Analysis and Quantification of Iron properties in relation to Health and Illness

Learn about the analysis and quantification of iron properties in relation to health and illness. Understand the importance of iron in living organisms and how it needs to be regulated. Explore the role of ferritin in balancing iron levels in the body. Discover the experimental methods used to analyze and determine iron ion concentrations. Study the results and discuss the implications for biological processes and overall health.

Analysis and Quantification of Iron properties in relation to Health and
Illness
Introduction
Iron is an important mineral ion in almost all living organisms. This is because iron is required in
active sites of enzymes that undertake activities that enable life enabling processes. These
processes include respiration, DNA synthesis, regulation of genetic materials, storage and
transport of oxygen, and metabolism of energy.
Although iron, like many other essential minerals and sources of living organisms, is needed in
the body, it needs to be regulated. (Templeton, 2002)This is because excess iron can initiate
series of reactions in the cell and lead to production of reactive oxygen species (ROS) and other
free radicles that are devastatingly harmful and can lead to permanent cell damage, organ failure
and eventually cause death. In addition, excess iron ions in the bloodstream can react in the
bloodstream with oxygen to form insoluble compound, like rust which is extremely harmful in
one’s body. This can lead to rusty kidneys. (Gregory J. Anderson, 2012)And thus it is important
for a body to device ways on how to have iron levels in the bloodstream consistent.
The body’s mechanism is by use of ferritin. Ferritin is a universal storage site of protein found in
living organisms both prokaryotic and eukaryotic. This compound is found plants and animals,
and even in bacteria. Ferritin works in both intracellular, that’s inside cells, and extracellular,
that’s in bloodstream, environments. Ferritin balances the amount of iron in the body by storing
excess iron, that is iron overload, and releases it when there is a deficiency or needed. Ferritin is
shaped like a hollow cage and this enables it to store Fe 2+. Ferritin is made up of subunits and
most of ferritin is made of 24 units. (Worwood, 1982)Ferritin has the ability to get rid of toxic
oxygen, ROS, through reacting with ferrous ion and thus reducing the overload. Ferrous ions are
transported to ferritin by a protein for transport called transferrin. In the site of this ferritin
compound, ferrous ions are oxidized to less soluble ferric ions at the ferroxidase site. Ferritin has
the ability to store high volumes of iron of up to 4500 iron ions. (Lauffer, 1992)
When cells require ion, for instance ion is required for synthesis of enzymes after blood loss or
during the time of embryonic development, iron stored in ferritin is released in regulated rates
into the cells. The process of releasing iron ions involves complex chain reactions which
encompasses a process called reverse mineralization. In this process ferric ion or oxy mineral is
dissolved from its solid state to aqueous ferric ions and these ions reduced to ferrous ions.
In the experiment, the analysis and determination of quantities of iron ion has been done. It has
been shown how iron (III) or ferric ion is reduced to ferrous or iron (II) ions which are essential
mineral ions in our cell metabolism. (M. Worrall, 1991)The experiments done in the laboratory
try to analyze how a defect in the chemical properties of ferratin or its reducing agents could
affect biological processes and functionality of living organisms and even affect health and
wellbeing of the organisms.

In reference to the experiment and the experimental results it can shown that using various
laboratory chemicals together with serum (to give us ferritin containing iron ions) conversion of
ferric ions to ferrous ion is possible yielding quantifiable results and thus giving us better
understanding of biological processes that occur in most of living organisms. (Worwood, 1990)
The whole experiment involved various chemicals and measurement devices which initiated
series of reactions and processes. They included processes from the use of dilute acid,
dihydroxyfumarate, DHF, recording quantities of concentration of iron ion s using spectrometer,
reduction and oxidation of ferric and ferrous ion through redox reactions, breaking of ferritin
protein and data analysis and calculations
Materials and Methodology
The main of this experiment is to determine the amount and concentrations of iron ions in the
ferritin compound in the blood serum. (David Dunaief, 2014) This analysis and determination is
done by trace analysis of iron ions using a measurement device called spectophotometer.
They procedure in the experiment involved putting ferritin compound of known quantity and
concentration into a solution, the protein ferritin was broken and iron ions were released by usin
an agent by the anme chromophoric complexing agent ferrozine (II) ions. (Worwood, 1979)The
iron released was the measured and recorded which were used by the students to determine
moles of iron ions that were present in a mole of apoferritin, which is an empty ferritin.
Part 1
In the first part of the experiment we used standard solutions of, that is known concentration,
ferrous ions (iron (ii)). The solution containing ferrous ions was prepared by adding the dilute
acid into the solution. The acid is used in breaking apart ferratin protein so that iron ions
enclosed inside the ferratin sac is freed. When the iron ions had been released, we used a
chemical agent called dihydroxyfumarate. The redox agent was used in reducing insoluble ferric
ions to ferrous ions (Fe III to Fe II).
The 7 test tubes contained different amounts of reagents iron, water, dihydroxyfumarate, and
ferrozine. The reagents were the mixed as instructed and then left for two minutes at room
temperature for changes (reactions) to take place. The results were observed by
spectrophotometer and as shown in the table below. The absorbance of these materials was
measured. The reading values observed in the in each of seven test tubes was recorded as shown
below and the concentration of the samples (Fe II) under observation was calculated and
recorded as shown below.

Results: Table 1
Blank 1 2 3 4 5 6 7
1 mM Fe(NH4)2(SO4)2 ·
6H2O in H2SO4 0.02
M(mL)
0 0.04 0.06 0.08 0.10 0.12 0.14 0.16
Water
(mL)
1.52 1.48 1.46 1.44 1.42 1.40 1.38 1.36
5 mM
Dihydroxyfumarate (mL)
0.16 0.16 0.16 0.16 0.16 0.16 0.16 0.16
2.5 M NaOAc
(mL)
0.16 0.16 0.16 0.16 0.16 0.16 0.16 0.16
12.5 mM Ferrozine(mL) 0.16 0.16 0.16 0.16 0.16 0.16 0.16 0.16
[Fe(NH4)2(SO4)2 · 6H2O]
(mM)
0 0.02 0.03 0.04 0.05 0.06 0.07 0.07
Absorbance
values(spectrophotometer
) readings at 562 nm)
0.522 00.716 0.987 1.139 1.399 1.582 1.782
The graph of absorbance against concentration of Ferrous

The graph shows clearly that the absorbance increases with increase in concentration of ferrous
ions. This is shown as the graph is a straight line thus obeying the Beer-Lambert law which states
that absorption of substances is directly proportional to the concentration of the same material.
Calculating using Beer’s Lambert law is given by;
A=ℇIC
Where A = the absorbance of solution
Molar Absorptivity; I = length of solution cell and c is concentration of solution.
Part 2
Determination of Total concentration in Ferritin
0.4 ml of Ferritin solution 0.5 g/ml was put in a volumetric flask and sodium acetate of
concentration 0.15 M was added. Then 0.8 ml 2M sulphuric acid (H2SO4) was added to the flask
which we labelled sample A. For sample B the solution containing sodium acetate and ferritin
was added into 0.8 ml 5mM dihydroxyfumarate. We waited for around 30 minutes and then we
added 1.6 ml 2.5 M NaOAc to sample 1 and 0.8 ml 12.5 mM ferrozine was added to sample B.
We then again waited for 30 minutes after both samples were diluted up t 10ml. The observation
and recording of the spectrophotometer was recorded.
Part 3
The concentration of ferrous ions which is unknown was determined by a linear fit equation from
the calibration curve. The concentration was then multiplied by 25 factor. This is due to the
reason that the initial ferritin solution was diluted 25 times.
Finally, the equation of a straight line y=mx + c was used to calculate the concentration of
ferrous ions which was the unknown concentration.
Letting X be the unknown concentration, therefore; X= (y-c)/m
Therefore, from the graph, it can be deduced that absorbance is;
y = 2.1157x + 0.1031 and
R2 = 0.9974
Where y is the absorbance and x is the concentration and from the formula;
M1V1 = M2V2

Results
Table 3 (The average concentration of unknown ferratin samples A and B)
Discussion
In the experiment the diluted acid, sulphuric acid was added to the standardized solution contain
ferratin protein in order to break the ferratin compound and release iron ions of ferric ions.
Thereafter dihydroxyfumarate was added, as a redox agent, to the other solution containing freed
iron ions. The redox agent reduced ferric ions to ferrous ions.
The main aim of the experiment was to help in analysis and determination of the concentration of
iron ions that are present in a mole of aporeferritin. This would then help in calculation and
establishment of relationship between the absorption and concentration of ferratin. The
Absorbance
(A)
Unknown of Fe (II)
concentration - A
(multiply by 25)
Absorbance (B) Unknown of Fe
(II) concentration
– B
(multiply by 25)
0.398 3.48 0.415 3.68
0.404 3.55 0.392 3.41
0.412 3.65 0.381 3.28
Average of
A: 0.407
3.59 Average of B: 0.396 3.31

experiment was testing lamberts law which states that absorption of substances is directly
proportional to the amount of concentration.
Moles = Concentration * Volume
Therefore;
Moles of Fe2+ =3.31*10mL
= 33.1/1000
= 0.0331 moles
The standard deviation for results obtained from
the calibration curve, Sc;
Where N = is the number of pared data used,
m = the gradient of the curve,
Sr = the standard deviation about regression,
Yc = average of the values of the unknown,
N = Approximate calibration points
Standard Deviation about Regression;
= 1-0.9974
= 0.0026
The gradient of the curve, m;
= 2.1157
The average of the values of the unknown (Fe III) for sample A
= (3.48+3.55+3.65)/3
= 3.56 mL
The average concentration in Sample B,
= (0.415+0.392+0.381)/3
= 0.396 mL
Therefore,

Sc = (1-0.9974)/2.1157 ((1/4) +(1/7) + (3.59-3.31)^2/(2.1157^2))^1/2
=0.0041+0.6406
= 0.6448
This is the standard deviation and it lower than the coefficient of determination which is 0.9974.
The difference may have arisen in the entry of data caused by factors mentioned in the
conclusion part.
Conclusion
In the calibration graphs it can be clearly noted that diluted sample solutions with higher
concentrations of ferratin yield a higher absorption value. And therefore it is clear that by
converting ferric ions to ferrous ions it becomes possible to establish the concentration of iron in
ferratin protein There is a slight deviation of R2 from the its ideally value and from the
experiment it can be deduced that the small deviation shown might be due to stray light. This is
the ray that is striking the spectrophotometer detector whose wavelength is outside the spectra
band pass which may have not passed through the samples. In some instances, as it has been
discussed by researchers it may lead to curving of the graph otherwise the effect is relatively
negligible many a times. (Barton, 2005) In addition, R2 of the graph above may have deviated
and the error might be due; unequal path lengths across a light beam, unequal absorber
concentrations across the light beam and changes in the chemical equilibria as functions of
concentration.
The blue dots in the graph above represent the measured absorbance for each standard solution.
Also on the plot is the equation of concentration and absorbance and finally the R2 value which
usually is the measure of then degree of correlation between absorbance concentration and it is
0.9974 which means that it is not the perfect correlation since ideal case has a 1.000 correlation
coefficient. This deviation might be due to concentration prediction error being relatively
measurable. (Nall, p. 2017)
In addition, basing discussions upon the experiments that were done, it can be deduced that since
ferric ions from ferratin can be reduced to ferrous ions experimentally in the laboratory, this
knowledge can be of great help for research and innovations which can help in treatments
associated to blood and iron deficiencies with a conclusive and evidence-based understanding of
chemical properties of ferratin and its concentration levels. (Babu, 2017)

References
Babu, V., 2017. Evaluation and association of serum iron and ferritin levels in children with dental
carries. [Online]
Available at: https://www.ncbi.nlm.nih.gov/m/pubmed/28492187
[Accessed 7 january 2019].
Barton, L. L., 2005. Structural and Functional Relationships In Prokaryotes. 1 ed. New Mexico: Springer.
David Dunaief, J. D., 2014. Iron Induce Retinal Damage. In: A. Cwanger, ed. Nutrition, Diet and the Eye.
New York: RELX Group, pp. 619-26.
Gregory J. Anderson, G. D. M., 2012. Iron Physiology and Pathophysiology in Humans. 1 ed. London :
Springer.
Lauffer, R. B., 1992. Iron and Human Disease. 1st ed. London: CRC Press.
M. Worrall, M. W., 1991. Immunological Properties of Ferritin. European Journal of Haematology, 1(12),
pp. 223-228.
Nall, R., 207. Ferritin Levels Blood Test. [Online]
Available at: https://www.healthline.com
[Accessed 12 March 2019].
Templeton, D., 2002. Molecular and Cellular Iron Transport. 1 ed. London: CRC Press.
Worwood, M., 1979. CRC Critical Reviews. Reviews in Clinical Laboratory Sciences, 10(Serum Feritin), pp.
171-204.
Worwood, M., 1982. Ferritin in Human Tissues. Clinics in Haematology, 3(46), pp. 770-772.
Worwood, M., 1990. Ferritin. Blood Reviews, 69(12), pp. 259-269.