Pharmaceutical Analytical Chemistry

Dr. M Afroz Bakht

Course Topics

1. Acid-Base Titration 6 hours

2. Precipitation and Complex-formation Titration 4 hours

3. Oxidation-reduction Titration4 hours

4. Electrochemical methods4 hours

5. Ultraviolet/visible spectrophotometry6 hours

6. Introduction to chromatographic separation 4 hours

ObjectivesThe Student will be able to : 1. Check the feasibility of titrimetric

reactions. The student shall be also able to choose the suitable indicator and derive the titration curve.

2. Calculate the percentage purity of solid (powdered) samples and the concentration of liquid samples. 

3. Obtain the absorption spectrum of the light-absorbing compound and define its max. 

4. Know the principles of chromatographic separation of pharmaceutical compounds in a mixture.

5. Check the feasibility of electrochemical methods of analysis.

Course Evaluation Continuous Assessment:

• First Assessment Test15%

• Second Assessment Test10%

• Term Activity* 10%• Laboratory Test 10%• Final Laboratory Test 15%

Final Examination:

Final Paper test Final Exam40%

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).

Principles 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)

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 .

Areas of Chemical Analysis and questions they answer

IdentificationWhat is the identity of the substance in the sample? QuantitationHow much of the substance x is in the sample? DetectionDoes the sample contain substance X or not? SeparationHow the species of interest can be separated?

Quantitative Chemical AnalysisClassification 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 methodsused for determination of quantities more than 100

mg.

Quantitative Chemical Analysisb-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

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).

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.

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).

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).

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)

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.

Molar Standard Solutions (M) -ExamplesMolar standard solution (M)

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 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 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.

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 OH-

Example Eq.Wt of NaOH = m.wt / 1 Eq.Wt of Ba(OH)2 = m.wt / 2

Eq. Wt For Salts = m. wt/( number of metal x its charge )

Examples NaCl eq. wt = m.wt / 1 CaCl2 eq. wt = m.wt / 2

-N.B. Equal volumes of equal normalities contain equal number of molecules, that means equal normalities react 1 to 1 ratio.

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 .

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.

Electrolytes and Non-electrolytes 2-Weak 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 non-conducting solutions.

Examples: Sugar , Glycerin , Ethyl acetate.

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.

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.

Molecular and Ionic Equations

Molecular equations represent the reaction species (reactant and products ) as molecules .

NaOH + HCl → NaCl + H2O 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+ → H2O 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

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 .

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 = C + D Vf α [A] [B] or Vf = K1 [A] [B]

Vb α [C] [D] or 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

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 right ) . solute + solvent = solution + heat

Theories of Acids and Bases1-Arrhenius theory: -An acid forms H+ in water (upon ionization) HCl → H+ + Cl- HNO3 → H+ + NO3

-

H2SO4 → 2H+ + SO4--

- A base forms OH- in water (upon ionization) Na OH → Na+ + OH-

Ca(OH)2 → Ca+ + + 2OH-

N.B. 1- Not all acid-base reactions involve water2- Many bases (NH3, and carbonate) do not contain

any OH-

Theories of Acids and Bases2-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

Theories of Acids and Bases3-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

Dissociation of Water

H2O = 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 OH-

Kw = [H+ ]2 =10-14 [H+ ] =√ 10-14 = 10-7

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

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

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)

pH of acids and Bases

ExamplesCalculate the pH of 0.1 M solution of acetic acid (Ka =1.75x10-

5) pH = ½ pKa + ½ pCa

= ½ (-log 0.1) + ½ (-log 1.75 x

10-5) = (0.5 x 1) + (0.5 x 4.757) = 2.88

Calculate the pH of 0.25 M solution of formic acid (Ka =1.76x10-4)

pH = ½ pKa + ½ pCa

= ½ (-log 0.25) + ½ (-log 1.76 x 10-4)

= (0.5 x 0.602) + (0.5 x 3.754) = 0.301 + 1.877 = 2.18

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 )

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 solutions1-Acidic buffer solutionsConsists 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

Buffer Solutions

2- Basic Buffer solutionsConsists 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 +

Cl-

Henderson Equation for calculation of pH of buffer solutions

pH of acidic buffer: pH = pKa + log Cs/Ca

pH of basic buffer: pOH = pKb + log Cs / Cb

pH = pKw - pOH pH = pKw - pKb - log

Cs/Cb

OR pH = pKw - pKb + log Cb/Cs

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 – 4.745 + 0.222 = 9.477

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

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

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

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 8.3-10 colorless to pink , and Methyl Red (M.R.) pH range 4.4-6.3 red to yellow

2 - Turbidity indicatorsPrecipitation or turbidity appears at the end point e.g. Isonitrosoacetyl-p-aminobenzene 3 – Fluorescence indicatorsCertain 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

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.

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 indicatorsSharper 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

Color Indicators

Screened indicatorsWhen 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

Neutralization Titration CurvesTitration curve is the plot of pH versus the

volume of titrant Titration curves are constructed to - study the feasibility of titration - choosing indicator1- 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 = pCa3- 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. pH = pKw - pCb

Strong Acid Vs Strong Base

Both M.O. and Ph.Ph. are suitable

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

Weak Acid Vs Strong Base

The suitable indicator is Ph.Ph. Not M.O.

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.

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.

Neutralization reactions in Non-Aqueous 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 Titration1– 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.

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.

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.