• we first react the sample with a color developing reagent
• now then make the measurements as above

Background--Absorption of Light and Absorbance

• When visible light passes through a colored sample
• Photons of light are absorbed
• An electron is excited from one orbital to a higher energy orbital
  • The color is fixed by the orbital energies involved
• The more concentrated the solution, the more photons absorbed
  • We usually measure I/Io
  • the ratio of light intensities into and emerging from the sample
• It is more useful to compute the Absorbance
  • defined, Absorbance = - log₁₀(I/Io)
• if a sample absorbs no light the absorbance is 0
• if a sample absorbs 90% of the light
  • I/Io = 0.1 and Absorbance is 1.0
• if a sample absorbs 99% of the light
  • I/Io=0.1 and Absorbance= 2.0

Beer's Law states

• Absorbance = (constant) x (concentration) x (path length)
  • we will always work with a path length of 1.00 (cm)
  • the constant is called the molar extinction coefficient
  • the constant varies with wavelength
  • we generally select the wavelength where the sample absorbs light most strongly

• We will begin by examining a sample using the spectrometers we used during the first week of the semester. These will display a graph of the absorbance vs. wavelength. This lets us find the wavelength that is best suited to this analysis.
• We then switch to a simpler spectrometer.
  • The brand name is the Spectronic 20
    • This instrument allows you to select one wavelength
    • It then displays absorbance on its meter
    • A few calibration steps are required

The Chemistry Involved

We want to determine the concentration of iron (Fe²⁺ or Fe³⁺) in a sample. Iron compounds are generally not strongly colored, so direct spectrophotometry is not practical. However, we can find a reagent that reacts with Fe²⁺ to produce a strongly colored molecular complex. Measuring the concentration of this species is equivalent to measuring the concentration of iron.

The reagent of choice is a colorless compound called 1,10 phenanthroline.

This species has two nitrogen atoms with available electron pairs. Each of the nitrogens can bind to an Fe²⁺ ion. In the language of week two, this is a bidentate ligand. If there is an excess of the reagent, we will end up with three phenanthroline molecules bound to each iron ion.

This complex is an intense red color and it easy to measure with a spectrophotometer. It is easy to measure the iron in very dilute solutions-- concentrations of 1 mg/liter (1 part per million) are easily measured. That's around 2 x 10^-5 Molar.

We have two complications when we analyze our K₃Fe(C₂O₄)₃ :

• first, there is no free iron-- it's all tightly bound into the Fe(C₂O₄)₃3- anion.
• treating the solution with acid can quickly free the iron from the oxalate.
• even then, it's Fe³⁺ and not Fe²⁺ and we won't get the red complex
• treating the solution with hydroxylamine will reduce iron(III) to iron(II)
• reaction: 4Fe³⁺ +2 NH₂OH ---> 4 Fe²⁺ + N₂O + 4H+ + H₂O
• even if the sample were Fe²⁺ we'd need this treatment since some Fe²⁺ would be air oxidized to Fe³⁺ before we made our measurements.

Experimental Procedure

• Solution with precisely known Fe content (will be about 0.050 mg/ ml)
• 10% Hydroxylamine Hydrochloride (to reduce Fe₃₊)
• 1 M Ammonium Acetate (to act as a buffer and control pH)
• 0.3% 1,10 phenanthroline
• 6M Sulfuric acid

Special Equipment

• Buret
• volumetric flasks, 100 l and 1000 ml
• pipet, 10 ml, and pipet bulb
• Spectrophotometer

Standard Solutions and the Calibration Plot

• Using a buret, carefully measure 2 ml of the iron standard into a 100 ml volumetric flask
  • the flask may be wet
  • measure the volume precisely, + 0.02 ml (it need not be exactly 2.00)
• Add 1 ml of the 1M Ammonium acetate
• Add 1ml of Hydroxylamine Hydrochloride
  • mix (swirl a few times)
• Add 10 ml of the 0.3% 1,10 phenanthroline
• Fill flask to the mark with distilled water
  • stopper, mix by inverting 5 times (there's no other way to mix in these flasks)
  • transfer to a clean dry flask, beaker or test tube
    • you will only need about 10 ml of this mixture
    • if the container is wet, rinse it several times with the rest of the sample
• Set this aside for 30-45 minutes before making measurements
• Repeat with samples of 4, 6, 8 and 10 ml of the iron solution.

Solution of the Iron Salt

• Start by accurately weighing about 300 mg of your green salt
• Transfer it to a 1000 ml volumetric flask
• Add 25 ml of 6M H₂SO₄
  • swirl carefully to dissolve
  • let this mixture stand for 15 minutes
• The fill the flask with distilled water (to the mark)
  • stopper and mix by inverting 5 times
• Now use a pipet and transfer 10 ml of this liquid to a 100 ml volumetric
• Treat it the same way as the iron standards above
  • (add ammonium acetate, hydroxylamine, 1-10 phenanthroline, dilute to volume)
  • wait 30-45 minutes before measuring the absorbance.

The Spectronic 20 Spectrophotometer

The basic design of the instrument is shown in the cartoon

• a light bulb supplies light
• a lens aims the light at a diffraction grating
• the grating disperses the light into the spectrum
• a small slit selects light of a very narrow spectral range
• a knob allows us to rotate the grating and select the wavelength
• the sample is placed in the light beam
• a light detector measures light intensity
• electronics converts intensity to absorbance

An external diagram of a Spectronic 20

• Button "a" switches the instrument from %T to Absorbance
• Knob "b" allows you to select the wavelength. the value is displayed on the panel meter-- we want this set to 510 nm for the iron measurements.
• Knob c is used to set the meter to 0%T when no sample tube is present and the light beam is blocked.
• Knob d is used to set the meter to read 100%T when a tube of water is inserted into the cell compartment ("e" on the drawing.)
• Pressing "a" will allow you to set meter to read absorbance.
• Each sample tube is placed into compartment "e" and the Absorbance is read from the meter.

Calculations

• for the standard solutions
  • determine the concentration of Fe in each of the standard solutions
  • Plot Abs (Y axis) vs. concentration (X axis)
  • The slope of the graph is the molar extinction coefficient

• For the sample
  • Divide the Absorbance by the molar extinction coefficient
  • this calculates the concentration of iron in this solution
  • compute back to determine the moles of iron per gram of sample
  • compare with the expected result of 1 mole/ 491 grams of salt.