1. Pharmaceutical Analytical Chemistry

2. Course Topics Acid-Base Titration
Precipitation and Complex-formation Titration
Oxidation-reduction Titration
Electrochemical methods
Ultraviolet/visible spectrophotometry
Introduction to chromatographic separation

3. 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)
The handouts are for guidance and studying must be from textbooks

4. 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 .

5. 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?

6. 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 mg.
3- Macro methods
used for determination of quantities more than 100 mg.

7. 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

8. 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).

9. 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.

10. 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).

11. 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).

12. 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)

13. 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.

14. 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 gm/L of solution.
2M solution of H2SO4 contains 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.

15. 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 cation or anion 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.

16. 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 .
نهاية محاضرة الاثنين20/04/1436

17. 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.

18. 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 non-conducting solutions.
Examples: Sugar , Glycerin , Ethyl acetate.

19. 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.

20. 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 = NH Cl-
The degree of dissociation characterizes the chemical activity of the respective substance.

21. 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

22. 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 .

23. 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

24. 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

25. Theories of Acids and Bases
1-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 OH- N.B. 1- Not all acid-base reactions involve water 2- Many bases (NH3, and carbonate) do not contain any OH-

26. 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
NH 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

27. 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

28. 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- ] = Since the dissociation of water gives rise to equal number of H+ and OH- Kw = [H+ ]2 =10-14 [H+ ] =√ = 10-7
نهاية محاضرة الاثنين 27 /4

29. 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 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 = pKw = 14 -In general : for acids pH = -log [H+] for bases pOH = -log [OH-] pKw = pH + pOH pH = pKw - pOH = 14 - pOH

30. 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 = 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 = pOH =14- 1= 13

31. 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)

32. 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)
= ½ (-log 1.76 x 10-4)+ ½ (-log 0.25)
= (0.5 x 3.754)+ (0.5 x 0.602) =
= 2.18

33. 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 )

34. Buffer Solutions 1-Acidic 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

35. 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 Cl-

36. 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

37. 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 log 1/1 = = 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 log 0.5/0.3 = 14 – = 8.967

38. 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 = = 1.1 Cs = = 0.9 pH = pKa + log Cs/Ca = -log 1.75 x log 0.9/1.1 = ( 0.087) = 4.67

39. 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

40. 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

41. 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 red to orange or yellow, Phenolphthalein(Ph.Ph.) pH range colorless to pink , and Methyl Red (M.R.) pH range 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
An indicator is a substance which is used to determine the end point in a titration. In acid-base titrations, organic substances (weak acids or weak bases) are generally used as indicators.

42. 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.
Considering two important indicators phenolphthalein (a weak acid) and methyl orange (a weak base), Ostwald theory can be illustrated as follows: Phenolphthalein: It can be represented as HPh. It ionises in solution to a small extent as: HPh ↔ H+ + Ph
Colourless Pink

43. 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 , M.R Ph.Ph
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

44. 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

45. 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. pH = pKw - pCb

46. Strong Acid Vs Strong Base
Both M.O. and Ph.Ph. are suitable

47. 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

48. Weak Acid Vs Strong Base
The suitable indicator is Ph.Ph. Not M.O.

49. 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.

50. This titration curve shown in the figure involves 1
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.

51. 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 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.

52. 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.

53. 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.
1- a protic solvent is a solvent that has a hydrogen atom bound to an oxygen (as in a hydroxyl group) or a nitrogen (as in an amine group). In general terms, any solvent that contains labile H+ is called a protic solvent. The molecules of such solvents readily donate protons (H+) to reagents. Conversely, aprotic solvents cannot donate hydrogen.