow, or when the analyte to be
assayed does not give a distinct sharp end point with an indicator by
direct titration.

Analyte is combined with excess standard solution
Burette is filled with second titrant
Indicator is added to the solution to be titrated
Second titrant is added to flask with swirling until the end point
Read and record the volume of titrant delivered

Direct Titration:
% Purity =
𝑁 𝑥 𝑚𝑙 𝑥 𝑚𝐸𝑞
𝑊𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑆𝑎𝑚𝑝𝑙𝑒
x 100 mEq = (
𝑀𝑊
𝑓 𝑥 1000
)
Residual Titration:
% Purity =
𝑁 𝑥 𝑚𝑙𝑎 − 𝑁 𝑥 𝑚𝑙𝑏 𝑥 𝑚𝐸𝑞
x 100
Acidimetric Analysis: Formulas

Direct Titration
Sample/Analyte Titrant Indicator
Assay of Sodium
Bicarbonate
1 N Sulfuric Acid Methyl orange
Assay of Sodium
Hydroxide
1 N Sulfuric Acid Methyl orange or
Phenolphthalein
Assay of Sodium
Salicylate Tablets
0.1 N HCl Bromophenol blue TS

• A 3-gram sample of sodium bicarbonate (NaHCO3) required
35.1 ml of 1 N sulfuric acid in titration to a methyl orange end
point. What is the % purity (w/w) of the analyte? (Sodium = 23
g/mol; Oxygen = 16; Hydrogen = 1; Carbon = 12)
Solution:
% Purity =
𝑤𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑠𝑎𝑚𝑝𝑙𝑒
x 100
=
1 𝑁 𝑥 35.1 𝑚𝑙 𝑥 (
84 𝑔/𝑚𝑜𝑙
1 𝑥 1000
)
3 𝑔𝑟𝑎𝑚𝑠
x 100 = 98.28%
Sample Problem

• If a 0.2800-g sample of Sodium Bicarbonate (96.5% NaHCO3)
is titrated with 0.9165 N sulfuric acid, what volume of the acid
should be required to produce an end point? (Sodium = 23
g/mol; Oxygen = 16; Hydrogen = 1; Carbon = 12)
Checkpoint
% Purity =
x 100

% Purity =
x 100
96.5% =
(0.9165)(𝑋)(0.084)
0.2800 𝑔𝑟𝑎𝑚𝑠
x 100
X =
96.5% 𝑥 0.2800
(0.9165)(0.084)
= 3.51 ml
Solution

• Twenty (20) sodium salicylate (NaC7H5O3) tablets labelled 325 mg were
dispersed in sufficient water to make 200 ml. A 15.0-ml aliquot of the filtrate
was titrated to a bromophenol blue end point in the usual way by 29.11 ml of
0.100 N hydrochloric acid. Calculate the % purity (w/w) of the sample.
Checkpoint
% Purity =
x 100

325 𝑚𝑔
1 𝑡𝑎𝑏𝑙𝑒𝑡
=
𝑥
20 𝑡𝑎𝑏𝑙𝑒𝑡𝑠
x =
(325 𝑚𝑔)(20 𝑡𝑎𝑏𝑙𝑒𝑡𝑠)
1 𝑡𝑎𝑏𝑙𝑒𝑡
= 6500 mg
6500 𝑚𝑔
200 𝑚𝑙
=
𝑥
15 𝑚𝑙
x =
(6500 𝑚𝑔)(15 𝑚𝑙)
(200 𝑚𝑙)
= 487.5 mg
Solution

% Purity =
x 100
=
0.1 𝑁 (29.11 𝑚𝑙)(
1 𝑥 1000
)
0.4875 𝑔
x 100 = 95.54%
Solution

Residual Titration
• Assay of Zinc Oxide
• Assay of Potassium Sodium Tartrate
• Assay of Milk of Magnesia
• Assay of Methenamine
• Assay of Ammonium Chloride

• A 1.250-gram sample of Zinc Oxide (ZnO) required 50 ml of 1.1230
N sulfuric acid and 50.81 ml of 0.9765 N sodium hydroxide were
required in the back titration. What is the % purity (w/w) of the
analyte? (Zinc = 65.4 g/mol; Oxygen = 16)
Sample Problem

% Purity =
x 100
% Purity =
1.1230 𝑥 50 −(0.9765 𝑥 50.81) (
81.4 𝑔/𝑚𝑜𝑙
2 𝑥 1000
)
x 100
% Purity = 94.98 or 95.0%
Solution

• If a sample of milk of magnesia weighing 5.2430 g when
dissolved in 25 ml of 0.9915 N sulfuric acid required 9.85
ml of 1.1402 N sodium hydroxide to titrate the excess acid,
what is the percent of Mg(OH)2 in the sample? (Magnesium
= 24.31 g/mol; Oxygen = 16; Hydrogen = 1)
Checkpoint

Solution

• A method for the quantitative determination of nitrogen in organic and
inorganic substances that is based on neutralization reaction.
• Introduced by Johan Kjeldahl in 1883
• Applications: Analysis of food, beverages, meats, feeds, grains,
wastewater, soil, fertilizers, and many other samples.
Nitrogen Determination: Kjeldahl Method

• Purpose: Break down peptide bonds to liberate nitrogen and convert them to
ammonium (NH4) ions.
• Heating a sample at 350-380°C with concentrated sulfuric acid (H2SO4), which
decomposes ("digests") the organic sample to liberate the nitrogen
as ammonium sulfate.
• Faster digestion is achieved through addition of a catalyst (Potassium Sulfate)
Step 1: Digestion

• Purpose: Convert Ammonium (NH4) into
Ammonia (NH3)
• Addition of conc. sodium hydroxide raises the
pH of the mixture and cause the conversion of
NH4 to NH3
• Sample solution is distilled to separate
ammonia and received in trapping flask
containing absorbing solution (HCl or H2SO4)
Step 2: Distillation

• When using sulfuric acid standard solution as absorbing
solution, the residual sulfuric acid is titrated with sodium
hydroxide standard solution and by difference the amount of
ammonia is calculated. The end-point is detected using methyl
red as indicator.
Step 3: Titration

Nitrogen Determination: Kjeldahl Method
% N =
x 100

A 6.1000-gram sample of beef extract was distilled into 50.0 ml of
0.1246 N sulfuric acid, and the mixture was titrated with 22.42 ml of
0.0962 N sodium hydroxide. Calculate the percentage of nitrogen
present in the sample.
Sample Problem
% N =
x 100
% N =
0.1246 𝑥 50 −(0.0962 𝑥 22.42) (
1 𝑥 1000
)
x 100
% N = 0.93%

Sample Problem
A 0.6325 g sample of a wheat flour was analyzed by the
Kjeldahl method. The ammonia formed by addition of
concentrated base after digestion with H2SO4 was distilled
into 25.00 mL of 0.02486 M H2SO4. The excess HCl was
then back titrated with 3.97 mL of 0.04012 M NaOH.
Calculate the percent protein in the flour.

Sample Problem

Direct Titration
• Assay of Hydrochloric Acid
• Assay of Diluted Phosphoric Acid
• Assay of Boric Acid
• Assay of Tartaric Acid
Residual Titration
• Assay of Aspirin Capsules
Alkalimetric Analysis

•Theory is the same as acid-base titration
•Reaction is carried out in non-aqueous medium
Non-aqueous Titrimetric Analysis

Zn + 2H+ → Zn2+ + H2
Zn is being
oxidized
H is being
reduced
• An INCREASE in oxidation number of an atom signifies
OXIDATION (Example: from 0 to +2)
• A DECREASE in oxidation number of an atom signifies
REDUCTION (Example: From +1 to 0)
Identifying Oxidized and Reduced Elements

Step 2: Balance each half equations atomically in this order:
a. Atoms other than H and O
Cu → Cu2+
NO3
- → NO2
Here in the example, Copper (Cu) and Nitrogen (N) are already
balanced in both half equations.
Balancing Redox Reactions (in Acidic Solution)

Balance the following equation in an acidic solution:
SO3
-2 + MnO4
- → SO4
-2 + Mn2+
Balance the following equation in a basic solution:
CN- + MnO4
- → CNO- + MnO2
Checkpoint: Balancing Redox Reactions

Oxidation-Reduction
Reactions

• Also called as Redox reactions
• A reaction that involves the transfer of electrons between chemical
species (the atoms, ions, or molecules involved in the reaction).
• Redox reactions are all around us: the burning of fuels, the
corrosion of metals, and even the processes of photosynthesis and
cellular respiration involve oxidation and reduction.
Oxidation-Reduction Reactions

•Oxidation: LOSS of electrons
•Reduction: GAIN of electrons
Undergoes OXIDATION; the
atom is OXIDIZED
Undergoes REDUCTION; the
atom is REDUCED
+ 3
+ 2
+ 1
0
-1
-2
-3
Adding
electrons (e-)
reduces the
charge

• Oxidation: LOSS of electrons in a reducing agent
• Reduction: GAIN of electrons in an oxidizing agent
Undergoes OXIDATION;
Reducing agent
Undergoes REDUCTION;
Oxidizing agent
LEORA = Loss of Electrons is Oxidation of a Reducing Agent
GEROA = Gain of Electrons is Reduction of an Oxidizing Agent

+ →
e-
→
Na is being
oxidized
Cl is being
reduced
Na → Na+ + e- Cl + e- → Cl-

Rules for assigning oxidation number:
#1: The oxidation number (ON) of an uncombined element
is always zero.
Example:
Na Cu N2
0 0 0
Assigning Oxidation Numbers

#2: A monatomic ion has an oxidation number
equal to its charge.
Example:
Na+ Cu2+ N3-
+1 +2 -3

#3: The oxidation numbers of hydrogen is +1 with
non-metals and -1 with metals.
Examples:
HCl NaH
+1 -1

#4: The oxidation number of oxygen is usually -2, except in
peroxides (X2O2 or compounds containing O2
2-) where
oxygen has an oxidation number of -1.
Example:
H2O H2O2
-2 -1

#5:
Group IA (Alkali Metals): always +1
Group 2A (Alkaline Earth Metals): always +2
Example:
CaO LiCl
+2 -2 +1

#6: Fluorine has an oxidation number of -1 in all
compounds; other halogens (Cl, Br, I) usually have ON
of -1 unless combined with oxygen or fluorine where it is
a positive number.
Example:
LiCl BaF2
-1
+1 +2 -1

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