The purpose of this laboratory is to understand the redox chemistry and kinetics of the oxidation of ferrous iron by measuring the kinetics over a narrow pH range. The rate coefficient for Fe(II) oxidation will be compared with a value published in the literature (Stumm W. and Lee G.F., 1961, "Oxygenation of Ferrous Iron," Ind. Eng. Chem. 53, 143-146).
The TA will provide the following materials:
Prepare a calibration curve for measuring Fe(II)
concentrations using the Ferro-Ver (1,10-phenanthroline) method.
a. turn on the spectrophotometer, wait 10 min, set the wavelength to 508
nm, and zero the spectrophotometer.
b. create a 3.0 mg L-1 Fe(II) solution as directed by the
teaching assistant.
c. use the 3.0 mg L-1 Fe(II) solution to create Fe(II)
standards of 0.5, 1.0, and 2.0 mg L-1.
d. create a blank standard (0 mg L-1 Fe(II)).
e. measure the absorbance of the blank and the four standards using the spectophotometer:
i. rinse a 25 mL pipette with deionized water.
ii. take a 25.0 mL sample with the pipette and place it
in a Hach sample cell.
iii. add the iron reagent powder pillow to the sample
cell, invert to mix, and time for three minutes.
iv. place the sample cell in the spectrophotometer and
read absorbance.
f. measure the absorbance of one of the standards two more times to assess
the reproducibility of the measurement.
g. check the calibration curve (Abs vs. [Fe(II)]) for linearity.
Calibrate the pH meter using pH 4 and 7 standards.
Prepare the buffer solution.
a. add 125 mL of the Tris buffer solution and 350 mL of the deionized water to a 1 L
beaker (use the balance for accuracy).
b. place the beaker on the stir plate, add the stir bar, and begin
stirring.
c. adjust the pH to a pH of 7.4 to 7.6 by adding 1 M NaOH (the volume will increase, but
insignificantly).
d. aerate the solution by bubbling compressed air into the beaker through
the gas dispersion tube for 5
min.
e. re-adjust the pH if necessary and measure the temperature of the
solution.
f. continue to aerate.
Measure the dissolved oxygen concentration of the
buffer solution using the Hach DR/890 high range dissolved oxygen test kits as
in Lab 5.
Add Fe(II) to the buffer solution to start the
experiment.
a. add 15.0 mL of the ferrous ammonium sulfate solution to the beaker and
continue to aerate. This should produce
an initial concentration of Fe(II) of 2.0 mg L-1.
b. MARK THIS AS TIME ZERO.
Sample the Fe(II) solution over at the times indicated in the table below.
a. follow the same method for measuring Fe(II) as in step 1e.
i. rinse a 25 mL pipette with deionized water
ii. take a 25.0 mL sample with the pipette and place it
in a Hach sample cell
iii. add the iron reagent powder pillow to the sample
cell (record this time as the
exact sample time because
reagent addition stops the Fe(II) oxidation),
invert to mix, and time for
three minutes.
iv. place the sample cell in the spectrophotometer and
read the absorbance
b. record the pH in the beaker at this time.
c. dispose of the
solutions properly and rinse the sample cells with deionized water.
| sampling times (min) | |||||||||
| 2 | 6 | 10 | 15 | 20 | 25 | 30 | 35 | 40 | 45 |
The kinetics of ferrous iron oxidation play a key role in acid mine drainage. Normally, the oxidation of Fe(II) to Fe(III) is very slow in the pH range encountered in acid mine drainage waters. However, chemolithoautotrophic bacteria (e.g., Thiobacillus ferrooxidans) can mediate the reaction and derive energy from ferrous iron oxidation even at very low pH values. One way to lessen the impact of acid mine drainage is to prevent the microbial mediation of this reaction by killing the microbes (with various poisons, including surfactants) or fostering competition by seeding waste rock piles with heterotrophic bacteria and a carbon source. Prof. JoAnn Silverstein and her students in the Civil, Environmental, and Architectural Engineering Department have been conducting research on this competition approach.
(Results) Prepare a plot of the Fe(II) calibration curve (absorbance versus [Fe(II)]) and provide the linear regression used to determine Fe(II) concentration as an equation in the Results section.
(Results) Report the data showing the reproducibility of the Fe(II) concentration measurement.
(Results) Prepare plots of [Fe(II)] versus time and ln [Fe(II)] versus time for the experiment. Provide the linear regression of ln [Fe(II)] versus time as an equation in the text.
(Results) Following the kinetic expression for ferrous iron oxidation introduced by Stumm and Lee (1961) and discussed in lecture, determine the pseudo-first order rate coefficient (k') for the experiment.
(Results) Taking into account the pH (averaged over time) and dissolved oxygen content and temperature (to determine PO2) measured during each experiment, determine the rate coefficient k for the experiment (remember, k' = k [OH-]2 PO2). Explain and show your calculation of PO2, including the measured [O2(aq)] and the temperature correction for HO2. Notice that PO2 needs to be in units of atmospheres.
(Discussion) Discuss factors that might affect the reproducibility and accuracy of the Fe(II) concentration measurement.
(Discussion) Compare the value of k calculated for your experiment with the value of k reported by Stumm and Lee (1961), 8.0(±2.5) x 1013 M-2 min-1 atm-1, for 20°C in a system buffered with bicarbonate. Suggest reasons for any difference between your experimental k value and the Stumm and Lee k value.
(Discussion) If your value of k does not match the k value of Stumm and Lee (1961), what would the pH (i.e., {OH-}) or [O2(aq)] have to be to make the k values match? How likely is it that an inaccurate measurement of pH or dissolved oxygen concentration could explain the difference between the k values?
Last updated on July 31, 2007 at 07:19 AM by Joe Ryan