Determination of Iron in a Hematite Ore

Abstract

This method development lab explored the redox titration of potassium dichromate with an iron
oxide ore dissolved in a concentrated nitric acid medium. While this experiment differed by the use of
only the acid, it became well apparent that the chemistry of the nitric acid prevents the adequate
determination of iron in the hematite ore. The use of nitric acid in the dichromate method of iron
determination results in a significant difference between the same method using concentrated
hydrochloric acid.

Introduction

     The determination of iron in oxide ores is determined by either precipitation of hydrous ferric oxide
and ignition to Fe₂O₃ or by the reduction of dissolved Fe(III) to Fe(II) followed by the titration with a
standard oxidizing agent such as the dichromate ion, Cr₂O₇^2-(Rice, 27). This lab will utilize the swift
and accurate properties of the dichromate method in order to analyze the iron ore.

     In the original dichromate method, hot concentrated hydrochloric acid dissolves iron in a hematite
ore according to the following reaction:
Fe₂O₃(s) + 6 HCl(aq) ---> 2 FeCl₃(aq) + 3 H₂O   (1)

Lack of stable reducing agents suitable for the direct titration of the ferric ion in solution requires the reduction to the ferrous ion. Stannous Chloride (namely the stannous ion, Sn(II)) is utilized to reduce
the iron prior to titration. The iron from hematite is reduced from Fe(III) to Fe(II) upon forming the
halogenated iron complex. This reaction is described as follows:

Sn²⁺ + 2 Fe³⁺ ---> 2 Fe²⁺ + Sn⁴⁺ (2)

Excess Sn²⁺ in solution is removed by adding mercuric chloride producing insoluble Hg₂Cl₂:

Sn²⁺ + 2 HgCl₂(aq) ---> Sn⁴⁺ + Hg₂Cl₂(s) + 2 Cl^- (3)

The ferrous ion can now be titrated with the standard potassium dichromate solution to yield the
following net ionic equation:

6 Fe²⁺ + Cr₂O₇^2- + 14 H⁺ ---> 6 Fe³⁺ + 2 Cr³⁺ + 7 H₂O (4)

An inside indicator, sodium diphenylamine sulfonate, is oxidized by the dichromate ion to produce a
purple product. This reaction, however, does not occur as quickly as that of the oxidation of ferrous ion (Rice, 27-28).

In designing this experiment to determine the iron in the oxide ore, the role that the hydrochloric acid plays in the reaction was looked upon with curiosity. In light of research and chemical intuition, it was
found that the role of excess hydrochloric acid is disputed by some chemists. Dr. James Carr from the
University of Nebraska believes that if the HCl method is used, the HCl must be removed so that the
Cr₂O₇^2- or MnO₄^1- (depending on method) ions will not react with the chloride during the titration
(Carr, internet source). Carr suggests removing the HCl by adding concentrated H₂SO₄ and heating
until dense white fumes of HCl come off. Meanwhile, authors Skoog and West suggest that moderate
amounts of hydrochloric acid do not affect the accuracy of the titration (Skoog and West, 125).

Recognizing the aforementioned arguments, we thought it would be interesting to introduce a new
ion into the picture. In much of the initial research, very little was noted about nitric acid's role in iron
ore determination. We then decided that an interesting element that could be introduced to the
experiment would be the use of nitric acid instead of the typical hydrochloric acid. Using an
electrochemical approach to the question, the potential for nitric acid was calculated to be +0.79 V or
+0.96 V, depending on which nitrogen oxide gas is liberated (please refer to Results and Discussion for
the derivation of these values; equations 5 and 6). Comparing this potential to that of hydrochloric acid
(-1.36 V), it is very noticeable which acid dominates as the oxidizing agent as the potential of nitric acid
surpasses that of hydrochloric acid. Recognizing the very strong oxidizing power of the acid, the
hypothesis that was formulated stated that there would be a significant difference between the
dichromate method involving hydrochloric acid and the same method involving nitric acid.

Materials and Methods

To adequately test our hypothesis, the following method is followed. A 0.1240 N potassium
dichromate solution was prepared by measuring 1.2160 grams of K₂Cr₂O₇ and transferring to a 200 mL volumetric flask. Four subsequent iron samples were weighed analytically to approximately 0.4 to 0.5 grams following being dried for one hour a 110 degrees Celsius. Each iron ore was then treated with 10 mL of concentrated nitric acid, covered with a watch glass, and allowed to heat gently for twenty to thirty minutes. After thirty minutes, red particles of undissolved iron oxide remained, so 0.5 M stannous chloride was added gradually using a medicine dropper, when the solution turned yellowish instead of light straw yellow (signifying to us the partial reduction of the ferric ion). Condensate from the watch glass was rinsed into the beaker as well as the inner walls of the beaker.

Treating one sample at a time, 0.5 M stannous chloride solution was added drop by drop from a
medicine dropper to the hot solution, but instead of becoming colorless, the solution would turn black
and then back to the yellow color. After considerable amounts of stannous chloride was added, it was
apparent that the solution would not decolorize from the yellow color. Recognizing the chemistry that
was taking place, it was agreed that the experiment would be continued. Approximately 60 mL of cold
water was added and prior to titration 10 mL of saturated mercuric chloride is added to remove stannous chloride. After waiting three to four minutes, a white precipitate of mercurous chloride did not appear. Subsequent steps including the addition of 10 mL 1:1 (V:V) phosphoric acid and 0.3 mL of 0.01 M sodium diphenylamine sulfonate indicator were taken. Each sample was then titrated with the standard potassium dichromate. No endpoint was attained in any sample. Instead, as the potassium dichromate was added, the solution became more and more yellow (resembling the color of K₂Cr₂O₇). No traces of purple color were observed in any of the titrations.

Once we encountered difficulty in dissolving the iron in the nitric acid and observing the black color
that occurred when adding stannous chloride, it was agreed that an interesting question may lie with the role that the acid aqua regia (3:1 hydrochloric and nitric acids) may play in the chemistry. The aqua regia was prepared and then 10 mL was mixed with another sample of iron ore. As was expected, the
ore dissolved readily. Upon dissolution, stannous chloride was added to reduce the ferric ions and again the sample formed a dark "cloud" that disappeared after swirling the beaker; this was identical to the
pure nitric acid samples above. When this sample was treated with water, phosphoric acid and
indicator, it also did not approach an endpoint when titrated with potassium dichromate.

Results and Discussion

The following data was recorded during the experiment:

**Weight of Iron Ore Samples**
Sample 1	0.4554 grams
Sample 2	0.4290 grams
Sample 3	0.4414 grams
Sample 4	0.4186 grams
Sample 5*	0.4592 grams*

* Treated with Aqua Regia

**Standardization of Potassium Dichromate**
1.2160 grams K₂Cr₂O₇	0.1240 N K₂Cr₂O₇

**Titration with Potassium Dichromate**
Sample 1	50.00 **
Sample 2	50.00 **
Sample 3	50.00 **
Sample 4	***
Sample 5*	50.00 **

                                       * = treated with aqua regia
                                       ** = no endpoint was reached; solution became yellow-
                                               orange; no purple detected during titration
                                       *** = no dichromate solution remained to titrate this
                                                 sample; aqua regia was titrated to explore the
                                                  possibility of reaching an endpoint

The preceding chart describing the titration of the sample conveys that there was definitely problems encountered by utilizing the nitric acid (and even the aqua regia which wasn't the main attention of this experiment). The results in themselves tell that this experiment failed, especially when an endpoint could not be obtained and the percentage iron consequently could not be calculated. However, the observations that were noted during the experiment tell us a great deal of what chemistry prevailed upon the use of nitric acid.

Firstly, when the nitric acid was added to the ore, it was apparent right away when red gas (namely
NO₂) was liberated and the remaining solution consisted of a heterogeneous mixture of undissolved iron and nitric acid that this experiment may run into difficulty. Secondly, after thirty minutes of continuous
heat, 0.5 M stannous chloride was added to the solution and the black cloud was encountered, which
disappeared quickly upon swirling the solution. This made us believe that something must be happening
to the added stannous ion.

To explain what went on in this experiment, our first goal was to explore if others who have tried this experiment in the same fashion experienced the same chemistry. Indeed, research prevailed and we found an array of professors who comment about nitric acid's role in the experiment. For example,
Lucio Gelmini, an apparent chemist, asks his chemistry education discussion list on the internet whether
or not anyone has a redox titration where dichromate or permanganate is used to titrate a sample of iron ore. In response, a professor responds with the following:

If memory serves me right, that is an analysis we did in my Analytical Chemistry class that I taught several years ago. 1. Dissolve the iron ore with a suitable acid such as nitric acid. 2. Reduce the Iron III cation to Iron II using Tin (II) chloride. 3. Oxidize the Iron II back to Iron III using a standardized solution of Potassium Dichromate or Permanganate (Gelmini, internet source).

Immediately, colleagues respond with suggestions not to use nitric acid. Two such follow:

A previous reply suggested dissolution of the sample in HNO₃, with the                         attendant problem of residual nitric acid cause an oxidation of the permanganate/
                        dichromate titrant or the pre-reduced iron. Concentrated HCl is a GOOD
                        solvent for iron oxide samples--but in excess chloride ion also reacts with
                        permanganate/dichromate. Rather than distilling either the excess HCl or
                        excess HNO₃ from the solution by concentrating a sulfuric acid solution to
                        formation of SO₃ (and all that entails)--addition of Zimmerman-Reinhardt
                        reagent to a solution in HCl inhibits the oxidation of chloride by the
                        permanganate/dichromate (Wiley, internet source).

                        Neither is it a good idea to have nitrate ion present as nitrate is a good
                        enough oxidizing agent that it will oxidize iron from Fe(II) to Fe(III)
                        during the titration. It might also oxidize Mn(II) back up to some higher
                        state. HNO₃ can be got rid of by the same technique with sulfuric
                        acid (Carr, internet source).

     From the above data, the question that our null hypothesis posed has been answered experimentally. However, in order to fully explain the failure of this lab, chemistry will play a role in describing what the nitric acid had to do with the experiment. The following suggestions have been derived. Electrochemistry can be used in many situations to describe situations in which oxidation and reduction are taking place. In the normal experiment, hydrochloric acid is used to dissolve the iron. The following two half reactions can be written for hydrochloric acid and nitric acid, respectively.

(reduction half-reaction) 2 H⁺ + 2 e^- ---> H₂ E_red = 0 V (oxidation half-reaction) Cl₂ + 2 e^- ---> 2 Cl^- E_ox = 1.36 V __________________________________________________
(overall) H+ + Cl- ---> (1/2) H2 + (1/2) Cl2 E = -1.36 V (5)

NO₃^-+ 2 H⁺ + e^- ---> NO₂ + H₂O E = + 0.79 V (6) OR NO₃^- + 4 H⁺ + 3e^- ---> NO + 2 H₂O E = + 0.96 V

From the above, the potential for hydrochloric acid is -1.36 V and that for nitric acid is +0.79 V or
+0.96 V, depending on what reaction prevailed with respect to the liberated nitrogen oxide. The more
positive the potential, the more powerful the substance is as an oxidizing agent. Therefore, it is apparent that nitric acid is a powerful oxidizing agent. This would explain that the nitric acid (namely the nitrate ion) would interfere with the dichromate ion in oxidizing the ferric ion to the ferrous ion during titration:

6 Fe²⁺ + Cr₂O₇^2- + 14 H⁺ ---> 6 Fe³⁺ + 2 Cr³⁺ + 7 H₂O (7)

Once again, one can refer to the potential of the dichromate ion in an acid environment in which
equation 8 below can be written:

Cr₂O₇^2- (aq) + 14 H⁺(aq) + 6 e^----> 2 Cr³⁺(aq) + 7 H₂O (l) E = + 1.33 (8)

This preceding potential reaction contains a potential greater than that of nitric acid, but if one compares it to hydrochloric acid, it is no wonder why hydrochloric acid is used as its potential does not compete
with dichromate's potential as an oxidizing agent.

Knowing what we have found above, we can then summarize the chemistry that occurred during the experimentation. With the hematite and the nitric acid being mixed together, it was evident that no iron went into solution because the only chemical observation that was noticed was the liberation of some sort of nitrogen oxide, whether it be NO, NO₂, etc. This part of the explanation, we believe, is not important because we were not able to get the iron to dissolve, so spending time trying to derive the
products from this reaction would be unimportant in the questions and discussions proposed by the
hypothesis of this lab.

When Fe³⁺ is reduced to Fe²⁺ using the stannous ion, the reaction of equation 4 above predominates.

For the stannous ion to go to the stannic ion, the following oxidation occurs:

Sn²⁺ ---> Sn⁴⁺ + 2 e^- E = -0.154 V (9)

Equation nine (9) is very fundamental to our argument because the negative sign reveals that this is
indeed a reducing agent. But when this is in a powerful oxidizing agent, such as nitric acid (E = +0.79
V), the reduction of the ferric ion will be almost impossible, as was in our experiment.

A final question lies with the existence of the black "cloud" that was formed when adding the
stannous chloride. Our first inclination at describing this phenomena was the oxidation of the stannous
ion to the stannic ion, Sn⁴⁺. Any metal in an oxidized form displays a dark color (Treadwell and Hall,
167). When HgCl₂ was added to the solution after adding stannous chloride, a white precipitate was
anticipated but did not prevail. Nor did black metallic mercury (Hg⁰) precipitate prompting too much
stannous chloride was added to the solution (Christian, 709), as in equation ten below.

Sn²⁺ + Hg₂Cl₂ ---> Sn⁴⁺ + 2 Hg⁰ + 2 Cl^- (10)

The most probable explanation of what happened when the stannous chloride was added to the sample
would be that the stannous ion is oxidized to the stannic ion by forming the oxide, SnO₂. When this was proposed by a group member, research prevailed and indeed stannous oxide is a black color (Treadwell
and Hall, 167).

Conclusions

Before this method development lab began, as a group we all felt that the use of nitric acid would
fail. Even though our null hypothesis conveyed this belief, we had no idea that the role that the nitric
acid would play could become so complex. Anticipating its failure, we were assured that the role of
nitric acid in this experiment would be an interesting one which would likely require a thorough
knowledge of this experiment as well as nitric acid chemistry. As shown in this report, one of nitric
acid's most profound characteristics is its strong oxidizing ability. In this lab, however, it is preferred to
use an acid that does not possess this characteristic (e.g.. HCl), as the reducing agent stannous chloride will be later used to reduce the ferric ion to ferrous ion. The role of nitric acid's oxidizing power
prevented the reduction of Fe³⁺ to Fe²⁺ and most likely interfered in the titration with dichromate, even
though the titration was "doomed" from the start. The failure of reaching an endpoint restricted us from performing the needed calculations required to determine the percent difference of iron. Obviously, the use of nitric acid in this method is not recommended and in our case, it was impossible to reach the
required point for iron determination. The use of nitric acid in place of hydrochloric acid results in a
significant difference of percentage iron determination in an iron oxide ore.

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