CHM 4130_ch.13_15_26-28_Spring11

CHM 4130_ch.13_15_26-28_Spring11 - An Introduction to...

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1 An Introduction to Ultraviolet-Visible Molecular Spectrometry Beer’s Law : A = -log T = -logP0 / P = ε x b x C See Table 13-1 for terms. In measuring absorbance or transmittance, one should compensate for reflections and scatter occurring at the interface of the cuvette. Compensation is always done by running the blank.
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2 Application of Beer’s Law to Mixtures Before applying these equations, you need to know the following: For compound 1: ε1 at λ 1 and λ 2 For compound 2: ε2 at λ 1 and λ 2 If their values are not know, you should obtain them via calibration curves. It is also common to obtain ε values from single standard solutions. You should always remember: ε is a constant that depends on: Wavelength Temperature Solvent Concentration, M Absorbance λ 1 λ2 ε1 ε2
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3 Limitations of Beer’s Law Three types of limitations cause deviations from linearity: a) Real Limitations b) Instrumental Limitations c) Chemical Limitations Real Limitations: Result from analyte-analyte interactions at high analyte concentrations (usually C > 0.01M). The upper limit of the LDR depends on the compound. For the same compound, it also depends on the solvent and the temperature. You should always determine experimentally the upper limit of the LDR. LDR Absorbance Concentration Solute-solute interactions are mainly dipole-dipole interactions Electrostatic interactions can also occur among strong electrolytes High analyte concentration causes average distances among analyte molecules to decrease
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4 Chemical deviations: Arise when an analyte dissociates, associates, or reacts with a concomitant (including solvent) to produce a product with different absorption spectrum than the analyte. Typical example: Acid-Base Indicators (HIn). HIn <=> H+ + In- (Color 1, λ 1) (Color 2, λ 2) Spectrum looks like: => There is a dissociation constant associated to the equilibrium above: Ka = [H+][In-] / [HIn] HIn and In- concentrations depend on the pH of solution: [H+] = Ka x [HIn] / [In-] => log [H+] = logKa + log [HIn] / [In-] => - log [H+] = - logKa - log [HIn] / [In-] => pH = pKa + log [In-] / [HIn] Assuming an indicator with pKa = 5 (Ka = 10-5): pH log[In-] / [HIn] [In-] / [HIn] [In-] = ? X [HIn] 1 - 4 10-4 0.0001 2 - 3 10-3 0.001 5 0 100 1 6 1 101 10 7 2 102 100 At any given wavelength, the total absorbance intensity of the solution is the sum of the individual intensities of HIn and In-:
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5 According to Beer’s law, the absorbance intensities of HIn and In- are the following: A HIn, λ = ε HIn, λ . b . [HIn] A In-, λ = ε In-, λ . b . [In-] Substituting in the total absorbance equation: A Total, λ = ε HIn, λ . b . [HIn] + ε In-, λ . b . [In-] => At any given wavelength, the total absorbance intensity of the solution (A Total) depends on the concentrations of HIn and In-. => Because [HIn] and [In-] depend on pH, A Total depends on pH. You should also remember that the indicator concentration is given by: C Indicator = C = [HIn] + [In-] HIn <=> H+ + In- C – x x x where x = [In-] = [H+]. From Ka: Ka = x2 / C – x => x2 = Ka (C – x) => x2 + Ka.x – Ka.C = 0 => x = -Ka ± {Ka2 + 4KaC}1/2 / 2 This equation demonstrates that the concentration of indicator varies non-linearly with the ionized fraction of the acid. The same is
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This note was uploaded on 11/07/2011 for the course CHM 4130 taught by Professor Staff during the Spring '11 term at University of Central Florida.

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CHM 4130_ch.13_15_26-28_Spring11 - An Introduction to...

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