Determination of Iron (II) in a Vitamin Tablet
The iron found in most vitamins is in the form of iron (II), the Fe2+ ion. In this experiment, you will use the
spectrophotometer to determine the amount of iron (II) in a vitamin tablet. The Beer-Lambert law provides a
relationship between the color of a solution and the concentration of the color-causing species. You will use this
relationship to determine the concentration of the iron(II)-phenanthroline complex, which in turn can be used to
determine iron.
Materials:
Spectrophotometer
Mortar and pestle
Cuvettes
Filter paper/funnel
Standard iron (ii) solution
50-mL graduated cylinder
1M sodium acetate
Stirring rod
1% o-phenanthroline
2 100-mL beakers
Pipet bulb
Ring stand/ring/gauze
0.10 M HCl
2 10-mL graduated
cylinders
Bunsen burner
Vitamin tablet
1% hydroxylamine
hydrochloride
Distilled water
2 10-mL pipets
100-mL volumetric flask
Safety
1. o-Phenanthroline, hydroxylamine hydrochloride, and HCl are corrosive and toxic. You must wear safety goggles
during this experiment. Report any spills or exposure to the instructor.
2. o-Phenanthroline must be handled in the fume hood only.
Procedure:
Preparation of the Unknown Iron solution
1. Using the analytical balance, weigh your vitamin tablet accurately. Use the mortar and pestle to crush the tablet into
a fine powder. In your notebook, record this mass along with the mass of iron contained in the tablet as shown on
the vitamin label.
2. Tare a 100-mL beaker on the balance and use a scoopula to remove about 0.2 g of vitamin. Try to avoid the crushed
bits of vitamin coating and get only crushed vitamin powder. To the beaker add 20 mL of the 0.010 M HCl solution and 10
mL of distilled water. Mix the solutions thoroughly.
3. Heat the beaker over the flame to a gentle boil (this will not take long!). Boil it gently for 5-7 minutes, stirring
frequently. After it has boiled, allow it to slowly cool to room temperature. Do not worry if there is still some
undissolved material.
4. Filter the solution through the filter paper into a 50-mL graduated cylinder. Rinse the beaker and filter paper with
distilled water. Fill the graduated cylinder to the 50-mL mark with distilled water.
Preparation of standards
1.
Label five cuvettes 1-5. Add to them the following reagents in the amounts listed. The total volume in each case
should be 2000 μL. Use the 200-μL micropipet provided to dispense both the standard iron solution and solutions
B, C, and D. Be sure to change pipet tips when adding a different solution. Mix each solution thoroughly. Cover
these with Parafilm until you need them.
Tube
1
2
3
4
5 (blank)
A
Standard Iron
(μL)
200
400
600
800
None
B
Hydroxylamine
HCl (μL)
200
200
200
200
200
C
Sodium Acetate
(μL)
200
200
200
200
200
D
o-Phenanthroline
(μL)
200
200
200
200
200
Distilled Water
(μL)
1200
1000
800
600
1400
Preparation of the standard curve
1.
2.
3.
Use the SmartSpec to perform a wavelength scan. From the printout (include in your notebook) find λmax.
Read the absorbance for each of your four samples. Read each one twice, blanking using the blank each time.
Plot absorbance (y) versus concentration (in mg/L) (x) for the four solutions using Excel. Once you have plotted
this graph, use the regression line function in Excel to determine the equation for the line that connects your four
points. This is your “standard curve” which can be used to determine the concentration (x) of any unknown
solution given its absorbance (y).
Determination of the Unknown
1.
2.
3.
4.
5.
Pipet 5.00 mL of your tablet solution into a clean 100-mL volumetric flask. Add 5.00 mL each of the sodium
acetate, o-Phenanthroline, and hydroxylamine hydrochloride solutions to the flask. Mix thoroughly and then add 30
mL of distilled water (a total volume of 50.0 mL) Fill a clean cuvette with a small amount of this solution. Set the
spec to λmax, then zero and blank the spec using your prepared blank cuvette. Read the absorbance of the solution
three times (blanking each time).
Using your regression equation from step (4) above and the measured absorbance for your unknown (‘y”), solve for
x (the concentration) using the absorbance you measured for your unknown.
Since the solution was diluted 1/10 in step 1, multiply your concentration by 10 to obtain the correct value in
mg/L. Then, since you used 50 mL of solution, multiply this vale by 0.05 L to obtain your amount of iron, in mg.
Note: if your absorption does not fall within the range of your standard absorptions, dilute it again and measure the
absorption again. Repeat until you get a value in the range.
Use the volume of the tablet solution to determine the number of mg of iron in your sample (include any further
dilutions). Use the following formula to determine the number of mg of iron in your tablet:
mg iron x __total tablet mass__ = mg iron in the tablet
mass of tablet used
6.
Determine the percent error in your calculated concentration. Is your calculation high or low? Offer suggestions
why your answer might be off (error analysis) in your conclusion.
Data:
Mass of tablet (±0.0001g):
Mass of tablet used (±0.0001g):
Mass of iron in tablet:
Tube
1
2
3
4
Unknown
Calculated mass of iron:
Iron concentration in the unknown solution:
Percent Error:
Concentration (mg/L)
2.00
4.00
6.00
8.00
Absorbance