289397846 Anodes AND Cathodes IN Corrosion Reactions PDF

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EXPERIMENT 5 : ANODES AND CATHODES IN CORROSIONREACTIONSOBJECTIVE : To perform corrosion testing on several metals and understand hoe corrosion occur. To show the existance and location of anodes and cathodes in corrosion processes. To demonstrate methods of effective corrosion protection. INTRODUCT...

Description

EXPERIMENT 5 :

ANODES AND CATHODES IN CORROSION REACTIONS

OBJECTIVE

:

•

To perform corrosion testing on several metals and understand hoe corrosion

•

occur. To show the existance and location of anodes and cathodes in corrosion

•

processes. To demonstrate methods of effective corrosion protection.

INTRODUCTION

:

In analyzing corrosion, the first thing that must be determined is whether a metal reacts with its environment. It is generally accepted that corrosion processes are caused by the formation of electrochemical cells. The electrochemical reactions in these cells can be divided into two reactions:- anodic reactions and cathodic reactions where each reaction is called a half-cell reaction. In the anodic reaction, metal goes into solution as an ion. In the cathodic reaction, electrons provided by the anode, flow through the metal until they reach the cathode where they can be combined with positively charged ions. Both the anodic and cathodic reactions must occur simultaneously for a corrosion process to proceed. The electrode potential in the electrochemical cell is when an ideal metal is placed in an electrolyte, an electrode potential develops that is related to the tendency of the material to give up electrons. To measure this tendency, we measure the potential difference between the metal and a standard electrode using a half-cell. The relative corrosion resistance of a material can be evaluated using the following criteria: Relative Corrosion Resistance mpy outstanding 200 Consider a car engine block with and without the addition of antifreeze, as an example of a corrosion rate problem. Engine blocks are generally made from low carbon steel that corrodes when exposed to tap water. This is the brown rust that is often seen in cooling systems containing water only. The corrosion reactions for low carbon steel are:

Fe --> Fe2+ + 2eFe2+ --> Fe3+ + eO2 + 2H2O + e- --> 4OH2Fe + 2H2O + O2 --> 2Fe2+ + 2OH- --> 2Fe(OH)2 2Fe(OH)2 + H2O + 1/2 O2 --> 2Fe(OH)3 (rust)

PROCEDURE

:

1. 200 mL of 0.1 M sodium nitrate was heated to boil in a 500 mL beaker. While stirring with a glass rod, 3 g of powdered agar was added. The mixture was heated and stirred until agar forms a suspension. 2. 10 drops of 0.1 M potassium ferricyanide and 10 drops of phenolphthalain was added to the agar suspension. 3. Three iron nail was cleaned with sand paper. 4. One iron nail was placed in the bottom of the first petri dish. 5. The 10 cm of copper wire and 10 cm of zinc wire was polished with sand paper. 6. The 10 cm copper wire was wrapped around the second iron nail. 7. The 10 cm zinc wire was wrapped around the third iron nail. 8. The second and third wrapped iron was placed in the bottom of the second petri dish. The two wrapped nail is not touch each other. 9. Enough warm agar was pour in both petri dishes to cover all of the iron strips to a depth of about 1 mm. 10.Both dishes was covered and let stand about 48 hours. 11. On day 1, the subsequent color developments was observed over a period of several hours. 12.On the day 3,the petri dishes was observed against both white and black backgrounds and the results was recorded. 13.On the day 4,the petri dishes was observed against both white and black backgrounds and the results was recorded.

RESULT AND OBSERVATIONS : Time

Item

Experimental Conditions

Observations

Sketch

Dish 1 (iron nail)

Iron in agar + phenolphthalein + potassium ferricyanide

Blue and brown colours formed at different places on the nail.

Dish 2 (zinc)

Iron wrapped with zinc + phenolphthalein + potassium ferricyanide

Blue coloration appears.

Iron wrapped with copper + phenolphthalein + potassium ferricyanide

The brown and blue colour develops along the copper surface and the blue color develops on the bare.

Day 1

Dish 3 (copper)

Dish 1 (iron nail)

Iron in agar phenolphthalein potassium ferricyanide

+ The iron + corrode

nail

Dish 2 (zinc)

Iron wrapped with zinc + phenolphthalein + potassium ferricyanide

The surface of iron nail that does not covered with zinc start to corrode

Dish 3 (copper)

Iron wrapped with copper + phenolphthalein + potassium ferricyanide

The iron nail corrode

Item

Experimental Conditions

Day 2

Time

Observations

Sketch

Dish 1 (iron nail)

Dish 2 (zinc) Day 4

Dish 3 (copper)

DISCUSSION

Iron in agar + phenolphthalein + potassium ferricyanide

Brown colours formed along the nail.

Iron wrapped with zinc + phenolphthalein + potassium ferricyanide

Blue coloration appears along the nail.

Iron wrapped with copper + phenolphthalein + potassium ferricyanide

The brown colour develops along the copper surface and the blue color develops on the unwrapped side.

:

In this experiments, the observation for dish 1(where the experiment conditions with iron in agar, phenolphtalein and potassium ferricyanide) at day 1 is the outer layer of nails is brown and slightly blue while at day 3 and 4 the outer layer

of nail is clearly brown in color. The piece of iron appears uniform on the large scale, but at the atomic level it is quite irregular. Regions of the iron which have been subjected to intense stress, like the stamped head and point of the nail, or the sheared ends of the wire, contain atoms that have a higher energy than there unstressed neighbors. These regions lose electrons or undergo oxidation slightly more readily than the unstressed regions. The reaction is shown below. Fe (s)

Fe2+(aq) + 2 e-

These electrons are readily taken up during the reduction of water, according to the reaction below. 2 H2O (l) + 2 e-

2 OH-(aq) + H2 (g)

Alternatively 2 H2O (l) + O2 (g) + 4 e-

4 OH-(aq)

If we can keep water and oxygen away from our iron surfaces, we can minimize the amount of corrosion that can occur, since oxidation cannot occur without reduction. The colors are created from two additional reactions as shown below. Phenolphthalein is a weak acid and will be denoted as HPh. HPh(aq) + OH-(aq) colorless

Ph-(aq) + H2O(l) pink

Wherever there are significant quantities of OH- the region will turn pink. H2O (l) + K+(aq) + Fe2+(aq) + Fe(CN)63-(aq)  KFe[Fe(CN)6] · H2O(s) pale yellow blue This last compound is call variously Turnbull’s blue or Prussian blue. It is somewhat unusual in that it is a mixed valence compound containing iron in both the +2 and +3 oxidation states. The agar solution will turn blue in any region with a supply of Fe2+. When iron is corroded, rust is formed in the reaction : 4Fe(s) + 3O2(g) + H2O(l) = 2Fe2O3H2O(s) The amount of rust that forms depends on the amount of water available for the iron to react with. The two electrochemical half-reactions are: Fe(s) = Fe2+(aq) + 2e- . This is the anode. O2(g) + 2H2O(l) + 4e- = 4OH-(aq). This is the cathode.

The anode and cathode are both on the same piece of iron but different regions of it. For this reaction to work, the electrons need a wire of some sort or some way of conducting the electron flow. In this case, the iron itself acts as the wire. Besides, the observation for dish 2 (where the experiment conditions with iron wrapped with zinc in agar, phenolphtalein and potassium ferricyanide) at day 1 white haze appeared at the port with zinc wire and the wrapped port then turn slightly pink. While, at day 3 and 4 the unwrapped port become more pinky while more white haze form at wrapped port. In this results, blue coloration appears because zinc is more active than iron and performs as the anode in the zinc-iron galvanic couple. Zinc ions form at the anode but they do not form a colored compound with the indicators used. Hence, no color develops on the zinc area but white haze was appeared. Lastly, the observation for dish 3(where the experiment conditions with iron wrapped with copper in agar, phenolphtalein and potassium ferricyanide) at day 1 the unwrapped port turn slightly blue and wrapped port turn slightly pink. At the day 3 and 4, dark blue formed at the unwrapped port and more pinky color formed at wrapped part. The slightly pink colour develops along the copper surface because it is performing as a cathode and the blue color develops on the bare iron surface that is the anode in the copper-iron galvanic couple. The slightly pink colour results from the accumulation of alkali on cathode areas and the blue color reveals the presence of ferrous ions at the iron anode surface.

CONCLUSION

:

Corrosion is when metal is gradually destroyed through chemical reactions. Because it is a chemical reaction, we can prevent it with chemical reactions, namely redox reactions. For dish 1, Fe act as anode, while for dish 2, Zinc ions form at the anode and for dish 3, the copper surface performing as a cathode. Therefore, if any

of this component which is an anode, a cathode, a conducting environment for ionic movement (electrolyte), an electrical connection between the anode and cathode for the flow of electron current is missing or disabled, the electrochemical corrosion process will be stopped. Clearly, these elements are thus fundamentally important for corrosion control.

QUESTIONS

:

1. Locate the anode and cathode in the three different conditions. For dish 1 : The anode and cathode are both on the same piece of iron but different regions of it. Fe(s) = Fe2+(aq) + 2e- . This is the anode. O2(g) + 2H2O(l) + 4e- = 4OH-(aq). This is the cathode. For dish 2 : Zinc ions form at the anode and zinc as a cathode. For dish 3 : Iron as an anode and Copper surface performing as a cathode. 2. Why does the iron behave differently in the three cases in this experiments? Because it is depend on the reactivity of iron and the other metal (wrapped metal: zinc znd copper) 3. Which atom have corroded? Iron, Fe 4. Write the possible electrochemical reactions for the rusting of iron nails until it become stable oxide compound (or rust). 1) Fe (s) Fe2+(aq) + 2 e2) 2 H2O (l) + O2 (g) + 4 e4 OH-(aq) Phenolphthalein is a weak acid and will be denoted as HPh. HPh colorless 3) H2O

(l)

+

(aq)

OH-(aq)

+

Ph-(aq)

+

H 2O

(l)

pink K+(aq)

+

Fe2+(aq)

+

Fe(CN)63-(aq)  KFe[Fe(CN)6] · H2O(s)

pale yellow 4) 4Fe(s) + 3O2(g) + H2O(l) = 2Fe2O3H2O(s)

blue

5. Rank the metals used in this experiment based on their reactivity.

Cu Fe  Zn More reactive 6. Explain how cathodic protection works.

Corrosion is the principal cause of leaks in underground storage tanks and pipelines. External corrosion of underground steel is an electrochemical process. The soil is an electrolyte, or conductor of corrosion currents. Corrosion current flow through the soil between different points on the buried metallic structure. When corrosion current enters or leaves the metal surface, a chemical reaction takes place. The results of the process is corrosion of the metal surface where current leaves the metal and enter the soil. Cathodic protection halts the corrosion process on the underground or submerged structure by changing the electrical condition and transferring the damaging chemical reaction away from the structure to an independent anode. Once the anode is installed on the soil, it will continue to provide a protective barrier to the structure for the design life of the system.

REFERENCES

:

Cathode and Anode Half-Cell Reactions (n. d.). Education portal. Retrieved November 8, 2014 from http://education-portal.com. Joseph Franek (2001). Iron corossion. The Regents of the University of Minnesota . Retrieved November 8, 2014 from https://www.chem.umn.edu K.R. Trethewey and J. Chamberlain: "Corrosion for Science and Engineering 2 nd Edn.", Longman (UK), 1995. Retrieved November 8, 2014 from http://www.corrosion club.com.

Laque, May, and Uhlig. (n. d.). A Classic Corrosion Experiment : Anodes and Cathodes in Corrosion Reactions. Corrosion in Action. Retrieved November 8, 2014 from http://www.corrosion-doctors.org...