S. REDOx TITRATION oF FERRICYANIDE ION TH ASCORBIC ACID! Introduction ferricyani
ID: 528088 • Letter: S
Question
S. REDOx TITRATION oF FERRICYANIDE ION TH ASCORBIC ACID! Introduction ferricyanide Fe(CN)6', by In this lab, We explore the redox titration of the ion, salt, which itself ascorbic acid (vitamin C). introduced as its potassium has found use in applications such as photography, preparation of blueprints, and as a metal ion colorimetric indicator. red K3Fe(CN)6 dissolved in water, the ion is a bright yellow complex. It is commonly used in electrochemical studies on account of its low toxicity, robustness, and easily reversible redox behavior. The reduced form, Fe(CNDe" or ferrocyanide, is colorless. Today, we will reduce ferricyanide with ascorbic acid (vitamin C). Ascorbic acid is used for several reasons, including the fact that it is a mild reducing agent and a weak acid. Exposure of the ferricyanide ion to strong acids can result in the formation of toxic cyanide gas, a highly undesirable situation. The two half reactions are shown below (written, of course, as reductions Fe(CN)6 er Fe(CN)s ferrocyanide C6H606, 2H 2e ascorbic acid dehydroascorbic acid We will be monitoring the reduction of ferricyanide ion both spectrophotometrically and potentiometrically. Ferricyanide on absorbs strongly at 420 nm (E 1.02x10 M cm i), and bic acid are transparent at this wavelength. ferrocyanide ion, ascorbic acid, and dehydroascor Therefore, the solution gradually becomes colorless with the addition of ascorbic acid. We can use Beer's Law (A g bc, where b 1.00 cm for our cuvettes to determine the concentration of ferricyanide at any point in the redox titration. We can use that calculated concentration of ferricyanide to determine the amount of ferrocyanide at the same point, since [ferricyanide]i [ferricyanide] volume doesn't change significantly during the course of the redox titration. Potentiometry is the field of electroanalytical chemistry in which potential is measured while no current is flowing through the reaction container (the reaction container being an electrochemical cell). The measured potential may then be used to quantify an ion of interest (in our case, the ferricyanide ion). The potential that develops in the electrochemical cell is the result of the free energy change that would occur if the chemical reaction were to proceed until equilibrium. To measure this potential, you need both a reference electrode and an indicator electrode. The reference electrode maintains a constant potential which the indicator electrod potential is measured against, while the indicator electrode develops an electric potential in the above redox reaction occurring at its met n this experiment, you will electrode (in saturated Experimentally, we will prepare reaction to perform both the and trophotometric measurements. initial will allow the calculation of your initial concentration of ferricyanide Then, y will perform 15 sets measurements. These data will calculate the number of electrons transferred and the potential of the electrde
Explanation / Answer
1.
a) The Nernst equation is:
E = E0 - (RT/nF)ln([red]/[ox])
So we need to identify for each half-reaction which species is the reduced one and which is the oxidized one.
Let's start with the first half-reaction:
Fe(CN)63- + e- -> Fe(CN)64-
From this reaction we can know Fe(CN)63- is the oxidized form while Fe(CN)64- is the reduced form of the compound. Tip: the oxidized form is the one present in the side where the electron appears.
The Nernst equation for this would be then:
E = E0 - (RT/nF)* ln([Fe(CN)64-]/[Fe(CN)63-])
For the second half-reaction:
C6H8O6 -> C6H6O6 + 2H+ + 2e-
We know C6H8O6 is the reduced form and C6H6O6 is the oxidized form.
So, the Nernst equation for this half-reaction is:
E = E0 - (RT/nF)*ln ([C6H8O6]/[C6H6O6 ][H+]2)
Notice we also included the H+ concentration because it appears on half-reaction as well.
b) Why do we only consider the ferricyanide/ferrocyanide couple for our calculations?
Because from all the species present in the solution (ferricyanide, ferrocyanide, ascorbic acid and dehydroascorbic acid), we can only measure ferricyanide using spectophotometry (the others are transparent to the wavelenght we'll use so we won't be able to measure them). The ferrocyanide concentration however, is inversely proportional to the concentration of ferricyanide and this way we can get both concentrations.
Note: If we wanted to use the ascorbic acid/dehydroascorbic acid pair for our calculation we would need to use another technique to follow the titration, possible UV spectophotometry.
Related Questions
Navigate
Integrity-first tutoring: explanations and feedback only — we do not complete graded work. Learn more.