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lec6_2011

# lec6_2011 - Applications of standard potentials The...

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Applications of standard potentials

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The electrochemical series For two redox couples Ox 1 /Red 1 and Ox 2 /Red 2 , Red 1 , Ox 1 || Red 2 , Ox 2 E θ = E θ 2 – E θ 1 The cell reaction: Red 1 + Ox 2 Ox 1 + Red 2 If E θ > 0 then Δ r G θ < 0 (Nernst equation), the reaction will take place spontaneously. In other words, if E θ 2 > E θ 1 , the Ox2 has the thermodynamic tendency to oxidize Red1.

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The determination of activity coefficients Once the standard potential of an electrode (E θ ) is known, one can use the following equation E + (2RT/vF)ln(b) = E θ - (2RT/vF)ln(γ ± ) to determine the mean activity coefficient of the ions at the concentration of interest via measuring the cell emf (E).
The determination of equilibrium constants Self-test 7.11 Calculate the solubility constant (the equilibrium constant for reaction Hg 2 Cl 2 (s) ↔ Hg 2 2+ (aq) + 2Cl - (aq)) and the solubility of mercury(I) chloride at 298.15K. The mercury(I) ion is the diatomic species Hg 2 2+ . Answer: This chemical process does not involve electron transfer, i.e. is not a redox reaction. Choosing cathode reaction as: Hg 2 Cl 2 (s) + 2e → 2Hg( l ) + 2Cl - (aq) from reference table 7.2, E θ = 0.27 V the anode reaction can be obtained through R – Cell Hg 2 2+ (aq) + 2e → 2Hg( l ) from reference table 7.2, E θ = 0.79 V Therefore the standard cell potential = 0.27 – 0.79 = -0.52 V

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Using equation ln K = E ө * v /25.7mV, here v = 2 ln K = -40.467 K = 2.62 x 10 -18 K = ( a Hg(I) * a 2 cl- )/ a Hg2cl2 = b*( 2b ) 2 /1 = 4*b 3 = 2.62 x 10 -18 therefore b = 8.68 x 10 -7 mol/kg
Species-selective electrodes: Measuring pH

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