Name: 拉凯
Student ID: 27190080
Title: Spectrophotometric determination of trace iron.
Abstract:
The goal of this lab experiment is to quantify the concentration of an unknown material using
spectrophotometric measurements and spectrochemical measurements. When iron (II) reacts
with 1,10-phenanthroline, it produces a red color in a solution with a pH ranging from 2 to 9. The
light intensities of the red color of this solution will be measured and then a calibration curve will
be plotted. This plot will consist of absorbance versus concentration of iron (II)+ and the
concentration of the unknown sample will be analyzed and then determined from the data
constructed.
This approach is known as colorimetric analysis, and it consists of a change in color intensity that
is used to determine the concentration of an unknown at various concentrations. The difference
in intensity of the light beam can be measured using two methods: percent transmittance and
the Beer-Lambert Law. With these absorbance and concentration connections, it is simple to
create a calibration curve, which is a significant tool in the scientific community when determining
unknowns.
During this lab experiment, a stock solution will be prepared with proper dilutions and
concentrations. Iron will be reduced from Fe3+ to Fe2+ by using the excess of hydroxylamine
hydrochloride yielding 4xferrous iron, nitrous oxide, 4H+, and H2O. The reduction of Fe3+ to Fe2+ is
necessary to increase solubility. Sodium acetate is used to create a buffer environment since the
pH range for this reaction is from 2-9 as mentioned before. Sodium acetate is going to be utilized
to neutralize the acid present and it changes the pH in the range that the complex produces. The
blank solution is utilized as a reference solution.
It is very important in the science community to measure unknown concentrations of iron and to
be able to determine the concentrations based on various spectroscopic techniques because it
relates to the measuring of iron in the human body. High concentrations of iron in the body trigger
oxidative stress and iron deficiency and this may cause adverse health issues. Even though
measuring iron in the body by utilizing colorimetry is not common, other methods are used for
this purpose.
Introduction:
The three-ringed nitrogen-based compound 1,10-phenanthroline (also known as C12H8N2 or oPhen) interacts with certain metals like iron, nickel, ruthenium, and silver, resulting in vibrant
color complexes. This distinctive quality provides a precise method for detecting these metal ions
in water-based solutions. In a typical experiment, o-Phen combines with ferrous ions to yield a
deeply red-colored complex, following the reaction:
Fe2+ + 3o-Phen → [Fe(phen)3]2+
Iron is capable of existing in two forms: Fe2+ (ferrous ion) and Fe3+ (ferric ion). For accurate total
iron estimation, any Fe3+ present must be reduced to Fe2+ before introducing phenanthroline to
initiate the complex formation. Hydroxylamine hydrochloride is selected as the reducing agent in
this procedure, which interacts with Fe3+ as shown in the equation:
2Fe3+(aq) + 2NH3OH+(aq) → 2Fe2+(aq) + N2 (g) + 2H2O(l) + 4H+(aq)
This is an equilibrium reaction that can be pushed >99.99% towards the product side, even with
a gentle reducing agent. The pH of the solution is moderated between 6 and 9 by incorporating a
sodium acetate buffer.
Several factors influence the amount of light absorbed by a specific compound: concentration,
path length, wavelength, and solvent. The relationship between absorbance and concentration is
governed by the Beer-Lambert law:
A = εbc
In this equation, 'A' represents the absorbance measured by the spectrophotometer, 'ε' is the
extinction coefficient (indicating the capacity of the compound to absorb photons at a specific
wavelength, usually given in L.cm-1.mol-1), 'b' signifies the path length of the cuvette in cm (a 2
cm path length cuvette is utilized in this experiment), and 'c' denotes the solution concentration
in mol.L-1.
By preparing and plotting a series of solutions with known concentrations against their
absorbance (concentration on x-axis, absorbance on y-axis), one can generate a Beer's Law plot.
The slope of this plot corresponds to 'ε'. Consequently, this graph can assist in determining the
concentration of an unknown solution."
Materials and apparatus:
For this experiment, we utilized the following materials: hydroxylamine hydrochloride, 1,10phenanthroline, sodium acetate (1 mol/L), and NH4Fe(SO4)2 (10 mg/L) as reagents. We employed
a volumetric flask (50 mL), cylinder, and pipette for the preparation and measurement of the
reagents..
Procedure:
Preparation of the standard and unknown solutions.
1. Preparation of solutions:
i.
ii.
iii.
iv.
Utilizing the volumetric flask, several standard solutions of NH4Fe(SO4)2 with
concentrations of 0.40 mg/L, 0.80 mg/L, 1.20 mg/L, 1.60 mg/L, and 2.00 mg/L were
prepared. The appropriate volume of the 10 mg/L NH4Fe(SO4)2 stock solution was added
to each flask using a pipette, then the flasks were filled to the 50 mL mark with distilled
water and mixed well.
To each of these flasks, 1 mL of hydroxylamine hydrochloride was added using the pipette.
This reduced any Fe(III) present to Fe(II).
Next, 5 mL of sodium acetate (1 mol/L) was added to each flask using the pipette. This
acted as a buffer to maintain an optimal pH for the formation of the Fe(II)-1,10phenanthroline complex.
Finally, 2 mL of 1,10-phenanthroline was added to each flask. This formed a colored
complex with Fe(II) that could be detected using the spectrophotometer.
2. Measurement of absorbance:
i.
ii.
iii.
iv.
The 721UV-Vis Spectrophotometer was set up according to the manufacturer's
instructions.
Before measuring the samples, the spectrophotometer was "zeroed" using a blank. The
blank was a solution containing everything the samples had except the analyte
(NH4Fe(SO4)2). This blank was prepared using the same method as for the samples, but
without adding NH4Fe(SO4)2.
The blank was inserted into the spectrophotometer, and the instrument reading was
adjusted to zero absorbance.
Once the spectrophotometer was zeroed, the absorbance of each prepared solution was
measured. The cuvette was rinsed with a small amount of the next solution to be
measured before each measurement to avoid cross-contamination.
3. Data analysis:
i.
ii.
iii.
The absorbance of each solution was recorded, and then a Beer-Lambert plot of
absorbance vs. concentration for the known solutions was created.
The slope of the best-fit line for the data, which represented the molar absorptivity, was
determined.
The equation of the best-fit line was used to calculate the concentration of any unknown
solutions based on their measured absorbance.
Prepare the Beer’s Law plot:
Pipet a certain amount of the solution from Flask 1~7 into cuvettes. Measurement of the
absorbance under 510 nm. Recorded the absorbance from the screen. Filled out the form
below.
Label of the
volumetric
flask(50mL)
1
2
3
4
5
6
7(unknown)
NH4Fe(SO4)2(mL)
0.00
2.00
4.00
6.00
8.00
10.00
5.00
Hydroxylamine
hydrochloride(mL)
1.00
1.00
1.00
1.00
1.00
1.00
1.00
NaAc(mL)
5
5
5
5
5
5
5
1,10phenanthroline(mL)
2.00
2.00
2.00
2.00
2.00
2.00
2.00
C(Fe2+)(mg/L)
0.00
0.40
0.80
1.20
1.60
2.00
1.052
Absorbance
0.00
0.092
0.158
0.233
0.328
0.399
0.202
Result
1. Prepare a plot of absorbance versus concentration of the known solutions. Draw the best-fitting straight
line through the points – this is called the Beer-Lambert Law plot.
0,45
0,4
0,35
Absorbance
0,3
0,25
0,2
0,15
0,1
0,05
0
0
0,5
1
1,5
2
Concentration(M)
2,5
Figure: Beer-Lambert Law plot
2. Place the best Absorbance value of each unknown solution onto this plot and determine their
concentrations.
0,45
0,4
0,35
Absorbance
0,3
0,25
0,2
0,15
0,1
0,05
0
0
0,5
1
1,5
Concentration(M)
2
2,5
Figure: Beer-Lambert Law plot
The best Absorbance value of each unknown solution is 0.202 and the concentration is 1.052.
3. Calculate the amount of iron in the unknown sample. Express this as mg of Fe per liter of the original
unknown solution (mg/L Fe).
First, We created a calibration curve from the known data using the Beer-Lambert law. The
equation for a linear regression line is y = mx + b, where y is the absorbance, m is the slope of
the line (which represents the molar absorptivity), x is the concentration of Fe2+ (mg/L), and b is
the y-intercept.
To find the slope m, we calculated the change in absorbance divided by the change in
concentration between two points on the line. Let's choose the points from solutions 2 and 6.
The slope m is calculated as follows:
m = (y2 - y1) / (x2 - x1) m = (0.399 - 0.092) / (2.00 - 0.40) m = 0.307 / 1.60 m = 0.192
The y-intercept b should be close to 0, as when the concentration is 0, the absorbance should
also be 0 according to the Beer-Lambert law. Looking at the data, it seems that our assumption
holds, as the absorbance for the solution with 0.00 mg/L Fe2+ is indeed 0. Therefore, b is
approximately 0.
Now, we used this equation and the absorbance of the unknown solution to solve for its
concentration:
x = (y - b) / m x = (0.202 - 0) / 0.192 x = 1.052 mg/L Fe2+
Therefore, the concentration of Fe2+ in the original unknown solution is approximately 1.052
mg/L.
Discussion:
1. What is the purpose of adding hydroxylamine hydrochloride and sodium acetate?
Hydroxylamine hydrochloride and sodium acetate are commonly used in iron analysis,
particularly in spectrophotometric methods for the determination of iron.
i.
Hydroxylamine Hydrochloride: This is typically used as a reducing agent in these kinds of
analyses. Iron in solutions can exist in two oxidation states - Fe(II), also known as ferrous
iron, and Fe(III), also known as ferric iron. For spectrophotometric determination, iron
needs to be in the Fe(II) state. However, in many cases, the iron in your sample may be in
the Fe(III) state. Hydroxylamine hydrochloride reduces Fe(III) to Fe(II), ensuring that all the
iron in the solution is in the correct state for the analysis. This is particularly important
when using 1,10-phenanthroline as the complexing agent because it selectively binds to
Fe(II), not Fe(III).
ii.
Sodium Acetate (NaAc): Sodium acetate is typically used as a buffer in these analyses. A
buffer is a solution that resists changes in pH. In the case of iron analysis with 1,10-
phenanthroline, the complex that forms between Fe(II) and 1,10-phenanthroline is pHdependent, with optimal formation occurring around pH 3-9. Sodium acetate helps to
maintain the pH of the solution within this range, ensuring that the complex forms
correctly and that the analysis is accurate.
Conclusion:
In the conducted experiment, the principles of spectrophotometry and the Beer-Lambert Law
were employed to detect and quantify trace amounts of iron in a solution. Using the compound
NH4Fe(SO4)2, a series of standard solutions with known concentrations were prepared. A colored
complex was generated with the addition of hydroxylamine hydrochloride and 1,10phenanthroline. This allowed for the measurement of the absorbance of each solution using a
721UV-Vis spectrophotometer. From these known solutions, a calibration curve was plotted
which established the relationship between absorbance and concentration.
The slope of the calibration curve denoted the molar absorptivity of the Fe(II)-1,10phenanthroline complex under the given conditions. Using this calibration curve, the
concentration of an unknown solution was determined based on its measured absorbance. The
determined concentration of iron in the original unknown solution was approximately 1.052
mg/L. This conclusion highlighted the sensitivity and effectiveness of spectrophotometric
methods in the detection and quantification of trace elements in solutions.
References:
1. E. Merdivan, M. L. Yilmaz, "Spectrophotometric Determination of Iron and Cobalt with
1,10-Phenanthroline and Nitrite", Analytica Chimica Acta, Volume 171, 1985, Pages 327331, ISSN 0003-2670, https://doi.org/10.1016/S0003-2670(00)84014-2.
2. Paul A. Flowers, Klaus Theopold, Richard Langley, "Chemistry", OpenStax, Rice University,
2015. https://openstax.org/details/books/chemistry.