Pharmaceutical Analysis (P. Chemistry | Ch-2): Latest D Pharma 1st year Notes

What is Pharmaceutical Analysis?

Pharmaceutical Analysis is a branch of medicinal chemistry that includes a series of experiments and procedures for the purification, identification, determination, quantitation, and characterisation of any chemical compound or component.

In simple terms, it is the science of testing drugs — confirming their identity, measuring their concentration, and ensuring they meet the required standards of purity, quality, and safety.

Pharmaceutical analysis and quality control together play a vital role in ensuring drug stability, compatibility, potency, efficacy, and safety in pharmaceutical manufacturing.

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Methods of Pharmaceutical Analysis

• Chemical / Titrimetric methods
• Biological methods
• Physical and physicochemical methods
• Pharmaceutical methods

Common Apparatus Used

Nessler Cylinders, Gas Detector Tubes, Sieves, Thermometers, Ultraviolet Ray Lamps, Volumetric Glassware, Weights and Balances.

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Key Terms and Formulas

Term	Formula
Concentration	Moles of solute ÷ Volume in litres
Mass % (w/w)	(Mass of component ÷ Total mass of solution) × 100
Volume % (v/v)	(Volume of component ÷ Total volume of solution) × 100
Mass by Volume % (m/v)	(Mass of solute ÷ Volume of solution) × 100
Molarity	Moles of solute ÷ Volume of solution in litres
Molality	Moles of solute ÷ Mass of solvent in kg
Normality	Gram equivalents of solute ÷ Litres of solution
PPM	(Parts of component ÷ Total parts of all components) × 10⁶
Mole Fraction	Moles of component ÷ Total moles of all components

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Chemical / Titrimetric Methods

Volumetric Analysis

Volumetric analysis is a method of quantitative chemical analysis in which the amount of a substance is determined by measuring the volume it occupies in chemical reactions. It is widely used in pharmaceutical analysis because of its robustness, low cost, and high precision.

Types of volumetric methods:

1. Neutralisation Titration
2. Non-aqueous Titration
3. Redox Titration
4. Complexometric Titration
5. Precipitation Titration

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1. Neutralisation Titration (Acid-Base Titration)

What is Titration?

Titration is the process of determining the unknown concentration of a solution (analyte) by reacting it with a standard solution of known concentration (titrant) until the reaction is complete.

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Acid-Base Theories

Acids:

Concept	Definition	Examples
Arrhenius	Produces free H⁺ ions in aqueous solution	HCl, HNO₃, H₂SO₄, CH₃COOH
Bronsted-Lowry	Donates a proton (H⁺) in any solvent	Same as above
Lewis	Accepts a lone pair/electron pair	AlCl₃, SF₆, SO₃

Bases:

Concept	Definition	Examples
Arrhenius	Produces free OH⁻ ions in aqueous solution	NaOH, KOH, Ca(OH)₂
Bronsted-Lowry	Accepts a proton (H⁺) in any solvent	SO₄²⁻, Cl⁻, O²⁻
Lewis	Donates a lone pair/electron pair	NH₃, amines

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Types of Neutralisation Titration

A. Strong Acid / Strong Base Titration

Strong acid and strong base react rapidly, breaking into ions and forming salt and water. The neutralisation point occurs at pH 7.

Titration curve characteristics (HCl vs NaOH):

• Initially low pH due to excess acid
• pH rises slowly as base is added
• Sharp rise occurs near the equivalence point
• pH above 7 after equivalence point due to excess alkali
• All strong acid/strong base curves have the same shape

B. Weak Acid / Strong Base and Weak Base / Strong Acid Titration

The weak component dissociates slowly while the strong component dissociates rapidly — making this titration different from A above.

Titration curve characteristics:

• pH rises gradually at first
• A flat zone (half neutralisation point) is observed where pH = pKa of the weak acid
• At this point a buffer solution is formed (mix of weak acid and its conjugate base)
• After the buffer zone, a sharp pH rise occurs above pH 7 at the equivalence point

C. Weak Acid / Weak Base Titration

The change in pH near the equivalence point is very gradual throughout the entire neutralisation. Ordinary indicators cannot detect the endpoint — mixed indicators must be used.

Back Titration

Used when direct titration is not possible — for insoluble substances, volatile substances, or reactions requiring excess reagent. A known excess of standard solution is added to the analyte, and after the reaction is complete, the excess is back-titrated with another standard solution.

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Important Neutralisation Titrations — Summary Table

Drug / Chemical	Titration Type	Titrant
Aspirin	Back	H₂SO₄
Ammonium chloride	Back	NaOH
Methyl salicylate	Back	HCl
Ephedrine	Back	NaOH
Furosemide	Back	NaOH
Pyrazinamide	Back	NaOH
Sodium hydroxide	Direct	H₂SO₄
Thiopentone sodium	Direct	H₂SO₄
Calamine	Direct	NaOH
Oxyphenbutazone	Direct	NaOH
Salicylic acid	Direct	NaOH
Boric acid	Direct	NaOH
Hydrochloric acid	Direct	NaOH
Nicotinic acid	Direct	NaOH
Saccharin	Direct	NaOH

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Neutralisation Indicators

Indicators are substances that give a visible sign — usually a colour change — indicating the presence or absence of a threshold concentration of acid or alkali in a solution.

Indicator	pH Range	Acid Colour	Base Colour
Thymol blue	1.2 – 2.8	Red	Yellow
Methyl red	4.2 – 6.3	Red	Yellow
Methyl orange	3.1 – 4.4	Red	Orange
Bromocresol green	3.8 – 5.4	Yellow	Blue
Bromocresol blue	3.0 – 4.6	Yellow	Blue
Phenol red	6.8 – 8.4	Yellow	Red
Phenolphthalein	8.3 – 11.0	Colourless	Red

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2. Non-Aqueous Titration

Non-aqueous titration is the most common titration used in pharmacopoeial assays. It is suitable for weak acids and weak bases that are poorly soluble in water and have low reactivity in aqueous conditions.

Why not use water? Water is amphoteric — it competes with weak acids and weak bases for protons. If the acid or base is relatively weak, it cannot compete effectively with water, and titration becomes unreliable.

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Types of Non-Aqueous Solvents

Solvent Type	Properties	Examples
Protogenic	Strong acidic character; high dielectric constant; levelling effect on bases	Formic acid, acetic acid, sulphuric acid
Protophilic	Basic nature; accepts protons; high dielectric constant	Dimethyl formamide, pyridine, ammonia
Aprotic	Chemically neutral; no O-H or N-H bonds; low dielectric constant; not ionised	Benzene, acetone, chloroform
Amphiprotic	Both acidic and basic properties; high dielectric constant; partially ionised	Water, ethanol, isopropyl alcohol

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Types of Non-Aqueous Titration

A. Titration of Weak Bases Acetic acid is used as solvent. Acetous perchloric acid is the titrant. Weak bases compete effectively with acetic acid for protons.

Pharmacopoeial drugs titrated by this method: Adrenaline, Metronidazole, Codeine, Amitriptyline HCl, Lignocaine HCl, Neostigmine bromide, Pancuronium bromide.

Indicators used: Oracet blue, Quinolidine red, Crystal violet.

B. Titration of Weak Acids Alcohol or aprotic solvent is used. Titrants used are lithium methoxide in methanol or tetrabutyl ammonium hydroxide in dimethyl formamide.

Pharmacopoeial drugs titrated by this method: Barbiturates, Uracils, Sulphonamides.

Indicator used: Thymol blue.

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3. Redox Titration

Redox titration involves both oxidation and reduction reactions.

Oxidation = addition of oxygen / loss of electron / removal of hydrogen Reduction = removal of oxygen / gain of electron / addition of hydrogen

Example redox reaction: MnO₄⁻ + I⁻ → Mn²⁺ + I₂

Oxidising agent — accepts electrons; decreases in positive valency or increases in negative valency.

Reduction potential (E⁰) — measures how readily a compound gains electrons. A higher positive value = stronger oxidising agent.

Standard Reduction Potentials Table

Reaction	E⁰ Value
Ce⁴⁺ + e⁻ → Ce³⁺	+1.61 V
Mn⁷⁺ → Mn²⁺	+1.51 V
Cl₂ + 2e⁻ → 2Cl⁻	+1.36 V
Br₂ + 2e⁻ → 2Br⁻	+1.065 V
Fe³⁺ + e⁻ → Fe²⁺	+0.771 V
I₂ + 2e⁻ → 2I⁻	+0.536 V
AgCl + e⁻ → Ag + Cl⁻	+0.223 V
2H⁺ + 2e⁻ → H₂	0 V
Fe²⁺ + 2e⁻ → Fe	−0.440 V
Ca²⁺ + 2e⁻ → Ca	−2.888 V

A substance with a higher reduction potential will oxidise one with a lower reduction potential. Example: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

Direct Redox Titration — Iodine (moderately strong oxidising agent) oxidises substances with lower reduction potential. End point detected using starch indicator (blue colour with excess iodine). Pharmacopoeial assays: Ascorbic acid, Sodium stibogluconate, Dimercaprol injection, Acetarsal.

Displacement Redox Titration — Used in assays of Liquefied phenol, Methyl hydroxybenzoate, Propyl hydroxybenzoate, Phenidione.

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4. Complexometric Titration

This titration involves the formation of a complex between the titrant and the analyte. It is used to estimate metals and their salts. The titrant is called a ligand.

EDTA (Ethylenediamine Tetraacetic Acid / Sodium Edetate) — the most widely used titrant. Forms stable 1:1 complexes with all metals except alkali metals (Na, K). The general reaction:

Mⁿ⁺ + Na₂EDTA → (MEDTA)ⁿ⁻⁴ + 2H⁺

Alkaline earth metals (Ca, Mg) form complexes unstable at low pH — titrated in ammonium chloride buffer at pH 10.

Indicators used: Calcein, Eriochrome Black T, Curcumin, Calcon, Murexide (ammonium purpurate), Hematoxylin, Fast Sulphon Black.

Pharmacopoeial drugs titrated: Calcium acetate, Calcium chloride, Calcium gluconate, Magnesium carbonate, Magnesium hydroxide, Magnesium trisilicate, Bismuth subcarbonate, Bacitracin zinc, Zinc chloride, Zinc undecanoate.

Types of Complexometric Titration

1. Direct complexometric titration
2. Back complexometric titration
3. Replacement complexometric titration
4. Indirect complexometric titration

Back titration with EDTA is used for insoluble metal salts: Aluminium glycinate, Aluminium hydroxide, Aluminium sulphate, Calcium hydrogen phosphate.

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5. Precipitation Titration

In precipitation titration, the titrant reacts with the analyte to form an insoluble precipitate. The end point is reached when no more precipitate forms.

Methods of Precipitation Titration

A. Mohr’s Method

Determines chloride ion concentration. Silver nitrate is slowly added to the chloride solution.

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

When all chloride is precipitated, excess silver ions react with potassium chromate indicator to form a red-brown precipitate of silver chromate — this is the end point.

2Ag⁺(aq) + CrO₄²⁻(aq) → Ag₂CrO₄(s)

Must be carried out at pH 6.5–9.0. Used for chloride determination in river water, sea water, and stream water.

B. Volhard’s Method

Used for chloride determination by back titration with potassium thiocyanate. Excess silver nitrate is first added; then Fe³⁺ is added as indicator and the solution is titrated with potassium thiocyanate.

Ag⁺(aq) + Cl⁻(aq) → AgCl(s) — first precipitation Ag⁺(aq) + SCN⁻(aq) → AgSCN(s) — pale yellow precipitate Fe³⁺(aq) + SCN⁻(aq) → [FeSCN]²⁺(aq) — dark red end point

Used under acidic pH conditions. Modified Volhard’s method is used specifically for NaCl and KCl determination.

C. Fajan’s Method

Uses adsorption indicators — dyes that adsorb or desorb on the precipitate surface at the equivalence point, producing a colour change. Indicators used: Fluorescein, Eosin (acid dyes); Rhodamine series (basic dyes).

Fluorescein forms a yellow-green colour in solution and is the most commonly used indicator for titrating chloride ions with silver.

Note: Sodium chloride can be estimated by Volhard’s method, Mohr’s method, or Fajan’s method.

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Gravimetric Analysis

Principle: Gravimetric analysis involves isolating and weighing a substance from a solution after precipitating it as an insoluble compound of known chemical composition. It is a quantitative analysis method based on weight.

Methods

1. Precipitation method — analyte is converted to an insoluble precipitate, filtered, washed, ignited, and weighed.
2. Volatilisation method — analyte is heated so decomposition products volatilise, leaving an unvolatilised residue of known composition for weighing.
3. Electro-analytical method — ions are deposited on an electrode by passing electric current. Based on Faraday’s second law. Avoids filtration and decomposition steps.

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Steps in Gravimetric Analysis

Step 1 — Sample preparation and dissolution Take a homogeneous, powdered sample. Dissolve completely in water in a clean beaker with stirring or gentle warming. Adjust conditions (temperature, pH, volume) to obtain a precipitate with low solubility suitable for filtration.

Step 2 — Precipitation Add dilute precipitant solution slowly with continuous stirring in a hot solution. After settling, add a few drops of precipitant to confirm complete precipitation. The ideal precipitant should form a stable, easily filtered precipitate with low solubility.

Step 3 — Digestion Heat to speed up the process and improve purity and filterability of the precipitate.

Step 4 — Filtration Separate the precipitate from the mother liquor using appropriate filter media:

• Filter paper
• Filter mats
• Filter pulp
• Permanent porous filter disc

Step 5 — Washing Remove surface impurities from the precipitate. Wash with distilled water (e.g., BaSO₄) or with dilute acid if peptization is a concern (e.g., lead sulphate washed with dilute HNO₃).

Step 6 — Drying and Ignition

• Drying = below 250°C (removes water and adsorbed solvent)
• Ignition = above 250°C and below 1200°C (converts precipitate to final weighable form)

Step 7 — Weighing and Calculation Cool in a desiccator and weigh on an analytical balance. Calculate the result:

%A = (Grams of analyte ÷ Grams of sample) × 100

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Colloidal States in Gravimetric Analysis

Colloidal suspensions are not suitable for gravimetric analysis because their particle size (0.1µ to 1mµ) prevents filtration through ordinary filter paper (which retains particles only above 10µ). Their stability can be reduced by stirring, heating, or adding electrolytes — causing colloidal particles to bind together into a filterable mass.

All colloidal particles carry either a positive or negative charge, stabilised by surface adsorption of ions. When a powerful beam of light is passed through a colloidal solution and viewed at a right angle, scattering is observed — this is called the Tyndall Effect.

Classification of Colloids

Property	Lyophobic Colloids (Suspensoids)	Lyophilic Colloids (Emulsoids)
Viscosity	Slightly viscous	Very viscous — almost jelly-like (gel)
Effect of adding water	No effect	Affected by water or solvent
Electrolyte needed for flocculation	Small concentration	Large concentration
Electrical charge	Definite sign; changed only by specific methods	Can change charge easily
Under ultramicroscope	Bright particles in vigorous Brownian motion (e.g., gold solution)	Only a diffuse light cone (e.g., gelatin)