Pharmaceutical Analytical
Chemistry
1
Course Topics
1.
2.
3.
4.
5.
6.
Acid-Base Titration
Precipitation and Complex-formation Titration
Oxidation-reduction Titration
Electrochemical methods
Ultraviolet/visible spectrophotometry
Introduction to chromatographic separation
2
References
• Fundamentals of Analytical Chemistry, Douglas A. Skoog and
Donald M. West. Fourth edition. Sanders College Publishing,
Philadelphia (1984).
• Analytical Chemistry, Douglas A. Skoog; Donald M. West, F.
James Holter, Standey R. Crouch, 7th ed. Harcourt College
Publishers (2000).
• Principales of Quantitative Chemical Analysis, Robert de
Levie. McGraw Hill, New York (1997).
• Vogel’s Textbook of Quantitative Inorganic Analysis, 4th ed. J.
Baisett, R.C. Denney, G.H. Jefferg and J. Mendham,
Longman, Essex (1978)
3
Pharmaceutical Analytical Chemistry
• Analytical chemistry deals with methods used
for determining the composition of various
materials.
• The process of material identification called
Qualitative Analysis .
• The process of material quantitation called
Quantitative Analysis .
4
Areas of Chemical Analysis and
questions they answer
• Identification
What is the identity of the substance in the sample?
• Quantitation
How much of the substance x is in the sample?
• Detection
Does the sample contain substance X or not?
• Separation
How the species of interest can be separated?
5
Quantitative Chemical Analysis
• Classification of Quantitative methods:
a-According to the quantity to be analyzed
• 1- Micro methods
used for the determination of quantities less than 1 mg.
• 2- Semi-micro methods
used for determination of quantities ranging from 1-100 mg.
• 3- Macro methods
used for determination of quantities more than 100 mg.
6
Quantitative Chemical Analysis
b-According to technique
• I- Volumetric or Titrimetric methods
Analysis by volume.
• II- Gravimetric methods
Analysis by weight.
• III- Instrumental methods (Physicochemical
methods)
•
Electrochemical methods
•
Spectroscopic methods
•
Separation methods
7
Volumetric Analysis
• It is the quantitative chemical analysis carried
out by determining the volume of a solution of
accurately known concentration which is
required to react quantitatively with measured
volume of solution of the substance to be
analyzed.
• The solution of accurately known concentration
is called the standard solution (titrant).
8
Volumetric Analysis
• The process of adding standard solution
gradually to the sample until the reaction is
just completed is termed as titration.
• The point at which the reaction is completed
is called end point or equivalence point.
• The concentration of the substance to be
analyzed is calculated from the volume of the
standard solution.
9
Detection of End Point
1- Physical change produced by the standard
solution itself (Self indicator).
2-The Addition of a substance known as
indicator.
(Compound which has different colors at
different conditions).
10
Requirements for Quantitative
Titrimetric Analysis
• The reaction between the sample and the
standard solution must be simple and can be
represented by a chemical equation.
• The reaction must be instantaneous (relatively
fast or rapid). Sometimes catalyst is needed.
• The substance to be determined should react
completely with the titrant in stoichiometric
manner (definite ratio).
• The end point of the reaction can be detected
easily. (indicator is available).
11
Reactions Used in Titrimetric Analysis
• I- Neutralization Reactions
(Acid-Base Reactions)
• II-The Precipitation Reactions
(Precipitimetry)
• III- Complex Formation Reactions
(Complexometry)
• IV-Electron-transfer Reactions
(Redoximetry)
12
Standard Solutions
• These are solutions of exact known
concentration
• Types of standard solutions
1-Molar standard solution (M)
• It is the solution which contains the gram molecular
weight of the substance in 1L of solution.
1M solution contains 1 x gm m.wt of substance/L of
solution.
2M solution contains 2 x gm m.wt of substance/L of
solution.
M/10 solution contains 0.1 x gm m.wt of substance/L
of solution.
13
Molar Standard Solutions (M)
• -Examples
Molar standard solution (M)
1M solution of NaOH contains 40 gm/L of solution.
2M solution of NaOH contains 80 gm/L of solution.
M/10 solution of NaOH contains 4 gm/L of solution.
1M solution of H2SO4 contains 98.07 gm/L of solution.
2M solution of H2SO4 contains 196.14 gm/L of solution.
M/10 solution of H2SO4 contains 9.8 gm/L of solution.
1M solution of Na2CO3 contains 106 gm/L of solution.
2M solution of Na2CO3 contains 212gm/L of solution.
M/10 solution of Na2CO3 contains 10.6 gm/L of solution.
14
Normal Standard Solutions (N)
• Solution which contains gm equivalent weight /L of solution.
Equivalent Weight
• Eq.Wt of acids = m.wt / no. of replaceable H+
Example
Eq.Wt of HCl = m.wt / 1
Eq.Wt of H2SO4 = m.wt / 2
• Eq.Wt of bases = m.wt / no. of replaceable OHExample
Eq.Wt of NaOH = m.wt / 1
Eq.Wt of Ba(OH)2 = m.wt / 2
• Eq. Wt For Salts = m. wt/( number of cation or anion x its charge )
Examples
-N.B.
NaCl eq. wt = m.wt / 1
CaCl2 eq. wt = m.wt / 2
Equal volumes of equal normalities contain equal number of
molecules, that means equal normalities react 1 to 1 ratio.
15
Neutralization Reactions
• Acid-Base Titrations In Aqueous Solution
• Solutions
• Solution is a homogenous mixture of two or more substances.
The component (solid, gas or liquid ) present in small quantity is called
the solute, while the one present in large quantities is called the solvent .
- Solutions may be
•
•
•
1- Saturated solutions .
2- Unsaturated solutions .
3- Supersaturated solutions .
16
Electrolytes and Non-electrolytes
• I-Electrolytes:
• Electrolytes are substances when dissolved in water undergo
dissociation and give electricity-conducting solutions.
- Electrolytes may be :
• 1-Strong Electrolytes :
• Substances when dissolved in water dissociate or ionize to a
high degree.
• Examples of strong electrolytes .
•
Acid: HCl , HNO3 , H2SO4 , HBr , HI .
•
Base: NaOH , KOH , Ca(OH)2 , Ba(OH)2.
•
Salt: NaCl , CH3COONa , NH4Cl.
17
Electrolytes and Non-electrolytes
• 2-Week Electrolytes:
Substances when dissolved in water dissociate or
ionize to a slight degree. Examples of week
electrolytes .
•
- Acid: CH3COOH , HCN , H2S , H3BO3 , HF .
•
- Base: NH4OH , N2H4 .
•
- Salt: HgCl2 , CdCl2 , HgBr2, CH3COONH4 .
• II-Non-electrolytes:
Non-electrolytes are substances when dissolved in
water do not undergo dissociation and give a nonconducting solutions.
•
Examples: Sugar , Glycerin , Ethyl acetate.
18
Electrolytic Dissociation Theory
• Pure water is a bad conductor for electricity.
• When an electrolyte is dissolved in water, it dissociates into
negatively charged ions (anions) and positively charged ions
(cations).
• Solutions conduct the electric current due to the presence of
ions.
• The degree of dissociation is directly proportional to the
degree of dilution.
19
Degree of Dissociation (α)
• It is the ratio of the ionized fractions to the total
amount of the dissolved solute.
• For each concentration there is a state of equilibrium
between the un-dissociated molecules and the
dissociated molecules (ions).
•
Molecule =
Cation (+ve) +
Anion (-ve)
•
CH3COOH =
H+
+
CH3COO•
NH4Cl
=
NH4+
+
Cl• The degree of dissociation characterizes the chemical
activity of the respective substance.
20
Molecular and Ionic Equations
• Molecular equations represent the reaction species (reactant
and products ) as molecules .
• NaOH
+ HCl →
NaCl
+
H 2O
• This equation shows that one mole of NaOH neutralize one mole of HCl
to form exactly one mole of NaCl and one mole of H2O .
• In ionic equations, strong electrolytes are represented as ions
•
•
•
•
•
•
while weak electrolytes represented as molecules .
In the above equation NaOH and HCl are strong electrolytes and
present as ions in the solution , so that , the equation can be written as
follows:
Na+ + OH- + H+ + Cl- → Na+ + Cl- + H2O
OH- + H+
→
H 2O
In the reaction of NaOH (strong electrolyte ) and CH3COOH (week
electrolyte ) ,the equation is written as follows:
Na+ + OH- + CH3COOH → Na+ + CH3COO- + H2O
OH- + CH3COOH →
CH3COO- + H2O
21
Chemical Equilibrium
• In reversible reactions products are formed from
the reactants and the reactants are being
produced from the products.
•
A + B =
C + D
•
Reactants ↔ products
• Under that condition the composition of the
reaction mixture becomes constant and the
system is said to be in a state of equilibrium
which is the state at which the rate of forward
reaction equal to the rate of backward reaction .
22
Law of Mass Action
• The rate of chemical reaction is directly proportional to the
product of the molar concentration of the reacting
substances.
•
•
•
•
•
•
•
•
•
For the reaction
A + B
Vf α [A] [B]
or
Vb α [C] [D]
or
=
C + D
Vf = K1 [A] [B]
Vb = K2 [C] [D]
- At equilibrium Vf = Vb
K1 [A] [B] = K2 [C] [D]
- K1 / K2 = k equilibrium (equilibrium constant)
Keq = [C] [D] / [A] [B]
- In case of :
aA + bB =
cC + dD
Keq = [C]c [D]d / [A]a [B]b
23
Displacement of Equilibrium
• Le Chatelier Principle:
• According to Le Chatelier principle, if a stress is applied to a system
in an equilibrium state , the equilibrium will be shifted in such
direction to minimize that stress.
• Applications of Le Chatelier Principle:
• In Precipitation :
•
A + B =
AB (precipitate)
• Precipitating agent B is used to precipitate the compound AB by
combining with A to form more AB and the equilibrium is shifted to
the right .
• In Solubility :
• In endothermic solution , the solubility of the solute increases by
heating (equilibrium shifted to right ) .
•
solute + solvent + heat =
solution
• In exothermic solution , the solubility of the solute decreases by
heating (equilibrium shifted to left ) .
•
solute + solvent
=
solution + heat
24
Theories of Acids and Bases
1-Arrhenius theory:
-An acid forms H+ in water (upon ionization)
HCl
→
H+
+
ClHNO3
→
H+
+
NO3H2SO4
→ 2H+
+
SO4--
- A base forms OH- in water (upon ionization)
Na OH
→
Na+
+
OHCa(OH)2 →
Ca+ +
+ 2OHN.B.
1- Not all acid-base reactions involve water
2- Many bases (NH3, and carbonate) do not contain any OH25
Theories of Acids and Bases
2-Bronsted - Lowry theory:
-Acid is a proton donor H+
Acid →
H+
+ Conjugate base
HCl
→
H+
+
Cl- Base is a proton acceptor H+
Base + H+ →
Conjugate acid
NH3+
+ H+ →
NH4+
The conjugate base of an acid is the acid minus the proton it
has donated
The conjugate acid of a base is the base plus the accepted
proton
26
Theories of Acids and Bases
3-Lewis theory:
-Base is a substance containing an atom that has
unshared pair of electrons e.g. N , O , P , S (base is
an electron donor e.g. NH3, amines like
triethylamine).
-Acid is a substance that can accept that pair of
electrons e.g. AlCl3 , BCl3 , BF3
example of Lewis acid Lewis base reaction:
H3N:
+
BF3 →
H3N: → BF3
Lewis Base + Lewis acid
27
Dissociation of Water
H2 O
=
H+
+
OH-Water molecules ionize in very slight degree.
-According to the law of mass action
K = [H+ ] [OH- ] / [H2O ]
K [H2O ] = [H+ ] [OH- ]
Kw = [H+ ] [OH- ]
Kw = ionic product of water
-It was found that under normal experimental conditions
and at 250c
Kw = [H+ ] [OH- ] = 10-14
-Since the dissociation of water gives rise to equal number
of H+ and OHKw = [H+ ]2 =10-14
[H+ ] =√ 10-14
= 10-7
28
Hydrogen Ion Exponent
pH is the measure of acidity or alkalinity solution .
-pH is the negative logarithm of the hydrogen ion concentration
pH =-log [H+]
pH range
0 1 2 3 4 5 6
7
8 9 10 11 12 13 14
Acidic
Neutral
Basic
-pH is a number obtained by giving a positive value to the negative
power of 10 in the expression .
[H+] = 10-n
pH =
n
[H+] = 10-5
pH =
5
Kw = 10-14
pKw =
14
-In general : for acids
pH = -log [H+]
for bases
pOH = -log [OH-]
pKw = pH + pOH
pH = pKw - pOH
= 14 - pOH
29
pH of acids and Bases
pH of strong acids and strong bases
-Strong acid and strong base are completely ionized so,
concentration of acid or base represents the
concentration of [H+] or [OH-] .
For acids
pH =-log [H+]
For bases pOH = -log [OH-]
pH = 14 - pOH
For examples:
pH of 0.1 M HCl (strong acid)
pH =-log [H+] = -log 10-1 = 1
pH of 0.1 M NaOH (strong base)
pOH = -log [OH-] = -log 10-1 = 1
pH = 14 - pOH =14- 1= 13
30
pH of acids and Bases
- pH of weak acids
-A Small quantity of weak acid is dissociated with the
formation of [H+] . e.g. CH3COOH
CH3COOH
= CH3COO- + H+
Ka = [CH3COO- ] [H+] / [CH3COOH]
Where:
[H+] = [CH3COO- ]
[CH3COOH]= Ca (concentration of acid)
Ka = [H+] 2 / Ca
[H+] 2 = Ka Ca
[H+] = √Ka Ca
pH = ½ pKa + ½ pCa
pH = ½ (pKa + pCa)
31
pH of acids and Bases
Examples
Calculate the pH of 0.1 M solution of acetic acid (Ka =1.75x10-5)
pH = ½ pKa + ½ pCa
= ½ (-log 1.75 x 10-5)+ ½ (-log 0.1)
= (0.5 x 4.757)+ (0.5 x 1)
= 2.88
Calculate the pH of 0.25 M solution of formic acid (Ka =1.76x10-4)
pH = ½ pKa + ½ pCa
= ½ (-log 1.76 x 10-4)+ ½ (-log 0.25)
= (0.5 x 3.754)+ (0.5 x 0.602)
= 1.877+ 0.301
= 2.18
32
pH of acids and Bases
pH of salts
-Salts of strong acids and strong bases e.g. NaCl is
neutral pH = 7
-Salts of strong acids and weak bases e.g. NH4Cl
pH = ½ (pKw - pKb + pCs )
-Salts of weak acids and strong bases e.g. CH3COONa
pH = ½ (pKw + pKa - pCs )
-Salts of weak acids and weak bases e.g. CH3COONH4
pH = ½ (pKw + pKa - pKb )
33
Buffer Solutions
Buffer solutions are solutions which resist the change in the pH of
solution upon addition of small amount of strong acid or strong base
- Types of buffer solutions
1-Acidic buffer solutions
Consists of weak acid and its salt of strong electrolyte.
e.g. acetic acid and sodium acetate (CH3COOH/CH3COONa)
-Upon addition of a strong acid:
sodium acetate react with it giving weakly ionized acetic acid
H+ + CH3COONa → CH3COOH + Na+
-Upon addition of a strong base:
acetic acid react with it and unionized water is formed
OH- + CH3COOH → CH3COO- + H2O
34
Buffer Solutions
2- Basic Buffer solutions
Consists of weak base and its salt of strong electrolyte.
e.g. ammonium hydroxide and ammonium chloride
(NH4OH/NH4Cl)
-Upon addition of a strong acid:
H+ + NH4OH → NH4+ + H2O
-Upon addition of a strong base:
OH- + NH4Cl → NH4OH
+
Cl35
Henderson Equation for calculation of
pH of buffer solutions
pH of acidic buffer:
pH = pKa + log Cs/Ca
pH of basic buffer:
pOH =
pH =
pH =
OR pH =
pKb
pKw
pKw
pKw
+
-
log Cs / Cb
pOH
pKb - log Cs/Cb
pKb + log Cb/Cs
36
Examples
Calculate the pH of a buffer solution consisting of 1 M
CH3COOH and 1 M CH3COONa where Ka=1.75x1 0-5
pH = pKa + log Cs/Ca
= -log 1.75 x 1 0-5 + log 1/1 = 4.76 + 0 = 4.76
Calculate the pH of a buffer solution consisting of 0.5 M
NH4OH and 0.3 M NH4Cl where Kb = 1.8 x 10 5
pH = pKw – pKb + log Cb/Cs
= 14 – log 1.8 x 1 0 5 + log 0.5/0.3
= 14 – 5.255 + 0.222 = 8.967
37
Examples
Calculate the pH of a buffer solution consists of 1 M CH3COOH and 1 M
CH3COONa after addition of 0.1 mol of HCl to one L of solution
where Ka=1.75x10-5
after addition of 0.1 mol HCl , it will react with an equivalent amount
of CH3COONa forming the same amount of CH3COOH
HCl + CH3COONa → CH3COOH + NaCl
0.1 mol
0.1 mol
0.1 mol
0.1 mol
Ca = 1+ 0.1 = 1.1
Cs = 1 - 0.1 = 0.9
pH = pKa + log Cs/Ca
= -log 1.75 x10-5 + log 0.9/1.1 = 4.757 + ( 0.087) = 4.67
38
Buffer Capacity
It is a magnitude of the resistance of a buffer to change
in the pH
B = ΔB / Δ pH
- B is a buffer capacity
- ΔB is a strong acid or base added
- Δ pH is the change in pH
Buffer capacity is directly proportional to concentration
of buffer components
Solution has equal concentration of acid or (base) and
its salt appears to have the maximum buffer capacity
- Buffer solution with high B is of high efficiency
39
Neutralization Indicators
- Neutralization indicators are weak acids or weak bases
which change their color according to the pH of the
solution
• The acid form (HA) of the indicator has one color, the
conjugate base (A–) has a different color.
• In an acidic solution, [H+] is high. Because H+ is a common
ion, it suppresses the ionization of the indicator acid, and
we see the color of HA.
• In a basic solution, [OH–] is high, and it reacts with HA,
forming the color of A–.
• The change of color is not sudden but takes
place within small interval of pH(2 pH units or
less)
• It is preferred to select an indicator which
exhibits color change at pH close to that of
salt formed at the end point
40
Types of Neutralization Indicators
1 – Color indicators
Organic dyes that exhibit different colors at different pH
values
e.g. Methyl Orange (M.O.) pH range 3.3-4.4 red to
orange or yellow, Phenolphthalein(Ph.Ph.) pH range 8.310 colorless to pink , and Methyl Red (M.R.) pH range
4.4-6.3 red to yellow
2 - Turbidity indicators
Precipitation or turbidity appears at the end point
e.g. Isonitrosoacetyl-p-aminobenzene
3 – Fluorescence indicators
Certain compounds emit visible radiations when exposed to
ultraviolet light stop or intensify when certain pH is reached
and used to detect end point when color or turbid
solutions are titrated e.g. Umbelliferone
41
Theories of Color indicators
1 – Ostwald Theory:
- Neutralization indicators are either weak acids or weak bases
- The color of ionized form differs from that of non-ionized
form
- In acidic medium basic indicators ionized and changed in
color e.g. M.O.
- In basic medium acidic indicators ionized and changed in
color e.g. Ph.Ph.
2 – Chromophore theory
- Indicators are Organic dyes which contain an unsaturated
group called chromophore group e.g. C=C , N=N , C=N , NO,
NO2 which is responsible for the color change.
- Accumulation of unsaturated groups leads to color
development
- Presence of auxochromes (-OH, -NH2 ) influence the color. 42
Effective Range of Color Indicators
It is the pH units over which the indicator changes its color.
The color change within the effective range is gradual.
Effective range for a good indicator shouldn’t exceed 2 pH
units. Example: M.O. 3.3-4.4, M.R. 4.4-6.3 Ph.Ph. 8.3-10
• Mixed indicators
Sharper color produced by using mixture of two indicators
have the similar pH range but contrasting color.
Example: mixture of thymol blue with cresol red has:
Violet color at pH 8.4
Blue color at pH 8.3
Rose color at pH 8.2
43
Color Indicators
Screened indicators
When the color change isn’t easily detectable particularly in artificial
light, addition of another indicator obtain Sharper and more
pronounced color change
Example screened mixture of M.O. and indogocarmine has:
At pH 4 yellowish green (alkaline) and violet (acidic).
Universal or multi-range indicators
The pH range can be extended By suitable mixing certain indicators.
Example: mixture of Bromothymol blue with Ph.Ph.
has:
Red color at pH 2
Orange color at pH 4
Yellow color at pH 6
Green color at pH 8
Blue color at pH 10
44
Neutralization Titration Curves
Titration curve is the plot of pH versus the volume of titrant
Titration curves are constructed to
- study the feasibility of titration
- choosing indicator
1- Titration curve of strong acid Vs strong base
- e.g. HCl against NaOH .
1- At beginning , pH of acid.
pH = - log [H+]
2- During titration, pH of strong acid.
pH = pCa
3- At the end point, pH of salt of strong acid and strong base
(neutral).
pH = pOH = ½ pKw
4- After end point , pH of strong base.
45
pH = pKw - pCb
Strong Acid Vs Strong Base
Both M.O. and Ph.Ph. are suitable
46
Weak Acid Vs Strong Base
- e.g. CH3COOH against NaOH.
1- At beginning , pH of weak acid.
pH = ½ pKa + ½ pCa
2- During titration, PH of acidic buffer.
pH = pKa + log Cs/Ca
3- At end point , PH of salt of weak acid and strong base.
pH = ½ pKw + ½ pKa - ½ pCs
4- After end point , PH of strong base .
pH = 14 -pCb
So that M.O. Isn’t suitable
The suitable indicator Ph.Ph. pH range 8.3-10
47
Weak Acid Vs Strong Base
The suitable indicator is Ph.Ph. Not M.O.
48
Weak Base Vs Strong Acid
- e.g. NH4OH against HCl.
1- At beginning , pH of weak base.
pH = pKw - ½ pKb – ½ pCb
2- During titration, pH of basic buffer.
pH = pKw - pKb + log Cb /Cs
3- At end point , pH of salt of strong acid and weak base.
pH = ½ pKw - ½ pKb + ½ PCs
4- After end point , PH of strong acid.
pH = pCa
So that M.O. or M. R. are used and Ph.Ph. Isn’t useful
N.B. Titration curve of weak acid against weak base and weak base
against weak acid.
Titration curves in both cases are smooth and change of pH at end
point is very small . So such titrations must be avoided.
49
This titration curve shown in the figure involves 1.0 M solutions of an
acid and a base. Identify the type of titration it represents.
50
Neutralization reactions in NonAqueous Medium
It means in a medium free of water and mainly used for determinations
of weak acids and weak bases
Solvent Properties and Role of Solvent in Non-Aqueos Titration
1– Relative acidity and basicity :
-According to Bronsted, acidity and basicity of substance are relative
to the solvent e.g. potassium acid phthalate when dissolved in water
acts as an acid while in glacial acetic acid acts as base.
-Similarly solvent behaves as an acid when the dissolved substance is
more basic e.g. acetic acid + pyridine and behaves as a base when
substance is more acidic e.g. acetic acid and perchloric acid.
2– Leveling effect:
-It is the ability of solvent to increase the strength of weak acids or
weak bases to reach that of strong acid or base .
-Acidic solvents have leveling effect on bases and basic solvents have
leveling effect on weak acids.
Example : acetic acid on amines and liquid ammonia on acetic acid.
51
Solvent Properties and Role of Solvent
in Non-Aqueous Titrations
3– Differentiating effect:
-It is the ability of solvent to differentiate
between the strength of acids or bases e.g.
glacial acetic and mixture of HNO3, HCl, HClO4.
4– Autoprotolysis effect:
-It is self dissociation of solvent
HA + HA
↔ A- +H2A+
-Two molecules of solvent interact ,one as
proton donor and one as proton acceptor.
52
Solvents used in Non-Aqueous
Titrations
1– Aprotic Solvents:
-These are neutral, inert, can't donate or accept protons e.g.
hexane, benzene , nitrobenzene, chloroform .
2– Amphiprotic Solvents:
-They act as acids or bases (may donate or accept protons):
a-Neutral solvents:
They have tendency to accept or donate proton e.g.
methanol, ethanol.
b-Protogenic solvents:
They are more acidic than water and have tendency to give
proton than accept proton e.g. acetic acid.
c-Protophillic solvents:
They are more basic than water and have higher tendency
to accept proton than to give proton e.g. ammonia.
53