Name of Student:
Date performed:
Section:
Date Submitted:
Group Number: SOME INVESTIGATIONS ON THE CORROSION OF IRON I.
Introduction
Electricity is one important product of the reactions in voltaic cells or batteries. However, this reactions also leads to unfavorable outcomes. The electrochemical basis of this reaction underlies the process of corrosion (Petrucci et al, 2011). The process of corrosion is one view of returning metals to its natural state from which they were originally obtained in ores (Zumdahl et al, 2010). Corrosion is defined as the deterioration of metals due to the electrochemical process it undergoes. In our surroundings, many manifestations of corrosion is taking place. Some of these are the formation of green patina on copper and brass, tarnishing of silver, and the most common is the rusting of iron. Corrosion is responsible in the immense damages in cars, buildings, bridges, ships, and other metal-involved infrastructures and equipment. (Chang et al, 2013). It involves oxidation or conversion of uncombined metals into oxides or other compounds (Torio et al, 2012). The corrosion of metals except gold is so because their standard reduction potentials is less positive compared to that of oxygen gas. Thus, when their half-reactions is reversed and combined with the half-reaction of oxygen, the total charge (Ecello) would be positive, which says that the reaction is spontaneous. However, the positive value does not tell how fast the oxidation of metals by oxygen will occur (Zumdahl et al, 2010). Since our experiment focuses mainly to rusting of iron, let’s take it first into consideration. The reactions involved in corrosion of iron is held quite complex and is not completely understood yet. However, main steps are believed to be as what really occurs in the reaction. In order for rust to form, oxygen gas and water must be present on iron. Oxidation occurs at a certain region of the metal’s surface which serves as anode: Fe(s)
Fe2+(aq)
+
2e-
+0.44 V
The charge of the reaction is positive since we reverse the reaction. In another region of the same metal’s surface, this electron product reduces atmospheric oxygen forming water at the cathode. O2 (g)
4H+(aq)
+
4e-
+
2H2O(l)
+1.23 V
The overall redox reaction is 2Fe(s)
+
O2 (g)
+
4H+(aq)
2Fe2+(aq)
Computing for the standard electromotive force (E ocell): Eocell
=
Eocathode
-
Eoanode
+
2H2O(l)
= 1.23 V
-
(-0.44 V)
=1.67 V By convention, the Eoanode would still be in negative sign since the reaction in the anode is oxidation but the formula has already factored the negative sign. Since H+ ions is supplied in part due to the reaction of atmospheric CO 2 with water forming H2CO3, this reaction occurs in an acidic medium. Further oxidation of Fe2+ by oxygen occurs again in the anode through the reaction: 4Fe2+(aq)
+
O2 (g)
+
(4+2x)H2O (l)
2Fe2O3 ∙ xH2O (s)
+8H+(aq)
Rust A product of the reaction which is a hydrated form of Iron (III) oxides is the one known as the rust. We represent the formula as Fe 2O3 ∙ xH2O since the amount of H2O associated with the Iron (III) oxide varies. Through migration of electrons and ions, electrical circuit is completed (Chang et al, 2013). Varying color of the rust from black to yellow to the familiar reddish brown is affected by the degree of hydration of the iron oxide (Zumdahl et al, 2010). Other than iron, other metals also undergo corrosion. Like for example aluminum, a metal used in building airplanes and making beverage cans. Since it has a more negative standard reduction potential compared to Fe, it has a greater tendency to oxidize than iron does. However, we do not see airplanes dissolving and corroding in rainstorms. This is so because the layer of insoluble Al2O3 that forms on its surface after it has been exposed to air shields the aluminum underneath from further corrosion. In the case of iron, rust is too porous to protect the iron underneath. Metals like copper and silver also corrode much more slowly. After being exposed in atmosphere, copper forms a verdigris which is basic Copper carbonate (CuCO 3), a green substance also called as patina, which protects metal underneath from further corrosion. Similarly, silverwares, after coming to contact with foods, develop a film of black silver sulfide (Ag2S) layer. In this particular experiment, the class stated the effects of three factors on the corrosion of iron. These factors are; the acidity or basicity of the solution in contact with the metal, mechanical stress applied on the metal and lastly, its contact with other metals. Objectives II.
Materials A. Reagents Part A 12 small bright nail NaOH NaCl
HCl KOH KNO3 HNO3 Na3PO4 NaSCN H2SO4 dH2O 12 - 1 to 2 drops 0.1 M K3Fe(CN)6 1 mL FeSO4 Part B and C 1 drop of 0.1 M K3Fe(CN)6 200 mL dH2O 4g agar 10 drops K3Fe(CN)6 2 drops phenolphthalein 4 big and bright nails Copper strip Zinc strip
B. Apparatus/Other Materials 12 centrifuge tubes 1 test tube Phenolphthalein Hot plate
3 -250 mL beaker Stirring rod Pliers Sand paper
III.
Procedure A. Reaction of Iron with various Aqueous Solutions To carry out the experiment, small nails was polished first with sandpaper in order to remove the rust it already has. Each nails was putted in each of the five centrifuge tubes. They were carefully slid down the bottom to avoid breaking of tubes. Each tubes was filled with solutions categorized in three sets A, B, and C. Our group was assigned with Set A containing NaOH, NaCl, HCl, and distilled H 2O solutions. The acidity of each solution was determined using phenolphthalein. Then, the nails was allowed to stand overnight in the solutions. Observations on changes that took place was recorded on Table _____. After standing overnight, one to two drops of 0.1 M potassium ferricyanide (K 3Fe(CN)6) solution was added into each test tube. The potassium ferricyanide dissociated to form K + and [Fe(CN)6]3-. Changes in the solution was observed and noted. In order to give descriptions to K3Fe(CN)6 solution, together with FeSO4 solution, a drop of 0.1 M K3Fe(CN)6 was added to 1mL FeSO4 solution placed in a test tube. Visible observations of the reaction was noted in Table ______.
Preparation of Agar Medium
About 200 mL of distilled water was heated to a gentle boil to be used in making agar medium. The heated water was removed from hot plate and was added with 4 g agar. After addition, the mixture was stirred and continued to be heated until the agar is spread all throughout. Then, 10 drops of K3Fe(CN)6 and 2 drops of phenolphthalein was added to the agar mixture. The mixture was allowed to cool until lukewarm. B. Reactions of Iron as affected by Mechanical Stress Two large nails were refurbished to be used in the experiment. One nail was placed in the beaker while the other was bent with a pair of pliers before putted in the same beaker but not allowing the two nails to touch each other. The cooled agar mixture was carefully poured into the beaker until the nails were covered to a depth of about 0.5 cm. Then, the beaker was allowed to stand overnight. Changes on the nails were observed and recorded in Table ______. C. Reactions Involving Metal Couples – Two Metals in Contact Same is done in having the large nails polished beforehand in this next experiment. A clean piece of copper strip was winded around the clean iron nail. Then the coil was allowed to be tightened by removing the nail and tightening the copper coil. Thus, when the nail was reinserted it made a tight contact with the coil. Same procedure was done using a zinc strip instead of copper on another nail. These nails together with its constricting coil were placed in the beaker. However, the nails was ensured not to touch each other. Then, the lukewarm agar mixture was poured into the beaker as what was done before in part B. The beaker was
allowed to stand overnight until changes is observable. Changes was recorded in Table ____. IV.
Data/Observations Table 12.1 Observations of the reaction of Iron with various aqueous sol’n after standing overnight and upon addition of K 3Fe(CN)6 tA OH Cl
Se
Cl
H
H 2O
d
tB OH
Se
Na Na
K K
NO3
After standing overnight No corrosion observed Brown ppt formed; nail corroded Brown ppt formed; dull texture; corroded tip Brown ppt formed; concentrated rusting on the tip After standing overnight No corrosion observed S l i g h t
Upon addition of K3Fe(CN)6 turned yellow green turned yellow
pH
Basic
Neutral
Turned blue
Acidic
turned yellow
neutral
pH
Basic
Neutral
Upon addition of K3Fe(CN)6 turned Yellow green Turned yellow
c o r r o s i o n
H
NO3 H 2O
d
Red brown ppt formed; dull color of nail Brown ppt formed;
Turned blue
Acidic
Turned yellow
neutral
Se
tC
Na
3
PO4
SCN
H2
SO4 H 2O
Na
concentrated rusting on the tip After standing overnight White ppt (error); cloudy mixture Brown ppt formed Gas evolved;
Upon addition of K3Fe(CN)6 Turned Yellow green
pH
basic
Turned Yellow
neutral
Turned blue
Acidic
d
Brown ppt Turned yellow neutral formed; concentrated rusting on the tip Table 12.2 Observations of the reaction of K3Fe(CN)6 with FeSO4
KFeSO4 Reaction b/w K3Fe(CN)6 Fe(CN) and FeSO 3 6 4 Light Rusty Bluish green with blue bright yellow colored, orange , precipitate green sol’n with rust Table 12.3 Observations of the reaction of Iron after standing overnight in agar set-ups
V.
Agar Set-ups Straight nail
Bent nail
Nail coiled with Cu strips
Nail coiled with Zn strips
Changes overnight Rusting is spreaded; pink color body of nail; dark blue tip and head Prominent rust on bent part; dark blue on tip, head, and bent portion; pink is the rest of the nail Rusting on the nail; pink color following traces on Cu strip, dark blue rust on nail Rusting on the Zn strips; pink nail; faint white color on Zinc
Discussion Based on our three objectives, the results were derive from the experiment. The tables tell the factors that could affect the corrosion of iron. The first factor to be tested in the experiment is the acidity and basicity of the solution in contact with the metal which is iron. From the results of the experiment found in Table ___, it shows that basic solutions have induced no observable sign of corrosion on the nail. These solutions are NaOH and KOH. Since the dissociation of base has a product OH -, O2 (g) + 2H2O (l) + 4e 4OH-(aq)
increasing its concentration hinders the reduction reaction above of the oxygen (explained by Le Chatelier’s Principle) thus, blocking the corrosion. However, in the acidic solutions, the results from Table ___ shows that the nail corroded varying on the acid used. The HCl reacted to form brown rust while the HNO3 reacted to form red brown rust, and lastly, H2SO4 producing orange brown rust. The acids’ dissociation forms H + ions and it is a reactant in the overall redox reaction below. 2Fe(s) + O2 (g) + 4H+(aq) 2Fe2+(aq) + 2H2O(l) Thus, we increase the concentration of H+ ions shifting the redox forward to form Fe2+ ions and water. The concentration of H + in the reaction:
4Fe2+(aq)
+
O2 (g)
+
(4+2x)H2O (l)
2Fe2O3 ∙ xH2O (s)
+8H+(aq)
is not increased since the this reaction occurs in the anode and the overall redox is in the cathode. In condition wherein salt solution is used, the reactions according to Table ___, shows that the iron nail generally corroded except Na 3PO4 which happen to be a basic salt. These salt solutions are NaCl, KNO 3, and NaSCN. This salts dissociates readily into its ions. This ions helps facilitate the corrosion process by providing electrons with a means of transportation between the anodic and cathodic regions of the iron nail. Thus, salts serves somehow as catalyst of the reaction. Water is a requisite in rusting, thus it serves as the control to show which reactions corroded most and which did not. Upon addition of K3Fe(CN)6 to each mixture with nail Table ___ suggest; basic mixtures just retained the color of the K 3Fe(CN)6 which is yellow green, acidic mixtures produces blue precipitate, and salt and water turned yellow. The reaction of K3Fe(CN)6 and FeSO4 produces a blue green colored mixture with a blue precipitate in it. This blue precipitate is an iron blue called Turnbull’s blue. It signifies the reaction when Fe 2+ is treated with K3Fe(CN)6 (Petrucci et al, 2011). The FeSO4 dissociate into Fe2+ and SO42-, and the Fe2+ reacts with K3Fe(CN)6 to produce the blue precipitate. Thus, acidic mixtures that turned blue has evidence of having Fe2+ ion in the solution showing oxidation. In the second part of the experiment, the set up was putted in an agar solution to retain the location of color and rust formation. K 3Fe(CN)6 and phenolphthalein was added to detect the presence of Fe 2+ ion and basic region of the nail respectively. The next factor which is the application of mechanical stress was tested. After overnight, the straight nail appeared to have rust all throughout it. The body of the nail is colored pink thus a basic region which is the site of reduction (O2 (g) + 2H2O (l) + 4e 4OH-(aq)) reaction due to the presence of OH in the product side. The tip and head of the nail appeared to be dark blue (presence of Turnbull’s blue) which specifies the oxidation site. Since the reduction occurs in the cathode and the oxidation at anode, the body of the nail serves as the cathode while its head and tip as the anode. However,
upon application of mechanical stress by bending the nail, the nail appeared to have a Turnbull’s blue at the bent portion thus the oxidations sites were increased to three (tip, head, and bent portion). Because the strained metal is more active thus more anodic than unstrained metal, the nail preferentially oxidized at portions like this. This is similar to the preferential rusting of a dented automobile fender (Petrucci et al, 2011). The third part of the experiment was about the effect of contact with other metals on the rusting of iron. After overnight stand of nail coiled with Cu strip, the nail was still rusted. It is so that iron can be electroplated with Cu and be protected from rust, but once the coating is cracked, the underlying iron starts to corrode. Since iron is more active than Cu, it undergoes oxidation and copper plating undergoes reduction half-reactions. That’s why the copper strip is pink whilst the iron is dark blue and rusted according to Table ___.
Reduction half reactions: Cu2+
+
2e-
Cu
Oxidation Half-reactions: Fe
Fe2+
+
2e-
Overall reaction:
Fe
Cu2+
+
Cu
Fe2+
+
While in the zinc strip coiled on iron, the zinc coats the iron which is known as galvanized iron. Since zinc is more active compared to iron, the iron is still protected from rusting even if the zinc plating is broken. Zinc is oxidized instead of iron and corrosion of zinc protects it from further corrosion. The faint white precipitate in the zinc strip mentioned in Table __ is characterized by zinc ferricyanide as the reaction of K3Fe(CN)6 with Zn. The part of the nail uncovered with zinc was colored pink thus a reduction site.
Reduction half reactions: Fe2+
+
2e-
Fe
Oxidation Half-reactions: Zn
Zn2+
+
2e-
Overall reaction:
Zn
+
Fe2+
Fe
+
Zn2+
The zinc serves as the anode and the oxidation site while the iron as the cathode and reduction site (Petrucci et al, 2011). The errors committed in the experiment may mainly concern to the uncertain concentration of the reagents that affected the rate of corrosion. The experiment may also be faulty due to the impure reagents like the white substance in the Na3PO4. The colors also in the 2nd and 3rd part of the experiment was not quite observable and this may be accounted in the insufficient amount of indicating substance like phenolphthalein and K 3Fe(CN)6. I may suggest from the experiment to add more time for the experiment in order to achieve most favorable result and give points on certain highlights like why did KNO 3 did not quite evidenced rusting.
VI.
Conclusion
Rusting of an iron is mainly affected by the three given factors. From the experiment on acidity and basicity of solution in contact with the metal, we could generalize that:
basic solutions induces unappreciable rusting on the metal acidic solutions induces the metal to rust further salt solutions catalyze the corrosion reaction of the metal and water solution induces rust on the iron typically
On the second part of the experiment, mechanical stress on metal was tested to affect rusting. It can be concluded that strained metal portion has evidenced of the oxidation site and serves as the anode while the unstrained portion serves as the cathode and the reduction site. Lastly, contact with other metals affect the rusting of metal concerned. The metal with the most active or most anodic (characterized by less standard reduction potential) will be the one to be corroded and the other will be the cathode of the reaction. VII.
Literature Cited/Bibliography -Chang, R., Goldsby, K. A, Chemistry 11th Edition, McGraw-Hill Publishers, McGraw-Hill Companies, Inc., 1221 Avenue of the Americas, New York, NY 10020, © 2013 -Petrucci, R. H., Herring F.G., Madura, J. D., Bissonnette, C., General Chemistry: Principles and Modern Applications 10 th Edition, Pearson Prentice Hall, Pearson Canada Inc., Toronto, Ontario, © 2011 -Torio, M.A. O., Revilleza, Ma. J. R., et al, Laboratory Instruction Manual for CHEM 17.1: General Chemistry II Laboratory, General Chemistry and Chemical Education Division, Institute of Chemistry, CAS, UPLB, Los Baños, Laguna, © 2012 -Zumdahl, S. S., Zumdahl S. A. et al, Chemistry, 8 th Edition, Brooks Cole, Cengage Learning, 10 Davis Drive, Belmont, CA, USA, ©2010
