UV-VIS Spectrophotometry Review

Fundamentals of UV-Vis Spectrophotometry

• Definition and Range: UV-Vis spectrophotometry involves pharmaceutical drugs being exposed to electromagnetic radiation in the Ultraviolet-Visible (UV-Vis) range, which spans wavelengths ($\lambda$) between $200$ and $800\,nm$.

• Energy-Wavelength Relationship: Different wavelengths within the UV-Vis range carry distinct energy levels. The relationship is inverse: a shorter wavelength corresponds to higher energy.

• Molecular Mechanism of Absorption:

  • Absorbed energy from radiation stimulates a molecule’s valence electrons, causing them to move from a lower energy orbital (ground state) to a higher energy orbital (excited state). This process is known as "excitation."

  • The specific energy levels associated with UV-Vis radiation trigger the temporary excitation of double bonds ($\pi$-electrons) and lone electron pairs ($n$-electrons).

• Quantitative Use: The fundamental principle for quantitative analysis is that a higher concentration of a drug leads to greater absorption of light ("more drug = more absorption").

Key Terminologies and Structural Components

• Conjugated System: This refers to interlinked unsaturated bonds, typically characterized by alternating saturated and unsaturated bonds. These systems allow for the free movement of "delocalized" $\pi$ and $n$ electrons.

• Auxochromes: These are functional groups possessing lone electron pairs (EP) that are attached to a conjugated system. They contribute to the delocalization of electrons and enhance the absorbance of light.

• Chromophore: This is the specific structural part of a molecule responsible for absorbing UV-Vis radiation. It is defined as the combination of the conjugated system and any attached auxochromes.

• Chromophore Extension and Energy:

  • The more extensive and complex a chromophore is, the less energy is required for electronic excitation.

  • As a result, the absorbance shifts toward higher wavelengths ($\lambda$), which represent lower energy levels.

• Molecular Color: If the absorbance shift moves from the UV range into the visible (VIS) range, the molecule appears colored to the human eye. The color perceived is the complementary color of the wavelength ($\lambda$) that is absorbed.

Practical Examples of Conjugation and Absorption

• Benzene:

  • Structure: Contains $3 \times C=C$ bonds.

  • Properties: Features $6$ delocalized electrons.

  • Absorption: Absorbs at $\lambda \approx 170\text{--}250\,nm$.

• Crotamiton:

  • Structure: Contains $2$ conjugated systems plus $3$ lone electron pairs.

  • Absorption: Absorbs at $\lambda \approx 242\,nm$.

• Vitamin A:

  • Structure: Contains $5 \times C=C$ bonds; it is highly conjugated.

  • Absorption: Absorbs at $\lambda \approx 326\,nm$.

• $\beta$-Carotene:

  • Structure: Contains $11 \times C=C$ bonds; it is very highly conjugated.

  • Absorption: Absorbs at $\lambda \approx 497\,nm$.

  • Color: It absorbs mainly blue-green light, resulting in a red-orange appearance.

The Beer-Lambert Law for Quantitative Analysis

• Measurement Principle: UV-Vis spectrophotometry measures the absorption of radiation by comparing the amount of radiation entering the solution to the amount of radiation leaving the solution.

• Correlations for Dilute Solutions:

  • Absorbance ($A$) is proportional to the concentration ($c$) of the solution (Beer's Law): $A \propto c$.

  • Absorbance ($A$) is proportional to the light's pathlength ($l$) (Lambert's Law): $A \propto l$.

• Mathematical Formula:

  • Using specific absorbance: $A = A^{1\%}_{1\,cm} \times c \times l$

  • Using molar absorbance: $A = \epsilon \times c \times l$

• Variable Definitions:

  • $A$: Absorbance (unitless) at a specific wavelength ($\lambda$).

  • $A^{1\%}_{1\,cm}$: Specific absorbance coefficient.

  • $\epsilon$: Molar absorbance coefficient.

  • $c$: Concentration. Expressed in $g/100\,ml$ when using $A^{1\%}_{1\,cm}$, or in $moles/l$ when using $\epsilon$.

  • $l$: Path length within the sample, typically measured in $cm$.

• Standardization: Spectrophotometers use standard cuvettes with a pathlength of $l = 1\,cm$, which simplifies the Beer-Lambert Law and allows for a direct linear relationship between absorbance and concentration.

External Standard Method and Calibration

• Typical Procedure:

  1. Prepare a stock solution of the target drug.

  2. Prepare $3\text{--}5$ dilute samples with known concentrations from the stock solution.

  3. Measure the absorbance of each dilute solution at a specific wavelength ($\lambda$).

  4. Plot a calibration curve (Standard Curve) of concentration ($c$) versus absorbance ($A$).

  5. Measure the absorbance of the drug sample solution with an unknown concentration (diluting if necessary).

  6. Determine the concentration of the unknown sample from the standard curve.

• Linearity and Limits:

  • The plot should be a line passing through the origin (where $c = 0$, $A = 0$).

  • The slope of the line is $A^{1\%}_{1\,cm}$ if concentration is in "$\%\,w/v$ ($g/100\,ml$)" and $\epsilon$ if concentration is in $moles/l$.

  • Linearity is lost at high concentrations (e.g., c > 0.01\,M).

  • The Beer-Lambert Law is only valid for dilute solutions where absorbance ($A$) is typically less than $1.5$.

  • Optimal Range: For reliable quantitative analysis, absorbance readings should ideally fall between $0.3$ and $0.8$ to ensure the linearity of the plot (the "Beer-Lambert Range").

Interpretation of UV-Vis Spectra

• Recording Spectra: Spectra can be recorded over the entire UV-Vis range ($200\text{--}800\,nm$).

• $\lambda_{max}$: The wavelength that produces the highest absorbance value is known as $\lambda_{max}$. This wavelength is commonly used for quantitative analysis.

• Structural Reflection: The UV-Vis spectrum is reflective of the specific chromophore present in the molecule. Changes to chemical structure that affect the chromophore (e.g., ionization) will alter the appearance of the spectrum ($A^{1\%}_{1\,cm}$ and $\lambda_{max}$).

Structural and Environmental Effects on Spectra

• pH Dependency: The presence of ionizable auxochromes such as phenols or anilines makes the UV-Vis spectrum dependent on the pH of the solution. This can be exploited for qualitative analysis to identify particular functional groups.

• Phenol Example (Auxochrome Effects):

  • Acidic/Neutral Media: The molecule exists as an unionized phenol with $\lambda_{max} = 270\,nm$ and $A^{1\%}_{1\,cm} = 172$.

  • Basic Media: The phenol turns into a phenoxide ion. This results in an extended chromophore.

  • Observed Shift:

    • Bathochromic (red) shift: $\lambda_{max}$ increases to $287\,nm$.

    • Hyperchromic shift: $A^{1\%}_{1\,cm}$ increases to $271$.

• Aniline Example (Auxochrome Effects):

  • Basic/Neutral Media: The molecule exists as unionized aniline with $\lambda_{max} = 286\,nm$ and $A^{1\%}_{1\,cm} = 179$.

  • Acidic Media: The aniline is protonated to form an anilinium ion. This results in the loss of the auxochrome (lone pair is occupied), creating a smaller chromophore.

  • Observed Shift:

    • Hypsochromic (blue) shift: $\lambda_{max}$ decreases to $255\,nm$.

    • Hypochromic shift: $A^{1\%}_{1\,cm}$ decreases to $16$.

Limitations of UV-Vis Spectrophotometry

• Linearity Constraints: The Beer-Lambert Law is only reliable for dilute samples.

  • If concentration is too low, detecting the small amount of radiation leaving the sample is inaccurate.

  • If concentration is too high, particles may experience aggregation or dimerization. Aggregates often possess different spectral properties than individual molecules.

• Pharmacopoeial Requirements: The British Pharmacopoeia (B.P.) requires that absorbance values used for assays range between $0.3$ and $0.8$. This often necessitates the development of specific dilution strategies.

• Chemical Homogeneity: To ensure accurate results, all molecules in a sample must have the same chemical structure (same chromophore). Analyses are often conducted at extreme pH values (fully ionized or fully unionized) to ensure consistency in the number of lone electron pairs and absorbance values.

• Specificity (The Biggest Limitation): UV-Vis spectrophotometry lacks high specificity.

  • Any molecule with a similar chromophore, such as impurities or degradation products, will contribute to the total absorbance.

  • This leads to the potential over-estimation of drug content.

  • Example: In an assay of Aspirin, the degradation product Salicylic Acid will also be partially measured, interfering with the results. It is necessary to remove such interferants before analysis whenever possible.

Fundamentals of Titrimetric Analysis

• Definition and Overview: Titrimetric analysis is a wet chemical analytical method utilized for the quantitative analysis of drugs. It involves the reaction of a drug analyte with a titrant.

• Key Terminologies:

  • Analyte: The drug sample being analyzed.

  • Titrant: A standard solution with a precisely known concentration used in the reaction.

  • Equivalence Point: The theoretical endpoint where the reaction is complete (total). This is determined via visual assessment (indicators) or electrochemical methods.

  • Equivalency Value: A value commonly used in the British Pharmacopoeia (BP) instead of manual stoichiometric calculations. It indicates the Correlation between titrant volume and the amount of analyte present.

• Analytical Process: The amount of titrant required to reach the equivalence point provides a direct estimate of the drug amount in the sample.

Computational Examples and Calculations

• Stoichiometric Purity Calculation (Ibuprofen Example):

  • Scenario: $0.2681\,g$ of Ibuprofen is titrated with $0.1M\,NaOH$. It requires $12.4\,mL$ of $NaOH$ to reach the endpoint.

  • Step 1: Calculate moles of titrant used:

    • $c = \frac{n}{V} \rightarrow n = c \times V$

    • $n = 0.1\,mol/dm^3 \times 0.0124\,dm^3 = 0.00124\,moles$

  • Step 2: Determine Stoichiometry: Assuming a $1:1$ reaction, there are $0.00124\,moles$ of Ibuprofen present.

  • Step 3: Calculate mass of analyte:

    • Molecular weight ($MW$) of Ibuprofen = $206\,g/mol$

    • $n = \frac{m}{MW} \rightarrow m = n \times MW$

    • $m = 0.00124\,mol \times 206\,g/mol = 0.2554\,g$

  • Step 4: Calculate Purity:

    • $\%\,purity = \frac{\text{Calculated mass}}{\text{Sample mass}} \times 100$

    • $\%\,purity = \frac{0.2554\,g}{0.2681\,g} = 95.2\%$

• Equivalency Value Calculation (Paracetamol Example):

  • Equivalency Definition: $1\,mL$ of $0.1M\,NaOH \equiv 20.63\,mg$ of Paracetamol.

  • Calculation: If the titre volume is $15.4\,mL$, then the amount of Paracetamol = $15.4 \times 20.63 = 317.70\,mg$.

Advantages and Disadvantages of Titrimetric Analysis

• Advantages:

  • Versatile, simple, and relatively inexpensive.

  • High accuracy.

  • Absolute method: Does not require instrumental calibration against a standard.

  • Longest-standing traditional technique still widely used in pharmaceutical analysis.

  • Can be automated.

• Disadvantages:

  • Lack of specificity.

  • Can be time-consuming.

  • Requires a relatively large sample mass (typically > 100\,mg).

  • Requires large solvent volumes, posing potential safety and disposal issues.

• Main Applications:

  • Assessment of drug purity.

  • Assessment of drug content within a formulation.

  • Analysis of excipients.

Visualization of Titration Endpoints

• Chemical Indicators:

  • Rely on a specific color change range of weak acids or weak bases.

  • The choice of indicator is strictly dependent on the pH at the titration's equivalence point.

  • Provides rapid endpoint determination.

  • Common Examples: Phenolphthalein, Bromothymol Blue, and Crystal Violet.

• Electrochemical Endpoint Determination:

  • Methodology: Uses a potentiometer, such as a glass pH electrode.

  • Mechanism: The electrode monitors variations in potential difference caused by the interaction of $H^+$ ions with the outer surface of a pH-sensitive glass membrane.

  • Analysis: The endpoint is determined by plotting the potential difference ($mV$) of the titre solution against the titrant volume ($mL$).

  • Characteristics: The endpoint is the mid-point between two volumes where the largest difference in $mV$ is measured.

  • Advantages: More accurate than visual assessment and can be used for colored sample solutions that would obscure chemical indicators.

  • Note: Usually slower than chemical indicators, but can be automated.

Types of Drug Titrations

• Selection Criteria: The choice of titration depends on the drug's chemical structure and physiochemical properties, such as acid-base nature and solubility.

• Common Types (ordered by popularity):

  • Acid-Base Titrations: Most popular, used for weak acids/bases.

  • Non-Aqueous Titrations: Also extremely popular, especially in the British Pharmacopoeia.

  • Redox-Titrations: Used for oxidation-reduction reactions.

  • Argentimetric Titrations: Silver-based titrations (applied to select groups).

  • Compleximetric Titrations: Involves complex formation (applied to select groups).

  • Diazotisation Titrations: Specific to certain nitrogenous groups (applied to select groups).

Acid-Base Titrations and Back-Titrations

• Acid-Base Titration:

  • Employed extensively by the B.P. because many drugs are weak acids or bases.

  • Weak acids are titrated with a strong base; weak bases are titrated with a strong acid.

  • Solubility Requirements: The drug must be completely soluble in the media.

  • Solvent Adjustments: For low water-solubility drugs, alcohols like Methanol ($MeOH$) or Ethanol ($EtOH$) may be added. Alternatively, pure alcohol or water-miscible solvents like acetone or dimethylformamide can be used.

  • Examples: Tiaprofenic acid (Weak acid titrated with $NaOH$), Acenocoumarol, Diflunisal, Fusidic acid, Ibuprofen, Mefenamic acid, Phenylbutazone, and Tolbutamide.

• Back-Titration:

  • An excess amount of strong acid/base is added to the drug.

  • The unreacted (excess) amount is then titrated with a strong base/acid.

  • Logic: If the drug content is high, the unreacted amount of primary reagent is low, resulting in a low titre volume. If the drug content is low, the unreacted amount is high, resulting in a high titre volume.

  • Reaction Examples:

    • Aspirin: Reacts in a $1:2$ ratio with $NaOH$. Excess $NaOH$ is back-titrated with $HCl$.

    • Clofibrate capsules: Also utilizes back-titration.

Non-Aqueous Titrations

• Prevalence: This is the most common titrimetric method in the B.P.

• Conditions: The reaction occurs under non-aqueous conditions where both the drug and titrant are in non-aqueous solvents. Glassware must be completely dry.

• Rationales for Use:

  • Low water solubility of the drug.

  • Low aqueous stability (drug decomposes in water, e.g., esters prone to hydrolysis).

  • Very weak acid (pKa > 7) or very weak base (pKa < 7). In water, $H_2O$ competes for the titrant (acting as a base or acid), leading to incorrect results or difficult endpoint determination.

• Chemical Logic (Competition):

  • Weak Acid: $H_2O + \text{BASE} \rightarrow OH^- + \text{BASE-H}^+$ (Water competes with the analyte for the base titrant).

  • Weak Base: $H_2O + \text{ACID} \rightarrow H_3O^+ + \text{ACID-BASE}$ (Water competes with the analyte for the acid titrant).

• Specific Drug Examples:

  • Aciclovir: Assay involving a weakly basic guanine unit.

  • Cocaine: Assay involving a strongly basic $3^{\circ}$ aliphatic amine; esters in cocaine are prone to aqueous hydrolysis.

  • Gliquidone: Reasonably acidic sulfonylurea unit with poor water solubility.

Case Study: Quality Control of Vitamin C Tablets

• Methodology: A Redox-Titration where Vitamin C is oxidized by $I_2$. The $I_2$ is generated in situ from iodate (the titrant) and iodide.

• Protocol: Due to complex stoichiometry, a correct equivalency value must be determined before tablet analysis. Titrations must be carried out at least in duplicate.

• Equivalency Determination (Data Example):

  • Sample 1: Mass $0.0998\,g$; Volume $18.9\,mL$; Equivalency $5.28\,mg/mL$.

  • Sample 2: Mass $0.1063\,g$; Volume $20.2\,mL$; Equivalency $5.26\,mg/mL$.

  • Average Equivalency Value: $5.27\,mg/mL$.

• Tablet Analysis (Data Example):

  • Constants: Average tablet weight = $956.72\,mg$.

  • Tablet 1:

    • Powder used: $194.9\,mg$

    • Titre: $19.7\,mL$

    • Vitamin C in sample: $103.82\,mg$

    • Vitamin C in tablet: $509.63\,mg$

    • Percentage stated content: $101.9\%$

  • Tablet 2:

    • Powder used: $205.2\,mg$

    • Titre: $20.3\,mL$

    • Vitamin C in sample: $106.98\,mg$

    • Vitamin C in tablet: $498.78\,mg$

    • Percentage stated content: $99.8\%$

  • Average Stated Content: