3. Gravimetric Determination of Iron as Fe2O3
3.
Quantitative Chemical Analysis
Gravimetric Determination of Iron as Fe2O315
A sample containing iron can be analyzed by precipitation of the hydrous
oxide from basic solution, followed by ignition to produce Fe2O3:
Fe3+ + (2 + x)H2O
Base
FeOOH . xH2O(s) + 3H+
Hydrous ferric oxide
Green Profile
See Section 0
900ºC
FeOOH . xH2O
Fe2O3(s)
The gelatinous hydrous oxide can occlude impurities. Therefore, the initial precipitate is
dissolved in acid and reprecipitated. Because the concentration of impurities is lower during the
second precipitation, occlusion is diminished. Solid unknowns can be prepared from reagent
ferrous ammonium sulfate or purchased from Thorn Smith.2
Procedure
1. Bring three porcelain crucibles and caps to constant mass by heating to redness for 15 min
over a burner (Figure 1). Cool for 30 min in a desiccator and weigh each crucible. Repeat
this procedure until successive weighings agree within 0.3 mg. Be sure that all oxidizable
substances on the entire surface of each crucible have burned off.
Figure 1. Positioning a crucible above a burner.
15. D. A. Skoog and D. M. West, Fundamentals of Analytical Chemistry, 3d ed. (New York: Holt, Rinehart and
Winston, 1976).
18
3. Gravimetric Determination of Iron as Fe2O3
Quantitative Chemical Analysis
2. Accurately weigh three samples of unknown containing enough Fe to produce ~0.3 g of
Fe2O3. Dissolve each sample in 10 mL of 3 M HCl (with heating, if necessary). If there are
insoluble impurities, filter through qualitative filter paper and wash the filter well with
distilled water. Add 5 mL of 6 M HNO3 to the filtrate, and boil for a few minutes to ensure
that all iron is oxidized to Fe(III).
3. Dilute the sample to 200 mL with distilled water and add 3 M ammonia16 with constant
stirring until the solution is basic (as determined with litmus paper or pH indicator paper).
Digest the precipitate by boiling for 5 min and allow the precipitate to settle.
4. Decant the supernatant liquid through coarse, ashless filter paper (Whatman 41 or Schleicher
and Schuell Black Ribbon, as in Figures 2-18 and 2-19 in the textbook). Do not pour liquid
higher than 1 cm from the top of the funnel. Proceed to step 5 if a reprecipitation is desired.
Wash the precipitate repeatedly with hot 1 wt% NH NO until little or no Cl- is detected in
4
3
the filtered supernate. (Test for Cl- by acidifying a few milliliters of filtrate with 1 mL of
dilute HNO and adding a few drops of 0.1 M AgNO . If precipitate is observed, Cl- is
3
3
present.) Finally, transfer the solid to the filter with the aid of a rubber policeman and more
hot liquid. Proceed to step 6 if a reprecipitation is not used.
5. Wash the gelatinous mass twice with 30 mL of boiling 1 wt% aqueous NH4NO3, decanting
the supernate through the filter. Then put the filter paper back into the beaker with the
precipitate, add 5 mL of 12 M HCl to dissolve the iron, and tear the filter paper into small
pieces with a glass rod. Add ammonia with stirring and reprecipitate the iron. Decant
through a funnel fitted with a fresh sheet of ashless filter paper. Wash the solid repeatedly
with hot 1 wt% NH4NO3 until little or no Cl- is detected in the filtered supernate. Then
transfer all the solid to the filter with the aid of a rubber policeman and more hot liquid.
6. Allow the filter to drain overnight, if possible, protected from dust. Carefully lift the paper
out of the funnel, fold it (Figure 2), and transfer it to a porcelain crucible that has been
brought to constant mass.
16. Basic reagents should not be stored in glass bottles because they slowly dissolve glass. If ammonia from a
glass bottle is used, it may contain silica particles and should be freshly filtered.
19
3. Gravimetric Determination of Iron as Fe2O3
1. Flatten
paper
2. Fold in
edges
Quantitative Chemical Analysis
Figure 2. Folding filter
paper and placing it inside
a crucible for ignition.
Continue folding paper so
entire package fits at the
bottom of the crucible. Be
careful not to puncture the
paper.
3. Fold over
top
4. Place inside crucible
with point pushed
against bottom
7. Dry the crucible cautiously with a small flame, as shown in Figure 1. The flame should be
directed at the top of the container, and the lid should be off. Avoid spattering. After the
filter paper and precipitate are dry, char the filter paper by increasing the flame temperature.
The crucible should have free access to air to avoid reduction of iron by carbon. (The lid
should be kept handy to smother the crucible if the paper inflames.) Any carbon left on the
crucible or lid should be burned away by directing the burner flame at it. Use tongs to
manipulate the crucible. Finally, ignite the product for 15 min with the full heat of the
burner directed at the base of the crucible where the Fe2O3 is located.
8. Cool the crucible briefly in air and then in a desiccator for 30 min. Weigh the crucible and
the lid, reignite, and bring to constant mass (within 0.3 mg) with repeated heatings.
9. Calculate the weight percent of iron in each sample, the average, the standard deviation, and
the relative standard deviation (s/ x ).
20
CHEM 321
Gravimetric Determination of Iron as Fe2O3
Report
Name:
Lab Partner:
Experiment Date:
Report Due Date:
*When typing your report, please 1.5 or 2.0 space your text. Remember to submit your report to
Safe Assign and turn in a hard copy to your instructor. The hard copy may be printed on both
sides of the paper.
*When you write a report, please remove all questions/text that we have included to guide your
responses. You may leave the questions at the end in the report. You may need to reformat your
tables to fit your data. Do not treat this as a “fill in” worksheet – it should appear like a formal
lab report when printed out.
Abstract
In one paragraph, summarize the experiment, why you did it, and the results. Be specific and
include chemical equations, unknown identifiers (letter/number), results, uncertainties, RSDs,
confidence intervals.
I. Data and Results
Table 1. Data and calculations for the determination of iron as Fe2O3 from Unknown ___
Sample
(g) (± ?)
Crucible
(empty, g)
(± ?)
Crucible
(full, g)
(± ?)
Fe2O3 (g)
(± ?)
Fe2O3 (mol)
(± ?)
Fe (mol)
(± ?)
Fe (g)
(± ?)
Fe wt %
(± ?)
Discuss the data in Table 1 in terms of the reaction taking place, the calculations, results,
experimental difficulties, and any other pertinent information that you recorded in your lab
notebook.
II. Results and Conclusions
Briefly summarize your results and describe your conclusions. This section should include
everything that appears in your abstract as well as a thoughtful analysis of the experiment.
III. Calculations
You may hand-write or scan your notebook calculations.
1
IV. References
Please format properly for a scientific report (Journal of the American Chemical Society
format). By now you should be using Zotero (available from GMU library, www.zotero.org). If
you need assistance, please see the science librarian.
2
Quantitative Chemical Analysis
Chemistry 321
Gravimetric Determination of Iron as Fe2O3
Individual Data Report: 25 points
Points
1
1
3
Report
Followed sample report format and typed answers.
Data tables typed and complete
Abstract complete
1
1
5
5
2
2
Data and Results
Table 1: Uncertainties for all data columns reported
Table 1: Mass recorded to four decimal places
Table 1: Calculations correct (must match what instructor gets)
Table 1: Error propagation correct
Discussion/Table 1: Balanced reaction included
Discussion/Table 1: Includes observations, results, difficulties, etc.
2
Results/Conclusion: Detailed, thorough, thoughtful
2
References (JACS format)
Gravimetric Determination of Iron as Fe2O3
Class Data Report: 25 points
Points
2
2
5
2
2
2
2
2
4
2
Report
Abstract complete
-3 if data tables are not typed or are incomplete
Statistical Analysis
Table 1: Uncertainties reported, masses reported to four decimal places
Table 1: Calculations correct (must match instructor)
Tables must be described and explained – not just pasted into report.
Table 2: Residuals
Table 3: Descriptive statistics
Figure 1: Residual plot
Figure 2: Residual histogram
Table 4: t-test and comparison of unknowns
Conclusion: Detailed, thorough, thoughtful
References (JACS format)
[Type here]
Chem 321: Titration and Gravimetric Analysis of Chloride Unknown
Experiment date February 22, 2018
Name
Abstract
Titration is a common method that is used to determine the concentration of an unknown
substance in a solution by adding a measured amount of known concentration of a standard
solution. The known concentration solution (reactant) reacts with unknown concentration solution
(analyte) give an indicative result. The concentration of unknown can be calculated by knowing
the stoichiometry of the reaction and the molarity of the standard solution. This experiment is a
medium to determine the concentration of potassium chloride and unknown soluble chloride that
were titrated with silver nitrate and another regent which was used as an indicator. Six titrations
were performed. In the first three titrations, a known amount of potassium chloride samples titrated
with 0.05 M silver nitrate solution to calculate Molarity of silver nitrate. The last three titrations,
unknown soluble chloride samples titrated with 0.05 M silver nitrate, then the percentage of
chloride was calculated using the average of molarity of silver nitrate as determined in the first
three titrations.
Table 1. Standardization of silver Nitrate AgNO3
[Type here]
Titration
KCl
KCl
AgNO3
V initial
V final
V total
Molarity
(g)
(mol)
(mol)
(mL)
(mL)
(mL)
AgNO3
(mol/L)
1
0.0789
0.0010583316
0.0010583316
0.01
31.05
31.04
0.0341
2
0.0783
0.0010502835
0.0010502835
0.05
24.01
23.96
0.0438
3
0.0775
0.00010395526 0.00010395526
0.21
28.08
27.87
0.0373
Average = 0.0384
StdDev= 0.0049
% RSD = 12.8%
95% Cl =
In this experiment, the silver nitrate was standardized initially. This was an important step
to protect the solution from the light and to prevent its decomposition to elemental silver. Then
using a burette measured quantities of silver nitrate solution added to potassium chloride solution
with a few drops of dichlorofluorescein indicator. Dextrin was also added to the solution. In the
process of titration, the solution turned to milky white when silver chloride percipient began to
form. After all the chloride ions were precipitated, the solution turned to pink because of the
reaction between excess of silver ions and dichlorofluorescein and thus singled the endpoint. The
use of dextrin was to prevent coagulation of the silver chloride particles.
Initially, this experiment worked as was expected, the solution turned to white as silver
chloride began to form, them to light pink which was the indication of the endpoint. However,
[Type here]
experimental errors were obviously made. The errors were observed initially during the lab
experiment by seeing a major difference of the volumes of silver nitrate that were delivered to the
anlyate solution. The volumes were 31.05ml, 24.01ml, and 28.08ml which confirmed the
possibility of errors. After, the relative standard deviation was calculated and obtained to be of
12.8% which is far greater than the %RSD of 0.5%. The errors possibly arose because of improper
preparation of the standard silver nitrate or most likely because of poor titration technique.
Table 2 Volumetric determination of chloride Weight % for unknown
Titrations Unknown
(g)
V initial
V final
V total
Molarity AgNO3
(mL)
(mL)
(mL)
AgNO3
= Cl-
(mol/L)
(mol)
Cl-
Weight
(g)
Prevent
1
0.1216
0.10
34.01
33.91
0.0384
0.0013
0.0462
37.9%
2
0.1160
0.20
30.10
30.10
0.0384
0.0012
0.0425
36.6%
3
0.1189
0.10
30.00
30.00
0.0384
0.0012
0.0425
35.7%
Average = 36.7
StdDev= 1.1068
% RSD = 3.0%
95% Cl
In the second part of the experiment, three samples of soluble chloride titrated with
standard silver nitrate solution following the same procedures that were used in part one. However,
this part was done to discover the percentage mass of chloride in the unknown salt using the
average molarity that was obtained in the first part of the experiment. The average percentage mass
[Type here]
determined to be 36.7% and the relative standard deviation was 3.0 % which was still greater than
the %RSD of 0.5%. However, comparing the RSD% of the first part of the experiment of 12.8%
to the RSD% of the second part of the experiment of 3.0% shows that the second experiment was
a little more accurate than the first one. It indicates a change in the knowledge of the experimental
procedure in the second experiment.
Conclusion
The molarity of the standard solution and the amounts of the solution titrated are the
indicators for finding the molarity and the identity of the unknown solution. There is a possibility
of error originates from the titration process and the preparation of the solution. Since the
percentage error reduced from the first to second experiment, there is a high possibility that the
error was in the titration process for which the second experiment was done with more accuracy
compared to the first one.
Questions
1.
Identify three factors that must be considered in choosing a proper adsorption
indicator for precipitation titration. Explain.
A proper adsorption indicator should have a relatively medium pH. This helps the
indicator to maintain its anionic form.
Particle size of the precipitate matters as they precipitate should not be allowed to form
large particles since they will settle. Small particle surface of course provides a larger
surface area.
[Type here]
The charge of the indicator that of the precipitate should be at equivalence point. If the
indictor is strong enough, it is likely to displace main adsorbent ion right before
equivalence point is achieved.
2.
What would happen if dextrin were not used in the titration? Explain.
The precipitate will coagulate since dextrin plays the role of protective colloid that
ensures that the precipitate is highly dispersed.
REFERENCES:
Vogel, A. I. (1939). Text book of quantitative inorganic analysis. Longmans, Green: London.
Erdey, L. (2013). Gravimetric Analysis: International Series of Monographs on Analytical
Chemistry (Vol. 7). Elsevier.
This experiment is done to determine the percentage of iron in an unknown sample. for the
determination we have used the precipitate technique in the part gravimetric analysis is done
by using filtration method
Procedure (exactly what we did in the lab)
We weight enough of unknown to produce 0.3 g of Fe2O3 (only one sample). 20 mL Of HCl was
added to the sample in a 400 mL beaker to dissolve the Iron and the sample was filtered to
remove any insoluble. 5 mL of 6 M nitric acid was added and was boiled for a few minutes then
diluted to 200 mL. Then 3 M ammonia was added until the solutions tested basic. The solution
was boiled for five minutes and the supernatant was decanted through ashless filter paper and
washed with warm 1% ammonia nitrate. 1 mL aliquot of filtrate was acidified using 0.1 M nitric
acid and them a few drops of Silver Nitrate were added. If the solution appeared cloudy the
precipitated was washed with more warm 1% ammonia nitrart and this was repeated until the
filtrate tested non cloudy. . the funnel was lifted to drain overnight until next session. After one
week (next lab) the filter paper was lifted out of the funnel and transferred carefully to
crucible and weighted then baked at 900 C ( to ensure the filter paper have current completely
to get the precipitate. After one week after baked and cooled the crucible with filtrate
weighted.
Some notes
Our result for this experiment was precise but not as accurate and we were able to recover 70%
of the iron that we believed was present
The error in this experiment probably was in transferring the product the crucible ( We did not
know the proper technique for handling crucible)
or most likely during the process of filtering the product through the ashes filter
Sample
Crucible
Crucible
Fe2O3
(empty,
(full,g)
(g)
20.9442
0.3056
FeO3 ( mol) Fe (mol)
Fe(g)
Fe wt %
0.00191372
0.2147
70%
g)
0.3006
20.6386
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