PHARMACEUTICAL INORGANIC CHEMISTRY

Complete Short Answer Questions (5 Marks Each)

University Level - High Yield Answers

---

UNIT - I: VOLUMETRIC ANALYSIS

---

Q1. List out various volumetric methods and explain back titration with examples.

Volumetric Methods at a Glance

VOLUMETRIC ANALYSIS
        |
        ├── 1. Acid-Base Titrations (Acidimetry & Alkalimetry)
        ├── 2. Redox Titrations (Permanganometry, Cerimetry)
        ├── 3. Precipitation Titrations (Argentometry)
        ├── 4. Complexometric Titrations (EDTA Titrations)
        └── 5. Non-Aqueous Titrations

---

Back Titration

• The analyte does not react directly with the titrant
• The analyte is insoluble in water
• The direct endpoint is difficult to detect
• The analyte is volatile

Step 1: Add EXCESS known reagent (R) to analyte (A)
        A  +  R (excess)  →  Product  +  R (remaining)

Step 2: Titrate the REMAINING excess R with standard titrant (T)
        R (remaining)  +  T  →  endpoint

Step 3: Amount of A = Amount of R added - Amount of R remaining

Amount of analyte = Meq of Reagent added - Meq of back titrant used

---

• Add excess HCl (known amount) to CaCO₃
• CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
• Back titrate remaining HCl with NaOH
• Amount of CaCO₃ = Meq HCl added - Meq NaOH used

• Add excess NaOH (known) to NH₄Cl, heat to expel NH₃
• NH₄Cl + NaOH → NaCl + NH₃↑ + H₂O
• Back titrate excess NaOH with standard HCl

• Aspirin hydrolysed with excess NaOH (known)
• Back titrate with HCl to find excess NaOH
• Amount of aspirin calculated from difference

---

Q2. Preparation and Standardization of 500 mL of N/10 Sodium Hydroxide (NaOH)

Preparation

1. Weigh approximately 2.1 g of NaOH pellets (slight excess taken as NaOH is hygroscopic)
2. Dissolve in small amount of freshly boiled and cooled distilled water (CO₂-free water)
3. Transfer to a 500 mL volumetric flask
4. Make up to the mark with CO₂-free distilled water
5. Mix well and store in a well-stoppered bottle (NOT glass stoppered - use rubber stopper)

Note: NaOH absorbs CO₂ from air forming Na₂CO₃, which introduces error. Hence CO₂-free water and rubber stoppers are used.

---

Standardization

• 6.3 g/L → dissolve 0.63 g in 100 mL

1. Pipette 10 mL of N/10 oxalic acid into a conical flask
2. Add 2-3 drops of phenolphthalein indicator
3. Fill burette with NaOH solution
4. Titrate until a permanent pale pink colour appears
5. Note the burette reading (V mL)

N₁V₁  =  N₂V₂
N(NaOH) × V(NaOH)  =  N(Oxalic acid) × V(Oxalic acid)
N(NaOH) = (N/10 × 10) / V(NaOH)

---

Q3. Uses of Dilute HCl and Oxalic Acid

Dilute Hydrochloric Acid (HCl)

• Dilute HCl BP contains approximately 10% w/v HCl
• Used as pharmaceutical aid and in estimation of antacids by back titration

---

Oxalic Acid (H₂C₂O₄·2H₂O)

2KMnO₄ + 5H₂C₂O₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 10CO₂ + 8H₂O
(Purple)                                    (Colourless)

---

Q4. Primary Standard - Definition, Examples, and Properties

Definition

Example

• Oxalic acid (H₂C₂O₄·2H₂O)
• Potassium dichromate (K₂Cr₂O₇)
• Sodium chloride (NaCl)
• Anhydrous sodium carbonate (Na₂CO₃)
• Potassium hydrogen phthalate (KHP)

Properties of a Primary Standard

PROPERTIES OF PRIMARY STANDARD
         |
         ├── 1. HIGH PURITY (≥99.9%)
         |       Should be available in pure form
         |
         ├── 2. STABLE
         |       Should not absorb moisture or CO₂
         |       Should be stable to heat and light
         |
         ├── 3. HIGH MOLECULAR WEIGHT
         |       Reduces weighing error
         |
         ├── 4. READILY SOLUBLE
         |       Should dissolve in water/solvent
         |
         ├── 5. NON-HYGROSCOPIC
         |       Should not absorb water from air
         |
         ├── 6. REACT STOICHIOMETRICALLY
         |       Reaction should be complete and definite
         |
         └── 7. COMMERCIALLY AVAILABLE
                 Should be easily obtainable

Key Point: Oxalic acid satisfies all above criteria and is widely used as a primary standard for NaOH and KMnO₄ standardization.

---

Q5. Classification of Determinate Errors and Steps to Minimize

Definition

Classification

DETERMINATE ERRORS
        |
        ├── 1. INSTRUMENTAL ERRORS
        |       - Faulty calibration of burette, pipette, flask
        |       - Worn out glassware
        |       - Defective balance weights
        |
        ├── 2. PERSONAL (OPERATIVE) ERRORS
        |       - Wrong reading of meniscus
        |       - Colour blindness affecting endpoint detection
        |       - Parallax error in reading burette
        |
        ├── 3. METHOD ERRORS
        |       - Incorrect indicator choice
        |       - Incomplete reactions
        |       - Co-precipitation of impurities
        |
        └── 4. REAGENT ERRORS
                - Impure reagents used
                - Degraded/expired standards
                - Wrong strength of reagent

Steps to Minimize Errors

1. Perform blank correction - titrate without analyte and subtract
2. Perform triplicate titrations and take mean
3. Use freshly prepared and standardized solutions
4. Ensure proper indicator selection
5. Calibrate all volumetric glassware before use

---

Q6. Note on Back Titration

• Analyte is insoluble (e.g., CaCO₃, BaSO₄)
• Analyte is volatile (e.g., NH₃)
• Direct endpoint is unclear
• Reaction with direct titrant is too slow

Analyte + Excess reagent A (known) → Products
Excess reagent A + Titrant B (standard) → Endpoint
Analyte = Meq(A) - Meq(B)

• Add 50 mL of 0.1N HCl to tablet
• Heat to dissolve completely
• Back titrate with 0.1N NaOH using phenolphthalein
• Calculate CaCO₃ from difference

---

UNIT - II: ACID-BASE AND REDOX TITRATIONS

---

Q1. Acidimetry and Alkalimetry

Definitions

H⁺ + OH⁻ → H₂O

---

Indicators Used

---

Example - Acidimetry

• Titrant: 0.1N HCl (standard)
• Analyte: NaOH solution
• Indicator: Phenolphthalein
• Reaction: NaOH + HCl → NaCl + H₂O
• Endpoint: Pink to colourless

• Titrant: 0.1N NaOH (standard)
• Analyte: HCl solution
• Indicator: Methyl orange or Phenolphthalein
• Reaction: HCl + NaOH → NaCl + H₂O

---

Titration Curve

pH
14 |                          ___________
   |                         /
 7 |________________________/
   |
 0 |_____________________________|_____
                         Equivalence Point
                    Volume of base added →

---

Q2. Oxidizing and Reducing Agents

Definitions

REDOX REACTION
OIL RIG Rule: Oxidation Is Loss, Reduction Is Gain (of electrons)

Reducing Agent → loses electrons → gets OXIDIZED
Oxidizing Agent → gains electrons → gets REDUCED

---

Common Oxidizing Agents

Common Reducing Agents

---

Q3. Theory of Redox Titrations, Iodimetry and Iodometry

Theory of Redox Titrations

E = E° + (0.0591/n) × log [Oxidized form / Reduced form]

• Reaction must be rapid and stoichiometric
• Large difference in electrode potentials (ΔE > 0.2 V)
• Sharp endpoint must be detectable
• No side reactions

---

Iodimetry vs Iodometry

As₂O₃ + 2I₂ + 2H₂O → As₂O₅ + 4HI

Step 1: 2KMnO₄ + 10KI + 8H₂SO₄ → 2MnSO₄ + 5I₂ + K₂SO₄ + 8H₂O
Step 2: I₂ + 2Na₂S₂O₃ → 2NaI + Na₂S₄O₆

---

Q4. Principle and Reactions in Iodometry Titrations

Principle

IODOMETRY PRINCIPLE
Step 1: Oxidizing agent + KI (excess) → Liberated I₂
Step 2: Liberated I₂ + Na₂S₂O₃ (standard) → Endpoint

Conditions Required:

• Excess KI (10-15% solution) must be used
• Solution must be slightly acidic (H₂SO₄ or acetic acid)
• Starch indicator added near endpoint (not at start - starch-iodine complex is hard to break)
• Titration done in the dark to prevent photo-oxidation of I⁻

Reactions - Estimation of Copper Sulphate

Step 1: Liberation of Iodine
2CuSO₄ + 4KI → 2CuI↓ + 2K₂SO₄ + I₂
(Blue)                                 (Brown)

Step 2: Titration with sodium thiosulphate
I₂ + 2Na₂S₂O₃ → 2NaI + Na₂S₄O₆
(Brown → colourless at endpoint)

Reactions - Estimation of KMnO₄

2KMnO₄ + 10KI + 8H₂SO₄ → 2MnSO₄ + 5I₂ + 6K₂SO₄ + 8H₂O
I₂ + 2Na₂S₂O₃ → 2NaI + Na₂S₄O₆

Endpoint

---

Q5. Preparation and Standardization of 0.1N KMnO₄

Preparation

• Weigh 3.2 g KMnO₄
• Dissolve in 1 litre of freshly boiled distilled water
• Boil for 15 minutes (to oxidize organic matter)
• Cool and filter through sintered glass funnel (NOT filter paper - KMnO₄ oxidizes it)
• Store in amber-coloured (dark) bottle

Caution: KMnO₄ is unstable in light; always store in dark bottles.

---

Standardization

1. Prepare 0.1N oxalic acid (dissolve 0.63 g in 100 mL)
2. Pipette 10 mL of 0.1N oxalic acid into conical flask
3. Add 10 mL of dilute H₂SO₄ (to provide acidic medium)
4. Heat to 70-80°C (NOT boiling - CO₂ loss and decomposition)
5. Titrate from burette with KMnO₄ solution
6. First few drops decolourise slowly, then rapidly (autocatalysis by Mn²⁺)
7. Endpoint: Permanent pink/violet colour for 30 seconds

2KMnO₄ + 5H₂C₂O₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 10CO₂ + 8H₂O
(Purple)                               (Colourless)

N₁V₁ = N₂V₂
N(KMnO₄) × V(KMnO₄) = 0.1 × 10
N(KMnO₄) = 1 / V(KMnO₄)

---

Q6. Preparation and Standardization of 0.1N Sodium Thiosulphate

Preparation

1. Weigh 24.8 g of Na₂S₂O₃·5H₂O
2. Dissolve in freshly boiled and cooled distilled water
3. Add small amount of Na₂CO₃ (0.2 g) to prevent decomposition by CO₂
4. Make up to 1 litre in volumetric flask
5. Store in dark amber bottle away from light

Na₂S₂O₃ is unstable - decomposes in acidic conditions and light: Na₂S₂O₃ + HCl → NaCl + S↓ + SO₂ + H₂O

---

Standardization

1. Weigh accurately about 0.35 g of dried KIO₃
2. Dissolve in distilled water, add excess KI and 5 mL H₂SO₄
3. Liberated iodine appears (brown colour)
4. Titrate with Na₂S₂O₃ until pale yellow
5. Add starch indicator - deep blue colour
6. Continue titration until blue disappears = endpoint

Step 1: KIO₃ + 5KI + 6HCl → 3I₂ + 6KCl + 3H₂O
Step 2: I₂ + 2Na₂S₂O₃ → 2NaI + Na₂S₄O₆

---

UNIT - III: NON-AQUEOUS TITRATIONS & ARGENTOMETRY

---

Q1. Non-Aqueous Titrations - Definition and Important Conditions

Definition

• Water suppresses ionization of weak acids/bases
• Endpoint is not sharp in aqueous medium
• Many organic compounds are insoluble in water

---

Important Conditions

• Solvent: Glacial acetic acid (enhances basicity)
• Titrant: 0.1M perchloric acid in glacial acetic acid
• Indicator: Crystal violet (violet → blue-green = endpoint)

---

Q2. Precautions in Preparing Perchloric Acid Titrant

Precautions:

⚠ SAFETY PRECAUTIONS FOR PERCHLORIC ACID
        |
        ├── 1. NEVER add water to perchloric acid
        |       Hot anhydrous perchloric acid is explosive
        |
        ├── 2. Use glacial acetic acid as solvent - not other solvents
        |       Acetic anhydride may be added to remove water
        |
        ├── 3. Cool the mixture before adding perchloric acid
        |       Heat generated on mixing can cause explosion
        |
        ├── 4. Avoid contact with organic matter
        |       Perchloric acid + organic material = fire hazard
        |
        ├── 5. Store in glass containers only
        |       Not plastic - perchloric acid dissolves plastics
        |
        ├── 6. Standardize before each use
        |       Strength changes on standing
        |
        └── 7. Preparation: Add perchloric acid slowly to glacial
                acetic acid, then add acetic anhydride if needed

• Take 8.5 mL of 70% perchloric acid
• Slowly add to 900 mL glacial acetic acid
• Cool, add acetic anhydride if moisture is present
• Make up to 1000 mL with glacial acetic acid
• Standardize with potassium hydrogen phthalate (KHP)

---

Q3. Types of Solvents in Non-Aqueous Titration

Classification of Solvents

SOLVENTS IN NON-AQUEOUS TITRATION
        |
        ├── 1. PROTOPHILIC (Basic solvents)
        |       Accept protons from analyte
        |       Examples: Pyridine, dimethylformamide, acetone
        |       Use: For estimation of weak acids
        |
        ├── 2. PROTOGENIC (Acidic solvents)
        |       Donate protons to analyte
        |       Examples: Glacial acetic acid, formic acid
        |       Use: For estimation of weak bases
        |
        ├── 3. AMPHIPROTIC (Both acid and base)
        |       Can donate and accept protons
        |       Examples: Ethanol, methanol, water
        |       Use: For moderately strong acids/bases
        |
        └── 4. APROTIC (Neutral solvents)
                Neither donate nor accept protons
                Examples: Benzene, chloroform, dioxane
                Use: As diluents; leveling medium

---

Q4. Argentometric Titrations

Ag⁺ + X⁻ → AgX↓ (precipitate)
Where X = Cl⁻, Br⁻, I⁻, SCN⁻

Three Methods of Argentometry:

ARGENTOMETRIC METHODS
        |
        ├── 1. MOHR'S METHOD
        |       Indicator: K₂CrO₄ (potassium chromate)
        |       pH: 6.5-10.5 (neutral to slightly alkaline)
        |       Endpoint: Red Ag₂CrO₄ precipitate forms
        |       Suitable for: Cl⁻, Br⁻
        |
        ├── 2. VOLHARD'S METHOD
        |       Back titration with KSCN
        |       Indicator: Ferric alum (ammonium ferric sulphate)
        |       pH: Acidic (HNO₃ medium)
        |       Endpoint: Red FeSCN²⁺ complex
        |       Suitable for: All halides including I⁻
        |
        └── 3. FAJAN'S METHOD
                Adsorption indicator (Fluorescein/Eosin)
                Direct titration
                Endpoint: Colour change on AgCl surface
                Suitable for: Cl⁻, Br⁻, I⁻, SCN⁻

---

Q5. Different Methods of Estimation of Halides

Ag⁺ + Cl⁻ → AgCl↓ (white) [primary reaction]
2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄↓ (red-brown) [endpoint reaction]

• First AgCl precipitates completely (lower Ksp)
• Then excess Ag⁺ reacts with CrO₄²⁻ to give red colour = endpoint

Ag⁺ (excess) + Cl⁻ → AgCl↓
Ag⁺ (remaining) + SCN⁻ → AgSCN↓
SCN⁻ (excess) + Fe³⁺ → [FeSCN]²⁺ (red) = endpoint

---

Q6. Standardization of Silver Nitrate (AgNO₃)

1. Dry NaCl at 120°C for 1 hour, cool in desiccator
2. Weigh accurately ~0.3 g NaCl
3. Dissolve in 50 mL distilled water
4. Add 1 mL of 5% potassium chromate indicator
5. Titrate with AgNO₃ from burette
6. Endpoint: Brick-red precipitate (Ag₂CrO₄) persists on shaking

AgNO₃ + NaCl → AgCl↓ + NaNO₃
(White precipitate)

N(AgNO₃) = (W × 1000) / (Eq.wt × V(AgNO₃))
           = (W × 1000) / (58.5 × V)

---

Q7. Fajan's Method

Principle

Before endpoint: AgCl surface adsorbs Cl⁻ → negative surface
                 Indicator (anionic dye) repelled → original colour

At endpoint: Excess Ag⁺ adsorbed → positive surface
             Indicator attracted → colour change

Indicators Used

Example - Estimation of NaCl using Fluorescein

Titrant: AgNO₃
Analyte: NaCl
Indicator: Fluorescein
Reaction: Ag⁺ + Cl⁻ → AgCl↓
Endpoint: Yellow-green → Pink-red

• Solution must be slightly acidic (pH 7-10 for fluorescein)
• No excess KNO₃ or KCl (interferes)
• Must be done in light (adsorption visible)

---

UNIT - IV: COMPLEXOMETRIC TITRATIONS

---

Q1. Complexometric Titrations - Principle and Examples

Definition

Principle

COMPLEXOMETRIC TITRATION PRINCIPLE

Metal ion (Mⁿ⁺) + Chelating agent (EDTA) → Stable Metal-EDTA complex

M⁴⁺ + H₂Y²⁻ → MY^(4-4)⁻ + 2H⁺
(Metal ion) + (EDTA)    → (Stable complex) + (Acid)

Where H₄Y = EDTA (Ethylene Diamine Tetraacetic Acid)

EDTA Structure

• 2 nitrogen atoms (N donor)
• 4 carboxylate groups (-COO⁻, O donor) = Hexadentate ligand (6-coordinate sites)

---

Types of Complexometric Titrations

Example - Estimation of Calcium (Ca²⁺)

pH = 12 (NaOH buffer)
Indicator: Murexide (ammonium purpurate)
Ca²⁺ + EDTA → CaEDTA complex
Endpoint: Pink → Purple (free indicator)

Example - Water Hardness (Ca²⁺ + Mg²⁺)

pH = 10 (ammonia buffer)
Indicator: Eriochrome Black T (EBT)
Wine red (Metal-EBT) → Blue (free EBT) = endpoint

---

Q2. Ligands - Definition and Types

Definition

Types of Ligands

TYPES OF LIGANDS (based on donor atoms)

├── 1. MONODENTATE (1 donor site)
|       Examples: Cl⁻, NH₃, H₂O, CN⁻, F⁻
|       "Mono" = one; "dentate" = tooth (bite at one point)
|
├── 2. BIDENTATE (2 donor sites)
|       Examples: Ethylenediamine (en), Oxalate (ox), 
|                 Acetylacetonate (acac)
|       Form 5 or 6 membered chelate rings
|
├── 3. TRIDENTATE (3 donor sites)
|       Example: Diethylenetriamine (dien)
|
├── 4. TETRADENTATE (4 donor sites)
|       Example: Triethylenetetramine (trien)
|
├── 5. HEXADENTATE (6 donor sites)
|       Example: EDTA (most important in pharmacy)
|       Forms most stable complexes
|
└── 6. AMBIDENTATE (can bind through different atoms)
        Example: SCN⁻ (can bond through S or N)
                 NO₂⁻ (can bond through N or O)

---

Q3. Preparation and Standardization of 0.05M Disodium EDTA

Preparation

1. Weigh approximately 9.3 g of disodium EDTA
2. Dissolve in small amount of distilled water with stirring
3. Transfer to 500 mL volumetric flask
4. Make up to mark with distilled water
5. Mix well

Note: EDTA dissolves slowly in cold water - use warm water to dissolve, then cool before making up to volume.

---

Standardization

1. Weigh accurately ~0.3 g anhydrous ZnSO₄
2. Dissolve in 50 mL water
3. Add 10 mL ammonia buffer (pH 10)
4. Add 2 drops of EBT indicator (wine red colour)
5. Titrate with EDTA from burette
6. Endpoint: Wine red → Blue

Zn²⁺ + Na₂H₂Y → Na₂ZnY + 2H⁺
(Zn-EBT complex)  (Zn-EDTA formed, EBT released)
Wine red → Blue

M(EDTA) = (W × 1000) / (MW × V)

---

Q4. Principle of Complexometric Titrations

• Ammonia buffer (NH₃ + NH₄Cl) maintains pH
• At pH 10, EDTA exists mainly as Y⁴⁻ (fully deprotonated) - most reactive form
• Too low pH - EDTA protonated, less reactive
• Too high pH - metal hydroxide precipitates

K'f = Kf × αY
Where αY = fraction of EDTA as Y⁴⁻ (depends on pH)

M-Indicator complex + EDTA → M-EDTA + free Indicator
(Coloured 1)                          (Coloured 2)

---

UNIT - V: GRAVIMETRIC ANALYSIS & LIMIT TESTS

---

Q1. Gravimetric Analysis - Definition and Importance

Definition

Importance

Analyte (solution) + Precipitating agent → Precipitate
Precipitate → Filter → Dry/Ignite → Weigh
% Analyte = (Weight of precipitate × Gravimetric factor × 100) / Weight of sample

---

Q2. Steps in Gravimetric Analysis

STEPS IN GRAVIMETRIC ANALYSIS

1. PREPARATION OF SOLUTION
   │── Dissolve sample in appropriate solvent
   │── Remove interfering ions if needed

2. PRECIPITATION
   │── Add precipitating agent slowly with stirring
   │── Heat (digest) to get coarse precipitate
   │── Use slight excess of precipitating agent

3. DIGESTION (Ageing)
   │── Heat precipitate in mother liquor for 30-60 min
   │── Coagulates colloidal particles (Ostwald ripening)

4. FILTRATION
   │── Use Whatman filter paper or sintered glass crucible
   │── Wash with cold dilute electrolyte solution
   │── Prevents peptization during washing

5. DRYING / IGNITION
   │── Dry at 105-120°C (if filter paper not used)
   │── Ignite at 800-1000°C in muffle furnace if needed
   │── Convert to stable weighing form

6. WEIGHING
   │── Cool in desiccator
   │── Weigh accurately on analytical balance

7. CALCULATION
   └── Calculate % analyte from weight of precipitate

---

Q3. Gutzeit's Apparatus for Arsenic Limit Test

Principle

ARSENIC LIMIT TEST - GUTZEIT'S METHOD

Reactions:
Step 1: Zn + H₂SO₄ → ZnSO₄ + 2H (nascent hydrogen)
Step 2: As₂O₃ + 6H → 2AsH₃↑ (arsine gas)
Step 3: AsH₃ + 3HgCl₂ → As(HgCl)₃ + 3HCl (yellow stain)
        (or)   2AsH₃ + 3HgCl₂ → 2As + 3Hg + 6HCl (brown stain)

Gutzeit's Apparatus Diagram

          ┌─────────────────────────────────────────┐
          │   GUTZEIT'S APPARATUS                   │
          │                                         │
          │  ┌──────────────┐                       │
          │  │ Cotton wool  │ ← Soaked in lead      │
          │  │    plug      │   acetate (removes H₂S)│
          │  └──────┬───────┘                       │
          │         │                               │
          │  ┌──────┴───────┐                       │
          │  │  Glass tube  │ ← 3-4 mm bore         │
          │  └──────┬───────┘                       │
          │         │                               │
          │  ┌──────┴───────┐                       │
          │  │HgCl₂ paper   │ ← Mercuric chloride   │
          │  │  (test strip)│   impregnated paper   │
          │  └──────────────┘                       │
          │         ↑                               │
          │    AsH₃ gas rises                       │
          │         ↑                               │
          │  ┌──────────────────────────────────┐   │
          │  │ Conical flask (250 mL)           │   │
          │  │ Sample + KI + HCl + SnCl₂       │   │
          │  │ Zinc granules added              │   │
          │  │ (H₂ generated → AsH₃ formed)    │   │
          │  └──────────────────────────────────┘   │
          └─────────────────────────────────────────┘

1. Wide-mouth conical flask (arsenic liberation vessel)
2. Lead acetate cotton plug (removes H₂S interference)
3. Glass tube
4. Mercuric chloride paper strip (test strip)

---

Q4. Limit Test for Sulphate

Principle

SO₄²⁻ + BaCl₂ → BaSO₄↓ (white turbidity) + 2Cl⁻

Procedure

1. Dissolve sample in distilled water (avoid sulphate-free water)
2. Add 2 mL of dilute HCl (acidify - prevents carbonate and phosphate interference)
3. Add 5 mL of Barium sulphate TS (25% BaCl₂ in 25% H₂SO₄)
4. Dilute to 50 mL
5. Allow to stand for 5 minutes
6. Compare turbidity with standard (containing known SO₄²⁻)

BaCl₂ + H₂SO₄ → BaSO₄↓ + 2HCl
(Acidic medium prevents Ba₃(PO₄)₂ precipitation)

---

Q5. Limit Test for Lead

Principle

Pb²⁺ + Na₂S → PbS↓ (brown-black) + 2Na⁺

Procedure

1. Dissolve sample in dilute acetic acid
2. Add 1 mL of potassium cyanide (to complex Cu, Fe, Zn and prevent interference)
3. Add sodium sulphide TS
4. Compare brown colour with standard lead solution treated identically

PbSO₄/Pb(NO₃)₂ (test) + CH₃COOH → Pb(CH₃COO)₂ (soluble lead acetate)
Pb²⁺ + S²⁻ → PbS↓ (brown)

• Acetic acid: Dissolves lead salts
• KCN: Masks Cu²⁺, Fe³⁺, Zn²⁺ interference
• Na₂S: Precipitating agent (forms brown PbS)

---

Q6. Role of Citric Acid, Thioglycolic Acid, and Ammonia in Iron Limit Test

Iron Limit Test Principle

Fe³⁺ + Thioglycolic acid → Purple-Red complex (ferric thioglycolate)

Role of Each Reagent

Steps:
1. Citric acid → keeps iron dissolved (solubilizer/masking agent)
2. Thioglycolic acid + Fe³⁺ → Fe²⁺ + purple-red complex
3. Ammonia → adjusts pH to alkaline (colour development)

---

Q7. Role of Acetic Acid and Ammonia in Limit Test for Heavy Metals

Heavy Metals Limit Test Principle

Role of Reagents

• Acidifies the solution to pH 3.5
• Dissolves the sample (converts insoluble salts to acetate form)
• Prevents interference from carbonates and phosphates
• Keeps metals in ionic form for reaction
• Provides the correct pH for thioacetamide reaction

• Used to adjust pH to slightly acidic range (pH 3.5) when combined with acetic acid
• Acts as buffer with acetic acid (acetate buffer)
• Prevents false positives by controlling pH precisely
• Too acidic → no precipitation; too alkaline → hydroxide precipitation (false positive)

Buffer system:
CH₃COOH + NH₃ → CH₃COONH₄ (pH ≈ 3.5)

At pH 3.5:
Heavy metal ions + Thioacetamide → Metal sulphide (brown)
S²⁻ + Pb²⁺ → PbS↓ (brown)

---

Q8. Basis for Fixing Limits for Impurities

Basis for Impurity Limits

BASIS FOR FIXING IMPURITY LIMITS
        |
        ├── 1. TOXICOLOGICAL BASIS
        |       Based on the known toxic dose of the impurity
        |       Limit = fraction of toxic dose / daily dose of drug
        |
        ├── 2. PHARMACOPOEIAL BASIS
        |       IP, BP, USP specify limits based on
        |       - Manufacturing feasibility
        |       - Analytical sensitivity
        |
        ├── 3. THERAPEUTIC USE
        |       Higher purity required for injectables,
        |       ophthalmic preparations than oral tablets
        |
        ├── 4. STABILITY STUDIES
        |       Some impurities increase on storage
        |       Limits account for maximum at shelf-life
        |
        ├── 5. ANALYTICAL METHOD SENSITIVITY
        |       Limit can't be set below detection limit
        |       of the test method
        |
        └── 6. MANUFACTURING PROCESS
                Economic feasibility of achieving purity
                Current Good Manufacturing Practices (cGMP)

The amount of impurity present in the maximum daily dose should be less than the amount known to cause any pharmacological or toxic effect.

---

UNIT - VI: PHARMACEUTICAL INORGANIC COMPOUNDS

---

Q1. Preparation and Uses of Aluminium Hydroxide Gel

Chemical Formula

Preparation

AlCl₃ + 3NaOH → Al(OH)₃↓ + 3NaCl
(Aluminium chloride) + (Sodium hydroxide) → (Aluminium hydroxide gel)

Al₂(SO₄)₃ + 6NH₄OH → 2Al(OH)₃↓ + 3(NH₄)₂SO₄

• React aluminium sulphate with sodium carbonate
• Wash precipitate thoroughly to remove sulphate
• Suspend in purified water to form gel
• Contain 3.5-4.5% Al(OH)₃ w/v (IP specification)

Uses

---

Q2. Requirements for Ideal Antacids

IDEAL ANTACID PROPERTIES
        |
        ├── 1. RAPID ONSET OF ACTION
        |       Neutralizes acid quickly
        |
        ├── 2. SUSTAINED ACTION
        |       Prolonged buffering effect (pH 3-5 maintained)
        |
        ├── 3. NON-ABSORBABLE
        |       Should not be absorbed systemically
        |       (prevents systemic alkalosis)
        |
        ├── 4. NO CO₂ PRODUCTION
        |       Carbonates produce CO₂ causing belching
        |       Ideal antacid should not produce gas
        |
        ├── 5. NO CONSTIPATION OR DIARRHOEA
        |       Al(OH)₃ causes constipation
        |       Mg(OH)₂ causes diarrhoea
        |       Combination used to balance
        |
        ├── 6. NO DRUG INTERACTIONS
        |       Should not chelate or adsorb other drugs
        |
        ├── 7. PALATABILITY
        |       Pleasant taste and easy to take
        |
        ├── 8. ECONOMICAL
        |       Low cost, widely available
        |
        └── 9. ACID NEUTRALIZING CAPACITY (ANC)
                High ANC (>5 mEq/dose) preferred

---

Q3. Composition and Uses of Talc and Kaolin

Talc

1. Dusting powder - base for medicated powders
2. Lubricant in tablets - reduces friction in punch/die
3. Glidant - improves powder flow in tablet manufacturing
4. Filtering aid - clarifies solutions
5. Cosmetics - talcum powder base

---

Kaolin (China Clay)

1. Antidiarrhoeal - adsorbs toxins, bacteria, and irritants from GI tract
2. Kaolin-Morphine mixture - official preparation for diarrhoea
3. Poultice - in kaolin poultice (BPC) for local heat/drawing action
4. Adsorbent - adsorbs drugs and toxins
5. Tablet filler - diluent in tablet formulations

---

Q4. Physiological Role of Copper and Chromium

Copper (Cu)

---

Chromium (Cr)

Chromium (Cr³⁺) + Nicotinic acid + Glutathione = GTF (Glucose Tolerance Factor)
GTF → potentiates insulin → enhances glucose uptake → normal blood sugar

---

UNIT - VII: BODY FLUIDS AND ELECTROLYTES

---

Q1. Physiological Acid-Base Balance

Definition

ACID-BASE BALANCE OVERVIEW

Normal pH = 7.35-7.45
Acidosis = pH < 7.35
Alkalosis = pH > 7.45

Buffer Systems Maintaining pH

1. BICARBONATE BUFFER SYSTEM (most important in blood)
   H₂CO₃ ⇌ H⁺ + HCO₃⁻
   Ratio HCO₃⁻ : H₂CO₃ = 20:1 (at pH 7.4)

2. PHOSPHATE BUFFER SYSTEM (important in kidney/cells)
   H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻

3. PROTEIN BUFFER SYSTEM
   Haemoglobin and plasma proteins

4. RESPIRATORY REGULATION
   ↑CO₂ → ↑ventilation → ↓pCO₂ → normalize pH

5. RENAL REGULATION (slower)
   Kidney excretes H⁺ or HCO₃⁻ to adjust pH

---

Q2. Functions of Bicarbonate and Sodium Ions

Bicarbonate (HCO₃⁻)

• Sodium bicarbonate (NaHCO₃) given in metabolic acidosis
• NaHCO₃ used to alkalinize urine in salicylate poisoning

---

Sodium (Na⁺) Ions

• Hyponatraemia (< 135 mEq/L): Oedema, brain swelling, confusion
• Hypernatraemia (> 145 mEq/L): Dehydration, CNS effects

---

Q3. Electrolytes Used in Replacement Therapy

Definition

REPLACEMENT FLUIDS
        |
        ├── CRYSTALLOIDS (ionic solutions)
        |    ├── Normal Saline (0.9% NaCl)
        |    ├── Ringer's Lactate (Na⁺, K⁺, Ca²⁺, Cl⁻, Lactate)
        |    └── Dextrose Saline
        |
        └── COLLOIDS (large molecules)
             ├── Dextran
             └── Plasma protein solutions

---

Q4. Composition and Uses of ORS (Oral Rehydration Salt)

WHO-ORS Composition (New Formula 2002)

• Na⁺ = 75 mEq/L
• K⁺ = 20 mEq/L
• Cl⁻ = 65 mEq/L
• Citrate = 10 mmol/L
• Glucose = 75 mmol/L
• Osmolarity = 245 mOsm/L (hypotonic - reduces stool output)

Uses

1. Acute diarrhoea (most important use)
2. Cholera treatment
3. Dehydration due to vomiting
4. Heat exhaustion - electrolyte replacement
5. Post-surgical fluid maintenance

---

Q5. Buffer Systems of the Body

BICARBONATE BUFFER (Most important):
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻
                    ↑ Regulated by lungs and kidneys

Henderson-Hasselbalch equation:
pH = 6.1 + log [HCO₃⁻] / [H₂CO₃]
pH = 6.1 + log 20/1 = 6.1 + 1.3 = 7.4

---

Q6. Preparation and Assay of Magnesium Sulphate

Magnesium Sulphate (MgSO₄·7H₂O)

Preparation

Method 1: MgO + H₂SO₄ → MgSO₄ + H₂O
Method 2: Mg(OH)₂ + H₂SO₄ → MgSO₄ + 2H₂O
Method 3: MgCO₃ + H₂SO₄ → MgSO₄ + H₂O + CO₂

Assay (Complexometric Method - IP)

1. Dissolve accurately weighed sample (~0.3 g) in water
2. Add 10 mL of ammonia buffer pH 10
3. Add 2 drops of EBT indicator (wine red colour)
4. Titrate with 0.05M disodium EDTA
5. Endpoint: Wine red → Blue

Mg²⁺ + Na₂H₂Y → Na₂MgY + 2H⁺
(at pH 10, ammonia buffer)

1 mL of 0.05M EDTA = 0.01232 g MgSO₄·7H₂O
% MgSO₄ = (V × 0.05 × 246.5 × 100) / (1000 × W)

Uses

---

UNIT - VIII: ANTIMICROBIALS

---

Q1. Antimicrobials - Definition, Mode of Action, Official Antimicrobials

Definition

Mode of Action

MECHANISMS OF ANTIMICROBIAL ACTION
        |
        ├── 1. CELL WALL DISRUPTION
        |       Break down bacterial cell wall
        |       Example: Benzalkonium chloride, Chlorhexidine
        |
        ├── 2. PROTEIN DENATURATION
        |       Denature bacterial proteins
        |       Example: Phenol, Formaldehyde, Alcohol
        |
        ├── 3. OXIDATION
        |       Release active oxygen → oxidize cell components
        |       Example: H₂O₂, KMnO₄, Halogens (Cl₂, I₂)
        |
        ├── 4. MEMBRANE DISRUPTION
        |       Disrupt cytoplasmic membrane permeability
        |       Example: Quaternary ammonium compounds
        |
        └── 5. INHIBITION OF ENZYMES
                Block key metabolic enzymes
                Example: Heavy metals (Hg, Ag)

---

Official Antimicrobials (IP)

---

Q2. Mechanism of Action of Antimicrobial Agents

Cl₂ + H₂O → HOCl (hypochlorous acid) → nascent O → oxidizes cell proteins

Ag⁺ + SH groups of proteins → Ag-S-protein (inactive enzymes → cell death)

Benzalkonium chloride → inserts into lipid membrane → disrupts membrane integrity → leakage of cell contents

Formaldehyde → alkylates NH₂, OH, SH groups of proteins → irreversible protein denaturation

---

Q3. Preparation and Uses of Chlorinated Lime

Preparation

Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O
(Slaked lime) + (Chlorine gas) → (Chlorinated lime)

• Dry slaked lime [Ca(OH)₂] is exposed to chlorine gas in rotating cylinders
• Temperature kept below 40°C (higher T → CaCl₂ + Ca(ClO₃)₂ formed)
• Product contains about 30-35% available chlorine

Mechanism of Action

CaOCl₂ + CO₂ + H₂O → CaCO₃ + 2HOCl
HOCl → HCl + [O] (nascent oxygen)
Nascent O → oxidizes bacterial protoplasm → kills bacteria

Uses

---

Q4. Role of Chloroform in Assay of Potassium Iodide

Step 1: KIO₃ + 5KI + 6HCl → 3I₂ + 6KCl + 3H₂O
        (KI is oxidized to I₂)

Step 2: I₂ partitions into chloroform layer
        I₂ (aqueous) → I₂ (CHCl₃ layer) - brown/violet colour

Step 3: Titrate I₂ in CHCl₃ with Na₂S₂O₃ until colour disappears

1. Extraction solvent - extracts I₂ from aqueous phase (CHCl₃ has high affinity for I₂)
2. Indicator function - I₂ in CHCl₃ gives violet/brown colour (detection of I₂)
3. Removes interference - organic impurities that dissolve in CHCl₃ are separated
4. Endpoint clarification - appearance/disappearance of violet colour in CHCl₃ layer gives sharp endpoint

---

Q5. Physiological Role of Copper and Chromium

---

UNIT - IX: DENTAL AND ZINC COMPOUNDS

---

Q1. Role of Fluorides in Dental Cement

Fluoride Mechanism in Teeth

FLUORIDE ACTION ON TEETH

1. ANTI-CARIES (Prevention of tooth decay):
   Ca₁₀(PO₄)₆(OH)₂ + 2F⁻ → Ca₁₀(PO₄)₆F₂ + 2OH⁻
   (Hydroxyapatite)           (Fluorapatite - harder, acid-resistant)

2. REMINERALIZATION:
   Fluoride promotes deposition of Ca²⁺ and PO₄³⁻
   back into demineralized enamel

3. ANTI-BACTERIAL:
   Inhibits enolase enzyme in Streptococcus mutans
   → Reduces acid production → Less enamel erosion

4. TOPICAL FLUORIDE in dental cement:
   GIC (Glass Ionomer Cement) releases F⁻ slowly
   → Sustained protection of adjacent teeth

Fluoride-Containing Dental Materials

---

Q2. Composition and Uses of Zinc Eugenol Cement

Composition

• Zinc oxide (ZnO) - 80-90%
• White rosin (resin) - 10%
• Zinc stearate (accelerator) - trace

• Eugenol (4-allyl-2-methoxyphenol) - from clove oil - 85%
• Olive oil - 15% (plasticizer)

Setting Reaction

ZnO + Eugenol → Zinc eugenolate chelate (amorphous matrix)
     (slow reaction over 24 hours)
     Water catalyzes the setting reaction

Uses

---

Q3. Fluorides Used in Dental Products

---

Q4. Official Zinc Compounds

Detailed: Zinc Oxide (ZnO)

Method 1 (French Process - indirect):
Zn metal → heated in air → Zn vapour + O₂ → ZnO (white powder)

Method 2 (Direct / American Process):
Zinc ore (ZnS/ZnCO₃) → roasted → ZnO formed directly

1. Dissolve ZnO in HCl, neutralize with ammonia
2. Add buffer pH 10, EBT indicator
3. Titrate with 0.05M EDTA until wine red → blue
4. 1 mL 0.05M EDTA = 0.004069 g ZnO

• Astringent and antiseptic in skin preparations
• Dusting powders, ointments, creams
• Sunscreen (physical UV blocker)
• Zinc oxide tape (adhesive plaster)
• Calamine lotion (with ferric oxide)

---

Q5. Pharmaceutical Aids

PHARMACEUTICAL AIDS
        |
        ├── 1. DILUENTS / FILLERS
        |       Calcium phosphate, Lactose, Talc
        |       (Bulk up tablet to manageable size)
        |
        ├── 2. BINDERS
        |       Acacia, Gelatin, Starch
        |       (Hold tablet ingredients together)
        |
        ├── 3. LUBRICANTS
        |       Magnesium stearate, Talc, Stearic acid
        |       (Reduce friction during tablet compression)
        |
        ├── 4. DISINTEGRANTS
        |       Starch, Sodium starch glycolate
        |       (Help tablet break apart in GI tract)
        |
        ├── 5. ANTIOXIDANTS
        |       Sodium metabisulphite, Ascorbic acid
        |       (Prevent oxidation of formulation)
        |
        ├── 6. PRESERVATIVES
        |       Benzalkonium chloride, Thimerosal
        |       (Prevent microbial growth)
        |
        ├── 7. SUSPENDING AGENTS
        |       Bentonite, Kaolin, Silica
        |       (Keep insoluble drugs suspended)
        |
        └── 8. COLOURING / FLAVOURING AGENTS
                Titanium dioxide, tartrazine
                (Appearance and palatability)

---

UNIT - X: RADIOPHARMACEUTICALS AND ANTIDOTES

---

Q1. Clinical Applications of Radiopharmaceuticals

Four Clinical Applications:

• Diagnosis: Thyroid uptake scan, thyroid imaging
• Therapy: Treatment of hyperthyroidism and thyroid cancer
• Basis: Thyroid actively concentrates iodine

• Technetium-99m methylene diphosphonate
• Diagnose bone metastases, fractures, osteomyelitis
• High sensitivity before X-ray changes appear

• Myocardial perfusion imaging
• Diagnose coronary artery disease, myocardial infarction
• Stress testing and rest imaging

• Fluorine-18 labelled fluorodeoxyglucose
• Cancer cells have higher glucose uptake → concentrate FDG
• Used in cancer staging and monitoring therapy response

---

Q2. Definitions of Key Terms

(a) Observed Activity

Observed activity = Raw count rate from detector

(b) Absolute Activity

• Background radiation
• Detector efficiency
• Geometry factor
• Dead time correction

Absolute Activity = Observed Activity / (Detector efficiency × Geometry factor)

(c) Nuclides

• Stable nuclides - do not undergo radioactive decay
• Radionuclides - unstable, undergo radioactive decay

(d) Isotopes

Example: Iodine isotopes
¹²⁷I (stable) - 53 protons, 74 neutrons
¹³¹I (radioactive) - 53 protons, 78 neutrons
¹²³I (radioactive) - 53 protons, 70 neutrons
All are ISOTOPES (same Z = 53, different N)

---

Q3. Antidotes - Definition, Official Antidote Compounds

Definition

MECHANISMS OF ANTIDOTE ACTION
        |
        ├── 1. CHEMICAL ANTAGONISM
        |       Antidote reacts with poison chemically
        |       Example: Sodium thiosulphate + CN⁻ → SCN⁻
        |
        ├── 2. PHYSIOLOGICAL/PHARMACOLOGICAL ANTAGONISM
        |       Antidote opposes effect of poison
        |       Example: Atropine (anticholinergic) vs Organophosphates
        |
        ├── 3. CHELATION
        |       Antidote chelates heavy metal ions
        |       Example: EDTA for lead poisoning
        |
        ├── 4. COMPETITIVE INHIBITION
        |       Competes for same receptor
        |       Example: Naloxone for opioid overdose
        |
        └── 5. METABOLIC BLOCKADE
                Blocks toxic metabolite formation
                Example: Ethanol for methanol poisoning

---

Official Antidote Compounds

1. Sodium thiosulphate - cyanide
2. Potassium permanganate - oxidizes alkaloids
3. Calcium gluconate - fluoride poisoning
4. Sodium EDTA - lead and cadmium
5. Activated charcoal - general antidote

---

QUICK REVISION SUMMARY CARD

---

Key Formulas to Remember

N₁V₁ = N₂V₂ (Normality equation for all titrations)

Normality = (Wt. in g × 1000) / (Eq. wt. × Volume in mL)

% w/v = (Wt. of solute in g × 100) / Volume of solution in mL

Gravimetric factor = Mol. wt. of analyte / Mol. wt. of precipitate form

Henderson-Hasselbalch: pH = pKa + log [A⁻/HA]
Blood pH = 6.1 + log [HCO₃⁻/H₂CO₃] = 7.4 (normal)

---

Important Reactions at a Glance

---

Top Mnemonics

• OSPRC for Primary standard properties: Obtainable pure, Stable, Pure, Reacts stoichiometrically, Commercially available
• OIL RIG: Oxidation Is Loss, Reduction Is Gain (electrons)
• EBT indicator: "End = Blue Transition" (wine red → blue at endpoint with EDTA)
• Gutzeit's: "Zinc + acid → H₂ → AsH₃ → Yellow stain"
• Buffer 20:1: Blood HCO₃⁻ : H₂CO₃ = 20:1 → pH 7.4

---