← B.Pharm Semester 1

BP106T — Pharmaceutical Inorganic and Analytical Chemistry

3 hrs/week · 3 credits · 45 hours · 75 marks (30 internal + 45 end-semester). Notes follow the PCI B.Pharm (NEP 2020) Semester 1 syllabus, unit by unit.

Units in this subject

1. Unit 1 — Pharmaceutical Analysis, Errors, Impurities
2. Unit 2 — Acid-Base Chemistry and Buffer Systems; Electrolytes
3. Unit 3 — Titrimetric Analysis
4. Unit 4 — Redox Titrations; Gastrointestinal Agents; Radiopharmaceuticals
5. Unit 5 — Miscellaneous Compounds

B.Pharm Semester 1 · NEP 2020 · 45 hours · 75 marks (30 internal + 45 end-sem)

This subject has two halves that examiners love to mix: analytical chemistry (how we test drugs — Units I and III) and inorganic pharmaceuticals (the actual compounds used as medicines — Units II, IV and V). Unit III alone is 14 hours, so expect the heaviest question weight from titrimetry. Work through the units in order; Unit I's limit tests and Unit III's titrations are the backbone of your practicals too (BP111P).

A note on the asterisks (*): compounds like Sodium chloride*, Calcium gluconate*, Sodium bicarbonate* appear with an asterisk in the official PCI syllabus itself. The syllabus does not define what the mark means — most teachers treat these as high-priority compounds (frequently asked, usually with IP monographs). The asterisks are preserved below exactly as in the syllabus.

Unit I — Pharmaceutical Analysis, Errors and Impurities (7 hours)

1.1 Introduction to pharmaceutical analysis

Pharmaceutical analysis answers two questions about any drug sample: what is in it (qualitative) and how much of it is there (quantitative). You do it to confirm identity, check purity, and verify strength — every IP monograph's "Assay" section is built on these techniques.

Techniques of analysis, in brief:

• Titrimetry (volumetric analysis) — measuring the volume of a standard solution needed to react completely with the sample. Fast, cheap, still the workhorse of IP assays.
• Gravimetry — converting the analyte into a pure, weighable precipitate. Slow but very accurate; used as a reference method.
• Instrumental methods — UV-Visible spectroscopy, IR spectroscopy, HPLC, gas chromatography, atomic absorption spectroscopy, and electroanalytical methods like pH-metry and potentiometry. You meet these properly in later semesters; here, just know what each one measures.

Methods of expressing strength of solutions (learn every definition — one-mark questions come straight from here):

• Molarity (M) — moles of solute per litre of solution.
• Normality (N) — gram-equivalents of solute per litre of solution. Normality = Molarity × valency (n-factor).
• Molality (m) — moles of solute per kg of solvent (temperature-independent).
• % w/v — grams of solute per 100 mL of solution. % v/v — mL per 100 mL. % w/w — grams per 100 g.
• Parts per million (ppm) — mg of solute per litre (for very dilute solutions, e.g. impurities).
• Ratio strength — e.g. 1:1000 means 1 g in 1000 mL.
• Milliequivalents (mEq) — used for electrolytes: mEq = (mg × valency) / atomic weight.

Primary and secondary standards:

1.2 Errors in analysis

Sources of errors: instrumental (uncalibrated burette, faulty balance), personal/operative (wrong endpoint judgement, parallax), and method errors (incomplete reaction, co-precipitation). Gross mistakes — like spilling sample — are not "errors" in this classification; they invalidate the analysis.

Types:

• Determinate (systematic) errors — reproducible, same direction every time; can be found and corrected. Subdivided into constant and proportional errors.
• Indeterminate (random) errors — small, unpredictable fluctuations; reduced by replicates, never fully removed.
• Absolute error = measured value − true value. Relative error = (absolute error / true value) × 100.

Minimising errors: calibrate instruments, run blank determinations, analyse replicates and take the mean, use control/standard samples, choose the right sample size (weigh larger quantities to shrink relative error).

Accuracy vs precision: accuracy is closeness to the true value; precision is reproducibility among repeats. Think of a target — tight cluster off-centre is precise but not accurate; scattered shots around the bullseye average out to accurate but not precise. A good analysis needs both.

Significant figures: every measurement is reported only up to the digit you can reasonably read. Rules that get asked: zeros between non-zero digits are significant; trailing zeros after a decimal point are significant; in multiplication/division the answer carries the fewest significant figures of the inputs; in addition/subtraction it follows decimal places, not digit count.

1.3 Impurities in pharmaceuticals

Definition: any unwanted foreign matter present in a drug substance or formulation. Types: organic impurities (starting materials, by-products, degradation products), inorganic impurities (heavy metals, residual catalysts, reagents like arsenic and lead), and residual solvents.

Why regulators care: impurities can be toxic (lead, arsenic), reduce efficacy, or destabilise the product. The IP therefore prescribes a limit test for specific impurities in each monograph — a semi-quantitative check, not a full assay.

Sources of impurities in pharmaceuticals:

1. Raw materials and starting materials
2. Reagents and solvents used in manufacture
3. The manufacturing process and equipment (metal leaching)
4. Containers and closures (glass alkali leaching)
5. Storage conditions (moisture, light, temperature)
6. Decomposition of the drug itself over time

Limit tests — principle and procedure. A limit test compares the opalescence, turbidity or colour produced by the sample with that produced by a standard solution containing a known amount of the impurity, under identical conditions. If the sample's effect is less than the standard's, it passes. Learn each test as: reagent → observation → comparison standard.

• Limit test for chloride: Sample solution + dilute nitric acid + silver nitrate → white opalescence of silver chloride (Ag⁺ + Cl⁻ → AgCl). Compared against a standard prepared from sodium chloride. The nitric acid keeps the medium acidic so that silver carbonate or phosphate don't precipitate and confuse the result.
• Limit test for sulphate: Sample + dilute hydrochloric acid + barium chloride → white turbidity of barium sulphate (Ba²⁺ + SO₄²⁻ → BaSO₄). Standard from potassium sulphate.
• Limit test for iron: Sample + citric acid + thioglycollic acid, made alkaline with ammonia → purple colour of ferrous thioglycollate. Citric acid prevents precipitation of iron by ammonia. Standard from ferric ammonium sulphate.
• Limit test for arsenic (Gutzeit test): The sample is placed in a Gutzeit generator bottle with zinc and hydrochloric acid. Nascent hydrogen reduces any arsenic to arsine gas (AsH₃), which rises and stains mercuric chloride paper yellow to brown. The stain is compared with a standard stain. Key reactions: As₂O₃ + 6Zn + 12HCl → 2AsH₃↑ + 6ZnCl₂ + 3H₂O. Lead acetate cotton in the tube traps H₂S that would give a false stain.
• Limit test for lead: The sample is extracted with dithizone (diphenylthiocarbazone) in chloroform at controlled pH → violet-red lead dithizonate in the chloroform layer, compared against a standard lead solution. Dithizone is chosen because it is highly selective for lead at the prescribed pH.
• Limit test for heavy metals: Sample solution + hydrogen sulphide (or thioacetamide as the H₂S source) → brown to black precipitate of metallic sulphides (PbS, HgS, etc.), compared against a standard lead solution. This is a group test for metals that precipitate as sulphides.

Modified limit test for chloride and sulphate: the normal tests fail when the substance itself interferes — for example a strongly coloured substance (like potassium permanganate) whose colour hides the opalescence, or a substance that reacts with silver nitrate or barium chloride. In such cases the IP monograph prescribes a modified procedure (decolourisation, filtration, or pH adjustment before adding the reagent). The exact pre-treatment differs from monograph to monograph — the exam point is why the modification is needed, not one fixed recipe.

Likely exam angles: principle of the Gutzeit test with reactions; why nitric acid is added in the chloride test; difference between a limit test and an assay; sources of impurities; accuracy vs precision with the target analogy; primary vs secondary standard with two examples each.

FAQs — Unit I

Q: What is the difference between a limit test and an assay? A: A limit test is a semi-quantitative check — it only tells you whether an impurity is below a permitted limit, by comparing with a standard. An assay quantitatively determines how much of the drug is present. Limit tests control impurities; assays confirm strength.

Q: Why is the arsenic limit test called the Gutzeit test? A: After the chemist who devised it. The distinctive feature is the Gutzeit generator bottle that converts arsenic to arsine gas, which stains mercuric chloride paper. Examiners ask for the gas (arsine, AsH₃) and the stain colour (yellow-brown).

Q: What is the difference between accuracy and precision in pharmaceutical analysis? A: Accuracy is closeness of a measurement to the true value; precision is the agreement between repeated measurements. You can be precise without being accurate (reproducible wrong answers, e.g. an uncalibrated burette), which is why calibration matters.

Q: Why can't sodium hydroxide be a primary standard? A: It is hygroscopic and absorbs carbon dioxide from air, so its weighed mass never corresponds to pure NaOH. It must be standardised against a primary standard like potassium hydrogen phthalate.

Key takeaways — Unit I

• Express solution strength in molarity, normality, % w/v, ppm and ratio strength — definitions are direct one-mark questions.
• Primary standards (Na₂CO₃, KHP, oxalic acid) are weighed directly; secondary standards (NaOH, HCl, KMnO₄) must be standardised.
• Determinate errors are correctable, indeterminate errors only reducible; accuracy ≠ precision.
• Limit tests are comparative, not quantitative: chloride (AgNO₃/opalescence), sulphate (BaCl₂/turbidity), iron (thioglycollic acid/purple), arsenic (Gutzeit/arsine stain), lead (dithizone), heavy metals (sulphide precipitate).
• Modified limit tests exist for substances that interfere with the normal test — know the reason, not a memorised recipe.

Unit II — Acid-Base Chemistry, Buffer Systems and Electrolytes (8 hours)

2.1 Acids, bases, buffers and pH

Three definitions of acids and bases get asked as a set — learn all three:

• Arrhenius: acid gives H⁺ in water, base gives OH⁻ in water. Simple, but limited to aqueous solutions.
• Bronsted–Lowry: acid is a proton donor, base is a proton acceptor. Explains conjugate acid–base pairs (e.g. CH₃COOH / CH₃COO⁻).
• Lewis: acid accepts an electron pair, base donates one. The broadest definition.

pH scale: pH = −log₁₀[H⁺]. At 25 °C, pH 7 is neutral; below 7 acidic, above 7 alkaline. Each unit is a tenfold change in [H⁺]. Why pH matters in pharmacy: drug solubility, stability (many drugs hydrolyse at extreme pH), absorption (weak acids absorb better from acidic stomach), and preservative efficacy all depend on pH.

Buffers: solutions that resist pH change when small amounts of acid or alkali are added. Two types — acidic buffers (weak acid + its salt, e.g. acetic acid/sodium acetate) and basic buffers (weak base + its salt, e.g. ammonia/ammonium chloride). Pharmaceutical examples: citrate and phosphate buffers in injections and eye drops.

Buffer equation (Henderson–Hasselbalch):

pH = pKa + log([salt] / [acid])

For a basic buffer: pOH = pKb + log([salt] / [base]). Maximum buffer capacity is at pH = pKa, when [salt] = [acid]. Buffer capacity is the moles of acid or base per litre needed to change the pH by one unit — it is highest when the ratio is 1:1 and increases with total buffer concentration.

Worked pattern for pH calculations (the numerical examiners repeat): for a buffer of 0.1 M acetic acid and 0.1 M sodium acetate (pKa 4.76), pH = 4.76 + log(0.1/0.1) = 4.76. Change the ratio and the log term shifts the pH — practice two or three such sums before the exam.

2.2 Isotonicity

An isotonic solution has the same osmotic pressure as body fluids (equivalent to 0.9% w/v NaCl, freezing point −0.52 °C). A hypertonic solution is more concentrated, hypotonic less.

Why it matters:

• IV fluids: a hypotonic injection causes water to rush into red cells → haemolysis; a hypertonic one draws water out → crenation (shrinking). Large-volume parenterals must be isotonic.
• Ophthalmic solutions: non-isotonic eye drops sting and irritate; isotonicity is adjusted for comfort.

Methods of adjusting isotonicity: freezing-point depression method, sodium chloride equivalent method, White–Vincent method, and Sprowls method. The freezing-point method is the one most often asked: calculate the freezing-point lowering of the drug solution and add enough NaCl to bring the total depression to 0.52 °C.

2.3 Electrolytes

Major extracellular ions: Na⁺, Cl⁻, HCO₃⁻. Major intracellular ions: K⁺, Mg²⁺, phosphate, protein. Functions: maintaining fluid balance and osmotic pressure, acid–base balance, nerve impulse transmission, and muscle contraction (Ca²⁺ in excitation–contraction coupling, Na⁺/K⁺ in the membrane potential).

Electrolytes used in replacement therapy:

• Sodium chloride* — 0.9% w/v "normal saline" for fluid and electrolyte replacement; also the reference for isotonicity and a primary standard in argentometric titrations. The asterisk marks it as a high-priority compound in the syllabus.
• Potassium chloride — treatment of hypokalaemia (low K⁺), e.g. after diuretic therapy or vomiting. Given diluted, never as a rapid IV bolus — concentrated KCl can cause cardiac arrest.
• Calcium chloride — hypocalcaemia, and in cardiac resuscitation; supplies Ca²⁺ rapidly.
• Oral Rehydration Salt (ORS) — for dehydration in diarrhoea. WHO reduced-osmolarity formula per litre of water: sodium chloride 2.6 g, potassium chloride 1.5 g, anhydrous glucose 13.5 g, trisodium citrate dihydrate 2.9 g (total osmolarity 245 mOsm/L). Glucose drives sodium–glucose co-transport in the gut, which is why sugar + salt works better than salt alone.

Physiological acid–base balance: blood pH is held at 7.35–7.45 by the bicarbonate buffer system (H₂CO₃/HCO₃⁻), phosphate and protein buffers, with the lungs (CO₂ removal) and kidneys (H⁺ excretion, HCO₃⁻ reabsorption) as the long-term regulators. Acidosis = pH below 7.35; alkalosis = above 7.45.

Likely exam angles: derive and apply the Henderson–Hasselbalch equation; define buffer capacity; why IV fluids must be isotonic (haemolysis vs crenation); ORS composition with the role of glucose; extracellular vs intracellular ions.

FAQs — Unit II

Q: What is the Henderson–Hasselbalch equation in pharmacy? A: pH = pKa + log([salt]/[acid]). It lets you calculate the pH of a buffer from the ratio of conjugate base to weak acid, and it is the basis for choosing buffer systems for injections, eye drops and liquid formulations.

Q: Why must eye drops be isotonic? A: Tears are isotonic with 0.9% NaCl. A hypotonic or hypertonic drop causes stinging, irritation and tearing, which washes the drug out. Isotonicity adjustment keeps the drop comfortable so the drug stays in contact with the eye.

Q: What is the composition of ORS? A: Per litre: NaCl 2.6 g, KCl 1.5 g, glucose (anhydrous) 13.5 g, trisodium citrate dihydrate 2.9 g. It replaces water and electrolytes lost in diarrhoea; glucose is essential because sodium is absorbed with glucose in the intestine.

Q: What is buffer capacity? A: The amount of strong acid or base (in moles per litre) that a buffer can absorb with a pH change of only one unit. It is maximum when pH = pKa and rises with buffer concentration.

Key takeaways — Unit II

• Three acid–base definitions (Arrhenius, Bronsted–Lowry, Lewis); pH = −log[H⁺].
• Henderson–Hasselbalch: pH = pKa + log([salt]/[acid]); best buffering at pH = pKa.
• Isotonic = same osmotic pressure as body fluids (0.9% NaCl); hypotonic IV → haemolysis, hypertonic → crenation.
• Extracellular: Na⁺, Cl⁻, HCO₃⁻; intracellular: K⁺, Mg²⁺, phosphate.
• Sodium chloride*, potassium chloride, calcium chloride and ORS are the replacement-therapy electrolytes; ORS formula and the glucose co-transport point are exam favourites.

Unit III — Titrimetric Analysis (14 hours)

This is the largest unit and the most scoring one — almost every technique here doubles as a BP111P practical. Learn each titration as: principle → titrant → indicator → endpoint observation → example assay.

3.1 Acid–base titrations

Theories of acid–base indicators:

• Ostwald's (ionisation) theory: the indicator is a weak acid or base whose ionised and unionised forms have different colours. Colour change follows the shift in ionisation with pH.
• Quinonoid theory: the indicator exists in two tautomeric forms — benzenoid and quinonoid — with different colours; pH shifts the equilibrium between them. Phenolphthalein (colourless benzenoid in acid, pink quinonoid in alkali) is the classic example.

Classification of acid–base titrations:

1. Strong acid vs strong base (HCl vs NaOH) — sharp pH jump; methyl orange or phenolphthalein both work.
2. Strong acid vs weak base (HCl vs NH₄OH) — equivalence below pH 7; methyl orange/methyl red.
3. Weak acid vs strong base (CH₃COOH vs NaOH) — equivalence above pH 7; phenolphthalein.
4. Weak acid vs weak base — no sharp pH jump at equivalence, so no direct visual titration; done by other methods.

Preparation and standardisation of titrants:

• Hydrochloric acid: standardised against weighed anhydrous sodium carbonate (primary standard) using methyl orange: Na₂CO₃ + 2HCl → 2NaCl + CO₂ + H₂O.
• Sodium hydroxide: standardised against potassium hydrogen phthalate (KHP) using phenolphthalein. Prepare with carbon-dioxide-free water and protect from air — NaOH absorbs CO₂ forming carbonate, which causes titration errors.

Neutralisation (titration) curves plot pH against volume of titrant. The steep vertical portion marks the equivalence point; an indicator is suitable only if its colour-change pH range falls on that steep part — which is exactly why indicator choice depends on the titration type above.

Assay of ammonium hydroxide: ammonia solution is assayed by titration with standard acid (methyl red indicator). Because ammonia is volatile, the assay is done quickly in a cooled, stoppered flask — a detail examiners like.

3.2 Non-aqueous titrations

Why non-aqueous? Water has a levelling effect — it is too good a proton donor/acceptor, so very weak acids and bases (many drugs) give no sharp endpoint in water. Titrating in a non-aqueous solvent sharpens the endpoint.

Types of solvents:

• Aprotic — inert, neither donate nor accept protons (benzene, chloroform, toluene).
• Protophilic — basic, accept protons; enhance the acidic strength of dissolved acids (dimethylformamide, pyridine, liquid ammonia).
• Protogenic — acidic, donate protons; enhance the basic strength of dissolved bases (glacial acetic acid — the most used solvent).
• Amphiprotic — can both donate and accept (glacial acetic acid, alcohols).

Acidimetry (titration of bases): titrant is perchloric acid in glacial acetic acid, standardised against potassium hydrogen phthalate; indicator crystal violet (violet → green endpoint). Used for weakly basic drugs.

Alkalimetry (titration of acids): titrant is sodium or lithium methoxide in methanol–benzene (or toluene–methanol), standardised against benzoic acid; indicator thymol blue. Used for weakly acidic drugs.

Estimation of sodium benzoate: the classic example — dissolve in glacial acetic acid and titrate with perchloric acid (crystal violet). Sodium benzoate is too weak a base to titrate in water, which is precisely why the non-aqueous route is needed.

3.3 Precipitation titrations and gravimetry

Gravimetric analysis — principle and steps: the analyte is converted into a pure compound of known composition that is precipitated, filtered, washed, dried or ignited, and weighed.

1. Precipitation (with digestion to grow filterable crystals)
2. Filtration
3. Washing (to remove adsorbed impurities)
4. Drying or ignition to constant weight
5. Weighing and calculation via the gravimetric factor

A good gravimetric precipitate must be highly insoluble, pure, easily filterable, and of definite composition after drying/ignition.

Estimation of barium sulphate by gravimetry: Ba²⁺ is precipitated with dilute sulphuric acid as BaSO₄, digested, filtered through ashless paper, ignited and weighed as BaSO₄. Calculation uses the gravimetric factor (Ba/BaSO₄).

Precipitation (argentometric) titrations — titrant is silver nitrate; learn the four named methods as a comparison:

Reactions: Ag⁺ + Cl⁻ → AgCl↓ (white); at Mohr's endpoint 2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄↓ (red). Volhard's: Ag⁺ + SCN⁻ → AgSCN↓; endpoint Fe³⁺ + SCN⁻ → [Fe(SCN)]²⁺ (red).

3.4 Complexometric titrations

EDTA (used as disodium EDTA) forms stable 1:1 complexes with metal ions regardless of the metal's valency — the basis of complexometry.

• Metal ion indicators (e.g. Eriochrome Black T, murexide): these form a coloured complex with the metal ion (EBT: wine red). At the endpoint EDTA strips the metal from the indicator, freeing it — EBT changes wine red → blue.
• Masking and demasking: masking agents (cyanide for heavy metals, fluoride for aluminium, triethanolamine for iron) tie up interfering ions so they don't react with EDTA; demasking releases a specific ion later when needed.
• Preparation and standardisation of EDTA: standardised against a primary standard such as zinc oxide or calcium carbonate (dissolved in acid, buffered to pH 10, EBT indicator).
• Estimation of magnesium sulphate: direct titration with EDTA at pH 10 (ammonia buffer), EBT indicator, wine red → blue endpoint.
• Estimation of calcium gluconate*: same principle — Ca²⁺ titrated with EDTA (murexide or Patton–Reeder indicator is typical for calcium). The asterisk marks it as syllabus-priority.

Likely exam angles: Ostwald vs quinonoid theory; why weak acid–weak base can't be directly titrated; standardisation reactions of HCl and NaOH; why glacial acetic acid is the solvent of choice in acidimetry; Mohr vs Volhard vs Fajans as a table; why AgCl is filtered in modified Volhard's; EBT colour change; masking vs demasking with one example each.

FAQs — Unit III

Q: What is the difference between Mohr's and Volhard's method? A: Mohr's is a direct titration of chloride/bromide with AgNO₃ using potassium chromate indicator in neutral solution. Volhard's is a back titration — excess AgNO₃ is added and the remainder is titrated with thiocyanate using ferric alum indicator in acidic solution. Volhard's suits coloured or acidic samples where Mohr's fails.

Q: Why is glacial acetic acid used in non-aqueous titration? A: It is protogenic — it donates protons and thereby enhances the basic strength of weak bases dissolved in it, giving a sharp endpoint where water's levelling effect would give none. It is also the solvent for the perchloric acid titrant.

Q: What is the endpoint colour change with Eriochrome Black T? A: Wine red (metal–indicator complex) to blue (free indicator) when EDTA removes the metal ion at the endpoint, at pH 10.

Q: What are masking and demasking agents in complexometry? A: Masking agents bind interfering metal ions so they don't consume EDTA (e.g. cyanide masks heavy metals, fluoride masks aluminium). Demasking selectively releases one masked ion afterwards so it can be titrated.

Q: Why is the precipitate digested in gravimetric analysis? A: Digestion (heating the precipitate in its mother liquor) grows larger crystals, which filter easily and adsorb fewer impurities — giving a purer, more accurate result.

Key takeaways — Unit III

• Indicator theories: Ostwald (ionisation) and quinonoid (benzenoid ⇌ quinonoid tautomerism).
• HCl ↔ Na₂CO₃ (methyl orange); NaOH ↔ KHP (phenolphthalein); use CO₂-free water for NaOH.
• Weak acid–weak base titrations have no sharp jump — no direct visual endpoint.
• Non-aqueous: protophilic/protogenic/aprotic/amphiprotic solvents; perchloric acid in glacial acetic acid (crystal violet) for bases; sodium methoxide (thymol blue) for acids; sodium benzoate is the textbook example.
• Mohr (direct, chromate, neutral), Volhard (back, thiocyanate, acidic), modified Volhard (filter AgCl for chlorides), Fajans (adsorption indicators).
• Gravimetry: precipitate → filter → wash → ignite → weigh; BaSO₄ estimation is the standard example.
• EDTA forms 1:1 complexes; EBT wine red → blue at pH 10; masking (cyanide, fluoride) handles interferences; MgSO₄ and calcium gluconate* are the syllabus estimations.

Unit IV — Redox Titrations, Gastrointestinal Agents and Radiopharmaceuticals (10 hours)

4.1 Redox titrations

Concepts: oxidation is loss of electrons, reduction is gain; they always occur together. The oxidising agent is reduced, the reducing agent is oxidised. Redox titrations need a way to see the endpoint — some titrants are self-indicators (KMnO₄), others need external or specific indicators (starch for iodine, ferroin/diphenylamine for cerimetry).

Types:

• Permanganometry: titrant KMnO₄, a self-indicator (purple colour persists at the endpoint). Done in acidic medium (sulphuric acid — never HCl, which gets oxidised to chlorine). Key half-reaction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. KMnO₄ is standardised against oxalic acid or sodium oxalate (heated to ~70 °C). Used to estimate Fe²⁺, hydrogen peroxide, oxalates.
• Cerimetry: titrant ceric sulphate, Ce⁴⁺ → Ce³⁺ (single-electron change, so calculations are simple and the solution is very stable). Indicator ferroin or diphenylamine. Used for Fe²⁺ estimations where permanganate's side reactions are a problem.
• Iodimetry vs iodometry — the pair examiners love to confuse:
• Iodimetry: direct titration with standard iodine solution; the analyte is a reducing agent (e.g. estimation of arsenious oxide, sodium thiosulphate).
• Iodometry: indirect — the analyte is an oxidising agent that liberates iodine from KI; the liberated I₂ is then titrated with standard sodium thiosulphate (e.g. assays of copper sulphate, potassium permanganate). Starch is added near the endpoint (deep blue starch–iodine complex disappears). Key reaction: I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻.
• Titrations with potassium iodate: KIO₃ is a stable primary-standard-grade oxidant (Andrews titration — iodine monochloride endpoint in strong HCl), used where iodine titrations are unsuitable.

4.2 Gastrointestinal agents

Acidifiers (given when stomach acid is deficient — achlorhydria):

• Dilute hydrochloric acid — supplements gastric HCl to aid digestion.
• Sodium acid phosphate — also used as a urinary acidifier.

Antacids — bases that neutralise excess gastric acid in hyperacidity.

• Ideal properties of antacids: rapid onset and reasonably sustained action; should not cause systemic alkalosis; should not be constipating or laxative on its own; palatable; cheap.
• Combinations of antacids: aluminium salts constipate while magnesium salts are laxative — combining them (e.g. aluminium hydroxide + magnesium hydroxide/trisilicate) balances the bowel effect and broadens the action. This "why combine" point is a favourite short question.
• Sodium bicarbonate* — systemic antacid, very rapid, but releases CO₂ (belching, flatulence) and absorbed sodium can cause systemic alkalosis — caution in hypertension and cardiac patients.
• Aluminium hydroxide gel* — non-systemic, slow-acting, constipating; also adsorbs phosphate (used in chronic kidney disease to control phosphate).

Agents that promote bowel movement:

• Magnesium hydroxide (milk of magnesia) — osmotic laxative that also acts as an antacid.
• Sodium orthophosphate — saline purgative; draws water into the bowel osmotically.
• Sodium potassium tartrate (Rochelle salt) — saline purgative.
• Magnesium trisilicate — antacid with adsorbent action; also mildly laxative, which is why it pairs well with aluminium hydroxide.

4.3 Antimicrobials (inorganic)

Inorganic antimicrobials work by oxidation, protein precipitation, or halogenation — a broader, less selective attack than organic antibiotics, which is why most are for external/antiseptic use.

• Potassium permanganate — strong oxidising agent; used as a gargle/mouthwash and wound wash in very dilute solution (around 1:10000). Stains skin brown — a practical point worth remembering.
• Boric acid — weak antiseptic; used in eye washes and dusting powders. Toxic if absorbed in large amounts, especially in infants — never take casually.
• Hydrogen peroxide* — 3% solution releases nascent oxygen with effervescence (2H₂O₂ → 2H₂O + O₂), which mechanically cleans wounds. Decomposes on storage, so it's kept in light-resistant, cool containers.
• Chlorinated lime* (bleaching powder) — releases chlorine on exposure to air/moisture; used as a disinfectant for water and surfaces.
• Iodine and its preparations — tincture of iodine, povidone-iodine (less staining, sustained release), Lugol's solution. Broad-spectrum antiseptic; the brown colour itself shows where it has been applied.

4.4 Radiopharmaceuticals

Basics of radioactivity: unstable nuclei decay by emitting alpha, beta or gamma radiation. Half-life (t½) is the time for half the atoms to decay — it decides whether an isotope suits diagnosis (short t½, e.g. hours) or therapy. Activity is measured in becquerel (Bq) and curie (Ci).

Applications of radioisotopes:

• Sodium iodide I-131 (t½ ~8 days) — thyroid diagnosis and therapy (thyroid cancer, hyperthyroidism); the thyroid selectively takes up iodine.
• Technetium-99m (t½ ~6 hours) — the workhorse of diagnostic imaging (bone, heart, kidney scans); obtained from a molybdenum-99 generator; short half-life keeps patient radiation dose low.
• Cobalt-60 (t½ ~5.27 years) — external beam radiotherapy for cancer (gamma rays).
• Phosphorus-32 (t½ ~14.3 days) — therapy of polycythaemia vera and some bone conditions (beta emitter).

Safe handling, storage and disposal: follow the ALARA principle (As Low As Reasonably Achievable) — minimise time, maximise distance, use shielding (lead containers). Store in labelled, shielded areas; dispose through authorised radioactive-waste channels. In India this is regulated under AERB (Atomic Energy Regulatory Board) guidelines — the exam expects you to name safe-handling principles and the regulatory adherence, not the full rulebook.

Likely exam angles: iodimetry vs iodometry with one example each; why H₂SO₄ and not HCl in permanganometry; ideal properties of antacids + why antacids are combined; half-lives and one use each of I-131, Tc-99m, Co-60, P-32; ALARA.

FAQs — Unit IV

Q: What is the difference between iodimetry and iodometry? A: In iodimetry you titrate directly with standard iodine — the substance being estimated is a reducing agent. In iodometry the substance is an oxidising agent: it liberates iodine from potassium iodide first, and you titrate that liberated iodine with standard sodium thiosulphate. Starch indicator goes in near the endpoint in both.

Q: Why is sulphuric acid used instead of hydrochloric acid in permanganometry? A: Hydrochloric acid is itself oxidised by permanganate to chlorine gas, which consumes extra KMnO₄ and gives a wrong, high result. Sulphuric acid is not oxidised under these conditions.

Q: Why are antacids combined instead of given singly? A: To balance side effects and improve action — aluminium compounds constipate while magnesium compounds loosen the bowel, so a combination neutralises acid without upsetting bowel habit, and may combine fast and slow onset.

Q: What is the half-life of Technetium-99m and why does it matter? A: About 6 hours. Short enough that the patient gets a low radiation dose, long enough to prepare and complete a scan — which is why it's the most widely used diagnostic radioisotope.

Key takeaways — Unit IV

• Redox: permanganometry (KMnO₄, self-indicator, H₂SO₄ medium), cerimetry (Ce⁴⁺→Ce³⁺, ferroin), iodimetry (direct, analyte reduces) vs iodometry (indirect, analyte oxidises KI; titrate liberated I₂ with thiosulphate; starch near endpoint), potassium iodate titrations.
• Acidifiers: dilute HCl, sodium acid phosphate. Ideal antacid = fast, sustained, non-systemic, bowel-neutral, palatable. Combinations balance Al (constipating) with Mg (laxative). Sodium bicarbonate* (systemic, CO₂, alkalosis risk); aluminium hydroxide gel* (non-systemic, slow, phosphate-binding).
• Bowel agents: magnesium hydroxide, sodium orthophosphate, sodium potassium tartrate, magnesium trisilicate — mostly osmotic/saline action.
• Inorganic antimicrobials act by oxidation/protein precipitation: KMnO₄, boric acid, hydrogen peroxide* (effervescent O₂), chlorinated lime* (chlorine release), iodine preparations.
• I-131 (thyroid, 8 d), Tc-99m (imaging, 6 h), Co-60 (radiotherapy, 5.27 y), P-32 (polycythaemia vera, 14.3 d). ALARA + shielding + AERB-regulated disposal.

Unit V — Miscellaneous Compounds (6 hours)

Short unit, but every compound here is a direct "write a short note" question. Learn each as: use → key property → one caution.

5.1 Expectorants

Expectorants increase bronchial secretion or make sputum less viscous so it can be coughed out.

• Potassium iodide — stimulates bronchial glands to increase respiratory tract fluid, thinning mucus. Also the classic reagent for liberating iodine in iodometry (Unit IV crossover — examiners like linking these).
• Ammonium chloride* — reflex expectorant (irritates gastric mucosa → vagal reflex → increased bronchial secretion). Also a systemic acidifier and diuretic. Large doses can cause systemic acidosis — the caution to quote.

5.2 Emetics

Emetics induce vomiting, used in certain poisonings to empty the stomach (before safer methods like gastric lavage took over).

• Copper sulphate* — prompt emetic; also an astringent and fungicide. Excess is toxic (copper poisoning), so the dose window matters.
• Sodium potassium tartrate (Rochelle salt) — listed here as an emetic; it is better known as a saline purgative (you also met it in Unit IV). Know both roles.

5.3 Haematinics

Haematinics supply iron for haemoglobin synthesis in iron-deficiency anaemia.

• Ferrous sulphate* — the standard oral iron salt; ferrous (Fe²⁺) iron is absorbed far better than ferric (Fe³⁺). Given with vitamin C to keep iron reduced and improve absorption; common side effects are constipation and black stools — warn the patient so they don't panic.
• Ferrous gluconate — milder on the stomach than ferrous sulphate, often preferred when the sulphate causes gastric irritation; slightly lower elemental iron content.

5.4 Poison and antidote

• Poison: any substance that causes injury, illness or death on entering the body in relatively small amounts.
• Antidote: a substance that counteracts a poison. Classification: (1) physiological/pharmacological antidotes — oppose the poison's action on the body (e.g. physostigmine for atropine poisoning); (2) chemical antidotes — react with the poison to form a harmless compound (e.g. sodium thiosulphate for cyanide); (3) mechanical antidotes — prevent absorption (e.g. activated charcoal adsorbing the poison, demulcents coating the gut).
• Sodium thiosulphate — the syllabus antidote: used in cyanide poisoning (with sodium nitrite — it converts cyanide to non-toxic thiocyanate, which is excreted) and to neutralise excess iodine/chlorine. Reaction to remember: it reduces iodine to iodide — the same chemistry as the iodometric titration endpoint (I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻).

Likely exam angles: short notes on any asterisked compound; classification of antidotes with one example each; why ferrous (not ferric) salts are used as haematinics; mechanism of sodium thiosulphate in cyanide poisoning.

FAQs — Unit V

Q: Why is ferrous sulphate preferred over ferric salts for anaemia? A: Iron is absorbed in the ferrous (Fe²⁺) state in the duodenum. Ferric iron must first be reduced before absorption, so ferrous salts like ferrous sulphate give better, more predictable absorption.

Q: What is the classification of antidotes? A: Physiological (oppose the poison's effect on the body), chemical (react with the poison to detoxify it, e.g. sodium thiosulphate for cyanide), and mechanical (prevent absorption, e.g. activated charcoal).

Q: How does sodium thiosulphate act as an antidote? A: In cyanide poisoning (given with sodium nitrite) it donates sulphur to convert toxic cyanide into thiocyanate, which the kidneys excrete. It also neutralises halogens like iodine and chlorine by reducing them.

Q: Why does ammonium chloride cause acidosis in large doses? A: It is a systemic acidifier — the liver converts the ammonium part to urea, releasing hydrochloric acid equivalents, which lowers blood pH if the dose is excessive.

Key takeaways — Unit V

• Expectorants: potassium iodide, ammonium chloride* (reflex expectorant; systemic acidifier in large doses).
• Emetics: copper sulphate*, sodium potassium tartrate (also a saline purgative).
• Haematinics: ferrous sulphate* (standard; Fe²⁺ absorbed best; give with vitamin C; black stools are normal), ferrous gluconate (gentler on the stomach).
• Antidotes: physiological, chemical, mechanical; sodium thiosulphate for cyanide (→ thiocyanate) and halogens.

End of BP106T notes. Pair these with your BP111P practical record — limit tests (Unit I) and titrations (Units III–IV) are performed, not just written, in the end-semester practical exam.

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