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📘 PHARMACEUTICAL CHEMISTRY - COMPLETE ANSWER GUIDE

(Units 1-10 | 10 Marks Each | University Level)

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✅ UNIT 1 - GENERAL PRINCIPLES OF VOLUMETRIC ANALYSIS

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Q1. Discuss the Different Methods to Minimize Errors

Definition of Error

Error = Difference between the observed value and the true value of a measurement.

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Types of Errors and Methods to Minimize Them

ERRORS IN ANALYSIS
│
├── 1. DETERMINATE (Systematic) Errors ──→ Can be detected & corrected
│       ├── Instrumental errors
│       ├── Operative errors
│       └── Errors due to reagents
│
└── 2. INDETERMINATE (Random) Errors ──→ Cannot be fully eliminated, only minimized

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Methods to Minimize Errors

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Flowchart: Steps to Minimize Error

Identify type of error
        ↓
   Systematic?       Random?
       ↓                ↓
Calibrate, blank   Repeat determinations
  correction       Take mean & average
       ↓                ↓
  Personal error?   Instrumental?
       ↓                ↓
  Train analyst     Use calibrated
  Follow SOP        glassware
       ↓
REPORT RESULT WITH STANDARD DEVIATION

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Q2. Primary & Secondary Standard Substances + Standardization of 0.1 N Perchloric Acid

Primary Standard Substance

• Has high purity (>99.9%)
• Is stable (does not absorb moisture, CO₂)
• Has a high molecular weight (to reduce weighing error)
• Is readily available and inexpensive

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Secondary Standard Substance

• Less stable, less pure
• Must be standardized before use

• HCl, NaOH, KMnO₄, AgNO₃ (when not of high purity), Iodine

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Standardization of 0.1 N Perchloric Acid (HClO₄)

1. Weigh accurately ~500 mg of KHP (previously dried at 120°C for 2 hrs)
2. Dissolve in 20 mL of glacial acetic acid
3. Add 2 drops of crystal violet indicator
4. Titrate with 0.1 N HClO₄ in glacial acetic acid
5. Endpoint: Blue-violet → blue-green

Normality of HClO₄ = Weight of KHP (g) × 1000
                    ─────────────────────────────
                    Equivalent weight × Volume (mL)

= (0.500 × 1000) / (204.2 × V mL)

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Q3. Importance of Significant Figures

Definition

Rules for Significant Figures

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Importance in Pharmacy

SIGNIFICANT FIGURES MATTER BECAUSE:
        ↓
┌───────────────────────────────────────┐
│ 1. Convey PRECISION of measurement   │
│ 2. Avoid false precision in results  │
│ 3. Minimize calculation errors       │
│ 4. Important in drug dosage calc.    │
│ 5. Required in pharmacopoeial tests  │
│ 6. Critical in gravimetric analysis  │
└───────────────────────────────────────┘

• Multiplication/Division → Answer has as many sig. figs as the least precise number
• Addition/Subtraction → Answer rounded to least number of decimal places

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Q4. Methods to Express Concentration of Solution

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Q5. Types of Errors in Volumetric Analysis and Ways to Minimize Them

Classification of Errors

ERRORS IN VOLUMETRIC ANALYSIS
│
├── A. DETERMINATE (Systematic) Errors
│     ├── 1. Instrumental Errors
│     │       - Uncalibrated burette/pipette
│     │       - Poorly graduated glassware
│     ├── 2. Operative (Personal) Errors
│     │       - Parallax error in reading meniscus
│     │       - Overtitration or undertitration
│     │       - Incorrect endpoint detection
│     └── 3. Errors due to Reagents
│             - Impure reagents
│             - Incorrect concentration
│
└── B. INDETERMINATE (Random) Errors
      - Uncontrollable fluctuations
      - Minimized by repetition

Minimization Strategies

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Q6. The Term "mEq" - Calculation of mEq of NaCl in 1 Litre of 0.76% Solution

Definition of milliequivalent (mEq)

1 milliequivalent (mEq) = 1/1000th of an equivalent of a substance

mEq = (Weight in mg) / Equivalent weight (in mg/mEq)

Equivalent weight = Molecular weight / Valency

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For NaCl:

• Molecular weight of NaCl = 58.5 g/mol
• Valency (charge) = 1
• Equivalent weight = 58.5/1 = 58.5 g/Eq = 58.5 mg/mEq

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Calculation:

Weight of NaCl in 1 litre = 7.6 g = 7600 mg

mEq of NaCl = Weight (mg) / Equivalent weight (mg/mEq)

            = 7600 / 58.5

            = 129.9 ≈ 130 mEq

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✅ UNIT 2 - ACID-BASE AND REDOX TITRATIONS

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Q1. Theories of Neutralization Indicators

What is a Neutralization Indicator?

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Theory 1: OSTWALD'S THEORY (Ionic Theory)

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Theory 2: CHROMOPHORE THEORY (Quinonoid Theory)

PRINCIPLE:
- Indicator exists in two structural forms
- Acidic form (benzenoid) ──→ one color
- Basic form (quinonoid) ──→ different color

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Theory 3: KOLTHOFF'S ADSORPTION THEORY

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Q2. Ostwald's Theory for Acid-Base Indicators

Statement

An acid-base indicator is a weak acid or weak base that has different colors in its ionized and un-ionized forms.

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For an Indicator that is a Weak Acid (HIn):

HIn ⇌ H⁺ + In⁻
(Acid form)   (Base form)
  Color 1       Color 2

KIn = [H⁺][In⁻] / [HIn]

Taking log:

pH = pKIn + log [In⁻]/[HIn]

Color Change Range:

• When [In⁻]/[HIn] = 1/10 → pH = pKIn - 1 (acid color visible)
• When [In⁻]/[HIn] = 10/1 → pH = pKIn + 1 (base color visible)

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Flowchart of Color Change:

pH increases (adding alkali)
        ↓
[In⁻] increases
        ↓
At pH = pKIn - 1: Acid color visible
        ↓
At pH = pKIn: Equal concentrations → Intermediate color
        ↓
At pH = pKIn + 1: Base color visible

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Examples:

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Q3. Iodometry and Iodimetry

Iodimetry (Direct Iodine Titration)

C₆H₈O₆ + I₂ → C₆H₆O₆ + 2HI
(Ascorbic acid)     (Dehydroascorbic acid)

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Iodometry (Indirect Iodine Titration)

Step 1: Cu²⁺ + I⁻ → CuI + I₂ (iodine liberated)

Step 2: I₂ + 2Na₂S₂O₃ → 2NaI + Na₂S₄O₆
(Liberated iodine titrated with sodium thiosulphate)

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Key Differences:

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Q4. Types of Redox Titrations

Definition

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Types:

1. Permanganometry

• Titrant: KMnO₄ (self-indicator - purple/violet)
• Medium: Acidic (H₂SO₄)
• Reaction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
• Example: Assay of FeSO₄, H₂O₂, oxalic acid
• Endpoint: Permanent pale pink (excess KMnO₄)

2. Dichromatometry

• Titrant: K₂Cr₂O₇
• Medium: Acidic
• Reaction: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O
• Indicator: Diphenylamine or potentiometric
• Example: Assay of iron ore, FeSO₄

3. Cerimetry

• Titrant: Ceric ammonium sulphate (Ce⁴⁺)
• Reaction: Ce⁴⁺ + e⁻ → Ce³⁺
• Advantage: Stable, accurate, used in presence of HCl
• Example: Assay of arsenious oxide, iron

4. Iodimetry

• Titrant: Standard I₂ in KI solution
• Example: Vitamin C assay (as above)

5. Iodometry

• Indirect: Oxidant liberates I₂ from KI → titrated with Na₂S₂O₃
• Example: Cu²⁺ estimation

6. Bromatometry

• Titrant: KBrO₃ (in presence of KBr)
• Reaction: BrO₃⁻ + 6H⁺ + 6e⁻ → Br⁻ + 3H₂O
• Example: Assay of phenol, 8-hydroxyquinoline

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Summary Table:

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Q5. Primary Standards for Redox Titration

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Q6. Preparation and Storage of Volumetric Solution of Iodine

Preparation of 0.1 M Iodine Solution (I₂/KI)

1. Dissolve 14 g KI in 20 mL distilled water
2. Add 12.7 g of iodine (I₂) in small portions with shaking
3. Stir until completely dissolved
4. Transfer to 1000 mL volumetric flask
5. Make up to volume with water

As₂O₃ + 2I₂ + 2H₂O → As₂O₅ + 4HI
Indicator: Starch (turns blue at endpoint)

STORAGE OF IODINE SOLUTION
│
├── Store in AMBER-COLORED glass bottles
│     (protects from light decomposition)
│
├── Use GLASS stopper (not rubber - I₂ reacts with rubber)
│
├── Store in COOL, DARK place
│
├── Avoid HEAT (I₂ is volatile)
│
└── RESTANDARDIZE frequently (concentration changes with time)

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Q7. Nernst Equation and its Importance in Redox Titration

Nernst Equation:

E = E° - (RT/nF) × ln [Products] / [Reactants]

At 25°C:

E = E° - (0.0592/n) × log [Oxidized form] / [Reduced form]

Explanation of Terms:

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Importance in Redox Titrations:

IMPORTANCE OF NERNST EQUATION
│
├── 1. Predicts FEASIBILITY of a redox reaction
│       (if ΔE is positive, reaction proceeds)
│
├── 2. Calculates ACTUAL electrode potential at any concentration
│
├── 3. Determines the ENDPOINT of titration
│       (sharp change in E at equivalence point)
│
├── 4. Helps in POTENTIOMETRIC TITRATIONS
│       (plotting E vs volume → endpoint)
│
├── 5. Explains effect of pH on oxidation potential
│       (KMnO₄ works better in acid medium)
│
└── 6. Calculates EQUILIBRIUM CONSTANT
        K = 10^(nΔE°/0.0592)

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Example: Fe²⁺ / Fe³⁺ System

Fe³⁺ + e⁻ → Fe²⁺     E° = +0.77 V

E = 0.77 - (0.0592/1) × log [Fe²⁺]/[Fe³⁺]

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✅ UNIT 3 - NON-AQUEOUS TITRATIONS & PRECIPITATION TITRATIONS

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Q1. Non-Aqueous Titration - Types of Solvents and Pharmaceutical Applications

Definition

Non-aqueous titration (NAT) is a titration carried out in a non-aqueous solvent to determine substances that are too weak to be titrated in water, or those that are insoluble in water.

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Why Non-Aqueous?

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Classification of Solvents Used in NAT:

SOLVENTS IN NAT
│
├── A. PROTOPHILIC SOLVENTS (basic)
│       - Accept protons from solute
│       - Used for titrating WEAK ACIDS
│       - Examples: Pyridine, acetone, dimethylformamide (DMF)
│
├── B. PROTOGENIC SOLVENTS (acidic)
│       - Donate protons to solute
│       - Used for titrating WEAK BASES
│       - Examples: Glacial acetic acid, formic acid, sulfuric acid
│
├── C. AMPHIPROTIC SOLVENTS
│       - Act as both acid and base
│       - Examples: Ethanol, methanol, isopropanol
│
└── D. APROTIC (INERT) SOLVENTS
        - Neither donate nor accept protons
        - Dissolve substances but don't participate
        - Examples: Benzene, chloroform, acetonitrile

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Most Used Solvent in NAT: Glacial Acetic Acid

• Acidic solvent → enhances basicity of weak bases
• Makes weak amines appear stronger
• Used with 0.1 N HClO₄ as titrant
• Indicator: Crystal violet (violet → blue-green)

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Applications in Pharmacy:

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Q2. Classification of Solvents in Non-Aqueous Titration

Additional Details:

Leveling effect: Strong solvents make all acids/bases appear equally strong. Differentiating effect: Weak solvents allow distinction between acids/bases of different strengths.

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Q3. Principle of NAT + Preparation and Assay of Sodium Benzoate

Principle of Non-Aqueous Titration:

Base (B) + HClO₄ → BH⁺ + ClO₄⁻

Glacial acetic acid enhances ionization:
B + CH₃COOH → BH⁺ + CH₃COO⁻ (base appears stronger)

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Assay of Sodium Benzoate (C₆H₅COONa, MW = 144.1)

C₆H₅COONa + HClO₄ → C₆H₅COOH + NaClO₄

• Primary standard used: Potassium hydrogen phthalate (KHP) to standardize HClO₄

1. Accurately weigh ~150 mg of Sodium Benzoate
2. Dissolve in 20 mL of glacial acetic acid (warming if needed)
3. Add 2 drops of crystal violet indicator
4. Titrate with 0.1 N HClO₄ in glacial acetic acid
5. Endpoint: Violet → blue-green (or green)
6. Each mL of 0.1 N HClO₄ = 14.41 mg of C₆H₅COONa

% Purity = (V × N × Equivalent weight × 100) / (Weight taken × 1000)

Equivalent weight of sodium benzoate = 144.1/1 = 144.1

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Q4. Precipitation Titration - Principle and Assay of NaCl

Definition

Titrations where the reaction produces an insoluble precipitate. Based on formation of slightly soluble salts.

Classification:

PRECIPITATION TITRATIONS
│
├── 1. ARGENTOMETRIC TITRATIONS (using AgNO₃)
│       ├── Mohr's method
│       ├── Volhard's method
│       └── Fajan's method
│
└── 2. MERCURIMETRIC TITRATIONS (using Hg(NO₃)₂)

Principle:

Ag⁺ + Cl⁻ → AgCl↓ (white precipitate)
Ksp of AgCl = 1.8 × 10⁻¹⁰ (very low → reaction complete)

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Assay of NaCl (Sodium Chloride) by Mohr's Method:

NaCl + AgNO₃ → AgCl↓ (white) + NaNO₃

At endpoint:
2AgNO₃ + K₂CrO₄ → Ag₂CrO₄↓ (brick-red) + 2KNO₃

1. Dissolve accurately weighed NaCl (~100 mg) in 50 mL water
2. Adjust pH to 6.5-10 (neutral to slightly alkaline)
3. Add 1 mL of 5% K₂CrO₄ indicator
4. Titrate with 0.1 N AgNO₃ solution
5. Endpoint: Yellow → brick-red/orange (permanent Ag₂CrO₄ precipitate)
6. Each mL of 0.1 N AgNO₃ = 5.844 mg NaCl

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Q5. Volhard's Method + Fajan's Indicator Mechanism

Volhard's Method (Back Titration):

Step 1: Ag⁺ + Cl⁻ → AgCl↓ (excess Ag⁺ added)

Step 2 (back titration):
Ag⁺(excess) + SCN⁻ → AgSCN↓ (white)

At endpoint:
Fe³⁺ + SCN⁻ → [Fe(SCN)]²⁺ (blood red color)

• Done in acidic medium (HNO₃) - prevents hydrolysis of Fe³⁺
• For chloride estimation: filter off AgCl precipitate first (or add nitrobenzene) to prevent back reaction AgCl + SCN⁻ → AgSCN + Cl⁻

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Fajan's Method - Mechanism of Adsorption Indicator:

BEFORE ENDPOINT:
AgCl↓ + Cl⁻ (excess) → AgCl·Cl⁻ surface (negative charge)
                          ↓
                  Indicator anion repelled from surface
                          ↓
                  Solution remains YELLOW-GREEN (fluorescein)

AT/AFTER ENDPOINT:
Excess Ag⁺ → AgCl·Ag⁺ surface (positive charge)
                          ↓
            Indicator anion (Fl⁻) ADSORBED on surface
                          ↓
        Precipitate turns PINK (AgCl·Ag·Fl complex)

• Neutral or slightly acidic pH (4-7)
• Use dextrin to prevent coagulation of precipitate
• Avoid bright sunlight (photodecomposition)

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Q6. Mohr's Method for Estimation of Halides

Principle:

Direct titration of halide (Cl⁻, Br⁻) with standard AgNO₃ using potassium chromate (K₂CrO₄) as indicator.

Reactions:

Primary reaction:
Ag⁺ + Cl⁻ → AgCl↓ (Ksp = 1.8 × 10⁻¹⁰)

Indicator reaction (at endpoint):
2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄↓ (Ksp = 1.1 × 10⁻¹²)
         (brick-red precipitate)

Conditions:

• pH must be 6.5 to 10 (neutral to slightly alkaline)
  • Too acidic: CrO₄²⁻ → Cr₂O₇²⁻ (indicator doesn't work)
  • Too alkaline: Ag₂O precipitates
• Indicator: 5% K₂CrO₄ (1-2 mL)

Procedure:

Sample (halide solution)
    ↓
Adjust pH 6.5-10 (with Na₂CO₃ or dilute HNO₃)
    ↓
Add K₂CrO₄ indicator (yellow solution)
    ↓
Titrate with standard 0.1 N AgNO₃
    ↓
ENDPOINT: Persistent brick-red/orange color
    ↓
Record volume → Calculate %Cl⁻

Limitations:

• Only for Cl⁻ and Br⁻ (NOT iodide or thiocyanate)
• Cannot use in acidic or strongly alkaline medium
• Interference by Pb²⁺, Ba²⁺ (precipitate chromate)

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Q7. Mohr's, Volhard's and Fajan's Methods - Comparison Table

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✅ UNIT 4 - COMPLEXOMETRIC TITRATIONS

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Q1. Complexometric Titrations - Types and Estimation of Calcium Gluconate

Definition

Titrations based on the formation of a stable, soluble complex between a metal ion and a complexing agent (ligand).

Types of Complexometric Titrations:

COMPLEXOMETRIC TITRATIONS
│
├── 1. DIRECT TITRATION
│       - Metal ion titrated directly with EDTA
│       - Example: Ca²⁺, Mg²⁺ titrated with Na₂EDTA
│
├── 2. BACK TITRATION
│       - Excess EDTA added; unused EDTA back-titrated
│         with standard metal salt
│       - Used when: Direct endpoint not sharp / metal reacts slowly
│       - Example: Al³⁺ (add excess EDTA, back-titrate with ZnSO₄)
│
├── 3. INDIRECT TITRATION
│       - Anion precipitated with excess metal → precipitate
│         dissolved in acid → metal titrated with EDTA
│       - Example: SO₄²⁻ precipitated as BaSO₄
│
├── 4. DISPLACEMENT (Substitution) TITRATION
│       - Metal displaces another metal from a weaker complex
│       - Displaced metal titrated with EDTA
│       - Example: Ca²⁺ displaces Mg from MgY² complex
│
└── 5. POTENTIOMETRIC TITRATION
        - Endpoint detected by metal-sensitive electrode

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Estimation of Calcium Gluconate by Complexometry:

Ca²⁺ + H₂Y²⁻ → CaY²⁻ + 2H⁺
(EDTA anion)  (stable complex)

1. Dissolve ~400 mg of calcium gluconate in 100 mL water
2. Add 5 mL of 2M NaOH to adjust pH to ~12
3. Add a pinch of murexide indicator (pink/red color)
4. Titrate with 0.05 M disodium EDTA
5. Endpoint: Pink/Red → Pure violet (blue)
6. Each mL of 0.05 M EDTA = 22.42 mg of Ca(C₆H₁₁O₇)₂·H₂O

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Q2. Methods in Complexometry + Masking and Demasking Agents

Masking Agents:

A masking agent is a substance added to prevent a metal ion from reacting with the titrant (EDTA), without removing it from solution.

Demasking Agents:

Demasking is the process of releasing the masked metal so it can react with EDTA.

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Q3. Ligands and their Types

Ligand Definition:

A ligand is a molecule or ion that donates electron pairs to a central metal atom/ion to form a coordination complex.

Types of Ligands:

Why EDTA is the Best Ligand in Complexometry:

EDTA (Ethylenediaminetetraacetic acid)
- 2 nitrogen donor atoms + 4 carboxylate oxygens = 6 donor atoms
- Forms STABLE 1:1 complex with almost all metal ions
- Reaction is RAPID and COMPLETE
- Forms WATER-SOLUBLE complexes
- pH can be controlled to achieve SELECTIVITY

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Q4. Types of EDTA Titrations + Preparation of 0.05 M Disodium EDTA

Types of EDTA Titrations:

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Preparation of 0.05 M Disodium EDTA:

0.05 M → 0.05 × 372.2 = 18.61 g per litre

1. Weigh 18.61 g of disodium EDTA dihydrate
2. Dissolve in small amount of water
3. Transfer to 1000 mL volumetric flask
4. Make up to mark with distilled water
5. Mix well

Standardization of 0.05 M Disodium EDTA:

1. Weigh accurately ~100 mg CaCO₃ (dried at 110°C)
2. Dissolve in minimum HCl
3. Neutralize with NaOH
4. Add 5 mL NH₄Cl/NH₃ buffer (pH 10)
5. Add Eriochrome Black T (EBT) indicator (wine-red)
6. Titrate with 0.05 M EDTA until wine-red → pure blue
7. Repeat and take mean

Molarity of EDTA = Weight of CaCO₃ (mg) / (MW × Volume of EDTA in mL)

= 100 / (100.09 × V mL)

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✅ UNIT 5 - GRAVIMETRIC ANALYSIS AND LIMIT TESTS

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Q1. Principle of Gravimetric Analysis with Example

Definition

Gravimetric analysis is a method of quantitative analysis in which the mass (weight) of an isolated pure substance is used to calculate the quantity of the analyte.

Principle:

GRAVIMETRIC ANALYSIS PRINCIPLE
│
Sample in solution
    ↓
Add precipitating agent (excess)
    ↓
Precipitate forms (insoluble compound)
    ↓
Filter (Whatman No. 42 filter paper / sintered glass crucible)
    ↓
Wash precipitate (remove impurities)
    ↓
Dry / Ignite at specific temperature
    ↓
Weigh the final product accurately
    ↓
Calculate % analyte from weight of precipitate

Example - Estimation of Sulphate (SO₄²⁻):

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

BaSO₄ is filtered, dried, ignited at 600°C, and weighed.

% SO₄²⁻ = (Weight of BaSO₄ × Gravimetric factor × 100) / Weight of sample

Gravimetric factor for SO₄²⁻ = MW of SO₄ / MW of BaSO₄
                                = 96/233 = 0.4115

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Q2. Steps in Gravimetric Analysis + Limitations

Steps:

1. PREPARATION OF SOLUTION
   ↓
2. PRECIPITATION
   - Add precipitating agent slowly, with stirring
   - Use hot, dilute solution (for large, pure crystals)
   ↓
3. DIGESTION (aging)
   - Heat precipitate in mother liquor for 30-60 min
   - Increases crystal size, reduces surface area
   ↓
4. FILTRATION
   - Filter through proper filter medium
   ↓
5. WASHING
   - Remove adsorbed impurities
   - Use small amounts of wash liquid repeatedly
   ↓
6. DRYING / IGNITION
   - Remove moisture (100-130°C for drying)
   - Ignite at 600-1000°C to convert to oxide if needed
   ↓
7. WEIGHING
   - Cool in desiccator, weigh to constant weight
   ↓
8. CALCULATION

Limitations:

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Q3. Limit Test for Arsenic - Principle, Procedure, Diagram

Principle (Gutzeit's Method):

Step 1: Arsenate → Arsenite (reduction with KI + SnCl₂)
As⁵⁺ + 2I⁻ → As³⁺ + I₂

Step 2: Arsine generation
As³⁺ + 3Zn + 3H₂SO₄ → AsH₃↑ + 3ZnSO₄ + 3H⁺ (simplified)

Step 3: Detection
AsH₃ + HgCl₂ → AsH(HgCl)₂ (yellow stain)
or
AsH₃ + 3HgCl₂ → AsCl₃ + 3Hg + 3HCl (brown stain)

Gutzeit's Apparatus:

         ┌─────────────────────────────┐
         │   HgCl₂ PAPER (stain)      │ ← Arsine detected here
         └──────────────┬──────────────┘
                        │
         ┌──────────────┴──────────────┐
         │   LEAD ACETATE COTTON       │ ← Removes H₂S interference
         └──────────────┬──────────────┘
                        │
                ┌───────┴────────┐
                │  GLASS TUBE    │
                └───────┬────────┘
                        │
         ┌──────────────┴──────────────┐
         │   CONICAL FLASK (100 mL)    │
         │   Sample + Zn + H₂SO₄      │
         │   + KI + SnCl₂             │
         │   (H₂ generated)            │
         └─────────────────────────────┘

Procedure:

1. Take sample solution in conical flask
2. Add 1g KI and 0.5 mL SnCl₂ (to reduce As⁵⁺ → As³⁺, and remove oxidizing agents)
3. Add 3g Zinc granules
4. Add 5 mL dilute H₂SO₄ (generates H₂ which reduces As to AsH₃)
5. Assemble Gutzeit's apparatus with lead acetate cotton (traps H₂S) and HgCl₂ paper
6. Allow to react for 40 minutes
7. Compare stain intensity with standard stain (prepared from arsenic standard solution)

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Q4. Limit Test for Iron

Principle:

Fe³⁺ + thioglycolic acid + NH₃ → Purple complex
(at pH 8.5-9.5)

Reagents and Roles:

Procedure:

1. Dissolve sample in dilute HCl / HNO₃
2. Add 2 mL citric acid (prevents iron hydroxide precipitation)
3. Add 0.1 mL thioglycolic acid
4. Make alkaline with concentrated ammonia (~10 mL)
5. Dilute to 10 mL
6. Compare purple color with iron standard (2 ppm Fe)
7. Color of sample should NOT be more intense than standard

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Q5. Sources of Impurities in Pharmaceuticals + Importance of Limit Tests

Sources of Impurities:

SOURCES OF IMPURITIES IN PHARMACEUTICALS
│
├── 1. RAW MATERIAL IMPURITIES
│       - Starting material impurities
│       - Reagent impurities used in synthesis
│
├── 2. PROCESS IMPURITIES (During Manufacturing)
│       - Intermediates
│       - By-products of synthesis
│       - Degradation during processing
│
├── 3. ENVIRONMENTAL IMPURITIES
│       - Water (heavy metals, chloride, sulphate)
│       - Air (dust, microorganisms)
│
├── 4. CONTAINER/PACKAGING impurities
│       - Leachables from plastic containers
│       - Metal contamination from equipment
│
└── 5. DEGRADATION PRODUCTS
        - Light, heat, humidity-induced degradation

Importance of Limit Tests in Quality Control:

LIMIT TESTS ARE IMPORTANT BECAUSE:
│
├── ✅ Ensure SAFETY of pharmaceutical product
│     (toxic impurities like As, Pb, Hg can be harmful)
│
├── ✅ Ensure QUALITY and PURITY of drug
│
├── ✅ Mandatory compliance with PHARMACOPOEIA
│     (IP, BP, USP standards)
│
├── ✅ Protect the PATIENT from adverse effects
│
├── ✅ Determine FITNESS of raw materials
│
└── ✅ Regulatory approval requirement

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Q6. Limit Test for Sulphate and Iron

Limit Test for Sulphate:

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

1. Dissolve sample (as specified) in water
2. Add 2 mL of dilute HCl (to prevent precipitation of BaCO₃/BaHPO₄)
3. Add 2 mL of 25% BaCl₂ solution
4. Allow to stand for 5 minutes
5. Compare turbidity with standard sulphate solution (10 ppm SO₄²⁻)

---

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Q7. Limit Test for Chloride

Principle:

Cl⁻ + AgNO₃ → AgCl↓ (white opalescence) + NO₃⁻

Procedure:

Take sample (as specified in monograph)
        ↓
Dissolve in 10 mL of water (if not already solution)
        ↓
Add 1 mL of dilute HNO₃ (to prevent precipitation of Ag₂CO₃, Ag₃PO₄)
        ↓
Add 1 mL of 2% AgNO₃ solution
        ↓
Allow to stand 5 minutes (in dark - AgCl photosensitive)
        ↓
Compare OPALESCENCE (turbidity) with STANDARD
(standard = known amount of NaCl giving 10 ppm Cl⁻ in same conditions)
        ↓
Sample turbidity should NOT EXCEED standard

• HCl would add extra chloride
• H₂SO₄ would precipitate BaSO₄ (interference)

---

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✅ UNIT 6 - INORGANIC PHARMACEUTICALS - GIT AGENTS

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Q1. Labelling and Storage of Oxygen

Oxygen (O₂) - Medical Grade

• Name: Oxygen / Medical Oxygen
• "FOR MEDICAL USE ONLY"
• Cylinder colour code: Black body + White shoulder (IP)
• Batch number, manufacturing date
• Name and address of manufacturer
• "CONTAINS COMPRESSED GAS" - Handle with care
• Flow rate / purity statement (NLT 99.0%)

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Q2. Antacids - Classification, Examples, Combination Therapy

Definition

Antacids are basic substances that neutralize excess gastric hydrochloric acid, providing relief from acidity, heartburn, and peptic ulcers.

Classification:

ANTACIDS
│
├── A. SYSTEMIC ANTACIDS (Absorbable)
│       - Absorbed into bloodstream → systemic alkalosis risk
│       - Examples: Sodium bicarbonate (NaHCO₃)
│
└── B. NON-SYSTEMIC ANTACIDS (Non-absorbable) ← PREFERRED
        ├── Magnesium salts
        │     - Magnesium hydroxide [Mg(OH)₂]
        │     - Magnesium carbonate [MgCO₃]
        │     - Magnesium trisilicate [Mg₂Si₃O₈·nH₂O]
        │
        ├── Aluminium salts
        │     - Aluminium hydroxide gel [Al(OH)₃]
        │     - Dried Al(OH)₃ gel
        │     - Aluminium phosphate
        │
        └── Calcium salts
              - Calcium carbonate (CaCO₃)

Ideal Properties of Antacid:

• Rapid onset, prolonged action
• Non-absorbed / non-systemic
• No CO₂ production
• No acid rebound
• No interference with drug absorption
• Palatable, inexpensive

Antacid Combination Therapy:

1. Balanced GI motility (no constipation or diarrhea)
2. Better neutralizing capacity
3. Reduced acid rebound
4. Reduced side effects

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Q3. Nitrous Oxide - Preparation, Labelling, Storage, Uses

Chemical formula: N₂O | Mol wt: 44 | Synonym: Laughing gas

Preparation:

Thermal decomposition of ammonium nitrate:

NH₄NO₃ ──Heat (170-230°C)──→ N₂O↑ + 2H₂O

Impurities removed:
- NO, NO₂ removed by passing through FeSO₄ solution
- NH₃ removed by passing through H₂SO₄
- Water removed by CaCl₂ drying tube
- Final gas compressed into cylinders

Labelling:

• Name: Nitrous Oxide / Nitrous Oxide for Medicinal Use
• Cylinder: Blue body + Blue shoulder (IP)
• "FOR ANESTHETIC USE ONLY"
• Purity: NLT 98.0% N₂O
• Volume, pressure

Storage:

• Blue cylinders stored in cool, ventilated area
• Away from heat and flammables
• Upright position
• Separate from other gases

Medicinal Uses:

---

Q4. Antacids - Ideal Properties + Sodium Bicarbonate (NaHCO₃)

Sodium Bicarbonate (Baking Soda)

Preparation (Solvay Process):

NH₃ + CO₂ + H₂O + NaCl → NaHCO₃↓ + NH₄Cl

NaCl + NH₃ + H₂O + CO₂ → NaHCO₃ (precipitation from Solvay tower)
    ↓
Filter, dry below 50°C (heating decomposes it)

Assay (IP Method):

NaHCO₃ + HCl → NaCl + H₂O + CO₂↑
(Equivalent wt = 84)

1. Weigh ~0.25 g NaHCO₃ accurately
2. Dissolve in 50 mL water
3. Add methyl orange indicator
4. Titrate with 0.5 M HCl until endpoint (yellow → orange-red)
5. Boil to remove CO₂, cool, continue if needed
6. Each mL of 0.5 M HCl ≡ 42 mg NaHCO₃

Uses:

• Antacid (systemic - rapid action)
• Urinary alkalinizer
• IV sodium bicarbonate for metabolic acidosis
• Baking (culinary use)
• Mouthwash (alkaline)

---

Q7. Aluminium Hydroxide Gel - Preparation and Uses

Chemical formula: Al(OH)₃ | Synonyms: Dried aluminium hydroxide gel

Preparation:

Method 1 (Chemical):
AlCl₃ + 3NaOH → Al(OH)₃↓ + 3NaCl
     (freshly precipitated gel washed, stabilized)

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

Properties:

• White gelatinous precipitate/suspension
• Acts by adsorption of pepsin (reduces peptic digestion)
• Has phosphate-binding property

Uses:

---

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✅ UNIT 7 - ELECTROLYTES AND BODY FLUID MANAGEMENT

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Q1. Acid-Base Balance + Role of Na⁺ and HCO₃⁻

Mechanism of Acid-Base Balance:

ACID-BASE BALANCE MECHANISMS
│
├── 1. BUFFER SYSTEMS (Fastest - seconds)
│       - Bicarbonate buffer: H₂CO₃/HCO₃⁻ (most important in blood)
│       - Phosphate buffer: H₂PO₄⁻/HPO₄²⁻ (important in urine)
│       - Protein buffer (hemoglobin, plasma proteins)
│
├── 2. RESPIRATORY REGULATION (Minutes)
│       - ↑CO₂ → hyperventilation → pH ↑
│       - ↓CO₂ → hypoventilation → pH ↓
│       - CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ (carbonic anhydrase)
│
└── 3. RENAL REGULATION (Hours - days - most powerful)
        - Kidneys excrete H⁺ or HCO₃⁻ as needed
        - Net acid excretion (NH₄⁺, titratable acid)

Henderson-Hasselbalch Equation:

pH = pKa + log [HCO₃⁻] / [H₂CO₃]

Normal: pH = 6.1 + log (24 mEq/L) / (1.2 mEq/L) = 6.1 + log 20 = 6.1 + 1.3 = 7.4

---

Role of Na⁺ (Sodium):

Role of HCO₃⁻ (Bicarbonate):

---

Q2. Major Intra and Extracellular Electrolytes + Dextrose and NaCl Injection

Electrolyte Distribution:

---

Dextrose and Sodium Chloride Injection:

1. Dissolve 50g glucose and 9g NaCl in water for injection
2. Make up to 1000 mL
3. Add activated charcoal, filter
4. Fill into glass ampoules/bottles
5. Sterilize by autoclaving at 121°C / 15 min
6. Check pH (3.5-6.5), clarity, particulate matter

• Fluid and electrolyte replacement
• Post-surgical hydration
• Dehydration (with sodium loss)
• Vehicle for IV drugs

---

Q4 & Q5. Cathartics and Saline Cathartics

Definition:

Cathartics (purgatives/laxatives) are substances that promote evacuation of bowel contents.

Classification:

CATHARTICS
│
├── SALINE CATHARTICS (Osmotic)
│       - Magnesium sulphate (Epsom salt)
│       - Magnesium hydroxide (Milk of Magnesia)
│       - Sodium sulphate (Glauber's salt)
│
├── BULK LAXATIVES
│       - Bran, psyllium, methylcellulose
│
├── STIMULANT CATHARTICS
│       - Castor oil, senna, bisacodyl
│
└── STOOL SOFTENERS
        - Docusate sodium

Mechanism of Saline Cathartics:

Mg²⁺ / SO₄²⁻ ions (poorly absorbed)
            ↓
Osmotic pressure in intestinal lumen increases
            ↓
Water drawn from tissues into intestinal lumen
            ↓
Increased intestinal volume
            ↓
Distension of intestinal wall
            ↓
Stimulates peristalsis
            ↓
Evacuation of bowel (within 1-3 hours)

---

Q6. Magnesium Sulphate - Preparation, Assay, Uses

Magnesium Sulphate

Preparation:

MgO + H₂SO₄ → MgSO₄ + H₂O
(Magnesium oxide reacts with sulphuric acid)

Or:
Mg(OH)₂ + H₂SO₄ → MgSO₄ + 2H₂O

Solution concentrated and crystallized at below 67°C to get heptahydrate

Assay (Complexometric - IP method):

Mg²⁺ + H₂Y²⁻ → MgY²⁻ + 2H⁺

1. Dissolve ~0.4 g MgSO₄ in water
2. Add 10 mL ammonia buffer (pH 10)
3. Add Eriochrome Black T (EBT) indicator (wine-red)
4. Titrate with 0.05 M EDTA
5. Endpoint: Wine-red → Pure blue
6. Each mL 0.05 M EDTA ≡ 12.32 mg MgSO₄

Uses:

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✅ UNIT 8 - HYDROGEN PEROXIDE AND ANTIMICROBIALS

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Q1. Hydrogen Peroxide - Preparation and Assay Principle

Chemical formula: H₂O₂ | MW: 34 | Synonym: Hydrogen dioxide

Preparation:

Industrial Method (Anthraquinone Process - Major method):
Anthraquinone + H₂ → Anthrahydroquinone (reduction)
Anthrahydroquinone + O₂ → Anthraquinone + H₂O₂ (oxidation)

H₂O₂ is extracted with water and concentrated.

Laboratory Method:
BaO₂ + H₂SO₄ → BaSO₄↓ + H₂O₂
(Barium peroxide + dilute H₂SO₄)

Strengths of H₂O₂:

• 3% w/v (10 vol) → medicinal/topical
• 6% w/v (20 vol) → hair bleaching
• 30% w/v (100 vol) → pharmacy stock solution
• "Volume strength" = volumes of O₂ released per volume of H₂O₂

Assay Principle (Permanganometry):

2KMnO₄ + 5H₂O₂ + 3H₂SO₄ → 2MnSO₄ + K₂SO₄ + 8H₂O + 5O₂

H₂O₂ acts as REDUCING AGENT (oxidized to O₂)

1. Dilute stock H₂O₂ appropriately
2. Add to dilute H₂SO₄
3. Titrate with standard 0.02 M KMnO₄
4. Endpoint: Permanent pale pink
5. Each mL 0.02 M KMnO₄ ≡ 0.68 mg H₂O₂

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Q2 & Q4. Mechanisms of Action of Inorganic Antimicrobial Agents

Classification and Mechanism:

INORGANIC ANTIMICROBIALS
│
├── 1. OXIDIZING AGENTS
│       H₂O₂, KMnO₄, Sodium hypochlorite
│       MECHANISM: Release active oxygen (nascent O)
│       → Oxidizes sulfhydryl (-SH) groups in enzymes
│       → Disrupts microbial protein structure
│       → Cell death
│
├── 2. HEAVY METAL SALTS (Oligodynamic action)
│       Silver nitrate, Zinc sulphate, Mercury compounds
│       MECHANISM:
│       Metal ions combine with sulfhydryl groups of enzymes
│       → Denature bacterial proteins
│       → Inhibit enzyme activity
│       → Bacteriostatic / Bactericidal
│
├── 3. SURFACE ACTIVE AGENTS
│       Quaternary ammonium compounds
│       MECHANISM: Disrupts bacterial cell membrane
│       → Loss of selective permeability
│       → Leakage of cell contents → death
│
└── 4. HALOGENS (Cl₂, I₂, iodophors)
        MECHANISM:
        Hypohalous acid (HOCl, HOI) formed
        → Oxidizes cell membrane lipids and proteins
        → Inhibits enzyme systems
        → Bactericidal

---

Q3. Boric Acid - Preparation and Assay

Chemical formula: H₃BO₃ | MW: 61.83 | Synonym: Orthoboric acid

Preparation:

Na₂B₄O₇ + H₂SO₄ + 5H₂O → 4H₃BO₃ + Na₂SO₄
(Borax + sulphuric acid → boric acid crystallizes out)

- Solution cooled → H₃BO₃ crystals filtered, washed, dried

Assay Principle:

H₃BO₃ + mannitol → mannitoboric acid complex
                           ↓
Titrate with 0.1 N NaOH (phenolphthalein indicator)

H₃BO₃ + NaOH → Na₂B₄O₇ + H₂O (after complex formation)

Uses:

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✅ UNIT 9 - DENTAL PRODUCTS AND MISCELLANEOUS COMPOUNDS

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Q1. Dentifrices - Classification + Role of Fluorides in Dental Caries

Definition:

Dentifrices are preparations used to clean teeth and maintain oral hygiene.

Classification:

DENTIFRICES
│
├── A. TOOTHPASTES (most common)
│
├── B. TOOTH POWDERS
│
├── C. TOOTH GELS
│
└── D. MOUTHWASHES (liquid dentifrices)

Composition of Dentifrice:

---

Role of Fluorides in Dental Caries:

Fluoride Mechanism:

MECHANISM OF FLUORIDE ACTION
│
├── 1. REMINERALIZATION
│       F⁻ + Ca²⁺ + PO₄³⁻ → Fluorapatite [Ca₁₀(PO₄)₆F₂]
│       Fluorapatite is HARDER and MORE ACID-RESISTANT than
│       hydroxyapatite → Strengthens enamel
│
├── 2. INHIBITION OF BACTERIA
│       Fluoride inhibits enolase enzyme of S. mutans
│       → Reduces lactic acid production
│       → Less demineralization
│
└── 3. SYSTEMIC EFFECT (water fluoridation)
        Fluoride incorporated into tooth structure
        during development (children)
        → Stronger enamel from inside

• Toothpaste: 1000-1500 ppm F⁻
• Water fluoridation: 0.7-1 ppm F⁻
• Dental gels: 5000 ppm F⁻ (professional)

---

Q2. Official Zinc Compounds + Zinc Sulphate Preparation and Assay

Official Zinc Compounds:

Zinc Sulphate - Preparation and Assay:

Zn + H₂SO₄ → ZnSO₄ + H₂↑
(Zinc metal dissolved in dilute H₂SO₄; purified by precipitation of Fe, Pb with ZnCO₃; crystallized)

Zn²⁺ + H₂Y²⁻ → ZnY²⁻ + 2H⁺ (at pH 5-6, xylenol orange indicator)

Or at pH 10 with EBT indicator:
Wine-red → Pure blue

• 0.25% eyedrops as astringent (zinc sulphate eye drops)
• Zinc supplement (dietary deficiency)
• Antidiarrheal (especially in children)
• Emetic (in poisoning)

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Q3. Bentonite and Magnesium Stearate

Bentonite:

Magnesium Stearate:

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Q5. Sodium Metabisulphite and Sodium Carboxymethylcellulose (CMC)

Sodium Metabisulphite:

---

Sodium Carboxymethylcellulose (CMC-Na):

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✅ UNIT 10 - AMMONIUM COMPOUNDS, RADIOPHARMACEUTICALS, ANTIDOTES

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Q1. Role of Ammonium Compounds as Respiratory Stimulants

Ammonium Compounds Used:

• Aromatic Ammonia Spirit (Sal Volatile)
• Ammonium Carbonate [(NH₄)₂CO₃]
• Ammonium bicarbonate [NH₄HCO₃]

Mechanism of Respiratory Stimulation:

MECHANISM:
│
Ammonium carbonate / smelling salts inhaled
            ↓
Decomposes and releases NH₃ gas (pungent, irritating)
            ↓
NH₃ irritates mucous membranes of nasal cavity
            ↓
Stimulates trigeminal nerve endings (sensory irritation)
            ↓
Reflex stimulation of respiratory center in medulla oblongata
            ↓
INCREASED RATE AND DEPTH OF RESPIRATION
            ↓
INCREASED AROUSAL (used in fainting)

(NH₄)₂CO₃ → 2NH₃ + CO₂ + H₂O
(ammonium carbonate decomposes rapidly)

• First aid for fainting/syncope (smelling salts)
• Expectorant (ammonium chloride in cough syrups)
• Aromatic ammonia spirit - reflex respiratory stimulant

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Q2. Radioactivity + Clinical Applications of Radiopharmaceuticals

Radioactivity:

Radioactivity is the spontaneous disintegration of unstable atomic nuclei with emission of radiation (alpha, beta, gamma).

Clinical Applications of Radiopharmaceuticals:

CLINICAL APPLICATIONS
│
├── 1. DIAGNOSTIC (Imaging)
│       - Tc-99m: Bone scan, lung scan, thyroid, renal scan
│         (Most widely used - 6 hr half-life, pure gamma)
│       - I-131: Thyroid scanning, uptake test
│       - Tl-201: Myocardial perfusion imaging (heart)
│       - Ga-67: Tumor/infection scanning
│       - F-18 (PET scan): Brain and cancer imaging
│
├── 2. THERAPEUTIC
│       - I-131: Treatment of hyperthyroidism, thyroid cancer
│       - P-32: Polycythemia vera, bone marrow suppression
│       - Sr-89: Bone pain palliation in cancer metastases
│       - Y-90: Liver cancer (microspheres - TheraSphere)
│
├── 3. RESEARCH
│       - Radiolabeled drugs for pharmacokinetic studies
│       - C-14/H-3: Metabolic pathway tracing
│
└── 4. IN VITRO DIAGNOSTICS
        - RIA (Radioimmunoassay) - hormone measurements
        - I-125 labeled antibodies

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Q3. Radioisotopes - Definition, Examples, Uses in Pharmacy

Definition:

Radioisotopes (radionuclides) are unstable forms of elements with the same atomic number but different atomic mass, that undergo spontaneous radioactive decay.

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Q4. Types of Antidotes + Official Antidote Compounds

Definition:

An antidote is a substance that counteracts the effects of a poison or drug overdose.

Classification of Antidotes:

ANTIDOTES
│
├── 1. MECHANICAL ANTIDOTES
│       - Work by physical means
│       - Prevent absorption of poison
│       - Examples:
│         • Activated charcoal (universal antidote) - adsorbs toxins
│         • White of egg - precipitates heavy metals
│         • Milk - dilutes and coats stomach wall
│         • Kaolin - adsorption
│
├── 2. CHEMICAL ANTIDOTES
│       - React chemically to detoxify poison
│       - Examples:
│         • Sodium thiosulphate → converts CN⁻ to thiocyanate
│         • Sodium bicarbonate → neutralizes acids
│         • Tannic acid → precipitates alkaloids/metals
│         • Potassium permanganate → oxidizes poisons (morphine, organophosphates)
│
├── 3. PHYSIOLOGICAL (PHARMACOLOGICAL) ANTIDOTES
│       - Counter the toxic effect by opposing pharmacological action
│       - Examples:
│         • Atropine → antidote for organophosphate poisoning (blocks muscarinic effects)
│         • Naloxone (Narcan) → opioid antidote (competitive antagonist at μ receptors)
│         • Flumazenil → benzodiazepine antidote (competitive antagonist)
│         • Physostigmine → atropine poisoning
│
└── 4. SPECIFIC ANTIDOTES (Chelating Agents)
        - Form stable, water-soluble complexes with metals
        - Examples:
          • EDTA (disodium) → lead poisoning (plumbism)
          • BAL (Dimercaprol) → arsenic, mercury, gold poisoning
          • Desferrioxamine (Deferoxamine) → iron poisoning
          • D-penicillamine → copper (Wilson's disease), lead, mercury
          • Prussian Blue → thallium and cesium poisoning
          • Pralidoxime (2-PAM) → organophosphate poisoning
            (reactivates acetylcholinesterase)

---

Official Antidote Compounds (IP/BP listed):

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Flowchart: Approach to Poisoning Management:

POISONING
    ↓
1. STABILIZE patient (ABC: Airway, Breathing, Circulation)
    ↓
2. IDENTIFY poison (history, symptoms, toxicology screen)
    ↓
3. DECONTAMINATION
   - Skin/eye: wash with water
   - GIT: activated charcoal (within 1-2 hrs), gastric lavage
    ↓
4. SPECIFIC ANTIDOTE (if available)
    ↓
5. SUPPORTIVE TREATMENT
   (IV fluids, O₂, vasopressors, dialysis)
    ↓
6. MONITORING and FOLLOW-UP

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📝 QUICK REVISION TABLE - KEY FORMULAS