B Pharmacy Sem 2: Pharmaceutical Analysis I

Subject 2. Pharmaceutical Analysis I

1. Principles of Quantitative Analysis & Expression of Concentration
2. Acid–Base Titrimetry
3. Complexometric & Precipitation Titrations
4. Redox Titrimetry
5. Gravimetric Analysis
6. Introduction to Instrumental Methods (UV–Vis, IR, Flame Photometry, AAS)

Unit 1: Principles of Quantitative Analysis & Expression of Concentration

This unit introduces the foundations of pharmaceutical analysis, focusing on accuracy, precision, sources of error, and the various ways to express the concentration of solutions—critical for formulation, quality control, and dosing.

1.1 Introduction to Quantitative Analysis

• Definition: Determination of the exact amount or concentration of an analyte in a sample.

• Applications: Assay of active pharmaceutical ingredients (APIs), detection of impurities, formulation verification.

• Key Terms:

  • Analyte: species being measured.

  • Matrix: sample components other than the analyte.

  • Method Validation: demonstrating accuracy, precision, specificity, linearity, range, limit of detection (LOD), limit of quantitation (LOQ).

1.2 Accuracy, Precision & Errors

• Minimizing Errors: proper glassware calibration, consistent technique, instrument maintenance, use of blanks and standards.

1.3 Significant Figures & Reporting Results

• Significant Figures: convey precision—retain only those digits known with certainty plus one estimated digit.

• Rounding Rules: round off at the last significant figure; avoid “false precision.”

• Reporting: include units, measures of dispersion (e.g., mean ± SD), and method conditions.

1.4 Expression of Concentrations

1.5 Solution Preparation & Dilutions

• Preparing Stock Solutions:

  1. Weigh precise mass of pure standard.

  2. Dissolve in volumetric flask to mark.

• Dilution Formula: C1V1=C2V2C_1V_1 = C_2V_2C1​V1​=C2​V2​

  • e.g., to prepare 100 mL of 1 M from 5 M stock: V1=1×1005=20 mLV_1 = \tfrac{1×100}{5} = 20\,\text{mL}V1​=51×100​=20mL.

• Serial Dilutions: stepwise dilutions to reach very low concentrations (e.g., for calibration curves).

1.6 Calibration & Standardization

• Calibration Curve: plot of instrument response (e.g., absorbance) vs. known concentrations.

  • Ensure linearity (R² ≥ 0.999), use at least five standards covering the expected range.

• Standardization: determining the exact concentration of a titrant by titrating against a primary standard (e.g., potassium hydrogen phthalate for acid–base).

1.7 Laboratory Glassware & Good Practices

• Volumetric Flasks: Class A flasks for accurate volumes.

• Pipettes & Burettes: measure or deliver liquids; calibrate periodically.

• Temperature Considerations: volumetric glassware is calibrated at 20 °C; correct for temperature deviations.

• Blanks & Replicates: run reagent blanks to zero instrument, perform measurements in triplicate for precision.

1.8 Key Points for Exams

• Define accuracy vs. precision, and list sources of systematic and random errors.

• State and apply concentration expressions: % w/v, molarity, normality, ppm.

• Perform dilution calculations using C1V1=C2V2C_1V_1=C_2V_2C1​V1​=C2​V2​.

• Explain steps in calibration curve preparation and interpretation (slope, intercept, linearity).

• Identify Class A vs. Class B glassware and their uses.

Unit 2: Acid–Base Titrimetry

This unit delves into volumetric determination of acids and bases. You’ll learn the principles of neutralization titrations, choice of indicators, titration curves, calculation of concentrations, and pharmaceutical applications.

2.1 Principle of Acid–Base Titration

• Neutralization Reaction:

  HA+BOH⟶BA+H2O \text{HA} + \text{BOH} \longrightarrow \text{BA} + \text{H}_2\text{O}HA+BOH⟶BA+H2​O

  where HA = acid, BOH = base.

• End Point vs. Equivalence Point:

  • Equivalence Point: stoichiometric completion (moles acid = moles base).

  • End Point: visible change (indicator color) approximating equivalence.

2.2 Types of Acid–Base Titrations & pH Curves

• Titration Curve Features:

  • Initial pH

  • Buffer region (for weak acid/base titrations)

  • Half‑equivalence point (pH = pKₐ for weak acids)

  • Equivalence point (vertical steep rise/drop)

  • Post‑equivalence plateau

2.3 Primary Standards

Characteristics of a primary standard:

• High purity and stability

• Non‑hygroscopic

• High equivalent weight

• Example: potassium hydrogen phthalate (KHP) for NaOH standardization.

2.4 Indicators

• Choice Criteria: indicator’s pKₐ ≈ pH at equivalence point.

• Common Indicators:

  Indicator	pH Range	Suitable Titration Type
  Methyl Orange	3.1–4.4	Strong acid vs. strong base; strong acid vs. weak base
  Phenolphthalein	8.3–10.0	Strong base vs. strong acid; weak acid vs. strong base
  Bromothymol Blue	6.0–7.6	Strong acid vs. strong base

2.5 Titration Apparatus & Technique

• Equipment:

  • Burette (class A), pipette, conical flask, pH meter (optional).

• Technique:

  1. Rinse and fill burette with titrant; note initial reading.

  2. Pipette measured volume of analyte into flask; add 2–3 drops of indicator.

  3. Titrate with stirring; approach end point slowly.

  4. Record final burette reading; calculate volume used.

  5. Repeat titration to concordant results (± 0.1 mL).

2.6 Calculations

• Molarity:

  MacidVacid=MbaseVbase M_\text{acid}V_\text{acid} = M_\text{base}V_\text{base}Macid​Vacid​=Mbase​Vbase​

• Normality (for polyprotic acids):

  N=eq of soluteL of solution=n×M,n=protons exchanged N = \frac{\text{eq of solute}}{\text{L of solution}} = n \times M,\quad n = \text{protons exchanged}N=L of solutioneq of solute​=n×M,n=protons exchanged

• Example: Titration of 25 mL 0.1 M HCl with 0.1 M NaOH at 24.8 mL at end‑point:

  MHCl=MNaOHVNaOHVHCl=0.1×24.825.0=0.0992  M M_{\rm HCl} = \frac{M_{\rm NaOH}V_{\rm NaOH}}{V_{\rm HCl}} = \frac{0.1 \times 24.8}{25.0} = 0.0992\;\text{M}MHCl​=VHCl​MNaOH​VNaOH​​=25.00.1×24.8​=0.0992M

2.7 Common Sources of Error & Precautions

• Systematic Errors: improperly calibrated glassware, incomplete reaction, indicator mismatch.

• Random Errors: parallax in readings, endpoint overshoot.

• Precautions:

  • Always rinse glassware with the solution it will contain.

  • Perform titrations at controlled temperature (20 ± 2 °C).

  • Use freshly standardized titrant and validated indicator.

2.8 Pharmaceutical Applications

• Assay of Aspirin: titrate phenolic –OH with standard alkali.

• Determination of Alkali in Antacid Tablets: titrate mixture of Mg(OH)₂ and Al(OH)₃ with HCl.

• Blood pH Buffer Capacity: titration of plasma against standard acid/base.

2.9 Key Points for Exams

• Distinguish equivalence vs. end point, and relate to indicator choice.

• Sketch titration curves for strong/weak acid–base pairs, labeling key points.

• Perform titration calculations using molarity and normality.

• List primary standard qualities and give examples.

• Describe common errors and best practices in titrimetry.

Unit 3: Complexometric & Precipitation Titrations

This unit covers volumetric methods based on formation of insoluble precipitates or metal–ligand complexes. You’ll learn the theories, titrants, indicators, titration curves, and pharmaceutical applications for determination of metal ions and anions.

3.1 Principle of Precipitation Titrations

• Definition: analyte reacts with titrant to form an insoluble precipitate, the disappearance of soluble ion marks the end point.

• General Reaction:

  Mx++y Y−⟶MyYx↓ \text{Mx}^+ + y\,\text{Y}^- \longrightarrow \text{M}_y\text{Y}_x\downarrowMx++yY−⟶My​Yx​↓

• End Point Detection: typically via adsorption indicators that precipitate with excess titrant or by back-titration.

3.2 Argentometric Titrations (Silver-Based)

• Mohr’s Method (using chromate indicator):

  • Reaction: Ag⁺ + Cl⁻ → AgCl↓

  • Indicator: K₂CrO₄; at end point, Ag₂CrO₄ (reddish) precipitates.

  • pH Range: 6.5–10 to prevent AgOH or Ag₂O formation.

• Volhard’s Method (back-titration):

  1. Add excess standard AgNO₃ to Cl⁻ sample.

  2. Centrifuge/filter off AgCl.

  3. Titrate remaining Ag⁺ with standard thiocyanate (SCN⁻) using Fe³⁺ indicator:

     Ag++SCN−→AgSCN↓ \text{Ag}^+ + \text{SCN}^- \rightarrow \text{AgSCN}\downarrowAg++SCN−→AgSCN↓

  4. At end point, Fe³⁺ + SCN⁻ → [FeSCN]²⁺ (blood‑red complex).

• Fajans’ Method (adsorption indicator):

  • Indicator: dichlorofluorescein adsorbs onto AgCl, changing from yellow to pink.

3.3 Complexometric Titrations

• Principle: formation of a stable, soluble complex between metal ions and a chelating agent, EDTA being most common.

• Chelating Agent: EDTA (ethylenediaminetetraacetic acid) forms 1:1 complexes with most divalent and trivalent metal ions.

3.4 Indicators in Complexometry

• Working Principle: indicator forms a weak complex with metal; free indicator color returns when EDTA binds all metal.

3.5 Titration Procedure & Calculations

1. Sample Preparation

   • Adjust pH to optimal range (e.g., pH 10 for Ca/Mg with ammonia buffer).

   • Add indicator.

2. Titration

   • Run EDTA solution from burette into sample under constant stirring.

   • Observe color change at end point.

3. Calculation

   MmetalVmetal=MEDTAVEDTA M_{\rm metal}V_{\rm metal} = M_{\rm EDTA}V_{\rm EDTA}Mmetal​Vmetal​=MEDTA​VEDTA​

   where $M$ = molarity, $V$ = volume.

3.6 Back-Titration for Poorly Reactive Metals

• Add known excess EDTA to sample.

• Boil or stand to ensure reaction.

• Back‑titrate unreacted EDTA with standard metal solution (e.g., Zn²⁺) using appropriate indicator.

• Calculate analyte concentration by difference.

3.7 Common Applications

• Water Hardness: determination of Ca²⁺ and Mg²⁺ via EDTA titration.

• Assay of Antacids: titrate Al³⁺ or Mg²⁺ content in formulations.

• Phosphate Estimation: precipitation as MgNH₄PO₄ followed by back‑titration of Mg²⁺ with EDTA.

• Trace Metal Analysis: e.g., Fe³⁺ in ferrous sulfate pills (after oxidation), Cu²⁺ in oral supplements.

3.8 Sources of Error & Precautions

• pH Control: inaccurate pH leads to side reactions or incomplete complexation/precipitation.

• Indicator Dosage: excess indicator can shift end point; too little gives poor color change.

• Titrant Standardization: EDTA solutions must be standardized periodically against primary standard metals.

3.9 Key Points for Exams

• Distinguish Mohr’s, Volhard’s, and Fajans’ methods, including indicators and pH conditions.

• Explain EDTA titration mechanism and select suitable indicator for given metal ion.

• Perform titration calculations for direct and back‑titrations.

• Describe common pharmaceutical applications of precipitation and complexometric titrations.

Unit 4: Redox Titrimetry

This unit focuses on titrations based on oxidation–reduction reactions. You’ll study the principles of redox chemistry, common redox titrants and indicators, titration curves, calculation of analyte concentration, and pharmaceutical applications.

4.1 Principles of Redox Titration

• Oxidation: loss of electrons

• Reduction: gain of electrons

• Redox Reaction:

  Oxidant+ne−⟶Reductant \text{Oxidant} + ne^- \longrightarrow \text{Reductant}Oxidant+ne−⟶Reductant

  or vice versa.

• Equivalence Point: amount of titrant provides exactly the stoichiometric number of electrons to oxidize or reduce the analyte.

4.2 Common Redox Titrants

4.3 Redox Indicators

• Permanganate Self‑Indicator: KMnO₄ is purple; end point when faint pink persists.

• Starch Indicator: forms blue complex with I₂; used in iodometric titrations.

• Diphenylamine Sulfonate: blue color at end point; for Ce⁴⁺ titrations.

• Ferroin: red (Fe²⁺) → blue (Fe³⁺); for iodometric titrations of copper.

4.4 Titration Techniques

1. Direct Titration: analyte reacted directly with titrant (e.g., Fe²⁺ with KMnO₄ in acidic medium).

2. Back Titration: excess oxidant added, then unreacted oxidant titrated (e.g., ASA assay via excess permanganate and back‑titration with oxalate).

3. Iodometric vs. Iodimetric:

   • Iodometric: analyte reduces I₂ to I⁻, then released I₂ titrated with thiosulfate.

   • Iodimetric: I₂ directly used as titrant.

4.5 Titration Curves & Calculations

• Half‑Cell Potentials: at equivalence, Ecell = E°cell

• Nernst Equation:

  E=E°−0.0591nlog⁡[Red][Ox] E = E° – \frac{0.0591}{n} \log\frac{\text{[Red]}}{\text{[Ox]}}E=E°−n0.0591​log[Ox][Red]​

• Example Calculation:
  Titration of 25 mL 0.01 M Fe²⁺ with 0.02 M KMnO₄:

  MnO4−+8 H++5 Fe2+⟶Mn2++5 Fe3++4 H2O \text{MnO}_4^- + 8\,\text{H}^+ + 5\,\text{Fe}^{2+} \longrightarrow \text{Mn}^{2+} + 5\,\text{Fe}^{3+} + 4\,\text{H}_2\text{O}MnO4−​+8H++5Fe2+⟶Mn2++5Fe3++4H2​O

  Equivalence when moles MnO₄⁻ × 5 = moles Fe²⁺.

  VKMnO4=0.01 mol/L×0.025 L0.02 mol/L×5=0.0025 L=2.5 mL V_{\rm KMnO_4} = \frac{0.01\,\text{mol/L} \times 0.025\,\text{L}}{0.02\,\text{mol/L} \times 5} = 0.0025\,\text{L} = 2.5\,\text{mL}VKMnO4​​=0.02mol/L×50.01mol/L×0.025L​=0.0025L=2.5mL

4.6 Sources of Error & Precautions

• Acid Strength: insufficient acid in permanganate titrations yields incomplete reaction.

• Light Sensitivity: I₂ solutions decompose; store in dark.

• Indicator Interference: excess indicator may react; use minimal drops.

• Temperature Control: reaction kinetics influenced by temperature.

4.7 Pharmaceutical Applications

• Assay of Ferrous Sulfate: direct titration with KMnO₄ in acid.

• Determination of Vitamin C: iodometric titration using I₂ and thiosulfate.

• Nitrite Estimation: diazotization followed by iodometric titration.

• Assay of Sulfite Preservatives: titrate with iodine.

4.8 Key Points for Exams

• List common redox titrants and their standard potentials.

• Describe iodometric vs. iodimetric methods.

• Apply stoichiometry to calculate titrant volume at equivalence.

• Identify redox indicators and their color changes.

• Recognize pharmaceutical assays using redox titrimetry.

Unit 5: Gravimetric Analysis

This unit covers methods in which the analyte is quantified by converting it into a pure, stable, weighed form. You’ll learn types of gravimetric procedures, precipitation techniques, requirements for a good gravimetric method, calculation of results, and pharmaceutical applications.

5.1 Principle of Gravimetric Analysis

• Definition: quantitative determination of an analyte by measuring the mass of a pure compound derived from the analyte.

• Key Steps:

  1. Conversion of analyte to an insoluble, well‑defined precipitate.

  2. Isolation by filtration or centrifugation.

  3. Washing to remove impurities.

  4. Drying or Ignition to constant weight.

  5. Weighing of the precipitate.

5.2 Types of Gravimetric Methods

5.3 Requirements for an Ideal Precipitate

• High Purity: precipitate must be free of co‑precipitated impurities.

• Known Stoichiometry: composition precisely defined (e.g., BaSO₄, AgCl).

• Particle Size: suitably large and easily filterable; avoid colloidal particles.

• Stability: stable at drying/ignition temperatures; no decomposition.

• Quantitative Recovery: complete precipitation under controlled conditions.

5.4 Precipitation Techniques & Control

• Controlled Precipitation:

  • Slow Addition of precipitant under stirring to form coarse crystals.

  • Digestion: gentle heating of slurry to promote crystal growth and ripening.

  • Aging: standing time to allow growth of larger crystals, reducing occlusion of impurities.

• Use of Digestion Aids: mild heating or addition of activators (e.g., small amount of precipitate seeds).

• Washing: use appropriate solvents (cold water, dilute reagents) to remove adhering impurities without dissolving precipitate.

5.5 Filtration & Isolation

• Filtration Methods:

  • Gravity Filtration: for coarse, rapidly settling precipitates.

  • Vacuum (Buchner) Filtration: faster, more efficient removal of mother liquor.

• Filter Medium: ash‑free filter paper or sintered glass crucibles to avoid contamination from the filter.

5.6 Drying and Ignition

• Drying: remove surface moisture at moderate temperature (e.g., oven at 105 °C) until constant weight.

• Ignition: heating to higher temperature to convert precipitate to a stable oxide form (e.g., BaSO₄ remains unchanged, but some carbonates are converted to oxides).

• Constant Weight: repeated cycles of heating, cooling in desiccator, and weighing until mass change < 0.5 mg.

5.7 Calculations

• Basic Gravimetric Equation:

  Mass of analyteSample mass=Mprecipitate×nanalyte per formulaMformula \frac{\text{Mass of analyte}}{\text{Sample mass}} = \frac{M_\text{precipitate} \times n_\text{analyte per formula}}{M_\text{formula}}Sample massMass of analyte​=Mformula​Mprecipitate​×nanalyte per formula​​

  where $M$ = molar mass, $n$ = stoichiometric coefficient.

• Example: Determination of sulfate as BaSO₄

  • Sample yields 0.1234 g BaSO₄.

  • Molar mass BaSO₄ = 233.39 g/mol; SO₄²⁻ part = 96.06 g/mol.

  • Sulfate mass = 96.06233.39×0.1234=0.0508 g\tfrac{96.06}{233.39} \times 0.1234 = 0.0508\,\text{g}233.3996.06​×0.1234=0.0508g.

5.8 Sources of Error & Precautions

• Incomplete Precipitation: affects quantitative recovery—ensure conditions favor complete reaction.

• Co‑precipitation: impurities occluded—control particle size and washing.

• Loss of Precipitate: during transfer or filtration—use proper technique, minimize handling.

• Filter Contamination: use ash‑free media and pre‑weigh filters/crucibles.

5.9 Pharmaceutical Applications

• Sulphate Determination: assay of sulfate salts (e.g., in effervescent tablets).

• Chloride Analysis: as AgCl in pharmaceutical excipients.

• Calcium Estimation: as CaC₂O₄ or CaSO₄; important in antacid formulations.

• Active Ingredient Assays: e.g., lead as PbSO₄ for certain formulations.

5.10 Key Points for Exams

• List types of gravimetric methods and give examples.

• Describe ideal precipitate characteristics and control techniques (digestion, aging).

• Outline filtration, drying, and ignition steps to constant weight.

• Perform stoichiometric calculations from precipitate mass.

• Identify common errors in gravimetry and how to avoid them.

Unit 6: Introduction to Instrumental Methods (UV–Vis, IR, Flame Photometry, AAS)

This unit introduces key spectroscopic and photometric techniques used for sensitive, specific quantitative and qualitative analysis of pharmaceutical compounds and trace elements.

6.1 UV–Visible Spectrophotometry

6.1.1 Principle

• Molecules absorb light in the ultraviolet (200–400 nm) and visible (400–700 nm) regions, promoting electrons from ground to excited states.

• Beer–Lambert Law: $A=\epsilon bc$

  • $A$: absorbance (unitless)

  • $\epsilon$: molar absorptivity (L mol⁻¹ cm⁻¹)

  • $b$: path length (cm)

  • $c$: concentration (mol L⁻¹)

6.1.2 Instrumentation

1. Light Source: deuterium lamp (UV), tungsten lamp (visible)

2. Monochromator: disperses light to select wavelength

3. Sample Cell: quartz cuvette for UV, glass for visible

4. Detector: photodiode or photomultiplier tube

5. Readout: display absorbance or transmittance

6.1.3 Applications & Examples

• Assay of Aspirin: measure absorbance of salicylic acid (λₘₐₓ ≈ 530 nm after diazotization)

• Protein Estimation: UV absorbance at 280 nm for aromatic residues

• Validation: linearity, limit of detection (LOD), limit of quantitation (LOQ)

6.1.4 Advantages & Limitations

• + Simple, rapid, inexpensive

• – Interference from overlapping spectra; requires clear solutions

6.2 Infrared (IR) Spectroscopy

6.2.1 Principle

• Molecules absorb infrared radiation (2.5–25 µm) causing vibrational transitions of bonds.

• Functional Group Identification: characteristic absorption bands (e.g., C=O ~ 1 700 cm⁻¹, O–H ~ 3 200–3 600 cm⁻¹).

6.2.2 Instrumentation

1. IR Source: Nernst glower or globar

2. Interferometer/Monochromator: Fourier-transform IR (FTIR) uses interferometer for rapid scanning

3. Sample Accessories: neat liquids (liquid cell), KBr pellet, ATR (attenuated total reflectance) crystal

4. Detector: deuterated triglycine sulfate (DTGS) or mercury cadmium telluride (MCT)

6.2.3 Applications & Examples

• Identification of APIs: confirming functional groups in ibuprofen, paracetamol

• Polymorphism Studies: detect solid-state forms via diagnostic peak shifts

• Excipients Characterization: identify polymers (e.g., HPMC)

6.2.4 Advantages & Limitations

• + Non‑destructive, minimal sample prep (ATR)

• – Overlapping bands in complex mixtures; requires spectral libraries

6.3 Flame Photometry

6.3.1 Principle

• Elemental atoms in a flame absorb and emit light at characteristic wavelengths.

• Emission Mode: measure emitted intensity proportional to concentration.

• Excitation: sample nebulized into flame (typically air–acetylene), free atoms formed.

6.3.2 Instrumentation

1. Nebulizer: converts sample to aerosol

2. Flame Burner: atomizes and excites atoms

3. Monochromator: selects element-specific wavelength

4. Detector: photomultiplier tube

5. Readout: emission intensity → concentration via calibration curve

6.3.3 Applications & Examples

• Sodium & Potassium Assay: in electrolyte formulations

• Calcium in Antacid Preparations

6.3.4 Advantages & Limitations

• + Simple, rapid multi-element capability for alkali/alkaline earth metals

• – Limited to easily atomizable elements; lower sensitivity than AAS

6.4 Atomic Absorption Spectroscopy (AAS)

6.4.1 Principle

• Ground‑state atoms absorb light of element‑specific wavelength from a lamp, reducing intensity.

• Beer–Lambert Law applies to absorbance vs. concentration of free atoms.

6.4.2 Instrumentation

1. Light Source: hollow‑cathode lamp for each element

2. Atomizer: flame (air–acetylene) or graphite furnace (GFAAS) for higher sensitivity

3. Monochromator: isolates analytical line

4. Detector: photomultiplier tube

5. Readout: absorbance → concentration via calibration

6.4.3 Techniques & Parameters

• Flame AAS: detection limits ≈ 0.1–1 ppm

• Graphite Furnace AAS: lower detection limits ≈ ppb; temperature program for drying, pyrolysis, atomization

• Background Correction: deuterium lamp or Zeeman effect to correct non‑specific absorption

6.4.4 Applications & Examples

• Trace Metal Analysis: Pb, Cd, Cu in raw materials and finished products

• Quality Control: ensuring levels below regulatory limits (e.g., heavy metals in herbal extracts)

6.4.5 Advantages & Limitations

• + High specificity and sensitivity; wide dynamic range with furnace mode

• – Single‑element analysis per run; lamp changeover time; higher cost

6.5 Key Points for Exams

• State Beer–Lambert Law and its applications in UV–Vis and AAS.

• Identify characteristic IR absorption bands for major functional groups.

• Compare flame photometry vs. AAS in terms of sensitivity and element scope.

• Describe basic instrument components for each technique: source, monochromator, detector.

• List pharmaceutical applications of UV–Vis, IR, flame photometry, and AAS.