The Understanding of the Whole Concept of Corrosion - Lab Report Example

Memorandum Report Megan Lynch, Corrosion Lab TA Kwame Owusu-abankwah CC: Dr. Guida Bendrich October 5, Re: CHE 2506 – Corrosion Lab Ms. Lynch, As per your October 5, 2010 request, our team comprising of Mhd Adiba Zahir, Megan O’Donnell and I have completed the Corrosion Lab at the University of New Brunswick’s Chemical Engineering Department. In this laboratory setup, galvanic corrosion was studied. With two metal strips electrically connected in a liquid electrolyte, the most noble and most corroded metals were discovered. The enclosed report studies the causes and mechanisms of corrosion. It also includes the results as well as the analysis of the data gathered. Regards, Kwame owusu Kwame Owusu-abankwah Encl: Report Corrosion Lab Prepared for Ms. Megan Lynch Corrosion TA CHE 2506- Material Science Prepared by Bismark Owusu October 5, 2010 Summary In an attempt to explore and delve deep into the understanding of the whole concept of corrosion, galvanic corrosion to be precise, the lab experiment did offer such worthwhile and insightful practical experiences. Galvanic corrosion occurs when two metals, or rather a galvanic metal couple, are placed in a common electrolyte and are electrically connected. As observed during the experiment, the behaviour of the two metals in the above conditions was dissimilar. One metal acted as the anode (corrodes) while the as the cathode. Two or more metal strips were immersed at different depths in a liquid electrolyte and the observations made are discussed below. One of the key concerns of the experiment was to determine the potential difference between the metal pairs; Zn-Al, Zn-Bx and Al-Bx. More so, the additive rule was put to test by adding a metal or two to the already established metal pair. As it turned out, the added value had a 0.5% difference with its theoretical value. With this factual finding, it can be hypothetically stated that the additive rule did not apply to this experimental set up. In addition, it was carefully noted that the area immersion ratio of the cathode/anode directly affected the current density consequently affecting the rate of corrosion. A constant cathode depth immersion with a decreased anode depth immersion saw Zinc exponentially lose its electrons. This ignited the flames that fuelled the corrosion of Zinc metal. On the other hand, when the anode immersion depth remained the same and the cathode immersion depth increased, the corrosion rate was gradual. Penultimate, the effects of stirring the liquid electrolyte was tested with empirical results. Stirring only left ephemeral results which were short lived. This implies that when the stirrer was turned on, the current shot up a few mille-Amperes and then started depreciating even with the stirrer still on at a similar rate to when the stirrer was off. This caused the electrolyte solution to become more uniform for a short while, which improved the current circulation. Lastly, a sacrificial anode was used to protect steel. The sacrificial anode gave up its electrons to the valuable steel thus protecting it from corrosion. Results The main effect of any form of corrosion is the deteriorative loss of a metal. These sentiments are echoed by Callister (2007), who defines corrosion as a process by which there is a deteriorative loss as a result of dissolution environmental reactions. There are different forms of corrosion such as rust but to define and limit the ambit of this study, galvanic corrosion suffices the objectives of this experiment. Galvanic corrosion is characteristically defined by its electrochemical reaction. As already indicated, two different strips of metals were electrically connected to a liquid electrolyte. A voltmeter was used to measure the potential difference that developed between the two metals. In experiment 1, the potential differences between three pairs of metal strips were measured: zinc-aluminum (Zn-Al), zinc-brass (Zn-Bx) and aluminum-brass (Al-Bx). The outcome of this experiment is tabulated bellow in Table 1. Table 1: Experiment 1-Potential difference Al (+) Bx (+) Zn (-) 0.422 0.848 Al (-) 0.430 In the second experiment, the effect of temperature and electrode areas on the corrosion rates of copper and zinc metal combination was investigated. With the same apparatus as in experiment 1, a multimeter was connected as an ammeter. Each current was measured at different area immersion depth ratios of the two metal strips. The current density, as well as the Zinc corrosion rate was then calculated. (See calculations in Appendix). The results for each area immersion depth ratios are as illustrated in table 2 below. . Table 2: Current measurements for varying ratios Current (mA) Current density ( Zinc corrosion rate 5 1 0.2 0.27 0.054 1.83 x 10 5 2 0.4 0.42 0.084 2.84 x 10 5 3 0.6 0.54 0.108 3.66 x 10 5 4 0.8 0.61 0.122 4.13 x 10 5 5 1.0 0.75 0.150 5.08 x 10 4 5 1.25 0.67 0.168 5.67 x 10 3 5 1.67 0.63 0.210 7.11 x 10 2 5 2.5 0.57 0.285 9.65 x 10 1 5 5 0.48 0.480 1.63 x 10 Experiment 2.2 was set with the main aim of determining the effects of different flow conditions on the corrosion rate of the zinc/copper galvanized couple. The key point of focus in this experiment was to establish the effects stirring the liquid electrolyte had on corrosion. The current was measured continuously with the stirrer being turned on and off. The results are as shown in Table 3 below. Table 3: Experiment 2.2- The effect of stirring on corrosion (Zn/Cu cell) t(s) Time (mm:ss) Stirrer Current (mA DC) t(s) Time (mm:ss) Stirrer Current (mA DC) -30 -0:30 On 0.97 260 4:20 on 0.75 0 0:0 On/Off 0.91 270 4:30 on 0.75 10 0:10 Off 0.87 280 4:40 on 0.74 20 0:20 Off 0.86 290 4:50 on 0.73 30 0:30 Off 0.84 300 5:00 on 0.73 40 0:40 Off 0.83 320 5:20 on 0.72 50 0:50 Off 0.82 340 5:40 on 0.71 60 1:00 Off 0.80 360 6:00 on 0.70 70 1:10 Off 0.79 380 6:20 on 0.70 80 1:20 Off 0.78 400 6:40 On/Off 0.69 90 1:30 Off 0.77 405 6:45 Off 0.69 100 1:40 Off 0.76 410 6:50 off 0.68 110 1:50 Off 0.75 415 6:55 off 0.68 120 2:00 Off 0.74 420 7:00 off 0.67 130 2:10 Off 0.74 440 7:20 off 0.66 140 2:20 Off 0.73 460 7:40 off 0.65 150 2:30 Off 0.72 480 8:00 off 0.64 160 2:40 Off 0.71 500 8:20 off 0.63 180 3:00 Off 0.70 520 8:40 off 0.62 200 3:20 Off 0.69 540 9:00 off 0.61 220 3:40 Off 0.67 560 9:20 off 0.61 240 4:00 Off/On 0.66 580 9:40 off 0.60 245 4:05 On 0.75 600 10:00 off 0.59 250 4:10 On 0.76 620 10:20 off 0.59 255 4:15 On 0.76 640 10:40 off 0.58 The next experiment (experiment 3) was set with much emphasis laid on trying to answer the question; how to protect steel electrochemically. There are two methods of protecting steel: a sacrificial anode and an imposed voltage are seen in experiments 3.2 and 3.3. The electrode immersion depth in each experiment is 5 in. The results are seen in table 4 below. Table 4: Experiment 3.x- Protection of steel by use of sacrificial anode Exp. 3.1 Zn not connected Exp. 3.2 Zn connected Exp. 3.3 Power off Exp. 3.3 Power on Meter A1 (mA) 0.0200 -0.5690 0.0260 0.00 Meter A2 (mA) 0.02 0.51 0.04 0.03 Power supply V when A1=0.0mA 0.02 Analysis and Discussion Experiment 1 Zn-Al, Zn-Bx and Al-Bx were the three metal pairs analyzed in this lab. Table 5 has the decreasing order of measured potentials of these pairs. Table 1: Metal pairs in decreasing order of measure potentials Metal Pair Measured Potential (V) Zn-Bx 0.848 Al-Bx 0.430 Zn-Al 0.422 The additive rule is used to estimate the potential difference of a third metal or metal alloy with two other metals or metal alloys, of which the potential differences are known. (Zn-Al) + (Al-Bx)  (Zn-Bx) (0.422 V) + (0.430 V)  0.848 V 0.852  0.848 V The percent difference between the measured potential difference (from table 1 - experimental) and the theoretical potential difference was calculated. % Difference = % % Difference =  = 0.5 With the result of 0.5% potential difference, the figure clearly shows that the additive rule is not accurate hence its significance and implication in this experiment is annulled. In understanding galvanic series a definition advanced by Callister (2007) suffices in advancing this argument. He defines galvanic series as a ranking of metals and alloys as to their relative electrochemical reactivity in seawater .It is a list of metals as well as metal alloys from the most noble to the most active. Table 6 lists the metals used in this lab experiment from the most cathodic to the most anodic for both the galvanic series and the standard emf series. Table 6: List of metals used in this lab set up in order of the most cathodic to the most anodic for the galvanic and standard emf series Galvanic Series Standard EMF Series (V) Brass Aluminum -0.633 Zinc -1.622 It was observed that zinc tends to corrode when it is tested electrochemically with the other metals. We can therefore deductively say that zinc is the most anodic element among brass and aluminum. Of the three Aluminum follows second as it tends to corrode more than brass. Brass thus ranks the lowest as it cannot corrode unless it reacts with a metal below its ranking in the galvanic series. This is the same for both experimental and theoretical galvanic series. For each pair of metal strip, one acts as an anode and the other as a cathode. The anode is the metal strip that lost the most electrons. The other metal strip acts as the cathode by gaining electrons. Zn acted as the anode for the Zn-Al and Zn-Bx pairs. Bx acted as the anode for the Al-Bx pair. This was determined by the recorded potential difference. If the potential difference was positive, then the metal strip with the negative lead touching it was the anode. Also, if the potential difference was negative, then the metal with the negative lead touching it was the cathode. Cold working, as defined by Callister (2007), is a phenomenon whereby a ductile metal becomes harder and stronger as it is plastically deformed. He further argues that Cold working affects the rate of corrosion because cold-worked metal is more vulnerable to corrosion than the same material in an annealed state. If a cold worked material is more susceptible to corrosion, then it will have a greater rate of corrosion. Cold working also affects the potential difference between two points. This lab experiment, however, was conducted under room temperature thus undergoing similar condition as the theoretical one. Experiment 2.1 Figure 1 represents the current density of Zinc versus the immersion area ratio. Figure 1: Current density of zinc versus immersion area ratio Figure 2 represents the corrosion rate of zinc versus immersion area ratio for the experiment Figure 2: Corrosion rate of zinc versus immersion area ratio The slope for the graph current density of zinc versus immersion area ratio is 0.075. Also, the slope for graph corrosion rate of zinc versus immersion area ratio is 2.5 x 10-8. From both graphs (figure 1 and 2), it is observed that the corrosion rate is higher when current density and immersion ratio is higher. This implies that corrosion rate is more influenced by the anodic area variation since putting a large surface of zinc in the liquid electrolyte produces more zinc ions. Consequently the current density will increase with an increasing concentration of Zn2+ in the solution. Temperature and the rate of corrosion exhibits a relationship expressed in the graph below. An increase in temperature results to a corresponding increase in the rate of corrosion, and as Bendrich, (2010) posits, Corrosion is based on a chemical reaction, and a chemical reaction’s rate increases with an increase in temperature. Experiment 2.2 Figure 3 represents the current versus time for the experiment that involves both the stirred and the unstirred case. Figure 3: Current versus time for both stirred and unstirred case. When the stirrer was turned off, for the first few minutes the current decreased over time. After some minutes when the stirrer was turned on, the current increased rapidly for a few seconds, and then began to decrease again at a similar rate to when the stirrer was turned off. Therefore, the stirrer does affect the cell current, but only for a short while. This is mostly due to the increase in uniformity of the liquid electrolyte, which is believed to allow current to flow more freely. Experiment 3.1 Rate of corrosion is calculated using the formulae: r = Where r: rate of corrosion (g/(s x unit area)); i: current density (A/unit area); M: atomic mass (g/mol); n: number of electrons transferred in reaction; F: Faraday constant; 96500 As/mol. Average Current (mA) = mA Immersion Area of Zinc = 5 in2 Current density (mA/in2) = 0.02 mA/ 5 in2 = 0.004 mA r = = = 1.36 x 10-6 g/in.s The anode in this experimental set up is the steel. Copper acts as the cathode. This is because the steel has more negative voltage. Experiment 3.2 Zinc has to be connected to the steel to provide protection to the steel. Zinc acts as a sacrificial anode to protect the steel by polarising its surface. This polarization stops corrosion of steel, although zinc continues to corrode. When Zinc is connected by either the wire or immersed in the salt the solution will provide protection for steel. Experiment 3.3 The DC power supply is a unit that supplies direct current. It maintains a fixed voltage power to the circuit. Without the DC supply being connected, the ‘scrap steel’ will have fewer electrons and so it will corrode faster. The corrosion rate is calculated as: Current (mA) = 0.03 mA Immersion Area of Zinc = 5 in2 Current density (mA/in2) = 0.03 mA/ 5 in2 = 0.006 mA/in2 r = = = 2.03x10-6 g/in.s In order to protect an underground steel pipeline from destruction, there is a must need to put in place measures to curb it from corrosion. One such important method is the use of a sacrificial anode – a metal with a more negative voltage than steel. The sacrificial anode in this case may be zinc or aluminum which provides the cathodic protection for the steel. Galvanising and electroplating are other methods of protecting steel from corrosion. As Callister (2007), illustrates, galvanizing is done completely immersing the steel pipe in a pool of molten zinc .This makes the zinc anodic, and cathodically, protects the steel. Electroplating on the other hand is done by passing an electrical current through a solution of dissolved metal ions and the metal object to be plated. In the above set up, the metal object acts as the cathode in an electrochemical cell, by attracting metal ions from the solution. Conclusion This laboratory experiment explored the electrochemical nature of corrosion and its control. Current density is defined as current per unit area and the corrosion rate is determined by the current density at the surface of the anode. It was observed that a metal will corrode faster when the anode area is smaller and vice versa. The concept of a sacrificial anode was successfully studied and demonstrated through the experimental set up. The use of zinc as an anode to protect a ‘valuable steel’ structure was compared with using ‘scrap steel’ with a DC power supply. It was also observed that the DC power supply increased the flow of electrons from the “scrap steel” to the “valuable steel”. This in turn, protected the “valuable steel”. Recommendation This lab was a good practice for concepts that were not yet taught in class. It would have been better and more beneficial if this experimental set up was conducted soon after the concepts were taught in class. References Bendrich, G. (2010) Materials Science in a Nutshell. New York: GD Publishing Bendrich, G. (n.d.). ChE 2506 -Materials Science Laboratory. Retrieved October 5, 2009, from Blackboard Learning System: https://learning.unb.ca Callister, D.W. (2007) Materials Science and Engineering an Introduction 7th Ed. New York: John Wiley & Sons Inc. Read More