This is a comprehensive pharmaceutical inorganic chemistry exam prep covering 10 units. Given the breadth of topics, I'll answer all questions directly from my knowledge base in clear, exam-ready language.
📚 Pharmaceutical Inorganic Chemistry — Short Answer Guide (2 Marks Each)
✨ UNIT 1 — General Principles of Analysis
1. Define Error
Error is the difference between the observed (measured) value and the true (accepted) value of a quantity.
Formula: Error = Observed value − True value
2. Difference Between Primary and Secondary Standard
| Feature | Primary Standard | Secondary Standard |
|---|
| Purity | Very high (99.9%) | Less pure |
| Stability | Very stable | Less stable |
| Example | Oxalic acid, K₂Cr₂O₇ | NaOH, KMnO₄ |
| Standardization | Not needed | Must be standardized against primary standard |
3. Define Precision and Accuracy
- Precision: Closeness of repeated measurements to each other (reproducibility). Example: Getting 10.1, 10.1, 10.2 mL in three readings.
- Accuracy: Closeness of a measurement to the true value. Example: True value is 10.0 mL and you get 10.0 mL.
Precision = repeatability; Accuracy = correctness.
4. Define Significant Figures
Significant figures are all the certain digits plus one uncertain (doubtful) digit in a measured value. They indicate the precision of a measurement.
Example: 25.34 mL has 4 significant figures.
5. Define Molality and Mole Fraction
- Molality (m): Number of moles of solute dissolved in 1 kg (1000 g) of solvent.
m = moles of solute / kg of solvent
- Mole Fraction (X): Ratio of moles of one component to the total moles of all components in a solution.
X_A = n_A / (n_A + n_B)
6. pH Range of Phenolphthalein (PHP) and Methyl Orange Indicator
| Indicator | pH Range | Color Change |
|---|
| Phenolphthalein | 8.3 – 10.0 | Colorless → Pink |
| Methyl Orange | 3.1 – 4.4 | Red → Yellow |
7. Types of Errors with Examples
- Constant Error: Same error in every reading. Example: Wrong weight of a standard.
- Proportional Error: Error increases with sample size. Example: Impure standard substance.
- Random (Indeterminate) Error: Due to chance, unpredictable. Example: Reading a burette slightly differently each time.
- Systematic (Determinate) Error: Due to a definite cause. Example: Uncalibrated balance.
✨ UNIT 2 — Acid-Base Titration & Redox Titration
1. What is Redox Titration?
Redox titration is a type of titration based on oxidation-reduction reactions between the titrant and analyte. One substance is oxidized and the other is reduced simultaneously.
Example: Titration of FeSO₄ with KMnO₄.
2. Define Oxidation and Reduction
- Oxidation: Loss of electrons (or gain of oxygen / loss of hydrogen). Oxidation state increases.
Fe²⁺ → Fe³⁺ + e⁻
- Reduction: Gain of electrons (or loss of oxygen / gain of hydrogen). Oxidation state decreases.
Mn⁷⁺ + 5e⁻ → Mn²⁺
3. What is Redox Potential?
Redox potential (electrode potential) is the tendency of a substance to gain or lose electrons compared to a standard hydrogen electrode (SHE = 0.00 V). Higher positive potential = stronger oxidizing agent.
4. What is Oxidation Number?
Oxidation number is the charge an atom would have if all bonds were completely ionic (electrons assigned to the more electronegative atom).
In KMnO₄: Oxidation number of Mn = +7.
5. Applications of Redox Titration
- Assay of hydrogen peroxide (H₂O₂) using KMnO₄.
- Estimation of iron (Fe²⁺) in ferrous sulfate tablets.
- Estimation of Vitamin C (ascorbic acid) using iodine.
- Assay of calcium gluconate using EDTA (indirect redox).
- Determination of bleaching powder (available chlorine).
6. One Theory of Acid-Base Indicator — Ostwald's Theory
According to Ostwald's theory, an acid-base indicator is a weak acid or weak base that has a different color in its ionized and unionized forms.
HIn (unionized, one color) ⇌ H⁺ + In⁻ (ionized, different color)
In acidic medium: equilibrium shifts left → color of HIn.
In alkaline medium: equilibrium shifts right → color of In⁻.
7. Define Alkalimetry Titration
Alkalimetry is the process of estimating the strength of an acid using a standard alkali (base) solution as the titrant.
Example: Titration of HCl with standard NaOH solution.
8. Define Normality and Molarity
- Normality (N): Number of gram equivalents of solute per liter of solution.
N = gram equivalents / volume in liters
- Molarity (M): Number of moles of solute per liter of solution.
M = moles / volume in liters
9. What is a Weak Base?
A weak base is a base that partially ionizes in water, producing a small amount of OH⁻ ions. It establishes an equilibrium in solution.
Example: NH₄OH (ammonium hydroxide): NH₄OH ⇌ NH₄⁺ + OH⁻
10. Different Types of Neutralization Indicators
- Phenolphthalein — used for strong acid vs strong base or weak acid vs strong base (pH range 8.3–10.0).
- Methyl Orange — used for strong acid vs weak base (pH range 3.1–4.4).
- Methyl Red — used for strong acid vs weak base (pH range 4.4–6.2).
- Litmus — general indicator (pH 5–8, red to blue).
✨ UNIT 3 — Precipitation & Non-Aqueous Titration
1. What is Precipitation Titration?
Precipitation titration is a type of titration in which the analyte reacts with the titrant to form an insoluble precipitate. The endpoint is detected when precipitation is complete.
Example: Argentometric titration — AgNO₃ + NaCl → AgCl↓ (white precipitate)
2. What is Non-Aqueous Titration?
Non-aqueous titration is a titration carried out in non-aqueous solvents (not water) for substances that cannot be titrated accurately in water due to weak acidic/basic nature or insolubility.
Example: Titration of aspirin (weak acid) in glacial acetic acid.
3. Principle of Non-Aqueous Titration
The principle is based on the leveling and differentiating effect of solvents. A solvent can enhance (level up) the acidity or basicity of a substance, making it possible to detect the endpoint sharply. Perchloric acid in glacial acetic acid is used for bases; sodium methoxide is used for acids.
4. Classification of Solvents Used in Non-Aqueous Titration
| Type | Example | Property |
|---|
| Protophilic (basic) | Pyridine, acetone | Accept protons |
| Protogenic (acidic) | Glacial acetic acid | Donate protons |
| Aprotic | Benzene, chloroform | Neither donate nor accept protons |
| Amphiprotic | Ethanol, methanol | Both donate and accept protons |
5. Fajans' Rules
Fajans' rules are used in argentometric titration with adsorption indicators (e.g., fluorescein).
- Before endpoint: Excess Cl⁻ ions are adsorbed on AgCl precipitate → indicator stays in solution (yellow-green).
- At/After endpoint: Excess Ag⁺ ions are adsorbed → indicator adsorbs on precipitate → color changes to pink/red.
Example: Fluorescein indicator in NaCl titration with AgNO₃.
6. Volhard's and Modified Volhard's Method
- Volhard's Method: Indirect method for estimation of Cl⁻, Br⁻, I⁻. Excess AgNO₃ is added, and the remaining Ag⁺ is back-titrated with standard NH₄SCN using ferric alum as indicator. Endpoint = red color (Fe(SCN)³).
- Modified Volhard's Method: Used for chlorides specifically. The AgCl precipitate is filtered or coated with nitrobenzene before back-titration to prevent it from reacting with SCN⁻ (since AgCl is more soluble than AgSCN).
✨ UNIT 4 — Complexometry
1. Different Methods in Complexometry
- Direct titration: Metal ion titrated directly with EDTA. Example: Ca²⁺ with EDTA.
- Back titration: Excess EDTA added, then back-titrated with a metal salt.
- Indirect titration: Metal not reacting with EDTA is precipitated, then the precipitate is dissolved and titrated.
- Substitution titration: Metal displaces another metal from its EDTA complex.
- Alkalimetric titration: H⁺ ions released on complex formation are titrated with NaOH.
2. Masking and Demasking Agents
- Masking agents: Substances added to prevent a metal ion from reacting with EDTA without removing it from solution. Example: KCN masks Zn²⁺, Cu²⁺, Ni²⁺.
- Demasking agents: Substances that release the masked metal back to react with EDTA. Example: Formaldehyde demasks Zn²⁺ from Zn-CN complex.
3. Sequestering Agents
Sequestering agents are substances that form stable, soluble complexes with metal ions, preventing them from participating in other reactions (keeping metal ions "hidden" in solution).
Examples: EDTA (ethylenediaminetetraacetic acid), NTA (nitrilotriacetic acid), citric acid, phosphates.
4. Chelating / Complexing Agents
Chelating agents are organic compounds that form ring-like (cyclic) stable complexes called chelates with metal ions by donating two or more electron pairs (polydentate ligands).
Example: EDTA forms a 6-membered ring with Ca²⁺. Used in complexometric titrations.
✨ UNIT 5 — Gravimetry & Limit Tests
1. Define Gravimetric Analysis
Gravimetric analysis is a quantitative analytical method in which the analyte is converted into a stable, insoluble precipitate, which is filtered, dried/ignited, and weighed to calculate the amount present.
It is based on the principle of measurement by weight.
2. Coprecipitation and Post-Precipitation
- Coprecipitation: Contamination of a precipitate by foreign ions that are carried down along with the desired precipitate during precipitation.
Example: BaSO₄ precipitate carrying down NO₃⁻ or Cl⁻.
- Post-precipitation: A second substance precipitates on top of the primary precipitate after the main precipitation is complete.
Example: MgC₂O₄ precipitating over CaC₂O₄.
3. Limitations of Gravimetric Analysis
- Time-consuming — requires filtration, washing, drying, and ignition.
- Not suitable for trace amounts — requires relatively large sample.
- Errors due to coprecipitation and post-precipitation.
- Requires skilled analyst — many steps can introduce error.
- Cannot be automated easily.
4. Sources of Impurities in Pharmaceutical Substances
- Raw materials used in manufacture
- Reagents and solvents used in synthesis
- Intermediates and by-products formed during reaction
- Degradation products during storage
- Packaging materials (heavy metals from containers)
- Environmental contamination (dust, microbes)
5. Reaction in Limit Test of Chlorides
The test is based on the reaction of chloride ions with silver nitrate (AgNO₃) in dilute nitric acid to form a white turbidity (AgCl precipitate):
Cl⁻ + AgNO₃ → AgCl↓ (white) + NO₃⁻
The turbidity of the test solution is compared with a standard chloride solution.
6. Reaction in Limit Test of Sulphates
Sulphate ions react with barium chloride (BaCl₂) in dilute HCl to form a white precipitate of barium sulphate:
SO₄²⁻ + BaCl₂ → BaSO₄↓ (white) + 2Cl⁻
The turbidity is compared with a standard sulphate solution.
7. Principle of Limit Test of Arsenic (Gutzeit Test)
Arsenic compounds are reduced to arsine gas (AsH₃) by zinc and dilute H₂SO₄. The arsine gas reacts with mercuric chloride paper, producing a yellow-brown stain. The intensity of the stain is compared with a standard.
As³⁺ + Zn + H₂SO₄ → AsH₃ (arsine gas) → reacts with HgCl₂ paper → yellow/brown stain
8. Chemical Reaction in Limit Test of Iron
Iron (Fe³⁺) reacts with thioglycolic acid (in ammoniacal solution) to form a purple-colored complex:
Fe³⁺ + thioglycolic acid + NH₃ → Purple color complex
The color is compared with a standard iron solution. (In IP, thioglycolic acid is used in ammoniacal medium.)
9. Apparatus Used in Limit Test of Arsenic
- Flat-bottomed flask (250 mL) — for reaction
- Scrubber tube — containing lead acetate cotton wool (to remove H₂S)
- Arsine generator tube with mercuric chloride paper at the top
- Zinc granules and H₂SO₄ — to generate AsH₃ gas
10. Define Limit Test and Its Types
Limit test is a quantitative/semi-quantitative test performed to check that the amount of an impurity does not exceed the prescribed (permissible) limit as per pharmacopoeia.
Types:
- Limit test for chlorides
- Limit test for sulphates
- Limit test for arsenic
- Limit test for iron
- Limit test for heavy metals (lead)
✨ UNIT 6 — Gastrointestinal Agents
1. What are Antacids? Give Examples.
Antacids are substances that neutralize excess hydrochloric acid (HCl) in the stomach, relieving acidity and heartburn.
Examples: Sodium bicarbonate (NaHCO₃), Aluminium hydroxide Al(OH)₃, Magnesium hydroxide Mg(OH)₂, Calcium carbonate (CaCO₃).
2. Systemic vs Non-Systemic Antacids
| Feature | Systemic Antacids | Non-Systemic Antacids |
|---|
| Absorption | Absorbed into blood | Not absorbed |
| Effect | May cause alkalosis | No systemic effects |
| Example | Sodium bicarbonate | Aluminium hydroxide, Mg(OH)₂ |
| Use | Short-term only | Preferred for chronic use |
3. What is Achlorhydria? Treatment?
Achlorhydria is a condition in which the stomach fails to produce hydrochloric acid (HCl), leading to poor digestion, especially of proteins.
Treatment: Administration of dilute hydrochloric acid (HCl) orally with water, or glutamic acid hydrochloride as an acidifier.
4. Agents Used as Acidifiers
- Dilute Hydrochloric Acid — given orally in achlorhydria
- Glutamic Acid Hydrochloride — safer oral acidifier
- Ammonium Chloride (NH₄Cl) — systemic acidifier used in metabolic alkalosis
5. Medicinal Gases and Their Uses
| Gas | Use |
|---|
| Oxygen (O₂) | Treatment of hypoxia, respiratory failure |
| Carbon dioxide (CO₂) | Respiratory stimulant, used in mixture with O₂ |
| Nitrous oxide (N₂O) | Anaesthetic gas (laughing gas) |
| Helium (He) | Diluent in respiratory therapy (mixed with O₂ for asthma) |
6. Acid-Neutralizing Reaction of Sodium Bicarbonate
NaHCO₃ + HCl → NaCl + H₂O + CO₂↑
Sodium bicarbonate neutralizes stomach HCl, releasing CO₂ gas (causes belching), NaCl (salt), and water.
7. Define Buffers. Two Official Buffers.
A buffer is a solution that resists changes in pH when small amounts of acid or base are added to it. It contains a weak acid and its conjugate base (or weak base and its conjugate acid).
Two Official Buffers (IP):
- Acetate buffer (pH 3.5–5.5) — made from acetic acid + sodium acetate
- Phosphate buffer (pH 5.8–8.0) — made from potassium dihydrogen phosphate + disodium hydrogen phosphate
8. Buffer Capacity and Isotonicity
- Buffer Capacity (β): The ability of a buffer to resist pH change. It is the moles of acid/base required to change the pH of 1 liter of buffer by 1 unit. Higher buffer capacity = more resistant to pH change.
- Isotonicity: A solution is isotonic when its osmotic pressure is equal to that of blood plasma (0.9% NaCl = isotonic). Used in IV fluids and eye drops to prevent cell lysis or crenation.
9. Storage and Labeling Conditions of CO₂
- Stored as a liquid under pressure in steel cylinders (gray/grey colored in IP).
- Labeled: "Carbon Dioxide — Medicinal Gas — For medical use only"
- Cylinders must be stored upright, away from heat, in well-ventilated areas.
- Used as a respiratory stimulant (5–7% CO₂ in O₂).
✨ UNIT 7 — GIT Agents & Electrolytes
1. Types of Cathartics
- Saline cathartics — MgSO₄, Na₂SO₄ (osmotic effect)
- Stimulant/Irritant cathartics — Castor oil, senna (stimulate gut motility)
- Bulk cathartics — Methylcellulose, ispaghula (increase stool bulk)
- Lubricant cathartics — Liquid paraffin (lubricate intestine)
- Emollient cathartics — Docusate sodium (soften stool)
2. Saline Cathartics
Saline cathartics are osmotically active salts that retain water in the intestinal lumen, increasing intestinal pressure and stimulating bowel movement.
Examples: Magnesium sulphate (Epsom salt — MgSO₄·7H₂O), Sodium sulphate (Glauber's salt — Na₂SO₄·10H₂O).
3. Extracellular and Intracellular Electrolytes
- Extracellular electrolytes: Found outside the cell, in plasma and interstitial fluid.
Examples: Na⁺, Cl⁻, HCO₃⁻, Ca²⁺
- Intracellular electrolytes: Found inside the cell (cytoplasm).
Examples: K⁺, Mg²⁺, HPO₄²⁻, proteins
4. Electrolyte Combination Therapy
This is the simultaneous administration of two or more electrolytes to correct multiple deficiencies or maintain electrolyte balance.
Example: Oral Rehydration Salts (ORS) — contains NaCl + KCl + NaHCO₃ + glucose in water. Used in diarrhea and dehydration.
5. Full Form of ORS
ORS = Oral Rehydration Salt(s) / Solution
Used to treat dehydration due to diarrhea and vomiting.
6. GIT Agents
GIT (Gastrointestinal Tract) agents are pharmaceutical substances that act on the gastrointestinal tract. They include:
- Antacids
- Cathartics/Laxatives
- Antidiarrhoeals
- Antiemetics
- Carminatives (relieve gas)
- Digestants (aid digestion)
7. Uses of Aluminium Hydroxide Gel
- Antacid — neutralizes excess stomach acid
Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O
- Phosphate binder — reduces phosphate absorption in renal failure
- Peptic ulcer treatment
- Astringent action — protects mucosa
- Used as adjuvant in vaccines
8. Major Physiological Ions
| Ion | Location |
|---|
| Na⁺ | Extracellular (major cation) |
| K⁺ | Intracellular (major cation) |
| Ca²⁺ | Bone, extracellular |
| Mg²⁺ | Intracellular |
| Cl⁻ | Extracellular (major anion) |
| HCO₃⁻ | Extracellular (buffer) |
| HPO₄²⁻ | Intracellular |
9. Natural Buffer Systems in Human Body
- Bicarbonate buffer system — H₂CO₃/HCO₃⁻ (most important in blood, pH 7.4)
- Phosphate buffer system — H₂PO₄⁻/HPO₄²⁻ (important in urine and cells)
- Protein buffer system — plasma proteins and hemoglobin (act as amphoteric buffers)
10. Role of Sodium, Calcium, Chloride, and Bicarbonate Ions
| Ion | Role |
|---|
| Na⁺ | Maintains extracellular osmotic pressure, nerve impulse transmission, fluid balance |
| Ca²⁺ | Bone/tooth formation, muscle contraction, blood coagulation, nerve function |
| Cl⁻ | Maintains acid-base balance, osmotic pressure, formation of HCl in stomach |
| HCO₃⁻ | Major blood buffer — maintains blood pH at 7.4; transports CO₂ from tissues to lungs |
✨ UNIT 8 — Antimicrobial Agents
1. Antimicrobial Agents — Examples
Antimicrobial agents are substances that kill or inhibit the growth of microorganisms (bacteria, fungi, viruses).
Examples:
- Boric acid (H₃BO₃) — antiseptic
- Hydrogen peroxide (H₂O₂) — antiseptic/disinfectant
- Iodine — antiseptic
- Chlorinated lime — disinfectant
- Potassium permanganate — antiseptic
2. Uses of Boric Acid
- Mild antiseptic — used for eye wash, skin infections
- Antifungal — for vaginal yeast infections
- Preservative in pharmaceutical preparations
- Buffer in ophthalmic solutions
- Used in boric acid lint for wound dressing
3. Why Sulphuric Acid is Added in the Assay of Hydrogen Peroxide
H₂SO₄ is added to:
- Acidify the medium — KMnO₄ works as an oxidizing agent only in acidic medium.
- Prevent decomposition of H₂O₂ in neutral/alkaline conditions.
- Provide Mn²⁺ for the reaction to proceed correctly.
Reaction: 2KMnO₄ + 5H₂O₂ + 3H₂SO₄ → 2MnSO₄ + K₂SO₄ + 8H₂O + 5O₂
4. Molecular Formula of Boric Acid
H₃BO₃ (also written as B(OH)₃)
Molecular weight = 61.83 g/mol. It is a weak, monobasic acid.
5. Preparation and Uses of Iodine
Preparation: Iodine is obtained by oxidation of iodide ions (from seaweed or iodide salts) using chlorine or MnO₂ in acidic medium:
2KI + Cl₂ → 2KCl + I₂
Uses:
- Antiseptic — Lugol's iodine, tincture of iodine (for wounds)
- Thyroid disorders — treatment of hyperthyroidism (pre-surgical)
- Povidone-iodine — surgical scrub
- Analytical reagent — iodometric titrations
6. Use of Glycerine in Boric Acid Assay
Boric acid is a very weak acid (pKa = 9.2) and cannot be directly titrated with NaOH. Glycerine is added to form a stronger glyceroboric acid complex, which can then be accurately titrated with NaOH using phenolphthalein indicator.
H₃BO₃ + Glycerol → Glyceroboric acid (stronger acid) → titrated with NaOH
✨ UNIT 9 — Miscellaneous Pharmaceutical Compounds
1. Dental Caries — Two Anticaries Agents
Dental caries is the decay/destruction of tooth structure caused by acid produced by bacteria (Streptococcus mutans) acting on sugars.
Two anticaries agents:
- Sodium fluoride (NaF) — hardens enamel, inhibits bacterial enzymes
- Stannous fluoride (SnF₂) — antibacterial and enamel-protective
2. Desensitizing Agents
Desensitizing agents are used to reduce tooth sensitivity (pain due to exposed dentin).
Examples:
- Potassium nitrate (KNO₃) — blocks dentinal tubules
- Strontium chloride (SrCl₂) — precipitates in tubules, blocks sensitivity
- Sodium fluoride (NaF)
3. Dentifrices (Dentifricing Agents)
Dentifrices are substances used in toothpaste/tooth powder to aid in cleaning and polishing teeth.
Examples:
- Calcium carbonate (CaCO₃) — mild abrasive
- Dicalcium phosphate (CaHPO₄) — abrasive
- Magnesium carbonate (MgCO₃) — abrasive
- Sodium lauryl sulphate — detergent/foaming agent
4. Two Examples of Antioxidants
Antioxidants prevent oxidative degradation of pharmaceuticals:
- Sodium metabisulphite (Na₂S₂O₅) — used in injections, liquid preparations
- Ascorbic acid (Vitamin C) — natural antioxidant
(Others: BHA, BHT, sodium bisulphite)
5. Water for Injection vs Sterile Water for Injection
| Feature | Water for Injection (WFI) | Sterile Water for Injection (SWFI) |
|---|
| Sterility | Not necessarily sterile (bulk use) | Sterile (filled in sealed containers) |
| Pyrogen test | Passes pyrogen test | Passes pyrogen test |
| Use | Used to make sterile preparations | Used directly as diluent/vehicle |
| Container | Large bulk containers | Sealed ampoules/vials |
6. Bentonite and Magnesium Stearate
| Compound | Formula | Uses |
|---|
| Bentonite | Al₂O₃·4SiO₂·H₂O (hydrated aluminium silicate) | Suspending agent, emulsifying agent, topical protective |
| Magnesium Stearate | Mg(C₁₇H₃₅COO)₂ | Tablet lubricant, anti-adherent in manufacturing |
7. Medicinal Uses of Calcium Carbonate (CaCO₃)
- Antacid — neutralizes stomach acid
- Calcium supplement — treats hypocalcemia, osteoporosis
- Dentifrice — mild abrasive in toothpaste
- Phosphate binder in chronic kidney disease
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
8. Sodium Metabisulphite and Sodium Carboxymethylcellulose (CMC)
| Compound | Formula | Uses |
|---|
| Sodium Metabisulphite | Na₂S₂O₅ | Antioxidant in injections, preservative, reducing agent |
| Sodium Carboxymethylcellulose (Na-CMC) | C₆H₇O₂(OH)₂·OCH₂COONa | Viscosity enhancer, suspending agent, tablet binder, laxative |
✨ UNIT 10 — Radiopharmaceuticals
1. What are Radiopharmaceuticals?
Radiopharmaceuticals are radioactive compounds or drugs used for the diagnosis and treatment of diseases. They contain a radioactive isotope (radioisotope) that emits radiation detectable by imaging equipment.
Example: Sodium Iodide I¹³¹, Technetium-99m.
2. Importance of Radioisotopes in Pharmacy
- Diagnosis — scanning of organs (thyroid, bone, kidney) using gamma emitters
- Therapy — destruction of cancer cells (e.g., I¹³¹ for thyroid cancer)
- Research — radioactive tracers to study drug metabolism and distribution
- Sterilization — gamma radiation sterilizes pharmaceutical products
- Quality control — checking integrity of sealed packages
3. Define Half-Life
Half-life (t₁/₂) is the time required for half (50%) of a radioactive substance to decay (disintegrate).
Example: Half-life of I¹³¹ = 8 days. After 8 days, only half the original activity remains.
4. Uses of Sodium Iodide (I¹³¹)
- Diagnosis of thyroid disorders — thyroid uptake scan
- Treatment of hyperthyroidism (Graves' disease)
- Treatment of thyroid cancer — destroys residual thyroid tissue
- Thyroid imaging — visualize size and function of thyroid gland
5. Define Isotope
Isotopes are atoms of the same element that have the same atomic number (same protons) but different mass numbers (different number of neutrons).
Example: Iodine-127 (stable) and Iodine-131 (radioactive) are isotopes of iodine.
6. Storage Conditions of Sodium Iodide (I¹³¹)
- Stored in lead-lined containers to absorb radiation
- Kept in a cool, ventilated, specially designated radioactive storage area
- Stored away from personnel (distance, shielding, time — radiation safety principles)
- Labeled with radioactive hazard symbols
- Use within the expiry date (short half-life = 8 days)
7. Precautionary Measures for Handling Radioactive Substances
- Distance — maintain maximum distance from source (inverse square law)
- Shielding — use lead aprons, lead glass screens, lead containers
- Time — minimize time spent near radioactive material
- Personal protective equipment (PPE) — gloves, lab coat, dosimeter badge
- No eating/drinking in radioactive areas
- Monitor with Geiger-Müller counter regularly
- Proper disposal of radioactive waste as per regulations
8. Two Properties of Gamma (γ) Radiation
- High penetrating power — can penetrate thick lead and concrete; most penetrating of all radiations (α < β < γ)
- No mass and no charge — gamma rays are electromagnetic waves (photons), not particles; they travel at the speed of light and are not deflected by electric or magnetic fields
9. What are Antidotes?
Antidotes are substances that counteract or neutralize the effects of a poison or drug overdose. They work by chemical binding, competitive antagonism, or physiological reversal.
Examples:
- Atropine — antidote for organophosphate poisoning
- Naloxone — antidote for opioid overdose
- Dimercaprol (BAL) — antidote for heavy metal poisoning
10. One Example of a Respiratory Stimulant
Nikethamide (Coramine) — stimulates the respiratory center in the medulla oblongata, increasing rate and depth of breathing. Used in respiratory depression and CO₂ narcosis.
(Caffeine and ammonia spirit are also used as respiratory stimulants.)
Exam Tip: For 2-mark answers, always give a definition + one example or a definition + one reaction/formula. Keep answers to 3–5 lines. Use chemical formulas wherever possible — examiners reward them!
