B Pharmacy Sem 1: Pharmaceutical Analysis

Subject 2. Pharmaceutical Analysis

1. 1. Fundamentals of Analytical Chemistry
   2. Acid–Base Titrations
   3. Precipitation & Complexometric Titrations
   4. Oxidation–Reduction Titrations
   5. Gravimetric Analysis
   6. Introduction to Instrumental Methods

Unit 1: Fundamentals of Analytical Chemistry

In this unit, we introduce the principles and terminology of analytical chemistry—the branch of chemistry concerned with the identification, separation, and quantification of chemical substances. Mastery of these basics is essential before performing any analytical procedure.

1.1 Definition and Scope

• Analytical Chemistry studies qualitative and quantitative determination of chemical components in samples.

• Qualitative Analysis answers “What is present?” (e.g., identifying anion or cation in a drug substance).

• Quantitative Analysis answers “How much is present?” (e.g., % assay of active pharmaceutical ingredient).

Analytical methods are used throughout the pharmaceutical lifecycle: raw material testing, in‑process control, finished product release, and stability studies.

1.2 Classification of Analytical Methods

1. Classical (Wet) Methods

   • Based on gravimetric or titrimetric procedures.

   • Typically simple equipment, direct measurement of mass or volume.

2. Instrumental Methods

   • Rely on measuring physical properties (absorbance, emission, mass‑to‑charge ratio, etc.).

   • Examples: UV‑Visible spectrophotometry, HPLC, GC, AAS, mass spectrometry.

3. Separation Techniques

   • Chromatographic (e.g., TLC, HPLC, GC) and electrophoretic methods isolate components prior to detection.

1.3 Key Analytical Terminology

• Analyte: Substance to be measured.

• Matrix: The sample’s background material (e.g., cream base, tablet excipients).

• Method: Protocol defining reagents, conditions, and measurements.

• Calibration: Establishing the relationship between instrument response and known analyte concentrations.

1.4 Sampling and Sample Preparation

Accurate analysis begins with a representative sample:

• Sampling Plan: Defines how to collect portions that reflect the whole batch.

• Sample Preparation: Includes drying, grinding, dissolution, filtration, and dilution to produce a homogeneous test solution.

Poor sampling or prep introduces bias and inaccuracy.

1.5 Error and Uncertainty

• Systematic Error (Bias): Consistent deviation from true value (e.g., imperfect calibration).

• Random Error (Precision): Scatter in repeated measurements due to uncontrollable fluctuations.

Accuracy describes closeness to true value; precision describes reproducibility under unchanged conditions.

1.6 Figures of Merit

• Sensitivity: Slope of the calibration curve; change in response per unit concentration.

• Limit of Detection (LOD): Lowest analyte level that can be distinguished from background noise (typically signal-to-noise ratio of 3:1).

• Limit of Quantification (LOQ): Lowest level that can be quantified with acceptable precision and accuracy (signal-to-noise ratio of 10:1).

• Linearity: Range over which response is directly proportional to concentration (expressed by correlation coefficient, r).

• Selectivity/Specificity: Method’s ability to measure analyte in presence of other components (matrix, impurities).

1.7 Calibration Techniques

1. External Standard Method

   • Prepare calibration standards of known concentrations.

   • Measure their instrument response to build a calibration curve.

   • Determine sample concentration by interpolating its response.

2. Internal Standard Method

   • Add a compound (internal standard) of known concentration to all standards and samples.

   • Corrects for sample loss and instrumental drift.

3. Standard Addition Method

   • Known increments of analyte are added directly to the sample.

   • Compensates for matrix effects that alter response.

1.8 Method Validation

Before routine use, analytical methods must be validated according to regulatory guidelines (e.g., ICH Q2[R1]):

1.9 Summary of Unit 1

• Analytical chemistry is divided into qualitative and quantitative methods.

• Proper sampling and sample preparation are critical to avoid bias.

• Understanding error types, figures of merit, calibration, and method validation ensures reliable, reproducible results.

• These fundamentals underpin every analysis you perform in pharmaceutical quality control.

Unit 2: Acid–Base Titrations

Acid–base titrations are quantitative methods used to determine the concentration of an acid or a base in solution by reacting it with a solution of known concentration (the titrant). This unit covers the theory, procedure, calculation, and choice of indicators for acid–base titrations.

2.1 Basic Concepts

• Acid: Substance that donates protons (H⁺) in solution (Bronsted–Lowry definition).

• Base: Substance that accepts protons (H⁺).

• pH: Negative logarithm of the hydrogen‑ion concentration, pH = –log₁₀[H⁺].

• pKa: Negative logarithm of the acid dissociation constant (Ka), pKa = –log₁₀Ka. A lower pKa indicates a stronger acid.

2.2 Principle of Titration

1. Titrant (standard solution) of known concentration is added gradually to a measured volume of analyte (unknown concentration).

2. The reaction proceeds until the stoichiometric (equivalence) point, where exactly equivalent moles of acid and base have reacted.

3. Indicator or pH meter is used to detect the end point, which approximates the equivalence point.

The basic reaction:

• For a monoprotic acid (HA) titrated with a strong base (MOH):
  HA+OH−→A−+H2OHA + OH⁻ → A⁻ + H₂OHA+OH−→A−+H2​O

2.3 Types of Acid–Base Titrations

1. Strong Acid vs. Strong Base

   • Example: HCl titrated with NaOH.

   • Equivalence pH ≈ 7.

   • pH curve is steep near equivalence; any indicator that changes color around pH 7 (e.g., phenolphthalein, pH 8.2–10; methyl orange, pH 3.1–4.4) may be used.

2. Strong Acid vs. Weak Base

   • Example: HCl titrated with NH₄OH.

   • Equivalence pH < 7 (acidic) because the conjugate acid of the weak base hydrolyzes.

   • Suitable indicators change color in an acidic range (e.g., methyl orange).

3. Weak Acid vs. Strong Base

   • Example: CH₃COOH titrated with NaOH.

   • Equivalence pH > 7 (basic) because the conjugate base (acetate) hydrolyzes.

   • Phenolphthalein (color change ~8.2–10) is commonly used.

4. Weak Acid vs. Weak Base

   • Example: CH₃COOH titrated with NH₄OH.

   • Equivalence pH ≈ intermediate (near neutral), but pH change is gradual.

   • Not commonly used analytically because end‑point detection is difficult.

2.4 pH Curve and Equivalence Point

• Titration Curve: Plot of pH versus volume of titrant added. Key regions:

  1. Initial pH: Depends on strength and concentration of analyte.

  2. Buffer Region (for weak acids or bases): pH changes gradually; can apply the Henderson–Hasselbalch equation.

  3. Steep Rise or Fall near equivalence.

  4. Post‑Equivalence: Excess titrant determines pH.

• Equivalence Point: Volume of titrant where moles of H⁺ = moles of OH⁻.

• End Point: Volume at which the indicator changes color; should coincide closely with the equivalence point.

2.5 Choice of Indicator

An indicator is a weak acid or base whose color differs between its protonated and deprotonated forms. Selection depends on the pH range over which the titration curve’s steep region occurs:

2.6 Titration Calculations

1. Before Equivalence

   • For strong acid/strong base, pH can be calculated directly from [H⁺] or [OH⁻] of the mixture.

   • For weak acid/strong base (or vice versa), use Henderson–Hasselbalch:
     pH=pKa+log⁡[A−][HA]\text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]}pH=pKa​+log[HA][A−]​

2. At Equivalence

   • All analyte is converted to its conjugate.

   • Calculate pH by considering hydrolysis of the conjugate (A⁻ or BH⁺):
     A−+H2O⇌HA+OH−A^- + H_2O \rightleftharpoons HA + OH^-A−+H2​O⇌HA+OH−
     Use Kb=KwKaK_b = \frac{K_w}{K_a}Kb​=Ka​Kw​​ and solve for [OH⁻], then pH = 14 – pOH.

3. After Equivalence

   • Excess titrant determines pH. For strong base in excess, calculate [OH⁻] = (moles excess)/(total volume).

Example Calculation
Determine the concentration of acetic acid if 25.00 mL of it requires 30.00 mL of 0.100 M NaOH to reach the equivalence point.

1. Moles NaOH = 0.03000 L × 0.100 mol/L = 0.003000 mol

2. At equivalence, moles CH₃COOH = moles NaOH = 0.003000 mol

3. Concentration CH₃COOH = 0.003000 mol / 0.02500 L = 0.120 M

2.7 Practical Procedure

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

2. Pipette a measured volume of analyte into a conical flask; add a few drops of chosen indicator.

3. Titrate by adding titrant dropwise, swirling flask constantly.

4. Near the expected end point, add titrant dropwise until the persistent color change.

5. Record final burette reading; calculate volume used.

6. Perform at least three titrations and take the average volume for accuracy.

2.8 Sources of Error and Precautions

• Air Bubbles in burette tip: lead to volume errors—always remove bubbles before titration.

• Indicator Overshoot: adding too much titrant past the end point—add dropwise and swirl well.

• Temperature Variations: can affect reaction equilibria and indicator color—conduct at constant temperature.

• Proper Calibration: regularly standardize titrant against a primary standard.

Unit 3: Precipitation and Complexometric Titrations

This unit covers two important classes of volumetric methods—precipitation titrations, which rely on the formation of insoluble precipitates, and complexometric titrations, which form soluble metal–ligand complexes. Both are widely used for assay of pharmaceutical ingredients, water hardness determination, and metal impurity analysis.

3.1 Precipitation Titrations

3.1.1 Principle

A precipitation titration involves adding a titrant that reacts with the analyte to form a sparingly soluble compound. At the equivalence point, stoichiometric amounts of reactants have combined, and further addition of titrant causes a detectable change (often in the appearance or disappearance of an indicator color).

3.1.2 Solubility Product (Kₛₚ)

For a salt $MX$ that dissociates as

MX(s)⇌M++X−MX_{(s)} \rightleftharpoons M^+ + X^-MX(s)​⇌M++X−

the solubility product constant is

Ksp=[M+] [X−]K_{sp} = [M^+]\,[X^-]Ksp​=[M+][X−]

Precipitation begins when the ionic product [M+][X−][M^+][X^-][M+][X−] exceeds KspK_{sp}Ksp​.

3.1.3 Argentometric Titrations (Silver Nitrate Titrations)

• Mohr’s Method

  • Titrate a halide (Cl⁻, Br⁻) solution with standard AgNO₃.

  • Use potassium chromate indicator: Ag⁺ + Cl⁻ → AgCl (white precipitate). At end point, excess Ag⁺ reacts with chromate to form red Ag₂CrO₄.

  • Applicable pH: 6.5–10 to prevent AgOH or HCrO₄⁻ interferences.

• Volhard’s Method (Back Titration)

  1. Add known excess of AgNO₃ to analyte containing halide.

  2. Filter off precipitated AgX.

  3. Titrate unreacted Ag⁺ with standard thiocyanate (KSCN) using ferric ion as an indicator (Fe³⁺ + SCN⁻ → [FeSCN]²⁺, red color).

  • Suitable for colored or turbid solutions where direct titration is difficult.

• Fajans’ Method

  • Direct titration of halides with AgNO₃ using adsorption indicators (e.g., fluorescein, dichlorofluorescein).

  • Indicator adsorbs on AgCl precipitate surface and changes color at the end point.

  • Works best at near-neutral pH.

3.1.4 Procedure and Calculations

1. Prepare and standardize AgNO₃ solution (e.g., against KCl primary standard).

2. Pipette analyte into a flask, adjust pH as required, and add indicator.

3. Titrate with AgNO₃ until the indicator color change persists.

4. Calculate analyte concentration from volume of AgNO₃ used:

   CX−=CAg+×VAgNO3VsampleC_{X^-} = \frac{C_{Ag^+} \times V_{AgNO_3}}{V_{\text{sample}}}CX−​=Vsample​CAg+​×VAgNO3​​​

3.2 Complexometric Titrations

3.2.1 Principle

Complexometric titrations quantify metal ions by forming a stable, water‑soluble complex with a ligand—most commonly ethylenediaminetetraacetic acid (EDTA). The titrant (EDTA) reacts with metal ions in a 1:1 molar ratio under controlled pH to form metal–EDTA complexes.

Mn++H2Y2−  ⟶  MY(n−2)−+2 H+\mathrm{M}^{n+} + \mathrm{H}_2\mathrm{Y}^{2-} \;\longrightarrow\; \mathrm{MY}^{(n-2)-} + 2\,\mathrm{H}^+Mn++H2​Y2−⟶MY(n−2)−+2H+

where H₂Y²⁻ represents the fully deprotonated form of EDTA.

3.2.2 EDTA and Its Forms

• EDTA exists in multiple protonated forms depending on pH.

• At pH 10 (for many divalent metals), the Y⁴⁻ form predominates and chelates metals most strongly.

3.2.3 Choice of Indicators

Complexometric indicators are dyes that change color when they bind to free metal ions but not to the metal–EDTA complex. Common indicators include:

• Eriochrome Black T (for Ca²⁺ and Mg²⁺)

• Calmagite (for hardness titrations)

• Murexide (for Ca²⁺ at pH 12)

At the equivalence point, all metal ions are chelated by EDTA; the free indicator then reverts to its unbound color.

3.2.4 pH Control

pH must be maintained within a narrow range to ensure selective complexation and accurate end‑point detection:

• Hardness (Ca²⁺/Mg²⁺): pH 10 (carbonate buffer)

• Zn²⁺, Cu²⁺: pH 5–6 (acetate buffer)

• Fe³⁺: pH 2–3 (buffered to prevent hydrolysis)

3.2.5 Masking and Demasking Agents

To titrate a specific metal in a mixture, other metals can be “masked” (temporarily prevented from reacting) with selective reagents (e.g., cyanide, tartrate, fluoride). After titration, the masking agent can be neutralized (“demasked”) to allow titration of the next metal.

3.2.6 Titration Procedure

1. Sample Preparation: Take a measured aliquot containing metal ions; adjust pH with appropriate buffer.

2. Indicator Addition: Add a few drops of complexometric indicator; solution develops the indicator–metal color.

3. Titration: Titrate with standard EDTA solution until the solution color changes to the indicator’s free form color.

4. Calculation:

   CMn+=CEDTA×VEDTAVsampleC_{\text{M}^{n+}} = \frac{C_{\text{EDTA}} \times V_{\text{EDTA}}}{V_{\text{sample}}}CMn+​=Vsample​CEDTA​×VEDTA​​

3.2.7 Applications

• Determination of Water Hardness (total, temporary, permanent)

• Assay of Active Pharmaceutical Metals (e.g., bismuth subsalicylate, ferrous sulfate)

• Metal Trace Analysis after appropriate sample digestion

3.3 Summary of Unit 3

• Precipitation titrations form insoluble salts; successful end‑point detection relies on appropriate indicators and control of pH to prevent side reactions.

• Complexometric titrations with EDTA enable precise determination of metal ions; require strict pH control, suitable indicators, and sometimes masking agents.

• Both methods yield accurate quantification when standardized titrants and careful technique are employed.

Unit 4: Oxidation–Reduction (Redox) Titrations

Oxidation–reduction titrations quantify analytes through electron‐transfer reactions. One reactant is oxidized (loses electrons) and the other is reduced (gains electrons). By measuring the volume of a standard oxidizing or reducing titrant required to reach the equivalence point, you determine the concentration of the analyte.

4.1 Fundamental Concepts

• Oxidation is loss of electrons; reduction is gain of electrons.

• Oxidizing agent accepts electrons (it is reduced).

• Reducing agent donates electrons (it is oxidized).

• Half‐reactions split the overall redox into its oxidation and reduction parts; each half includes electrons to balance charge.

  Example with iron(II) and potassium permanganate:

  yaml

  CopyEdit

  Oxidation half: Fe²⁺ → Fe³⁺ + e⁻
Reduction half: MnO₄⁻ + 8 H⁺ + 5 e⁻ → Mn²⁺ + 4 H₂O


• Balancing redox in acidic or alkaline medium ensures atom and charge balance by adding H⁺/OH⁻ and H₂O as needed.

4.2 Standard Electrode Potentials

• Each half‐reaction has a standard reduction potential (E°), measured in volts relative to the standard hydrogen electrode (SHE, E° = 0 V).

• Cell potential (Ecell) is the difference between the reduction potential of the cathode and that of the anode:

  Ecell=Ecathode∘−Eanode∘E_{\text{cell}} = E^\circ_{\text{cathode}} – E^\circ_{\text{anode}}Ecell​=Ecathode∘​−Eanode∘​

• A positive Ecell indicates a spontaneous reaction.

• Tables of standard potentials help select appropriate titrant and predict titration feasibility.

4.3 Common Redox Titrations

1. Permanganometry (KMnO₄ titration)

   • Reagent: Standardized potassium permanganate (strong oxidizing agent).

   • Analytes: Ferrous salts, oxalate, hydrogen peroxide, nitrite.

   • Medium: Acidic (usually H₂SO₄); MnO₄⁻ is purple, reduced Mn²⁺ is nearly colorless.

   • End point: Appearance of a faint pink color that persists for 30 seconds indicates slight excess MnO₄⁻.

2. Iodometry and Iodimetry

   • Iodometry (back‐titration): Titrate I₂ liberated by oxidizing analyte with standard thiosulfate.

   • Iodimetry (direct titration): Titrate I₃⁻ (from adding excess I⁻ to analyte oxidizer) directly with thiosulfate.

   • Reagent: Sodium thiosulfate; I₂ + 2 S₂O₃²⁻ → 2 I⁻ + S₄O₆²⁻.

   • Indicator: Starch forms deep blue complex with I₂; disappears at end point.

3. Cerimetry (Ce⁴⁺ titration)

   • Reagent: Cerium(IV) sulfate or ceric ammonium sulfate (strong oxidizer).

   • Analytes: Sulfites, ferrous ions, organic compounds.

   • Medium: Acidic; Ce⁴⁺ (yellow) reduces to Ce³⁺ (colorless).

   • End point: Disappearance of yellow color.

4. Dichrometry (K₂Cr₂O₇ titration)

   • Reagent: Standard potassium dichromate in acidic medium.

   • Analytes: Ferrous ions, oxalate, arsenic(III).

   • Indicator: Diphenylamine (turns blue at end point) or Mohr’s salt back‐titration with permanganate.

4.4 Procedure and Calculations

1. Preparation and Standardization

   • Accurately prepare titrant solution; standardize against a primary standard (e.g., Mohr’s salt for KMnO₄).

2. Sample Preparation

   • If solid, dissolve in appropriate acid or alkaline medium; dilute to known volume.

3. Titration

   • Pipette aliquot of sample into conical flask; add acid or buffer to set correct pH; add indicator if required.

   • Add titrant from burette until permanent color change (direct titrations) or perform back‐titration and titrate excess (iodometry).

4. Volume Measurement

   • Record initial and final burette readings; calculate volume of titrant used.

5. Stoichiometry

   • Use balanced redox equation to relate moles of titrant to moles of analyte.

   • Calculate analyte concentration:

   Canalyte=Ctitrant×Vtitrant×ntitrantVsample×nanalyteC_{\text{analyte}} = \frac{C_{\text{titrant}} \times V_{\text{titrant}} \times n_{\text{titrant}}}{V_{\text{sample}} \times n_{\text{analyte}}}Canalyte​=Vsample​×nanalyte​Ctitrant​×Vtitrant​×ntitrant​​

   where $n$ are stoichiometric coefficients.

Example
Determine the concentration of Fe²⁺ in 25.00 mL of solution if 20.00 mL of 0.0200 M KMnO₄ (E° = +1.51 V) titrates to the end point in acidic medium. Reaction:

5Fe2++MnO4−+8H+→5Fe3++Mn2++4H2O5 Fe^{2+} + MnO_4^- + 8 H^+ → 5 Fe^{3+} + Mn^{2+} + 4 H_2O5Fe2++MnO4−​+8H+→5Fe3++Mn2++4H2​O

• Moles KMnO₄ = 0.02000 L × 0.0200 mol/L = 4.00 × 10⁻⁴ mol

• From equation, 1 mol MnO₄⁻ oxidizes 5 mol Fe²⁺; so moles Fe²⁺ = 5 × 4.00 × 10⁻⁴ = 2.00 × 10⁻³ mol

• Concentration Fe²⁺ = 2.00 × 10⁻³ mol / 0.02500 L = 0.0800 M

4.5 Sources of Error and Precautions

• Acid Strength: Insufficient acidity can slow reaction or form precipitates (e.g., MnO₂ in permanganometry).

• Indicator Choice: Wrong indicator leads to false end point; choose based on redox potential.

• Temperature Control: Redox potentials are temperature‐dependent; titrate at constant temperature.

• Stabilization of Titrant: KMnO₄ solutions decompose on standing; standardize immediately before use.

Unit 5: Gravimetric Analysis

Gravimetric analysis is a quantitative method based on the measurement of mass. The analyte is converted into a pure, stable, insoluble compound (precipitate) that can be isolated, dried, and weighed. From the mass of the precipitate and known stoichiometry, the amount of analyte is determined.

5.1 Principle of Gravimetric Analysis

1. Precipitation: Convert the analyte into an insoluble form by adding a precipitating reagent.

2. Filtration: Separate the precipitate from the solution using filtration.

3. Washing: Remove impurities and adhering ions by washing with suitable solvents (often deionized water, sometimes with reagents to prevent loss of the precipitate).

4. Drying or Ignition:

   • Drying: Remove moisture at a specified temperature to a constant weight when the precipitate is stable and does not decompose.

   • Ignition: Heat to convert the precipitate into a known composition (e.g., convert barium sulfate to barium sulfate ash) at high temperature.

5. Weighing: Weigh the dried or ignited precipitate on an analytical balance.

6. Calculation: Use the mass of the precipitate and its chemical formula to calculate the amount of the analyte.

5.2 Types of Gravimetric Methods

1. Direct Gravimetry

   • Precipitate of the analyte itself is weighed.

   • Example: Precipitation of barium sulfate (BaSO₄) to determine sulfate.

2. Indirect (Back) Gravimetry

   • An excess of a standard reagent reacts with the analyte; the remaining reagent is then precipitated and weighed.

   • Example: Excess silver nitrate added to chloride solution; unreacted Ag⁺ is precipitated as AgCl and weighed.

3. Volatilization Gravimetry

   • Analyte is volatilized (as a gas or vapor) and the mass loss is measured.

   • Example: Loss on drying to determine water content; ignition to determine organic content by weight loss.

4. Electrogravimetry

   • Metal ions are deposited electrochemically onto an electrode, which is weighed before and after deposition.

   • Example: Silver plating from a silver nitrate solution.

5.3 Common Gravimetric Precipitates

5.4 Procedure and Best Practices

1. Selection of Precipitant and Conditions

   • Choose a reagent that forms a precipitate with high purity and low solubility product (Kₛₚ).

   • Control pH, temperature, and concentration to ensure complete precipitation.

2. Seeding and Digestion

   • Seeding: Add a small crystal of the pure precipitate to induce uniform crystal growth.

   • Digestion: Maintain the suspension at elevated temperature for a period to promote formation of larger, purer crystals that filter more easily.

3. Filtration Techniques

   • Büchner Funnel with vacuum for rapid filtration.

   • Gooch Crucible or sintered glass crucible for high-temperature ignition and weighing.

4. Washing the Precipitate

   • Use cold wash solutions to minimize solubility losses.

   • Final washes may include reagents (e.g., dilute solution of the precipitant’s salt) to prevent dissolution.

5. Drying vs. Ignition

   • Drying: In an oven at specified temperature (e.g., 105 °C) until constant weight.

   • Ignition: In a muffle furnace at high temperature (e.g., 600–800 °C) to convert precipitate to a stable oxide or ash.

6. Weighing

   • Cool in a desiccator to room temperature before weighing.

   • Record mass to four decimal places for accuracy.

5.5 Calculations

For a precipitate of known formula AxByA_xB_yAx​By​:

1. Determine the moles of precipitate:

   nprecipitate=mprecipitateMAxByn_{\text{precipitate}} = \frac{m_{\text{precipitate}}}{M_{\text{A}_x\text{B}_y}}nprecipitate​=MAx​By​​mprecipitate​​

2. Use stoichiometry to relate moles of precipitate to moles of analyte.

Example
Determine sulfate content in water sample by precipitating as BaSO₄.

• Mass of BaSO₄ precipitate = 0.5234 g

• Molar mass BaSO₄ = 233.39 g/mol

• Moles BaSO₄ = 0.5234 g / 233.39 g/mol = 2.243 × 10⁻³ mol

• Each mole of BaSO₄ corresponds to 1 mol SO₄²⁻; mass of SO₄²⁻ = 2.243 × 10⁻³ mol × 96.06 g/mol = 0.2155 g

5.6 Sources of Error

• Co-precipitation: Unwanted ions co-precipitate with the analyte; minimized by controlling conditions and digestion.

• Occlusion: Impurities trapped within precipitate crystals; reduce by digestion and proper washing.

• Losses during Filtration/Washing: Use appropriate wash solutions and techniques to minimize dissolution.

• Thermal Decomposition: Some precipitates decompose or lose water on drying/ignition; choose proper conditions.

Unit 6: Introduction to Instrumental Methods

Instrumental methods use sophisticated equipment to measure physical or chemical properties of analytes. They offer higher sensitivity, specificity, and speed compared to classical techniques. This unit provides an overview of key instrumental techniques used in pharmaceutical analysis.

6.1 Classification of Instrumental Methods

1. Spectroscopic Methods

   • Measure interaction between electromagnetic radiation and matter.

   • Examples: UV–Visible (UV–Vis), Infrared (IR), Atomic Absorption Spectroscopy (AAS).

2. Chromatographic Methods

   • Separate analyte components before detection.

   • Examples: High‑Performance Liquid Chromatography (HPLC), Gas Chromatography (GC), Thin‑Layer Chromatography (TLC).

3. Electrochemical Methods

   • Measure electrical properties (current, potential) resulting from redox reactions.

   • Examples: Potentiometry (pH meters, ion‑selective electrodes), Voltammetry.

4. Mass Spectrometry (MS)

   • Ionizes analytes and measures mass‑to‑charge ratios.

   • Often coupled to GC or HPLC for compound identification and quantification.

5. Hyphenated Techniques

   • Combine separation with detection (e.g., GC–MS, HPLC–MS, LC–MS/MS).

6.2 UV–Visible Spectrophotometry

Principle

• Molecules absorb UV or visible light, promoting electrons from ground to excited states.

• Absorbance (A) measured at specific wavelength (λ) follows the Beer–Lambert law:

  $$A=\epsilon bc$$

  where ε is molar absorptivity, b is path length, and c is concentration.

Instrument Components

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

2. Monochromator: Prism or grating to select wavelength.

3. Sample Holder: Quartz or glass cuvette.

4. Detector: Photodiode or photomultiplier tube.

5. Read‑out: Digital display of absorbance or transmittance.

Applications

• Assay of drug substances (e.g., paracetamol at 243 nm).

• Kinetic studies of reaction rates.

• Purity assessment of raw materials.

6.3 Infrared (IR) Spectroscopy

Principle

• Molecules absorb infrared radiation, causing vibrational transitions.

• Characteristic absorption bands correspond to functional groups.

Instrument Components

1. IR Source: Nernst glower or Globar.

2. Interferometer or Monochromator: Fourier transform IR uses an interferometer.

3. Sample Accessory: Pellet press (KBr), ATR (attenuated total reflectance).

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

Applications

• Identification of functional groups in drug molecules.

• Verification of excipients and impurities.

• Polymorph screening in solid dosage forms.

6.4 High‑Performance Liquid Chromatography (HPLC)

Principle

• Separation based on partition between mobile phase and stationary phase under high pressure.

Key Components

1. Pump: Delivers mobile phase at high pressure (up to 400 bar).

2. Injector: Introduces sample (manual or autosampler).

3. Column: Packed with stationary phase particles (e.g., C18 silica).

4. Detector: UV–Vis, photodiode array (PDA), refractive index (RI), or MS.

5. Data System: Processes chromatographic peak data.

Parameters

• Retention Time (tR): Time analyte spends in column.

• Resolution (Rs): Degree of separation between peaks.

• Plate Number (N): Measure of column efficiency:

  N=16(tRW)2N = 16 \left(\frac{t_R}{W}\right)^2N=16(WtR​​)2

  where W is peak width.

Applications

• Quantitative assay of pharmaceuticals.

• Impurity profiling.

• Dissolution testing of dosage forms.

6.5 Gas Chromatography (GC)

Principle

• Separation of volatile analytes between a gas mobile phase and a liquid or solid stationary phase.

Components

1. Carrier Gas: Inert gas (e.g., helium, nitrogen).

2. Injector: Vaporizes sample (split or splitless mode).

3. Column: Capillary or packed column within an oven to control temperature.

4. Detector: Flame ionization detector (FID), thermal conductivity detector (TCD), or MS.

5. Data System: Records chromatogram.

Applications

• Analysis of volatile organic impurities in drugs.

• Residual solvents testing.

• Flavor and fragrance profiling.

6.6 Atomic Absorption Spectroscopy (AAS)

Principle

• Atoms in the ground state absorb light of element‑specific wavelength. Absorbance is proportional to concentration.

Components

1. Light Source: Hollow cathode lamp for each element.

2. Atomizer: Flame (air–acetylene) or graphite furnace.

3. Monochromator: Selects analytical wavelength.

4. Detector: Photomultiplier tube.

5. Read‑out: Digital display of absorbance.

Applications

• Determination of trace metals (e.g., lead, cadmium) in pharmaceuticals and raw materials.

• Quality control of water for injection (heavy metal limits).

6.7 Method Selection and Validation

• Selectivity: Choose technique based on analyte properties (e.g., volatility, UV chromophore).

• Sensitivity: Ensure method can detect analyte at required limits.

• Robustness: Assess impact of small variations (flow rate, temperature).

• Validation Parameters: Linearity, accuracy, precision, LOD/LOQ, specificity, and stability.