L-Ascorbic acid, also known as vitamin C, is an essential human nutrient. It is believed to play a b — Physical Chemistry — Kinetics Chemistry Question
Kinetics of Ferricyanide Oxidation of Ascorbic Acid
L-Ascorbic acid, also known as vitamin C, is an essential human nutrient. It is believed to play a biochemical role as an antioxidant, protecting against damage from reactive oxidants by virtue of its ability to be easily oxidized itself. In this experiment, you will investigate the kinetics of oxidation of ascorbic acid by hexacyanoferrate(III) ion, Fe(CN)6 3–, also known as ferricyanide, running the reaction in the presence of more than 10-fold excess of the reducing agent. The bright yellow color of ferricyanide ion (λmax = 416 nm) is lost on its reduction to colorless ferrocyanide ion , allowing one to monitor the progress of the reduction of ferricyanide spectrophotometrically.
[VISUAL]
Chemicals and reagents
• L-Ascorbic acid (abbreviated HAsc)
• Potassium hexacyanoferrate(III) (potassium ferricyanide), K3[Fe(CN)6]
• Aqueous hydrochloric acid solution, (c = 0.120 mol dm–3)
• Deionized water
Equipment and glassware
• Analytical balance (± 0.0001 g)
• Volumetric flasks (2), 10 cm3 or 25 cm3
• UV-visible spectrophotometer capable of measuring absorbance at 416 nm
• Spectrophotometric cuvette, 1 cm path length
• Plastic Beral pipettes, 1 cm3 (4), graduated in increments of 0.25 cm3
Procedure
1. Prepare stock solutions of ascorbic acid (~0.060 mol dm–3) and of potassium ferricyanide (~6.0 ⋅ 10–3 mol dm–3) (10 or 25 cm3 each). The concentrations need not be exactly as stated, but you should record the exact concentrations of the stock solutions.
2. Using the Beral pipettes to dispense the solutions, mix 0.75 cm3 deionized H2O, 1.50 cm3 aqueous HCl, and 0.50 cm3 of the ascorbic acid stock solution and place the solution in a cuvette. If you have a single-beam spectrophotometer, blank the spectrophotometer using this solution. If you have a double-beam spectrophotometer, make up a second identical solution and use this as the reference sample.
3. Initiate the reaction by adding 0.25 cm3 of the ferricyanide stock solution to the above mixture and mixing thoroughly. If your cuvette has a lid that seals tightly, you can mix the solution in the cuvette itself. If the cuvette does not have a tight-fitting lid (or has a volume less than 3 cm3), you will need to mix the solution in a small vial, then transfer a portion of the mixed solution into the cuvette. As quickly as possible, replace the cuvette in the spectrophotometer and begin measuring the absorbance at 416 nm as a function of time.
4. Record absorption at 416 nm, A416, as a function of time over the course of 10 minutes. In the early part of the reaction (when the absorbance is changing rapidly), you should record the absorbance frequently (every 10 seconds or so), but as the reaction slows, you can make less frequent readings if you wish (every 30 seconds or so).
5. Repeat steps 2 – 4 as needed to explore the effect on the rate of varying the ascorbic acid concentration in the range [HAsc] = 0.005 – 0.015 mol dm–3 and of the acidity in the range [H+] = 0.01 – 0.10 mol dm–3. If the reaction is slower than the initial experiment, you may need to extend the monitoring period to 15 or 20 minutes in order to allow the reaction to go nearly to completion (the absorbance, A416, should fall below 0.02).
Give a balanced chemical equation for the oxidation of ascorbic acid by hexacyanoferrate(III) ion. Include a structural formula for the oxidation product of ascorbic acid.
Model Answer
Balanced chemical equation:
L-Ascorbic acid + 2 Fe(CN)6^3- → Dehydroascorbic acid + 2 Fe(CN)6^4- + 2 H+
[VISUAL] showing the skeletal structure of L-ascorbic acid reacting with 2 Fe(CN)6^3- to form dehydroascorbic acid (with two ketone carbonyls in place of the enediol system), 2 Fe(CN)6^4-, and 2 H+.
Determine the reaction order in Fe(CN)6 3–, and justify your determination.
Model Answer
All sample data were obtained at 18.2 °C; stock [HAsc] = 0.0608 mol dm–3; stock [Fe(CN)6 3–] = 6.26 ⋅ 10–3 mol dm–3.
In these reactions, all reagents are in at least 10-fold excess over the ferricyanide, so the only concentration that changes significantly over the course of a kinetics run is [Fe(CN)6 3–] ("pseudo-first-order conditions"). The order in ferricyanide can be determined by analysis of the time-dependence of the absorbance (which is related to the concentration). The modern way to do this is to fit the A416 vs. time plot directly to the mathematical expression appropriate for the given order (using a computer to do the nonlinear least-squares fitting). Thus, if the reaction is first-order in [Fe(CN)6 3–], one would expect to observe:
A(t) = Af + (A0 - Af) * e^(-kobs * t)
where A0 and Af are the initial and final (t = ∞) absorbance values, respectively, and kobs is the apparent first-order decay constant. This is in fact invariably observed here [VISUAL].
One can also analyze the data in the old-fashioned way, by plotting ln(A – Af) vs. t. This plot will be linear for four half-lives, and the negative slope of this plot gives kobs [VISUAL]. For this reaction, the products are nearly colorless, so Af is close to zero, but in fact a small positive Af is generally observed. Neglecting this causes a systematic error of ~5% in the rate constants measured here. The direct fit above is preferred because it is less sensitive to errors in measuring Af (since Af is treated as an adjustable parameter) and since it statistically weights the early data (where the most change is taking place) more heavily than the later data. In contrast, the natural log fit overemphasizes the late data, where the A – Af difference is small and where small errors cause large changes in the logarithm.
Either linearity of the ln(A – Af) plot or conformity to exponential decay in the direct plot is satisfactory evidence that the reaction is 1st-order in Fe(CN)6 3–.
Determine the reaction order in HAsc, and justify your determination.
Model Answer
By changing the concentration of HAsc, one can measure the variation of kobs with [HAsc] and deduce the order in HAsc. At constant [H+],
rate = k [Fe(CN)6 3–] [HAsc]^n
kobs = k [HAsc]^n
The concentration of HAsc was varied systematically (at [H+] = 0.0600 mol dm–3) as indicated in the table below. A plot of kobs vs. [HAsc] [VISUAL] gives a straight line with an intercept of zero within experimental error. This clearly indicates that the reaction is 1st-order in HAsc.
Run data (all with 0.25 cm3 K3Fe(CN)6 soln added):
• Run 1: 0.75 cm3 H2O, 1.50 cm3 HCl, 0.50 cm3 HAsc, kobs = 0.00899 s–1
• Run 2: 1.00 cm3 H2O, 1.50 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00444 s–1
• Run 3: 1.00 cm3 H2O, 1.50 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00441 s–1
• Run 4: 0.50 cm3 H2O, 1.50 cm3 HCl, 0.75 cm3 HAsc, kobs = 0.01333 s–1
Ascorbic acid readily ionizes to form the ascorbate anion, Asc–, with a pKa = 4.10 (Ka = 7.9 · 10–5). Indicate which proton in ascorbic acid is readily ionized and explain why it is so acidic.
Model Answer
Ascorbic acid is a vinylogous carboxylic acid. Loss of the proton from the 3-OH group (on the enediol system) forms a conjugate base (ascorbate monoanion) that is strongly stabilized by resonance delocalization of the negative charge across the conjugated carbonyl system:
[VISUAL] showing resonance structure delocalization.
The dependence of the reaction rate on [H+] is somewhat complex (it does not exhibit a simple, integer order). A plausible explanation for this is that both ascorbic acid (HAsc) and ascorbate anion (Asc–) can be oxidized by hexacyanoferrate(III) ion, but that they have different reactivities. Use this model to analyze your data quantitatively to determine the relative reactivity of ascorbate anion and ascorbic acid toward Fe(CN)6 3–.
Model Answer
The concentrations of HAsc and Asc– are related by the ionization equilibrium of HAsc:
Ka = [Asc-][H+] / [HAsc]
[Asc–] = Ka * [HAsc] / [H+]
Since [H+] >> Ka in this experiment, almost all the ascorbic acid is in the form of HAsc, and the concentration of Asc– is small. This means that we can use [HAsc] ≈ concentration of added ascorbic acid.
If the two forms of ascorbic acid react at different rates, then:
rate = k1[HAsc][Fe(CN)6^3-] + k2[Asc-][Fe(CN)6^3-]
kobs = k1 [HAsc] + k2 [Asc–] = k1 [HAsc] + k2 * Ka * [HAsc] / [H+]
kobs = [HAsc] * (k1 + (k2 * Ka / [H+])) [VISUAL]
A linear relation can be obtained by multiplying by [H+]:
kobs * [H+] = [HAsc] * (k1 * [H+] + k2 * Ka) [VISUAL]
Analysis of reactions conducted at constant [HAsc] confirm that the rates do indeed depend on [H+] in this manner:
• Run 2: 1.00 cm3 H2O, 1.50 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00444 s–1
• Run 3: 1.00 cm3 H2O, 1.50 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00441 s–1
• Run 5: 1.50 cm3 H2O, 1.00 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00503 s–1
• Run 6: 2.00 cm3 H2O, 0.50 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00618 s–1
• Run 7: 2.25 cm3 H2O, 0.25 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00848 s–1
• Run 8: 0.00 cm3 H2O, 2.50 cm3 HCl, 0.25 cm3 HAsc, kobs = 0.00418 s–1
Using the values from the linear relationship:
0.00367 s–1 = k1 [HAsc]
→ k1 = 0.724 ± 0.012 dm3 mol–1 s–1
4.92 ⋅ 10–5 mol dm–1 s–1 = k2 * Ka * [HAsc]
→ k2 = 123 ± 8 dm3 mol–1 s–1
Since k2 / k1 ≈ 170, the ascorbate anion is 170 times more easily oxidized by ferricyanide than is L-ascorbic acid itself.