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UNIT – III: Non-Aqueous Titrations & Precipitation Titrations

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① VERY SHORT ANSWER TYPE [2 Marks Each]

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(i) What is Precipitation Titration?

Example: Ag⁺ + Cl⁻ → AgCl↓ (white precipitate)

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(ii) What is Non-Aqueous Titration?

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(iii) Write the Principle of Non-Aqueous Titration

• When a weak base is dissolved in an acidic non-aqueous solvent (e.g., glacial acetic acid), the solvent enhances the basic strength of the analyte by accepting its proton, making it more readily titrated.
• Similarly, when a weak acid is dissolved in a basic non-aqueous solvent (e.g., dimethylformamide), the solvent enhances the acidic strength.
• The levelling effect: A solvent levels all acids stronger than its own conjugate acid to the same apparent strength.
• The differentiating effect: A solvent differentiates between acids/bases of different strengths, giving separate endpoints.

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(iv) Classify Solvents Used in Non-Aqueous Titration

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(v) Discuss Fajan's Rule

However, in context of Fajan's method (precipitation titration), the term refers to the rules governing the use of adsorption indicators:

1. The indicator must be an organic dye whose colour changes when adsorbed onto the precipitate surface.
2. The indicator ion must be opposite in charge to the primary adsorbed ion on the precipitate.
3. The sensitivity of adsorption depends on the surface area of the precipitate — a larger surface area gives a sharper colour change.
4. The pH of the solution must be controlled so the indicator is in its ionic (dissociated) form.
5. Coagulation of the precipitate must be prevented (e.g., by adding dextrin) so the precipitate remains colloidal with a large surface area.

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(vi) What is Volhard's and Modified Volhard's Method of Estimation?

• Developed by Jacob Volhard (1874).
• A known excess of standard AgNO₃ is added to the halide sample.
• The unreacted Ag⁺ is back-titrated with standard ammonium thiocyanate (NH₄SCN) using ferric alum [Fe₂(SO₄)₃·(NH₄)₂SO₄] as indicator.
• Endpoint: Formation of blood-red ferric thiocyanate complex: Fe³⁺ + SCN⁻ → [Fe(SCN)]²⁺ (red)
• Used for: Cl⁻, Br⁻, I⁻, SCN⁻

• In original Volhard's method, AgCl (Ksp > AgSCN) re-dissolves as SCN⁻ is added, causing a premature endpoint.
• In the modified method, after adding excess AgNO₃, the AgCl precipitate is removed by filtration, or nitrobenzene/chloroform is added to coat and protect the AgCl precipitate from reacting with SCN⁻.
• This gives a more accurate endpoint, especially for chloride estimation.

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② LONG ANSWER TYPE [10 Marks Each]

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(i) What are Non-Aqueous Titrations? Explain in Detail the Types of Solvents Used in NAT. Write a Note on Applications in Pharmacy.

Definition

Principle

Types of Solvents

• Neither acidic nor basic; chemically inert.
• Do not participate in proton transfer; act as diluents.
• They show the differentiating effect — can distinguish between acids/bases of different strengths, giving separate endpoints.
• Examples: Chloroform (CHCl₃), Benzene (C₆H₆), Dioxane.
• Use: Titration of mixtures; when individual components are to be quantified separately.

• Donate protons; enhance the strength of weak bases by accepting electrons.
• Show levelling effect for bases — all bases are levelled to the strength of the solvent's conjugate base.
• Examples: Glacial acetic acid (most commonly used), Formic acid, Sulphuric acid, HCl.
• Use: Titration of weak bases (e.g., alkaloids, antihistamines) with perchloric acid as titrant.

• Accept protons; enhance the strength of weak acids.
• Examples: Liquid ammonia, Primary/secondary amines, Dimethylformamide (DMF), Pyridine, Acetone, Ethylenediamine.
• Use: Titration of weak acids (e.g., phenols, sulfonamides, barbiturates) with sodium methoxide as titrant.

• Exhibit both protogenic and protophilic properties; can act as both acid and base.
• Examples: Water, Alcohols (methanol, ethanol, isopropanol), Weak organic acids.
• Use: Titration of soaps, salts of organic acids (alcohols used); alkalimetry in anhydrous conditions.

Key Solvents in Detail

Indicators Used

• Crystal violet (0.5% in glacial acetic acid) — violet in base, yellowish-green in acid; most widely used.
• 1-Naphtholbenzein — used in assay of sodium benzoate.
• Oracet Blue B — blue in base, pink in acid.
• Thymol Blue (0.5% in methanol) — pink to blue; used in alkalimetry.
• Methyl violet, Malachite green, p-Naphtholbenzein.

Titrants

• Acidic titrant: 0.1 M Perchloric acid in glacial acetic acid (most common).
• Basic titrants: Sodium methoxide, Potassium methoxide, Tetrabutylammonium hydroxide.
• Standardization of HClO₄: Against Potassium hydrogen phthalate (KHP) — endpoint: blue → blue-green with crystal violet.

Applications in Pharmacy

1. Assay of weak bases: Ephedrine, Codeine phosphate, Atropine, Antihistamines, Piperazine.
2. Assay of weak acids: Nalidixic acid, Acetazolamide, Fluorouracil, Allopurinol, Mercaptopurine.
3. Assay of salts: Sodium benzoate (weak base salt).
4. Estimation of halogen acid salts of bases (e.g., ephedrine HCl — mercuric acetate is used to replace Cl⁻ with acetate before titration).
5. Quality control of pharmaceutical dosage forms where water-soluble titration is not feasible.

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(ii) What is the Principle Involved in Precipitation Methods of Titration? Briefly Explain with One Example.

Introduction

Requirements for a Successful Precipitation Titration

1. The precipitate formed must be practically insoluble (very low Ksp).
2. The reaction must be stoichiometric and quantitative.
3. The reaction must reach equilibrium rapidly.
4. A suitable method for endpoint detection must be available.
5. Co-precipitation and adsorption errors must be minimal.

Principle

Analyte ion + Titrant ion → Insoluble precipitate

• Pre-equivalence: Excess analyte ion; Ag⁺ is low.
• Equivalence point: Sharp change in pAg.
• Post-equivalence: Excess Ag⁺; pAg rises steeply.

Example: Argentometric Titration of NaCl

• Titrant: 0.1 M AgNO₃
• Indicator: 5% potassium chromate (K₂CrO₄)
• Endpoint: First permanent brick-red/reddish-brown precipitate of Ag₂CrO₄

2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄↓ (reddish brown) — appears only after all Cl⁻ is precipitated.

1 mL of 0.1 M AgNO₃ ≡ 5.845 mg of NaCl

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(iii) Write in Detail the Principle and Procedure Involved in Mohr's, Volhard's, and Fajan's Method.

A. MOHR'S METHOD (Argentometry)

• Developed by K.F. Mohr (1865).
• Used for estimation of Cl⁻ and Br⁻ (directly) using AgNO₃ as titrant.
• Uses potassium chromate (K₂CrO₄) as indicator.
• Ag⁺ first precipitates AgCl (white, Ksp = 1.8 × 10⁻¹⁰) completely. When all Cl⁻ is consumed, excess Ag⁺ reacts with CrO₄²⁻ to form Ag₂CrO₄ (brick-red), signalling the endpoint.

Ag⁺ + Cl⁻ → AgCl↓ (white) — during titration 2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄↓ (brick-red) — at endpoint

• pH: 6.5–10 (neutral to slightly alkaline). In acidic medium, chromate is converted to dichromate (CrO₄²⁻ → Cr₂O₇²⁻); in basic medium, Ag₂O forms.
• Concentration of K₂CrO₄ indicator: ~0.005 M (2–3 drops of 5% solution).
• Not applicable for I⁻ and SCN⁻ (AgI and AgSCN adsorb chromate).
• Interfering ions: PO₄³⁻, S²⁻, Ba²⁺, Pb²⁺ must be absent.
• Solution should not contain ammonia (forms Ag-ammonia complex).

1. Weigh accurately ~0.6 g of dried NaCl.
2. Dissolve in 50 mL distilled water.
3. Add 1 mL of 5% K₂CrO₄ indicator.
4. Titrate with 0.1 M AgNO₃ until brick-red/reddish-brown precipitate persists.
5. Record the volume — 1 mL 0.1 M AgNO₃ ≡ 5.845 mg NaCl.

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B. VOLHARD'S METHOD (Thiocyanometry)

• Indirect (back titration) method developed by Volhard (1874).
• Suitable for Cl⁻, Br⁻, I⁻, SCN⁻ and also for determination of Ag⁺.
• A known excess of standard AgNO₃ is added to precipitate the halide completely.
• The unreacted excess Ag⁺ is back-titrated with standard ammonium thiocyanate (NH₄SCN) in acidic medium.
• Indicator: Ferric alum [Fe₂(SO₄)₃·(NH₄)₂SO₄] or ferric nitrate.
• Endpoint: Blood-red colour of [Fe(SCN)]²⁺ complex.

Ag⁺ (excess) + Cl⁻ → AgCl↓ (white) Ag⁺ (unreacted) + SCN⁻ → AgSCN↓ (white) Fe³⁺ + SCN⁻ → [Fe(SCN)]²⁺ (blood-red) — at endpoint

• Carried out in strongly acidic medium (dilute HNO₃) — prevents hydrolysis of Fe³⁺.
• Not affected by pH constraints like Mohr's method.
• Applicable even in the presence of many interfering anions.

• AgCl (Ksp = 1.8 × 10⁻¹⁰) is more soluble than AgSCN (Ksp = 1.0 × 10⁻¹²).
• SCN⁻ displaces Cl⁻ from AgCl → AgCl + SCN⁻ ⇌ AgSCN + Cl⁻ → premature endpoint.

• After precipitation of AgCl, add nitrobenzene or chloroform (2–3 mL) and shake well to coat AgCl particles.
• This prevents the back-dissolution of AgCl and gives an accurate endpoint.

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C. FAJAN'S METHOD (Adsorption Indicator Method)

• Developed by K. Fajan.
• Uses adsorption indicators — organic dye molecules (anionic in nature) that are adsorbed on the surface of the precipitate at the endpoint, resulting in a colour change.

1. Before equivalence point: AgCl particles adsorb excess Cl⁻ ions → negatively charged surface. Counter-ions are Na⁺. The dye anions (e.g., fluorescein → fluorescein⁻) are repelled from the negatively charged surface → solution appears yellow-green (free fluorescein colour).
2. At/after equivalence point: Excess Ag⁺ is present. AgCl particles adsorb Ag⁺ → positively charged surface. Now the dye anion (fluorescein⁻) is attracted and adsorbs onto the precipitate surface.
3. The adsorbed dye undergoes a structural change → colour changes from yellow-green to pink/red.

1. The indicator must be ionized (correct pH must be maintained).
2. The precipitate must be colloidal (large surface area) — dextrin is added to prevent coagulation.
3. The indicator ion must have opposite charge to excess titrant ion (so it can adsorb after equivalence point).
4. Avoid direct sunlight — silver halides are photosensitive; the adsorbed dye sensitizes them to light.
5. Concentration of analyte must not be too low — insufficient precipitate would not show colour change.

1. Take NaCl solution in conical flask.
2. Add 0.5 g dextrin (to prevent coagulation).
3. Add 5–6 drops of fluorescein indicator.
4. Titrate with 0.1 M AgNO₃ with constant stirring.
5. Endpoint: Yellow-green → pink/red (dye adsorbed on AgCl).

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(iv) Write the Mohr's Method for the Estimation of Halides

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(v) Give the Brief Classification of Solvents Used in Non-Aqueous Titration

1. Aprotic solvents — inert, differentiating, e.g., benzene, chloroform, dioxane.
2. Protogenic solvents — acidic, levelling for bases, e.g., glacial acetic acid, formic acid.
3. Protophilic solvents — basic, levelling for acids, e.g., DMF, liquid ammonia, pyridine.
4. Amphiprotic solvents — both acidic and basic, e.g., alcohols, water.

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(vi) Write a Principle of Non-Aqueous Titration. Write the Preparation and Assay of Sodium Benzoate

Principle of NAT

Estimation of Sodium Benzoate by Non-Aqueous Titration

C₆H₅COO⁻Na⁺ + HClO₄ (in glacial AcOH) → C₆H₅COOH + NaClO₄

1. Measure 8.5 mL of 72% perchloric acid.
2. Add slowly to ~900 mL of glacial acetic acid with constant stirring.
3. Add 30 mL of acetic anhydride (to remove residual water — reacts as: (CH₃CO)₂O + H₂O → 2 CH₃COOH).
4. Make up to 1000 mL with glacial acetic acid.
5. Allow to stand for 24 hours (for complete dehydration).

⚠️ Caution: HClO₄ must be diluted in glacial acetic acid BEFORE adding acetic anhydride — direct contact of HClO₄ with acetic anhydride is explosive.

1. Weigh accurately ~500 mg of potassium hydrogen phthalate (KHP).
2. Dissolve in 25 mL of glacial acetic acid.
3. Add a few drops of crystal violet (0.5% in glacial acetic acid) as indicator.
4. Titrate with the prepared HClO₄ solution.
5. Endpoint: Colour changes from blue → blue-green.
6. Factor: 1 mL of 0.1 M HClO₄ ≡ 0.02042 g KHP.

1. Weigh accurately ~0.25 g of sodium benzoate.
2. Dissolve in 20 mL of anhydrous glacial acetic acid (warm if necessary, then cool).
3. Add 0.05 mL (1 drop) of 1-naphtholbenzein solution as indicator.
4. Titrate with 0.1 M perchloric acid until the colour changes from greenish-yellow → orange-red.
5. Carry out a blank titration and subtract.

% Sodium Benzoate = (V × M_calculated × 0.01441 × 100) / w(g)

• V = volume of HClO₄ consumed (corrected for blank), mL
• M_calculated = actual molarity of HClO₄ (from standardization)
• 0.01441 = equivalent weight factor (1 mL 0.1 M HClO₄ ≡ 0.01441 g C₇H₅NaO₂)
• w = weight of sample in grams

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(vii) Define & Classify Precipitation Titration & the Principle and Ions Involved

A⁺ + B⁻ → AB↓ (precipitate)

• Very low Ksp of the precipitate (quantitative reaction).
• Rapid attainment of equilibrium.
• Stoichiometric reaction.
• Reliable endpoint detection method.

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(viii) Explain Volhard's Method of Estimation of Halides — Write the Mechanism of Indicator in Fajan's Method

Mechanism of Indicator in Fajan's Method (Detailed)

• Solution has excess Cl⁻.
• AgCl particles selectively adsorb Cl⁻ from solution (primary adsorption layer) → surface becomes negatively charged (AgCl·Cl⁻).
• Counter-ions (Na⁺) form the secondary layer.
• Fluorescein anions (Fl⁻) are repelled from the negative surface → remain in solution → yellow-green colour of free fluorescein.

• All Cl⁻ has been precipitated.
• Excess Ag⁺ is now present in solution.
• AgCl particles now adsorb Ag⁺ → surface becomes positively charged (AgCl·Ag⁺).
• The anionic dye (Fl⁻) is now electrostatically attracted and adsorbs onto the positively charged precipitate surface.
• Adsorption causes a structural rearrangement of the dye (quinonoid form) → colour change: yellow-green → pink/red.

Before EP:  [AgCl]·Cl⁻ : Na⁺  +  Fl⁻ (yellow-green, in solution)
After EP:   [AgCl]·Ag⁺ : Fl⁻  (pink/red — adsorbed on precipitate)

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③ SHORT ANSWER TYPE [5 Marks Each]

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① Define Non-Aqueous Titrations. Explain about the Important Conditions of Non-Aqueous Titration.

1. Choice of solvent: Must dissolve the analyte, must enhance its acid/base character, must be compatible with the titrant and indicator.
2. Exclusion of water: All glassware, solvents, and samples must be anhydrous; water interferes severely by acting as a levelling solvent.
3. Temperature control: Non-aqueous titrations are temperature-sensitive; correction factors must be applied if temperature varies significantly from the standardization temperature.
4. Proper indicator: Must be soluble in the non-aqueous solvent, must exhibit a sharp colour change at the endpoint.
5. Standardization of titrant: Titrants (e.g., HClO₄) cannot act as primary standards; must be standardized fresh against KHP.
6. Freshly prepared solutions: Non-aqueous solvents absorb moisture from the atmosphere; solutions must be used promptly.
7. Carbon dioxide exclusion: CO₂ from air dissolves in solvents and can interfere; avoid prolonged exposure.
8. Blank correction: A blank titration must be run to correct for solvent and indicator consumption.
9. Avoid explosive conditions: HClO₄ must be diluted in glacial acetic acid before adding acetic anhydride.

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② What are Precautions to be Taken While Preparing Perchloric Acid Titrant?

1. Order of addition: Always add HClO₄ (72%) slowly to glacial acetic acid (in ~900 mL) with constant stirring — NEVER add acetic anhydride directly to concentrated HClO₄ (explosive reaction).
2. Dilute first: Dilute HClO₄ with glacial acetic acid thoroughly before adding acetic anhydride.
3. Gradual addition of acetic anhydride: Add ~30 mL of acetic anhydride slowly to remove water.
4. Allow 24 hours to stand: For complete dehydration and equilibration.
5. Protect from moisture: Store in tightly stoppered amber bottles.
6. Standardize before use: Against potassium hydrogen phthalate (KHP).
7. Temperature correction: A correction of 0.001 mL per degree change in temperature must be applied.
8. Avoid contact with organic matter in concentrated form — HClO₄ is a strong oxidizer; fire/explosion hazard.
9. Use glass-stoppered burettes — rubber should not be used.

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③ Explain the Types of Solvents Used in Non-Aqueous Titration

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④ What are the Argentometric Titrations?

Ag⁺ + X⁻ → AgX↓ (X = Cl, Br, I, SCN)

• AgCl: 1.8 × 10⁻¹⁰
• AgBr: 5.0 × 10⁻¹³
• AgI: 8.5 × 10⁻¹⁷
• AgSCN: 1.0 × 10⁻¹²

• Estimation of NaCl, KCl, KBr in pharmaceutical preparations.
• Assay of mercaptans, fatty acids.
• Determination of Ag⁺ in photographic industry.
• QC of halide-containing drugs.

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⑤ Explain Different Methods of Estimation of Halides

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⑥ How is Silver Nitrate Standardized?

1. Dry NaCl at 120°C for 1 hour. Weigh accurately ~0.6 g.
2. Dissolve in 50 mL distilled water.
3. Add 1 mL of 5% potassium chromate indicator.
4. Titrate from the burette with the AgNO₃ solution with constant stirring.
5. Endpoint: Brick-red permanent precipitate of Ag₂CrO₄.

Ag⁺ + Cl⁻ → AgCl↓ (white, during titration) 2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄↓ (brick-red, at endpoint)

Molarity of AgNO₃ = (Weight of NaCl × 1000) / (Mol. wt. of NaCl × Volume of AgNO₃ in mL) = (w × 1000) / (58.45 × V)

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⑦ What is Fajan's Method? Explain by Taking Suitable Examples.

AgCl·Ag⁺ + Fl⁻ (yellow-green) → AgCl·Ag⁺·Fl⁻ (pink)

• Eosin (tetrabromofluorescein) is used for Br⁻, I⁻, and SCN⁻.
• Eosin can be used at more acidic pH (2.0–10.3) than fluorescein.
• Endpoint: orange → deep red/magenta adsorption on AgBr precipitate.

• Add dextrin to prevent coagulation of precipitate.
• Titrate under subdued light (avoid photodecomposition).
• Maintain correct pH for indicator ionization.
• Use freshly prepared fluorescein solution.

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UNIT – IV: Complexometric Titrations — Theory of Indicators

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① VERY SHORT ANSWER TYPE [2 Marks Each]

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(i) List Out Different Methods of Complexometric Titrations (CT)

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(ii) What are Masking and Demasking Agents?

• Add thiourea to mask Cu²⁺ → titrate Bi³⁺ with EDTA
• Then add excess EDTA → demask with HNO₂ → titrate Pb²⁺

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(iii) What are Sequestering Agents? Give Examples.

• EDTA (Ethylenediaminetetraacetic acid) — sequesters Ca²⁺, Mg²⁺, Fe³⁺, etc.
• Citric acid — sequesters Fe³⁺ in limit test for iron; prevents precipitation of Fe(OH)₃.
• Sodium hexametaphosphate (Calgon) — sequesters Ca²⁺ and Mg²⁺; softens hard water.
• EGTA — selectively sequesters Ca²⁺ over Mg²⁺.
• DTPA (Diethylenetriaminepentaacetic acid) — used in radiopharmaceuticals to sequester heavy metals.

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(iv) What are Chelating and Complexing Agents?

• Monodentate (one donor atom): NH₃, Cl⁻, CN⁻
• Bidentate (two donor atoms): Ethylenediamine (en)
• Polydentate (multiple donor atoms): EDTA, DTPA

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② LONG ANSWER TYPE [10 Marks Each]

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(i) What are Complexometric Titrations? List Out Different Types of CT with Examples. How Do You Estimate (Calcium) Gluconate?

Definition

Principle

M²⁺ + H₂Y²⁻ → MY²⁻ + 2H⁺ M³⁺ + H₂Y²⁻ → MY⁻ + 2H⁺

• 4 carboxylic acid groups (–COOH) + 2 amine groups (–N) = 6 donor atoms
• Forms highly stable 5-membered chelate rings with metals

Types of Complexometric Titrations

• Metal ion solution + buffer (pH 10) + metal ion indicator → titrate directly with 0.05 M EDTA.
• Endpoint: Indicator-metal complex (coloured) + EDTA → free indicator (different colour).
• Examples: Ca²⁺ (calcium gluconate), Mg²⁺ (magnesium sulphate), Zn²⁺, Cu²⁺, Pb²⁺.

• Excess standard EDTA is added to the metal solution; unreacted EDTA is back-titrated with standard MgSO₄ or ZnSO₄.
• Used when: reaction with EDTA is slow, indicator not available, or precipitation occurs at direct titration pH.
• Example: Al³⁺ (precipitates as Al(OH)₃ at pH 10 → back titration used).

• Analyte metal (M) is added to excess Mg-EDTA complex → M displaces Mg²⁺:

  M + MgY → MY + Mg²⁺

• Released Mg²⁺ is titrated with EDTA.
• Example: Ca²⁺ estimation when indicator gives sluggish endpoint.

• For anions not directly reactive with EDTA.
• Analyte precipitated with a metal, and the metal is titrated.
• Example: SO₄²⁻ → precipitated as BaSO₄ → excess Ba²⁺ titrated with EDTA.

Metal Ion Indicators

M-Indicator (coloured) + EDTA → M-EDTA + free Indicator (different colour)

1. Must form coloured complex with the metal ion.
2. Colour of M-indicator complex must differ clearly from free indicator.
3. M-indicator complex must be less stable than M-EDTA complex.
4. Should be soluble and chemically stable.
5. Reaction should be selective and reversible.

Masking and Demasking Agents (summary)

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Estimation of Calcium Gluconate by Complexometric Titration

Ca²⁺ + Indicator → Ca-Indicator (wine-red/pink) Ca-Indicator + EDTA → Ca-EDTA (stable) + free Indicator (blue) ← endpoint

1. Weigh accurately ~0.12 g of previously ignited calcium carbonate (CaCO₃, primary standard).
2. Dissolve in minimum dilute HCl.
3. Boil to expel CO₂, cool, adjust pH to ~10 with NH₃ buffer.
4. Add 0.1 mL of mordant black 11 indicator → wine-red.
5. Titrate with 0.05 M EDTA until colour changes to pure blue.
6. Factor: 1 mL 0.05 M EDTA ≡ 2.504 mg Ca²⁺.

1. Weigh accurately ~0.4 g of calcium gluconate (tablets/powder) into a conical flask.
2. Dissolve in 50 mL of warm water (calcium gluconate is sparingly soluble — warm water needed).
3. Add 5 mL of 0.05 M disodium EDTA (to prevent Ca²⁺ precipitation).
4. Add 10 mL of strong ammonia–ammonium chloride buffer (to adjust pH to ~10).
5. Add 2–3 drops of mordant black 11 (or EBT) indicator → wine-red colour.
6. Titrate slowly with 0.05 M disodium EDTA with constant swirling.
7. Endpoint: Wine-red → pure blue (free indicator colour).

Each 1 mL of 0.05 M EDTA ≡ 0.02242 g of C₁₂H₂₂CaO₁₄

% Calcium Gluconate = (V × M_calculated × 0.02242 × 100) / Weight of sample (g)

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(ii) What are the Different Types of EDTA Titration? How Do You Prepare and Standardize 0.05 M Disodium EDTA?

Preparation of 0.05 M Disodium Edetate (Disodium EDTA)

1. Weigh accurately 18.61 g of disodium edetate dihydrate.
2. Dissolve in approximately 900 mL of distilled (CO₂-free) water.
3. Make up to 1000 mL in a volumetric flask.
4. Mix well. Label and store in a well-closed amber bottle.

Note: Disodium EDTA cannot act as a primary standard (hygroscopic nature, variable water content) → must be standardized.

Standardization of 0.05 M Disodium EDTA

1. Accurately weigh ~0.12 g of dried CaCO₃.
2. Dissolve in minimum 2 M HCl.
3. Boil to remove CO₂, cool.
4. Add sufficient water and adjust pH to 10 with ammonia–ammonium chloride buffer.
5. Add 0.1 mL of mordant black 11 indicator → wine-red.
6. Titrate with 0.05 M EDTA until colour changes to pure blue (persistent for 30 seconds).

Molarity of EDTA = (Weight of CaCO₃ × 1000) / (100.09 × Volume of EDTA in mL)

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(iii) List Out Different Methods in Complexometric Analysis. Add a Note on Masking and Demasking Agents.

• Total hardness: titrate both Ca²⁺ + Mg²⁺ at pH 10 with EBT.
• Ca²⁺ alone: mask Mg²⁺ with KOH (pH 12 → Mg(OH)₂ precipitates) → titrate Ca²⁺ with murexide indicator.

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(iv) Write the General Principle Involved in CT. What are Ligands and Their Types?

General Principle

M^n+ + EDTA^4⁻ ⇌ [M-EDTA]^(n-4)

K_f = [MY^(n-4)] / [M^n+][Y^4-]

• Before equivalence: [M] decreases gradually.
• At equivalence: sharp rise in pM.
• After equivalence: pM plateau (excess EDTA controls [M]).

Ligands and Their Types

1. O-donor ligands: H₂O, OH⁻, carboxylate groups (–COO⁻)
2. N-donor ligands: NH₃, amines, pyridine, CN⁻
3. S-donor ligands: thiourea, dimercaprol (BAL)
4. Mixed donor ligands: EDTA (both N and O donors)

M(H₂O)₆ + EDTA → M-EDTA + 6H₂O (ΔG more negative — entropy driven)

• EDTA: Antidote for heavy metal poisoning; analytical titrant.
• Citric acid: Sequesters Ca²⁺/Fe³⁺ in formulations and limit tests.
• Dimercaprol (BAL): Antidote for arsenic, mercury, gold poisoning.
• DTPA: Chelation therapy; diagnostic imaging.
• Deferoxamine: Iron chelation in thalassaemia.

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③ SHORT ANSWER TYPE [5 Marks Each]

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① What are Complexometric Titrations? Explain Its Principle with Suitable Examples.

M²⁺ + H₂Y²⁻ → MY²⁻ + 2H⁺

• Correct pH (using buffer — pH 10 NH₃/NH₄Cl for most metals)
• Suitable metal ion indicator
• Absence of interfering ions (or masking if present)

1. Ca²⁺ estimation (Calcium gluconate) — pH 10, EBT indicator, endpoint: wine-red → blue.
2. Mg²⁺ estimation (Magnesium sulphate) — pH 10, EBT indicator, endpoint: red → blue.
3. Water hardness — total Ca²⁺ + Mg²⁺ titrated with EDTA using EBT.
4. Zn²⁺ — direct titration at pH 10, EBT indicator.

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② What are Ligands? Explain Types of Ligands with Examples.

• Monodentate: NH₃, CN⁻ — one bond with metal.
• Bidentate: Ethylenediamine (en), Oxalate — two bonds (chelate ring).
• Polydentate/Chelating: EDTA (hexadentate) — six bonds → very stable complex.
• Chelate effect: Polydentate ligands form more stable complexes due to entropy.

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③ How Do You Prepare and Standardize 0.05 M Disodium EDTA?

• Preparation: Dissolve 18.61 g Na₂H₂EDTA·2H₂O in water → make to 1000 mL.
• Standardization: Against ~0.12 g CaCO₃ (primary standard), at pH 10 (NH₃ buffer), EBT indicator, endpoint wine-red → pure blue.
• Factor: 1 mL 0.05 M EDTA ≡ 2.504 mg Ca ≡ 5.005 mg CaCO₃.

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④ Write in Detail the Principle for Complexometric Titration with Suitable Examples.

• 1:1 metal:EDTA stoichiometry
• Stability constant (K_f) governs completeness
• pH control essential (higher pH → EDTA more ionized → reacts better)
• Metal indicator theory: M-Indicator (weak complex) + EDTA → M-EDTA (strong) + Indicator
• Sharp pM change at equivalence point on titration curve

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UNIT – V: Gravimetric Analysis

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③ SHORT ANSWER TYPE [5 Marks Each]

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① Define Gravimetric Analysis. Write the Importance of It.

Gravimetric factor (GF) = (Molar mass of analyte × stoichiometric ratio) / Molar mass of precipitate

1. Precipitation gravimetry — analyte precipitated by a reagent; most common type.
2. Volatilization gravimetry — analyte converted to a gas and weighed (e.g., CO₂ from carbonate determination).
3. Electrodeposition gravimetry — metal deposited electrolytically on a weighed electrode.
4. Physical adsorption — rare; used for surface area measurement.

1. High accuracy and precision — considered one of the most accurate classical methods.
2. No calibration standards required — absolute method; based only on atomic weights and stoichiometry.
3. Error detection is easy — the precipitate can be examined for purity.
4. Applicable to diverse analytes — ions, gases, organic/inorganic compounds.
5. Pharmacopoeial use — IP/BP/USP use gravimetry for assay of inorganic salts, limit tests (BaSO₄ for sulphate), and residue on ignition.
6. Permanent record — the weighed precipitate can be re-examined.
7. Foundation for titrimetric methods — gravimetric calibration is used to validate titrimetric methods.

• Time-consuming (precipitation, digestion, filtration, ignition can take hours).
• Not applicable for trace quantities.
• Co-precipitation errors can reduce accuracy.

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② Enumerate the Different Steps Involved in Gravimetric Analysis

• Accurately weigh the sample.
• Dissolve completely in appropriate solvent (usually water or dilute acid).
• Adjust volume and pH appropriately.

• Add the precipitating agent (reagent) slowly and in slight excess to the hot, dilute sample solution to form a quantitative precipitate.
• Key factors:
  • Precipitation from dilute, hot solutions to get large crystals (easier to filter).
  • Slow addition with stirring prevents supersaturation.
  • Slight excess of reagent ensures completeness.
• Supersaturation = a state where concentration of dissolved salt exceeds its solubility → nucleation begins. High supersaturation → many small crystals (colloidal, hard to filter); low supersaturation → fewer, larger crystals (preferred).

• Allow the precipitate to stand in the hot mother liquor (at 80–90°C) for 30 min to several hours.
• Small crystallites dissolve and re-precipitate on larger crystals (Ostwald ripening).
• Result: larger, purer, more filterable crystals.

• Filter through ashless filter paper (for ignition) or a sintered glass crucible (for drying).
• Ashless filter paper leaves no residue on ignition.
• Use a porcelain crucible (ignition up to 1000°C) or a glass crucible (drying at 100–120°C).
• Transfer the precipitate quantitatively — rinse beaker 3 times.

• Wash the precipitate on the filter with a cold, dilute solution of an appropriate electrolyte (not water alone — to prevent peptization).
• Peptization: Washing with pure water may redisperse the precipitate into a colloidal state → passes through the filter paper. Prevented by washing with dilute electrolyte (e.g., dilute HNO₃ for AgCl).
• Wash until the filtrate is free of impurities (test with appropriate reagent).

• Drying (110–120°C, 1–2 hrs): When the precipitate is already in a weighable form with known composition (e.g., BaSO₄).
• Ignition (strong heating at 500–1200°C): When the precipitate must be converted to a more stable, well-defined form for weighing.
  • Example: CaC₂O₄·H₂O → ignited → CaO; Fe(OH)₃ → ignited → Fe₂O₃.
• Crucible is cooled in a desiccator before weighing.

• Weigh using a high-precision analytical balance (sensitivity to 0.0001 g).
• Calculate:

Weight of analyte = Weight of precipitate × Gravimetric Factor

GF = (nA × Mwt_analyte) / (nP × Mwt_precipitate)

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③ Draw a Neat and Labelled Diagram of Gutzeit's Apparatus Used for Arsenic Limit Test and Give the Reaction

H₃AsO₄ + KI + HCl → H₃AsO₃ + I₂ + H₂O

H₃AsO₃ + 3[H] → AsH₃↑ + 3H₂O (nascent H from: Zn + HCl → ZnCl₂ + 2[H])

AsH₃ + 3HgCl₂ → AsH(HgCl)₂ + 2HCl (yellow stain) Excess: As(HgCl)₃ → further darkening (brown/black)

     ┌─────────────────────────────┐
     │  HgCl₂ impregnated paper    │  ← Yellow stain appears here
     ├─────────────────────────────┤
     │  Lead acetate cotton wool   │  ← Traps H₂S (from sample) — prevents false positive
     ├─────────────────────────────┤
     │  Glass tube (6 mm bore)     │
     │  ↑ AsH₃ gas passes upward   │
     ├─────────────────────────────┤
     │  Conical flask / Woulff     │  ← Contains:
     │  bottle (125–250 mL)        │    - Test/standard As solution
     │                             │    - Stannated HCl (stannous chloride)
     │                             │    - KI solution
     │                             │    - Zinc (granulated)
     └─────────────────────────────┘

• Stannated HCl (stannous chloride in HCl): reduces As⁵⁺ → As³⁺; removes traces of Sb.
• Potassium iodide (KI): reduces As⁵⁺ → As³⁺ and ensures complete reduction.
• Zinc (granulated, As-free): generates nascent hydrogen.
• Lead acetate cotton/paper: absorbs H₂S (from any sulphide impurities) — prevents false positive brown stain.
• Mercuric chloride paper: reacts with AsH₃ to form yellow stain.
• Standard arsenic solution (10 ppm As): prepared from As₂O₃.

1. Set up the test solution in the Gutzeit bottle with KI, stannous chloride, and HCl.
2. Add zinc dust and quickly close the apparatus.
3. Immerse the bottle in a water bath (25°C) to maintain uniform gas evolution.
4. After 40 minutes, compare the test stain with the standard stain.

• Test stain not more intense than standard → sample passes the limit test.
• Test stain more intense → sample fails the limit test.

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④ Write the Principle and Reactions Involved in the Limit Test for Lead

Pb²⁺ + H₂S → PbS↓ (brown precipitate) + 2H⁺

1. Prepare Test Solution: Dissolve a specified weight of sample in water and adjust pH to 3–4 with acetic acid buffer. Make up to 25 mL in Nessler cylinder.
2. Prepare Standard Solution: Take 2 mL of Lead Standard Solution (10 ppm Pb) → 20 µg Pb → make to 25 mL with the same buffer.
3. To both cylinders, add 10 mL of hydrogen sulphide solution (freshly prepared).
4. Allow to stand for 5 minutes.
5. View the cylinders transversally in daylight against a white background.

• Test colour is not more intense (not darker) than standard → passes limit test.
• Test colour is more intense → fails limit test.

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⑤ Write the Procedure and Principle with Reactions for the Limit Test for Sulphate

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

• Dilute HCl: To prevent precipitation of BaCO₃ or BaHPO₄ (acidification ensures specificity).
• Barium chloride solution (25% w/v): Precipitating agent.
• Standard sulphate solution (100 ppm SO₄²⁻): Prepared from potassium sulphate (K₂SO₄).

1. Prepare Test Solution: Dissolve specified weight of sample in 10 mL of distilled water. Add 1 mL of dilute HCl. Make to 15 mL.
2. Prepare Standard Solution: Take 1.5 mL of 100 ppm sulphate standard → 150 µg SO₄²⁻ → add 1 mL dilute HCl, make to 15 mL.
3. To both cylinders, add 0.5 mL of 25% BaCl₂ solution.
4. Mix immediately and allow to stand for 5 minutes.
5. Compare the turbidity transversally in diffused daylight against a black background.

• Test turbidity not more than standard → sample passes limit test.
• Test turbidity greater than standard → sample fails limit test.

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⑥ Write the Use of Citric Acid, Thioglycolic Acid, and Ammonia in the Iron Limit Test

Fe³⁺ + Thioglycolic acid (HSCH₂COOH) + NH₄OH → Fe-thioglycolate complex (reddish-purple/pink)

1. Dissolve sample in 10 mL of iron-free water in Nessler cylinder.
2. Add 2 mL of citric acid solution → prevents Fe(OH)₃ precipitation.
3. Add 0.1 mL of thioglycolic acid → reacts with Fe³⁺.
4. Make alkaline with strong ammonia solution → reddish-purple colour develops.
5. Make to 20 mL. Allow to stand 5 min.
6. Compare colour with standard (20 ppm Fe) transversally against white background.

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⑦ Give the Role of Acetic Acid and Ammonia in the Limit Test for Heavy Metals

Pb²⁺ (as representative of heavy metals) + H₂S → PbS↓ (brown) + 2H⁺ Cu²⁺ + H₂S → CuS↓ (black) + 2H⁺ Bi³⁺ + H₂S → Bi₂S₃↓ (dark brown) + 6H⁺

1. Dissolve sample in water → add 2 mL dilute acetic acid → adjust pH to 3–4 with dilute ammonia → make to 25 mL.
2. Standard: take 2 mL of 10 ppm Pb standard → treat identically.
3. Add 10 mL freshly prepared H₂S solution to both.
4. Stand 5 min → compare colours transversally.

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⑧ What is the Basis for Fixing the Limits for Impurities?

• Limits are set based on the toxic threshold of the impurity for human use.
• Example: Arsenic limit (≤2–5 ppm) — set based on chronic toxicity; lead limit (≤10 ppm) based on neurotoxicity data.
• Heavy metals expressed as lead equivalent (ppm Pb).

• If a drug is given in large doses chronically, the limit for an impurity is stricter (lower).
• If given in small/infrequent doses, higher impurity levels may be acceptable.
• Formula used: Limit (ppm) = (Permissible daily intake of impurity × 1000) / (Maximum daily dose of drug)

• Limits are set at practically achievable levels using Good Manufacturing Practice (GMP).
• An impurity limit that cannot be consistently achieved by any known manufacturing process is not set.

• Limits must be within the detection range of the specified official test.
• Example: Gutzeit test detects As down to ~1 µg; sulphate test detects ~50 µg SO₄²⁻.

• Limits established over decades of clinical use, where no adverse effects were observed.
• Regulatory agencies (CDSCO, MHRA, FDA) review post-marketing data to update limits.

• For process impurities: 0.10% (1000 ppm) → identify; 0.15% (1500 ppm) → qualify.
• For residual solvents: Class 1 (carcinogens) — very low/avoid; Class 2 — limited daily exposure; Class 3 — low toxicity.

• In pharmacopoeial limit tests, the test result is always compared with a standard prepared simultaneously under identical conditions — ensuring fair, reproducible, objective comparison.
• Limits expressed in ppm (µg/mL or µg/g).

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