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UNIT-I: Errors in Volumetric Analysis — Pharmaceutical Inorganic Chemistry

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① VERY SHORT ANSWER TYPE [2 Marks Each]

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(i) Define Error?

Absolute Error = Observed Value − True Value

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(ii) Write down differences between 1° (Primary) and 2° (Secondary) Standards

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(iii) Define Precision & Accuracy

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(iv) Define Significant Figures

1. All non-zero digits are significant. (e.g., 1.25 → 3 sig. fig.)
2. Zeros between non-zero digits are significant. (e.g., 1.05 → 3 sig. fig.)
3. Leading zeros are NOT significant. (e.g., 0.025 → 2 sig. fig.)
4. Trailing zeros after decimal are significant. (e.g., 2.50 → 3 sig. fig.)
5. Trailing zeros without decimal may or may not be significant.

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(v) Define the terms Molality & Mole Fraction

m = moles of solute / mass of solvent (in kg)

• Unit: mol/kg
• Advantage: Not affected by temperature (uses mass, not volume)

Xₐ = nₐ / (nₐ + n_b)

• Dimensionless (no units)
• Sum of mole fractions = 1

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(vi) Give the pH range of Methyl Orange & Phenolphthalein

• Methyl Orange is used for strong acid vs weak base titrations
• Phenolphthalein is used for strong base vs weak acid titrations

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(vii) Give the types of errors with examples

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② LONG ANSWER TYPE [10 Marks Each]

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① Discuss the Different Methods to Minimize Errors

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② Primary & Secondary Standards with Example + Standardization of 0.1N Perchloric Acid

1. Available in high purity (99.9%+)
2. Stable composition (does not react with air, CO₂, moisture)
3. High equivalent weight (reduces weighing error)
4. Easily soluble in the titration solvent
5. Reacts stoichiometrically with the titrant

1. Cannot be obtained in a pure state
2. React with atmospheric CO₂, water (e.g., NaOH absorbs CO₂)
3. Standardized against a primary standard before use

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1. Accurately weigh ~204 mg of dried KHP.
2. Dissolve in glacial acetic acid (10 mL) with gentle warming.
3. Add 1–2 drops of crystal violet indicator.
4. Titrate with 0.1N HClO₄ in glacial acetic acid until colour changes from violet to blue-green.
5. Each mL of 0.1N HClO₄ ≡ 20.42 mg of KHP.

Normality = (Weight of KHP × 1000) / (Equivalent weight × Volume of HClO₄ in mL) Eq. wt. of KHP = 204.2 g/eq

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③ Explain the Importance of Significant Figures

1. Reflects Measurement Precision: Tells the reader how precisely a measurement was made. E.g., 25.0 mL (3 sig. fig.) is more precise than 25 mL (2 sig. fig.).
2. Avoids False Precision: Reporting more digits than the instrument can measure is misleading. E.g., a balance reading to 0.01 g should not report 5.1234 g.
3. Rounding in Calculations: When multiplying or dividing, the result should have the same number of sig. fig. as the least precise measurement. This prevents compounding errors.
4. Compliance with Pharmacopoeial Standards: IP/BP/USP require values to be reported to specified decimal places. Significant figures ensure compliance.
5. Quality Control: In pharmaceutical manufacturing, correct sig. fig. ensures drug dosage accuracy. An error in sig. fig. can mean a 10× or 100× dosing error.
6. Prevents Compounding of Errors: In multi-step calculations, tracking sig. fig. limits the accumulation of rounding errors.

• Addition/Subtraction → answer has fewest decimal places
• Multiplication/Division → answer has fewest significant figures

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④ Give Different Methods to Express Concentration of Solutions

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⑤ Explain Different Types of Errors in Volumetric Analysis & Ways to Minimize Them

Errors
├── Determinate (Systematic)
│   ├── Operational errors
│   ├── Instrumental errors
│   ├── Reagent errors
│   └── Method errors
├── Indeterminate (Random)
└── Gross Errors (Blunders)

• Calibrate instruments regularly
• Run blank titrations
• Use freshly prepared, standardized reagents
• Select correct method and indicator

• Cause: fluctuation in temperature, pressure, human variation in reading
• Cannot be eliminated but can be reduced by:
  • Taking multiple readings
  • Statistical analysis (mean, SD, confidence interval)

• Recording wrong reading, using wrong reagent, calculator error
• Minimization: Double-checking, following SOP, proper training

1. Calibration of apparatus
2. Blank titrations
3. Replicate titrations
4. Proper indicator selection
5. Use of primary standards
6. Temperature control
7. Proper cleaning of glassware
8. Statistical analysis of results

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⑥ Explain the Term Milliequivalent (mEq) + Calculate mEq of NaCl in 1 L of 0.769 Solution

1 mEq = 1 gram equivalent / 1000

• Used to express concentration of electrolytes in body fluids and IV solutions
• Allows calculation of ionic balance (Na⁺, K⁺, Cl⁻, HCO₃⁻)
• IV fluid labels express electrolytes in mEq/L (e.g., Normal saline: Na⁺ = 154 mEq/L)

mEq = (Weight in mg × Valency) / Molecular weight OR mEq/L = Normality × 1000

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• Volume = 1 L = 1000 mL
• Normality = 0.769 N
• NaCl: Molecular weight = 58.5 g/mol; Valency = 1; Equivalent weight = 58.5 g/eq

Gram equivalents = Normality × Volume (L) = 0.769 × 1 = 0.769 g-eq

mEq = 0.769 × 1000 = 769 mEq

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③ SHORT ANSWER TYPE [5 Marks Each]

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① List out Various Volumetric Methods & Explain Back Titration with Example

1. Acidimetry & Alkalimetry — acid-base titrations (e.g., HCl vs NaOH)
2. Precipitation Titrations — e.g., Argentometry (AgNO₃ for Cl⁻)
3. Complexometric Titrations — e.g., EDTA titrations for metal ions
4. Oxidation-Reduction (Redox) Titrations:
   • Permanganometry (KMnO₄)
   • Dichromatometry (K₂Cr₂O₇)
   • Iodometry/Iodimetry
5. Non-aqueous Titrations — HClO₄ in glacial acetic acid for weak bases

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• The analyte is insoluble or reacts too slowly with the titrant
• The endpoint of direct titration is not sharp
• The analyte decomposes during direct titration

1. Add excess known amount of reagent (A) to the analyte
2. Allow complete reaction
3. Titrate the unreacted (residual) excess of A with a standard solution of B
4. Amount of analyte = Amount of A added − Amount of A remaining (as titrated by B)

1. Add excess standard HCl (50 mL, 0.1N) to chalk
2. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
3. Excess HCl is back-titrated with standard NaOH (0.1N)
4. Amount of HCl reacted with CaCO₃ = (HCl added) − (HCl titrated by NaOH)
5. Amount of CaCO₃ calculated from this difference

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② How do you Prepare & Standardize 500 mL of N/10 Sodium Hydroxide Solution?

Weight = Normality × Equivalent weight × Volume (L) = 0.1 × 40 × 0.5 = 2.0 g

1. Weigh approximately 2.0 g of NaOH pellets using a watch glass (not paper — NaOH absorbs moisture).
2. Dissolve in CO₂-free distilled water.
3. Transfer to a 500 mL volumetric flask.
4. Make up to 500 mL with CO₂-free distilled water.
5. Mix well. Store in a polythene bottle (NaOH etches glass and absorbs CO₂).

• Eq. wt. of oxalic acid = 63 g/eq
• Dissolve accurately weighed oxalic acid (~630 mg per 100 mL) in water
• Titrate with NaOH using phenolphthalein as indicator
• Endpoint: colourless → faint pink (permanent for 30 sec)

N₁V₁ (oxalic acid) = N₂V₂ (NaOH) N₂ = (N₁ × V₁) / V₂

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③ Give the Uses of Dilute HCl & Oxalic Acid

1. Used as a standardizing agent in alkalimetry
2. Dilute HCl BP — used to adjust gastric pH in achlorhydria
3. Used in the preparation of official preparations
4. Solvent for many inorganic substances
5. Used in back titration of carbonates
6. Preparation of chloride salts in synthesis
7. Used in pH adjustment in pharmaceutical formulations

1. Primary standard in alkalimetry (standardizing NaOH)
2. Primary standard in permanganometry (standardizing KMnO₄)
3. Used as a reducing agent in chemical reactions
4. Used in rust removal and cleaning of metal surfaces
5. Used as a precipitating agent for calcium ions (Ca²⁺ + C₂O₄²⁻ → CaC₂O₄↓)
6. Used in leather tanning and textile industries
7. In pharmacy, used in the assay of calcium-containing preparations

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④ What is a Primary Standard? Give an Example. What are the Properties of a Primary Standard?

1. High Purity — Available in pure form (≥99.9%); free from impurities
2. Stability — Stable in air, not affected by moisture, CO₂, or light during storage
3. High Equivalent Weight — Reduces the relative error in weighing
4. Anhydrous or Definite Composition — Exact molecular formula known; should not be hygroscopic
5. Solubility — Readily soluble in the titration solvent (usually water)
6. Non-toxicity (preferred) — Safe to handle
7. Availability & Low Cost — Economically viable for laboratory use
8. Stoichiometric Reaction — Reacts exactly and completely with the substance being titrated

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⑤ Classify Determinate Errors? Explain the 5 Steps to Minimize the Error

• Due to analyst's technique
• Examples: Incorrect weighing, failure to read meniscus at eye level, not mixing solution properly, loss of precipitate during washing

• Due to faulty instruments or uncalibrated equipment
• Examples: Burette not calibrated, balance not zeroed, pipette delivering incorrect volume

• Due to impurities in chemicals used
• Examples: NaOH containing Na₂CO₃, KMnO₄ decomposing before titration

• Inherent to the analytical method itself
• Examples: Coprecipitation of foreign ions, incomplete reaction, wrong endpoint

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⑥ Write a Note on Back Titration

• Analyte is insoluble in the titrant
• Analyte reacts slowly (equilibrium not reached quickly)
• No suitable indicator available for direct titration
• Analyte is volatile and would be lost during direct titration
• Endpoint of direct titration is not sharp

1. Add a measured excess volume (V₁) of standard solution A (of known normality N₁) to the analyte
2. Allow the reaction to go to completion (heat if necessary)
3. Cool and titrate the residual unreacted A with a standard solution B (normality N₂, volume V₂)
4. Calculate amount of analyte:

Equivalents of analyte = (N₁ × V₁) − (N₂ × V₂)

• Aspirin is hydrolyzed with excess NaOH (heated)
• Excess NaOH back-titrated with standard H₂SO₄

• Chalk dissolved in excess standard HCl
• Excess HCl titrated back with standard NaOH
• Indicator: phenolphthalein

• Applicable to insoluble/slow-reacting substances
• Can determine multiple components simultaneously
• Avoids problems with volatile or unstable analytes

• Two standard solutions required
• Slightly more complex calculation
• Accumulation of errors from two titrations

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Exam Tips:

• For 2-mark questions: Define precisely + give one example or formula
• For 5-mark questions: Add classification table + 2 examples
• For 10-mark questions: Use headings, tables, formulas, examples, and a diagram if possible
• Always write units with numerical answers
• For calculation questions: Show each step clearly — marks are given for method

UNIT-II: Acid-Base Titrations & Redox Titrations — Pharmaceutical Inorganic Chemistry

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① VERY SHORT ANSWER TYPE [2 Marks Each]

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(i) What is Redox Titration?

2KMnO₄ + 10FeSO₄ + 8H₂SO₄ → 2MnSO₄ + 5Fe₂(SO₄)₃ + K₂SO₄ + 8H₂O

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(ii) Define Oxidation & Reduction

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(iii) What is Redox Potential?

• Also called electrode potential or oxidation-reduction potential (ORP)
• Measured in Volts (V)
• A higher (more positive) E° = stronger oxidizing agent
• A lower (more negative) E° = stronger reducing agent

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(iv) What is Oxidation Number?

1. Oxidation number of free element = 0 (e.g., O₂, Fe = 0)
2. For monoatomic ions = charge on ion (e.g., Na⁺ = +1)
3. O = −2 (except in peroxides = −1)
4. H = +1 (except in metal hydrides = −1)
5. Sum of oxidation numbers in a neutral compound = 0

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(v) Write Applications of Redox Titrations

1. Assay of iron preparations (FeSO₄ tablets) using KMnO₄
2. Determination of vitamin C (ascorbic acid) by iodimetry
3. Assay of calcium gluconate by permanganometry
4. Determination of copper by iodometric method
5. Assay of hydrogen peroxide using KMnO₄
6. Determination of chlorine in water by iodometry
7. Assay of sodium thiosulphate — iodometric method
8. Quality control of oxidizing/reducing pharmaceutical agents

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(vi) Write about Any One Theory of Acid-Base Indicator

HIn ⇌ H⁺ + In⁻ (Colour A) → (Colour B)

• In acidic medium (excess H⁺): equilibrium shifts LEFT → HIn form predominates → Colour A
• In alkaline medium (H⁺ removed): equilibrium shifts RIGHT → In⁻ form predominates → Colour B

• HIn (colourless) ⇌ H⁺ + In⁻ (pink)
• Acid → colourless; Base → pink

pH = pKin + log [In⁻]/[HIn] Colour change: pH range = pKin ± 1

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(vii) Define Alkalimetry Titration

• The base (alkali) acts as the titrant in the burette
• The acid is in the conical flask
• Indicator: phenolphthalein or methyl orange depending on the type of acid/base

NaOH + HCl → NaCl + H₂O

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(viii) Define Normality & Molarity

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(ix) What is Weak Base?

B + H₂O ⇌ BH⁺ + OH⁻

• Degree of ionization is small (Kb << 1)
• pH < 14 but > 7 in solution
• Examples: Ammonia (NH₃), pyridine, aniline, calcium carbonate

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(x) Write about Different Types of Neutralisation Indicators

1. One-colour indicators: Only one form is coloured (e.g., phenolphthalein)
2. Two-colour indicators: Both forms have different colours (e.g., methyl orange)

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② LONG ANSWER TYPE [10 Marks Each]

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① Explain Briefly Theories of Neutralisation Indicators

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HIn ⇌ H⁺ + In⁻ (Colour 1) (Colour 2)

KIn = [H⁺][In⁻] / [HIn]

pH = pKIn + log([In⁻]/[HIn])

• This gives a pH range of pKIn ± 1 for colour change

• Methyl Orange (weak acid type): pKIn = 3.46; range 3.1–4.4
• Phenolphthalein (weak acid type): pKIn = 9.3; range 8.3–10.0

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• In acidic form: lactoid form (colourless, closed ring)
• In alkaline form: quinoid form (pink, open ring structure with conjugated double bonds)

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② Explain Ostwald's Theory for Acid-Base Indicators

1. Acid-base indicators are weak acids (HIn) or weak bases (InOH)
2. The un-ionized form and ionized form have different colours
3. The ratio of [In⁻]/[HIn] determines which colour is predominant
4. This ratio is governed by the pH of the solution

KIn = [H⁺][In⁻] / [HIn] [H⁺] = KIn × [HIn]/[In⁻] pH = pKIn + log([In⁻]/[HIn])

• When [HIn]/[In⁻] ≥ 10 → eye sees Colour of HIn (acid colour)
• When [In⁻]/[HIn] ≥ 10 → eye sees Colour of In⁻ (base colour)
• Transition range = pKIn − 1 to pKIn + 1

• HIn (Red) ⇌ H⁺ + In⁻ (Yellow)
• pKIn = 3.46
• Transition range = 2.46 to 4.46 ≈ reported as 3.1–4.4

• HIn (colourless) ⇌ H⁺ + In⁻ (pink)
• pKIn = 9.3
• Transition range = 8.3 to 10.3 ≈ 8.3–10.0

• Mathematically sound
• Predicts transition pH range correctly
• Explains why different indicators suit different titrations

• Cannot explain phenolphthalein's structural change (quinoid vs. lactoid)
• Does not address mixed indicators fully
• Assumes only ionization — ignores tautomerism

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③ Explain Iodometry & Iodimetry with Examples

I₂ + 2e⁻ → 2I⁻ (Iodine is reduced; analyte is oxidized)

C₆H₈O₆ + I₂ → C₆H₆O₆ + 2HI (Ascorbic acid reduces I₂ → colourless; endpoint = first permanent blue)

2Na₂S₂O₃ + I₂ → Na₂S₄O₆ + 2NaI

• pH should be slightly acidic (pH 3–4); alkaline conditions cause I₂ to disproportionate
• Add starch near the endpoint (not at beginning — strong starch-I₂ complex makes endpoint hard to reverse)

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Step 1: Oxidant + 2KI → I₂ (liberated) Step 2: I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ + 2NaI

Step 1: 2Cu²⁺ + 4KI → 2CuI + I₂ Step 2: I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ + 2NaI

Cl₂ + 2KI → I₂ + 2KCl I₂ titrated with Na₂S₂O₃

H₂O₂ + 2KI + H₂SO₄ → I₂ + K₂SO₄ + 2H₂O

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④ Name 4 Primary Standards for Redox Titrations

• High purity, stable, high equivalent weight, soluble, reacts stoichiometrically

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⑤ Give Nernst Equation — Explain Important Terms — Importance in Redox Titrations

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Fe³⁺ + e⁻ → Fe²⁺; E° = +0.77 V

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

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1. Predicts feasibility of titration: If E°cell = E°oxidant − E°reductant > 0.2 V, the reaction proceeds quantitatively (>99.9% complete) — suitable for titration
2. Locates equivalence point: At the equivalence point, E_cell can be calculated using Nernst equation to identify the sharp potential jump
3. Potentiometric titrations: Nernst equation is the basis of potentiometric redox titrations (using platinum electrode + calomel reference electrode)
4. Effect of concentration: Shows how electrode potential changes with concentration of reactants/products — important for understanding indicator range
5. pH effect: In permanganometry, H⁺ appears in the half-reaction. Nernst equation shows potential decreases as pH increases → acidic conditions required
6. Identifies suitable indicator: Helps choose a redox indicator whose E°In lies within the steep part of the titration curve (E_equivalence ± 0.059/n V)

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⑥ Explain Different Types of Redox Titrations with Examples

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• Titrant: 0.1N KMnO₄ (self-indicating — purple/violet)
• Medium: Acidic (H₂SO₄) — never HCl (it reduces KMnO₄)
• Reaction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (colourless)
• Endpoint: Faint permanent pink/violet colour
• No external indicator needed (self-indicator)
• Example: Assay of H₂O₂, FeSO₄, oxalic acid, calcium gluconate

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• Titrant: 0.1N K₂Cr₂O₇ (orange) — primary standard
• Medium: Acidic (H₂SO₄)
• Reaction: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O (green)
• Indicator: Diphenylamine or N-phenyl anthranilic acid
• Advantages over KMnO₄: Stable, primary standard, can be used in HCl medium
• Example: Assay of iron in FeSO₄ tablets

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• Titrant: Standard I₂ solution (0.1N)
• Indicator: Starch solution
• Endpoint: Appearance of blue colour
• Used for: Reducing agents (ascorbic acid, Na₂S₂O₃)
• Example: Vitamin C assay

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• Titrant: Standard Na₂S₂O₃ (0.1N)
• Intermediate: KI + oxidant → I₂ liberated
• Indicator: Starch
• Endpoint: Disappearance of blue
• Used for: Oxidizing agents (Cu²⁺, Cl₂, H₂O₂, KMnO₄)

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• Titrant: Ceric sulphate [Ce(SO₄)₂], 0.1N
• Medium: Acidic (H₂SO₄)
• Reaction: Ce⁴⁺ + e⁻ → Ce³⁺ (yellow → colourless)
• Advantages: Very stable standard, can be used in HCl medium, no MnO₂ precipitate
• Indicator: Ferroin (1,10-phenanthroline-ferrous complex)
• Example: Assay of iron, As₂O₃, organic compounds

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• Titrant: KBrO₃ (0.1N) — primary standard
• In acid + KBr: BrO₃⁻ + 5Br⁻ + 6H⁺ → 3Br₂ + 3H₂O
• Br₂ generated in situ reacts with analyte
• Indicator: Methyl orange or methyl red (destroyed by Br₂ at endpoint — colour disappears)
• Example: Assay of phenol, 8-hydroxyquinoline, antimony compounds

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③ SHORT ANSWER TYPE [5 Marks Each]

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① Write a Note on Preparation & Storage of Volumetric Solution of Iodine

Weight = 0.1 × 127 × 1 = 12.7 g of I₂

1. Weigh 12.7 g of iodine and 20 g of potassium iodide (KI) (to increase solubility as KI·I₂ → KI₃ complex).
2. Dissolve KI in a small volume of water, then add iodine crystals.
3. Stir until completely dissolved (I₂ is poorly soluble in water; soluble in KI solution via: I₂ + I⁻ → I₃⁻).
4. Transfer to a 1000 mL volumetric flask.
5. Make up to volume with distilled water.

• Standardize against 0.1N sodium thiosulphate (primary standard: As₂O₃ or pure iodine)
• Indicator: starch; endpoint = disappearance of blue colour

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1. Store in amber-coloured (brown) glass bottles — prevents photodecomposition
2. Use glass-stoppered bottles — iodine vapour is corrosive to rubber stoppers
3. Store in a cool, dark place — iodine is volatile; heat increases evaporation
4. Do not store in plastic containers — iodine reacts with plastic
5. Standardize frequently — iodine solution changes strength over time
6. Keep excess KI in the solution to prevent volatilization of I₂

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② Define & Explain Acidimetric & Alkalimetric Titrations

• Acid is in burette; base is in flask
• Indicator: Methyl orange (for weak base) or phenolphthalein (for strong base) depending on titration type
• Example: Determination of Na₂CO₃ using standard HCl

Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂ Indicator: methyl orange; endpoint: yellow → red

• Base (NaOH) is in burette; acid is in flask
• Example: Determination of acetic acid using standard NaOH

CH₃COOH + NaOH → CH₃COONa + H₂O Indicator: phenolphthalein; endpoint: colourless → pink

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③ Explain Theory of Redox Titration + Note on Iodimetry & Iodometry

1. The reaction is quantitative (proceeds >99.9% to completion)
2. Reaction is fast (rapid equilibrium)
3. A suitable indicator or potentiometric method exists

ΔE° = E°(oxidant) − E°(reductant) > 0.2 V ensures >99% reaction completion

E = E°In ± 0.059/n V (transition range)

• E.g., Ferroin (E° = +1.06 V): colourless (ox) ↔ red (red)
• Starch-iodine: specific for I₂/I⁻ system (blue ↔ colourless)

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④ Explain Principle & Reactions Involved in Iodometry Titrations

Oxidant + 2KI (excess) + H₂SO₄ → I₂ + reduced product + K₂SO₄

I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ (sodium tetrathionate) + 2NaI

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2CuSO₄ + 4KI → 2CuI + I₂ + 2K₂SO₄ I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ + 2NaI

H₂O₂ + 2KI + H₂SO₄ → I₂ + K₂SO₄ + 2H₂O I₂ titrated with Na₂S₂O₃

Cl₂ + 2KI → I₂ + 2KCl I₂ titrated with Na₂S₂O₃

Same as above; used in water analysis

• Reaction done in acidic medium (dilute H₂SO₄) — prevents disproportionation of I₂
• Excess KI added to ensure complete liberation of I₂
• Starch indicator added just before endpoint (near pale yellow)
• Titration done quickly — I⁻ can be oxidized by atmospheric O₂ if left too long
• Bottles should be stoppered to prevent oxidation by air

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⑤ Write the Preparation & Standardization of 0.1N Potassium Permanganate

• Molecular weight = 158 g/mol
• Equivalent weight = 158/5 = 31.6 g/eq (in acid medium, gains 5e⁻: Mn⁷⁺ → Mn²⁺)
• Cannot be used as a primary standard (it oxidizes dust, rubber, etc. during preparation)
• Deep purple solution; self-indicator

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Weight = 0.1 × 31.6 × 1 = 3.16 g

1. Weigh approximately 3.2 g of KMnO₄
2. Dissolve in 1000 mL of freshly boiled and cooled distilled water (boiling removes organic matter + CO₂ that would reduce KMnO₄)
3. Allow to stand for 24 hours (Mn₂O₇ and organic matter are oxidized and precipitate as MnO₂)
4. Filter through a sintered glass crucible (NOT filter paper — filter paper is organic and would be oxidized)
5. Store in a brown bottle away from light

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• Eq. wt. of oxalic acid = 63 g/eq (MW = 126, n-factor = 2)
• Weigh accurately ~630 mg; dissolve in water with dilute H₂SO₄

5H₂C₂O₄ + 2KMnO₄ + 3H₂SO₄ → 10CO₂ + 2MnSO₄ + K₂SO₄ + 8H₂O

1. Pipette 20 mL of oxalic acid solution into a conical flask
2. Add 20 mL of dilute H₂SO₄ (to acidify)
3. Heat to 70–80°C (reaction is slow at room temperature; first few drops of KMnO₄ are decolorized slowly, then autocatalyzed by Mn²⁺)
4. Titrate with KMnO₄ from burette
5. Endpoint: Faint permanent pink/violet colour that persists for 30 seconds (self-indicator)

N₁V₁ (oxalic acid) = N₂V₂ (KMnO₄) N(KMnO₄) = (N₁ × V₁) / V₂

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⑥ Write the Procedure & Reactions in Standardization of 0.1N Sodium Thiosulphate

• Molecular weight = 248 g/mol (hydrated)
• Equivalent weight = 248 g/eq (n-factor = 1; S₂O₃²⁻ → S₄O₆²⁻, each S₂O₃ loses 1e⁻)
• Cannot be used as a primary standard (absorbs CO₂, decomposes slowly)
• Used as titrant in iodometry

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Weight = 0.1 × 248 × 1 = 24.8 g

1. Weigh 24.8 g of Na₂S₂O₃·5H₂O
2. Dissolve in freshly boiled and cooled distilled water (remove CO₂; prevents bacterial decomposition)
3. Add a small amount of Na₂CO₃ (~0.1 g) to maintain slightly alkaline pH (prevents bacterial decomposition and acid decomposition)
4. Make up to 1000 mL; store in a brown bottle

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• Eq. wt. of KIO₃ = 214/5 = 35.67 g/eq (IO₃⁻ → I₂, n = 5 per iodine = 10 for 2I) (More precisely: KIO₃, MW = 214, n-factor = 5 in iodometric context)

IO₃⁻ + 5I⁻ + 6H⁺ → 3I₂ + 3H₂O

I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ + 2NaI

KIO₃ + 5KI + 3H₂SO₄ → 3I₂ + 3K₂SO₄ + 3H₂O 3I₂ + 6Na₂S₂O₃ → 3Na₂S₄O₆ + 6NaI

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1. Accurately weigh ~357 mg of dried KIO₃ (primary standard)
2. Dissolve in water; transfer to 100 mL volumetric flask
3. Pipette 10 mL into a conical flask
4. Add 10 mL of 10% KI solution
5. Add 10 mL of dilute H₂SO₄ — brown/yellow colour appears (I₂ liberated)
6. Titrate with Na₂S₂O₃ from burette until solution turns pale yellow
7. Add 1 mL starch indicator — solution turns blue
8. Continue titrating until blue disappears (endpoint = colourless)

N₁V₁ (KIO₃) = N₂V₂ (Na₂S₂O₃) N(Na₂S₂O₃) = (N₁ × V₁) / V₂

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Exam Tips for Unit II:

• Always include half-reactions with electron counts in redox questions — earns easy marks
• For Nernst equation: write the formula, define every symbol, give one numerical example
• For iodimetry vs iodometry: the contrast table alone scores 2–3 marks
• In standardization questions: always show weight calculation + procedure + endpoint colour + equation for N
• Write reaction equations wherever possible — they demonstrate understanding and are specifically marked

UNIT-V: Gravimetry & Limit Tests — Pharmaceutical Inorganic Chemistry

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① VERY SHORT ANSWER TYPE [2 Marks Each]

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(i) Define Gravimetric Titration (Gravimetric Analysis)?

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(ii) What is Co-precipitation & Post-precipitation?

1. Surface adsorption — ions adsorbed on precipitate surface
2. Occlusion — foreign ions trapped inside growing crystal
3. Isomorphous inclusion — foreign ions replace lattice ions

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(iii) What are the Limitations of Gravimetric Analysis?

1. Time-consuming — requires precipitation, filtration, washing, drying, ignition — takes hours to days
2. Applicable to macro-amounts only — not suitable for trace/micro quantities
3. Co-precipitation & post-precipitation — cause errors in weighing
4. High analytical skill required — errors in filtration or ignition directly affect results
5. Single analyte determination — not suitable for simultaneous multi-component analysis
6. Volatile precipitates — some precipitates may decompose or volatilize during ignition
7. Non-specific precipitation — interfering ions may co-precipitate, giving false results
8. Requires expensive equipment — analytical balance, muffle furnace, desiccator

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(iv) List Different Sources of Impurity in Pharmaceutical Substances

1. Starting materials — unreacted raw materials carried forward
2. Synthetic intermediates — partially reacted or by-products of synthesis
3. Degradation products — due to hydrolysis, oxidation, photolysis, or thermal decomposition
4. Reagents used in synthesis — solvents, catalysts, heavy metals (e.g., Pb, As, Hg)
5. Process-related impurities — contamination from equipment (metals like Fe, Cu)
6. Microbial contamination — endotoxins, microbial by-products
7. Packaging materials — leaching from containers (plasticizers, heavy metals)
8. Environmental contamination — dust, water, air-borne particles
9. Isomers/enantiomers — unwanted stereoisomers from synthesis
10. Residual solvents — methanol, benzene, chloroform from manufacturing

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(v) Give Rxn (Reaction) Involved in Limit Test of Chloride

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

• Dilute nitric acid (HNO₃) — acidifies the medium, prevents precipitation of Ag₂CO₃ or Ag₃PO₄
• Silver nitrate solution (AgNO₃, 2% w/v)

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(vi) Give Rxn Involved in Limit Test for Sulphate

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

• Dilute hydrochloric acid (HCl) — acidifies medium, prevents BaCO₃ precipitation
• Barium chloride solution (BaCl₂, 25% w/v)

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(vii) Principle Involved in Limit Test for Arsenic

As⁵⁺ + SnCl₂ → As³⁺ + SnCl₄

2As³⁺ + 3Zn + 6HCl → 2AsH₃↑ + 3ZnCl₂

AsH₃ + HgCl₂ → AsH(HgCl)₂ (yellow stain — arsenic trichloride) (Deeper brown with higher arsenic)

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(viii) Write Chemical Reactions Involved in Limit Test for Iron

Fe³⁺ + Thioglycolic acid (reducing agent) → Fe²⁺ + oxidized thioglycolic acid

Fe²⁺ + 2 Thioglycolate → [Fe(HSCH₂COO)₂]²⁻ (pink/red colour)

• Citric acid (removes phosphate interference)
• Thioglycolic acid (HSCH₂COOH) — complexing + reducing agent
• 13.5M ammonia — alkaline medium

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(ix) Write the Apparatus Used in Limit Test for Arsenic

1. Wide-mouthed conical flask (250 mL) — contains the test/standard solution + Zn + HCl + SnCl₂ for arsine generation
2. Rubber stopper — fits the flask, holds the tube
3. Delivery tube (glass) — leads arsine gas upward; packed with:
   • Lead acetate cotton wool — at the bottom, absorbs H₂S (from sulphide impurities which would interfere)
4. HgCl₂ paper strip — held at the top of the tube; reacts with AsH₃ to produce yellow-brown stain
5. Comparison tube — standard run simultaneously for visual comparison

[HgCl₂ Paper] ← stain appears here
      |
[Glass tube — narrow]
      |
[Lead acetate cotton — removes H₂S]
      |
[Rubber stopper]
      |
[250 mL conical flask]
[Contains: Zn + HCl + SnCl₂ + sample]

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(x) Define Limit Test & Its Types

• Not a quantitative determination — only checks if impurity is above or below a limit
• Result: PASS (test ≤ standard) or FAIL (test > standard)
• Based on comparison of colour or turbidity of test with a standard

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② LONG ANSWER TYPE [10 Marks Each]

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(i) Explain the Principle Involved in Gravimetric Analysis with One Example

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1. Precipitation — The analyte is converted into a sparingly soluble compound (precipitate) of definite composition by adding an appropriate precipitating agent
2. Filtration — The precipitate is separated from the solution
3. Washing — Impurities (co-precipitated substances) are removed
4. Drying or Ignition — Convert precipitate to a form of constant, known composition (the weighing form)
5. Weighing — The mass of the weighed form is measured on an analytical balance
6. Calculation — Using the gravimetric factor, amount of analyte is calculated

GF = (Atomic/molecular weight of analyte × stoichiometric factor) / Molecular weight of weighed form

Weight of analyte = Weight of precipitate × GF

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1. Sparingly soluble — losses due to solubility should be minimal (<0.1 mg)
2. Easily filterable — large crystals (not colloidal) for easy separation
3. Pure — free from co-precipitated impurities
4. Stable — chemically stable during drying/ignition
5. Known composition — exact stoichiometric formula

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Ba²⁺ + H₂SO₄ → BaSO₄↓ (white crystalline precipitate) + 2H⁺

1. Dissolve sample in water; acidify with dilute HCl
2. Heat to near-boiling; add hot dilute H₂SO₄ slowly with stirring
3. Digest (heat gently for 1–2 hours) to grow large crystals → reduces co-precipitation
4. Filter through Whatman No. 42 filter paper or sintered crucible
5. Wash with hot dilute H₂SO₄ (to prevent peptization), then hot water
6. Dry and ignite at 800°C in muffle furnace (converts to pure BaSO₄)
7. Cool in desiccator; weigh accurately

GF = Molecular weight of Ba / Molecular weight of BaSO₄ = 137.3 / 233.4 = 0.5884 Weight of Ba = Weight of BaSO₄ × 0.5884

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(ii) Enumerate the Different Steps Involved in Gravimetric Analysis & What are the Limitations of Gravimetry

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• Dissolve the sample in appropriate solvent (usually water or dilute acid)
• Remove interfering ions if necessary (masking agents or separation techniques)
• Adjust pH to optimum for precipitation

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• Add precipitating reagent slowly, dropwise with constant stirring to a hot, dilute solution
• Hot solution → larger particles (less co-precipitation)
• Dilute solution → minimizes co-precipitation (von Weimarn principle)
• Add excess precipitant (~10% excess) to ensure complete precipitation
• Check for complete precipitation: add a drop of reagent to clear supernatant — no turbidity = complete

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• Heat the precipitate with the mother liquor for 30–60 minutes just below boiling
• Ostwald ripening occurs — small particles dissolve and redeposit on larger crystals
• Result: larger, purer, more filterable crystals with less co-precipitation

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• Filter through appropriate medium:
  • Filter paper (Whatman 40, 41, 42) — for precipitates to be ignited
  • Sintered glass crucible (G3, G4) — for precipitates only dried (not ignited)
• Decant the bulk of supernatant first; then transfer precipitate

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• Wash precipitate with cold/hot wash solution (dilute electrolyte or water)
• Purpose: remove impurities and mother liquor
• Avoid excess washing — may dissolve some precipitate (peptization)
• Test for complete washing: test last washing with appropriate reagent (e.g., AgNO₃ for Cl⁻)

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• Drying: at 105–120°C to remove water → used when precipitate is thermally stable
• Ignition: at 600–1200°C in muffle furnace → converts to oxide or anhydrous form of known composition
• Example: BaSO₄ — ignited at 800°C; Fe(OH)₃ → Fe₂O₃ at 900°C

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• Cool in a desiccator containing silica gel or CaCl₂ to prevent moisture absorption

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• Weigh on a calibrated analytical balance (accurate to 0.1 mg)
• Weigh to constant weight — repeat drying and weighing until two consecutive weighings agree within 0.2 mg

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% Analyte = (Weight of precipitate × GF / Weight of sample) × 100

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(iii) Write the Principle, Rxn, Procedure Involved in Limit Test for Arsenic

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As₂O₅ + SnCl₂ + 2HCl → As₂O₃ + SnCl₄ + H₂O

2As³⁺ + 3Zn + 6HCl → 2AsH₃↑ + 3ZnCl₂

AsH₃ + 2HgCl₂ → AsH(HgCl)₂ + HCl (pale yellow stain) With more As: AsH₃ + HgCl₂ → AsHg₂Cl (brown stain)

H₂S + Pb(CH₃COO)₂ → PbS (black) + 2CH₃COOH (Lead acetate cotton wool removes H₂S before it reaches HgCl₂ paper)

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• 250 mL wide-mouth conical flask
• Rubber stopper with narrow glass tube
• Lead acetate cotton wool packed in tube (removes H₂S)
• HgCl₂ paper at the top of tube

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1. Dilute HCl (7N)
2. Stannous chloride solution (SnCl₂ in HCl)
3. Zinc granules (arsenic-free)
4. Lead acetate cotton wool
5. Mercuric chloride (HgCl₂) paper strips
6. Standard arsenic solution (1 ppm As — from As₂O₃)

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1. Cut HgCl₂ paper strips (0.5 cm × 7 cm); soak in 5% HgCl₂ solution; dry — set aside
2. Roll lead acetate cotton wool lightly; place in glass tube

1. Take 1 mL of standard arsenic solution (≡ 1 µg As = 1 ppm in 1 mL) in conical flask
2. Add 1 mL of SnCl₂ solution, 5 mL of 7N HCl
3. Add 2 g of arsenic-free zinc granules
4. Immediately insert the delivery tube assembly with HgCl₂ paper
5. Allow reaction for 40 minutes at room temperature

1. Prepare test solution as directed for individual monograph
2. Run simultaneously with standard under identical conditions

• Compare the stain on test HgCl₂ paper with the standard stain
• Pass: Test stain ≤ standard stain
• Fail: Test stain > standard stain (darker brown/yellow)

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1. Temperature must be same for test and standard
2. Only arsenic-free Zn and reagents must be used
3. Lead acetate wool must be present — H₂S causes false positive
4. HgCl₂ paper must be freshly prepared
5. Reading done in good light; compare with white background

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③ SHORT ANSWER TYPE [5 Marks Each]

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(i) Draw a Neat & Labelled Diagram of Gutzeit's Apparatus

        ┌──────────────────┐
        │  HgCl₂ Paper     │  ← Yellow-brown stain (proportional to As)
        │  (Strips at top) │
        └────────┬─────────┘
                 │  Glass tube (narrow bore)
        ┌────────┴─────────┐
        │  Lead Acetate    │  ← Absorbs H₂S (prevents false positive)
        │  Cotton Wool     │
        └────────┬─────────┘
                 │
        ┌────────┴─────────┐
        │   Rubber Stopper │  ← Ensures air-tight seal
        └────────┬─────────┘
                 │
   ┌─────────────┴───────────────┐
   │   250 mL Wide-Mouth Flask   │
   │                             │
   │  Sample/Standard + HCl      │
   │  + SnCl₂ + Zn granules      │
   │                             │
   │  → AsH₃ gas generated       │
   └─────────────────────────────┘

1. Wide-mouth conical flask — reaction vessel
2. Rubber stopper — airtight seal
3. Narrow glass tube — guides AsH₃ gas upwards
4. Lead acetate cotton — removes H₂S interference
5. HgCl₂ paper — reacts with AsH₃ to form yellow-brown stain

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(ii) Give Principle, Procedure, Role of Reagents Involved in Limit Test for Iron

• Citric acid — prevents precipitation of iron as Fe(OH)₃; chelates phosphate/arsenate interference
• Thioglycolic acid (HSCH₂COOH) — reduces Fe³⁺ → Fe²⁺ AND forms the pink complex
• Ammonia — provides alkaline medium (pH ~9) required for complex formation

Fe³⁺ + Thioglycolic acid → Fe²⁺ (reduced) + oxidized product Fe²⁺ + 2 Thioglycolic acid + NH₃ → [Fe(SCH₂COO)₂]²⁻·NH₄⁺ (pink/red colour)

1. Take 2 mL of standard iron solution (1 ppm → 2 µg Fe) in Nessler cylinder
2. Add 2 mL of 20% citric acid solution
3. Add 0.1 mL of thioglycolic acid
4. Add 10 mL of 13.5M ammonia solution
5. Dilute to 20 mL with water; mix well
6. Allow to stand for 5 minutes

1. Dissolve required quantity of sample (as per monograph) in water
2. Follow same steps as standard

• View Nessler cylinders from top against white background
• Pass: Test colour ≤ Standard colour (pale pink/pink)
• Fail: Test colour deeper than standard

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(iii) Explain the Various Sources of Impurities in Pharmaceuticals + Importance of Limit Test in Quality Control

1. Raw materials / Starting materials — Impure precursors or reagents used in synthesis may carry metal impurities (e.g., arsenic from sulphates, iron from acids)
2. Manufacturing process — Reactions produce by-products; equipment made of metal can leach Fe, Cu, Pb into the product
3. Degradation during storage — Chemical breakdown due to light, heat, moisture, or oxygen produces degradation products (e.g., hydrolysis of aspirin → salicylic acid + acetic acid)
4. Solvent residues — Organic solvents used in synthesis (methanol, chloroform) may remain as residual impurities
5. Environmental contamination — Microbial growth produces toxic by-products; dust, heavy metals from air/water

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1. Patient Safety — Toxic impurities (As, Pb, Hg) even in trace amounts can cause poisoning. Limit tests ensure safety.
2. Pharmacopoeial Compliance — IP/BP/USP specify maximum permissible limits for each impurity. Limit tests confirm compliance.
3. Batch Release — Pharmaceutical manufacturers must pass limit tests before releasing each batch for sale.
4. Simple & Rapid — Limit tests are semi-quantitative, quick, and cost-effective for routine quality control (unlike full quantitative analysis).
5. Cumulative toxicity check — Some impurities (e.g., heavy metals) accumulate in body over time — limit tests ensure chronic exposure stays below safe threshold.
6. Process Validation — Limit tests help identify if a manufacturing process is producing unacceptable impurity levels.
7. Stability studies — Regular limit testing during stability studies detects degradation products forming over time.

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(iv) Define Limit Test? List Out Different Limit Tests — Detail on Sulphate & Iron

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Ba²⁺ + SO₄²⁻ → BaSO₄↓ (white turbidity)

• Dilute HCl (prevents BaCO₃/BaHPO₄ precipitation)
• BaCl₂·2H₂O solution (25% w/v)

1. Dissolve sample in 15 mL distilled water
2. Add 1 mL of dilute HCl; mix
3. Add 3 mL of BaCl₂ solution; mix
4. Allow to stand for 1 minute; compare turbidity with standard

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• Citric acid (20%) — prevents precipitation
• Thioglycolic acid — reducing + complexing agent
• Ammonia (13.5M) — alkaline medium
• Endpoint/comparison: Pink colour — test ≤ standard = PASS

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(v) Explain the Various Sources of Impurities & Importance of Limit Tests in Quality Control

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(vi) How Do You Carry Out the Limit Test for Chloride in a Given Sample?

Cl⁻ + AgNO₃ → AgCl↓ (white opalescence/turbidity) Medium: dilute HNO₃ (prevents Ag₂CO₃, Ag₃PO₄ precipitation)

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1. Dilute nitric acid (HNO₃)
2. Silver nitrate solution (AgNO₃, 2% w/v)
3. Standard chloride solution (0.0085% w/v NaCl = 0.05 mg Cl/mL)

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1. Pipette 1 mL of standard NaCl solution (0.0085% ≡ 0.05 mg Cl) into Nessler cylinder A
2. Add 15 mL of distilled water
3. Add 1 mL of dilute HNO₃
4. Add 1 mL of AgNO₃ solution
5. Mix; dilute to 50 mL with distilled water
6. Allow to stand for 5 minutes in diffused light

1. Dissolve the required quantity of sample in 15 mL distilled water in Nessler cylinder B
2. Add 1 mL of dilute HNO₃
3. Add 1 mL of AgNO₃ solution
4. Mix; dilute to 50 mL
5. Allow to stand for 5 minutes (same time as standard)

• View from the top against a black background in diffused light
• Compare the degree of opalescence (turbidity) of test with standard

• PASS: Opalescence of test ≤ opalescence of standard (Cl⁻ within limit)
• FAIL: Opalescence of test > standard (Cl⁻ exceeds limit)

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1. Use matched (identical) Nessler cylinders
2. Both test and standard prepared simultaneously
3. View in diffuse (not direct) light
4. Use HNO₃ (not HCl!) as acidifying agent
5. Reagents must be chloride-free
6. Time of standing must be same for both

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Exam Tips for Unit V:

• For Gutzeit apparatus — always draw a labelled diagram, even rough — gets 3–4 easy marks
• Always write reactions for each step in arsenic limit test — each reaction carries marks
• For gravimetric analysis: name the weighing form and give the gravimetric factor formula
• In limit test questions, always state: principle → reaction → standard → procedure → comparison → pass/fail criteria
• Include units (ppm, mg, µg) in all limit test calculations/discussions