CHEMISTRY 221 LABORATORY—QUANTITATIVE ANALYSIS

EXPERIMENT:  SPECTROPHOTOMETRIC DETERMINATION OF IRON

Spectrophotometry:
For chemical species that appear to have color, it is a logical assumption that the intensity of the color is proportional to the concentration of the species in solution. We see color as a complement of the visible wavelength being absorbed by the sample. Things that appear red absorb blue visible light and reflect other visible colors to our eyes. Conversely, things that appear blue are absorbing red light. Absorption of light or more precisely electromagnetic radiation is related to available energy levels in the molecule or ion. A molecule in its "ground state" or lowest energy level can absorb energy to jump to an "excited state" or higher energy state. The amount of energy and therefore the wavelength of radiation involved in this transition is a function of the electronic structure of the molecule or ion.

The eye can only see a limited range of electromagnetic radiation, from approximately 400 to 700 nm. However, molecules, atom, and ions are capable of absorbing many different energies of radiation ranging from ultraviolet (UV) to microwaves depending on the specific energy levels being excited. For some types of energy changes, the wavelengths of light are very specific for certain types of chemical structure resulting in a method of qualitatively identifying chemical species. Other types of energy absorption may be less qualitative since it may relate only to bond types. In both cases however, our initial premise that intensity of absorption is related to concentration can be used for quantitative analysis.

Since our vision if not quantitatively calibrated, an electronic instrument called a spectrophotometer is used to precisely measure light intensities at given energy (wavelength) settings. A spectrophotometer is an instrument that measures the amount of transmission of light through a substance. The drawing below illustrates a simple spectrophotometer system consisting of a light (energy) source, a monochromator to select a given energy range, a sample, and a light intensity detector.

When light is absorbed by a sample, the radiant power or intensity of the light beam decreases. Radiant power, I, refers to the energy per second per unit area of the beam. In the figure, light passes through a monochromator that selects one wavelength. Light of this wavelength, with radiant power I0, passes through a sample of pathlength b. The radiant power of the beam emerging from the other side of the sample is I. Mathematically, the amount of light that is absorbed (A) is given by

Note that if no light is absorbed, A = 0 and if all the light is absorbed ( I = 0) then A = ¥. The amount of light absorbed by the sample should be proportional to the probability that the molecule or ion will absorb the electromagnetic radiation (a), the number of absorbing molecules or ions per unit volume that the light beam passes through (C), and the length of the light path (b). This relationship is quantified in the Beer-Lambert (or Beer's) Law which is 

A = a × b × C 

Note that this equation is in the form of Y = m × X + b where the intercept, b, is zero when X or the concentration, C, is zero. If we measure a series of solutions of known C at a given wavelength in a cuvet or sample cell with a constant pathlength, b, then we can determine the proportionality constant, m, which is a × b. This procedure generates a "calibration curve" which allows the determination of an unknown concentration, C_unk, from the measurement of the absorbance of the unknown, A_unk. Determination of the slope, m, and intercept, b, of the calibration curve gives

Many of the transition metal ions such as copper, nickel, cobalt, and chromium exhibit color in solution. However, this color can be made more intense by reacting the metal ion with a molecule that increases the absorbance of the metal ion. Iron(II), Fe²⁺, exhibits little color in solution. When Fe²⁺ reacts with the ligand o-phenanthroline (or 1,10-phenathroline), a stable, intensely colored red complex is formed that can be used to determine iron. The intensity of the color varies over the pH range of 2 to 9. In this procedure an ammonium acetate buffer will adjust the pH to between 6 and 9.

The iron must be in the +2 oxidation state, requiring a pre-reduction step before formation of the colored complex. Hydroxylamine is used as a reducing agent 

2 Fe³⁺  +  2 NH₂OH  +  2 OH^-      2 Fe²⁺  +  N₂  +  4 H₂O

EXPERIMENTAL PROCEDURE 

Preparation of Standards and Determination of a Calibration Curve
You will be divided into groups to prepare the following solutions: 

1. Prepare a stock Fe solution by accurately weighing to the nearest 0.1 mg approximately 0.07 g of pure iron (II) ammonium sulfate hexahydrate and quantitatively transferring to a 1 L volumetric flask. 

2. Add 200 mL water and shake to dissolve any remaining solid. 

3. Add 5 mL of 1:1 sulfuric acid. 

4. Dilute to the mark with distilled water and homogenize thoroughly. 

5. Calculate the concentration of the solution in mg Fe/L.  

6. Prepare a series of standards by pipetting into each of five 100 mL volumetric flasks, 1.00, 5.00, 10.00, 25.00, and 50.00 mL aliquots of the stock Fe²⁺ solution (a buret can be used for this addition). 

7. Into a sixth 100 mL volumetric flask pipet 50 mL of distilled water to serve as a blank. 

8. To all of the solutions add, in sequence   
         1 mL of hydroxylamine hydrochloride solution, 
         10 mL of 1,10-phenanthroline, and 
         8 mL of sodium acetate buffer.

9. Dilute to the mark, mix thoroughly, and allow to stand for 10 minutes. 

It is important to add all reagents in proper sequence.

Determination of Unknown Ferrous Ammonium Sulfate:

In order to prepare your unknown in the desired concentration range for the spectrophotometric measurement it will be necessary to do a serial dilution. Each student will individually prepare their own unknown. Be sure to record your unknown number in your lab notebook. 

1. Accurately weight 0.70-0.75 g of your unknown to the nearest 0.1 mg. 

2. Quantitatively transfer the solid to a 1 L volumetric flask. 

3. Add 200 mL of water, 50 mL of 1:1 H₂SO₄ and dissolve. 

4. Dilute to the mark and mix thoroughly 

5. Pipet a 10.00 mL aliquot into a 100 mL volumetric flask. 

6. Dilute to the mark and mix thoroughly 

Finally prepare the actual sample for analysis. 

1. Take a 20 mL aliquot (2 × 10 mL) of this second solution and place it in a 100 mL volumetric flask. 

2. Treat this as you did your standards by adding · 
         1 mL of the hydroxylamine solution, 
         10 mL of the 1,10- phenanthroline solution, and 
          8 mL of sodium acetate. 

3. Dilute to the mark, mix thoroughly and allow to stand for 10 minutes.

Operating Procedure for the Spectronic 20:

1. Turn on the instrument with the zero adjust/on-off control and allow about five minutes warm-up time. Turn the wavelength control knob to select the desired wavelength of 512 nm.

2. With no sample tube in place and the cover closed, turn the zero adjust knob to bring the meter to 0% transmittance.

3. Fill the spectrophotometer tube to about 3/4 with distilled water or your blank solution. Make sure that the outside of the tube is clean and dry (do not handle the tube on the sides so as to avoid fingerprints) and insert it in the cuvet holder. Push the cuvet all the way down and aligning it so that the reference line on the cuvet holder matches a line (or other mark) on the tube. Close the cover of the cuvet holder. 

4. Adjust the light control knob on the spectrophotometer so that the meter reads 100% transmittance. Steps (2) - (4) should be repeated before each absorbance measurement. 

5. Fill (to 4/5) a tube with the solution to be measured. Wipe the outside of the tube to remove water and fingerprints and insert into the cuvet holder, aligning as above. Close the cover and read the % transmittance from the meter scale. 

Notes: Ideally, the same sample tube should be used for all absorbance measurements. Care must be taken to align the tube in exactly the same way each time. A separate tube may be used for the distilled water blank, but again always align it in the same way. Because the sample may be heated in the light beam or photochemical decomposition may occur, the absorbance reading should be made quickly after the tube is inserted. It is good practice to read and record the % transmittance meter reading and to convert to absorbance later. The conversion factor is A = 2 - log(%T). The transmittance scale is linear and thus less subject to reading error than the logarithmic scale.

Calculations:

In order to calculate the concentration of iron in your unknown sample, a calibration plot must be calculated. A calibration plot or working curve is a plot of the analytical signal (the instrument or detector response; in this case the absorbance (A) on the y axis) as a function of known analyte concentration (C) on the x axis. These calibration plots are obtained by measuring the signal from a series of standards of known concentration. The calibration plots are then used to determine the linearity of response of an analytical method.

Use a spreadsheet to generate a plot of your data. Perform a least-squares regression analysis of your data to determine the slope and intercept. If you are using Excel, you can use the "add trendline" feature to draw your best regression line and choose the "show equation" option to obtain the slope and intercept. Use the standards to prepare a calibration curve as directed by your instructor. Plot absorbance vs. concentration. Check the linearity of the curve to see if Beer's Law is obeyed.

Calculate the percentage of iron in your unknown at this point:    

• Report the result in terms of the percentage of iron in your unknown.  The range of unknown values should be 5% to 15%.

Return the CHEM 221L Laboratory Schedule