The Anoxic Corrosion of Copper in Pure Water and Chloride Rich Brines

Transcription

1 The Anoxic Corrosion of Copper in Pure Water and Chloride Rich Brines by Emilija Ilic A thesis submitted in conformity with the requirements for the degree of Master of Applied Science Graduate Department of Chemical Engineering and Applied Chemistry University of Toronto Copyright by Emilija Ilic 2015

2 The Anoxic Corrosion of Copper in Pure Water and Chloride Rich Brines Emilija Ilic Master of Applied Science Graduate Department of Chemical Engineering and Applied Chemistry University of Toronto Abstract 2015 The Nuclear Waste Management Organization (NWMO) is developing an approach for the permanent geological disposal of nuclear waste. The waste will be encased in copper coated used fuel containers (UFCs) and placed in a deep geological repository (DGR). To support the NWMO in their investigations on the long-term corrosion of copper a lab scale simulation of the DGR environment was created. Copper wires were placed in glass electrochemical cells and exposed to one of two environments; pure anoxic water or chloriderich anoxic brine. The systems were allowed to freely corrode and accumulate hydrogen within their headspaces over extended durations at 30 to 75 C. The hydrogen was periodically purged and subsequently analyzed using a highly sensitive amperometric sensor; these measurements were utilized to calculate the corresponding copper corrosion rates. Corrosion with hydrogen evolution was demonstrated in both pure water and brines at slow rates below 1 and 10 nm/year, respectively. ii

3 Acknowledgments I would like to express my gratitude to my supervisor Professor Roger Newman for welcoming me into the Corrosion Group and for his guidance over the duration of this research. I would also like to send my appreciation to the NWMO for their financial support, particularly to Peter Keech whose feedback during meetings greatly benefited the progress of this work. Thank you to Nicholas Senior and Bahar Abgahri for their collaboration on this project and for training me in all aspects of the experimental work. Dorota Artymowicz this thesis would have been difficult to complete without your intelligent insights and opinions during our numerous discussions, for this I give you great thanks. Anatolie Carcea thank you for ensuring the lab was always stocked with the equipment and materials necessary for my experimental work. Finally, thank you to the remainder of my lab mates; Amir, Ayman, Ali, Suraj and Mariusz for your support, enthusiasm, and for creating a pleasurable work environment over these past 2 years. iii

4 Table of Contents Abstract... ii Acknowledgments... iii Table of Contents... iv List of Tables... vi List of Figures... vii 1 Introduction Background Information Used Nuclear Fuel in Canada Canada s Strategy for the Long-term Storage of Nuclear Waste Used Fuel Container (UFC) Effect of the Repository Environment on the UFC Temperature Redox Conditions Degree of Saturation within the Repository Effect of Salinity in Groundwater Copper Considerations Possible Corrosion Mechanisms General Corrosion Principles Uniform Corrosion Localized Corrosion Microbial Corrosion Stress Corrosion Cracking (SCC) Research Objectives Experimental Procedures Hydrogen Evolution over Freely Corroding Copper Hydrogen Monitored Pure Water Cells Hydrogen Monitored High Salinity Brine Cells Preparation of Copper Wires Electrolyte and Cell Preparation iv

5 4.1.5 Replenishing the Electrolyte Hydrogen Analysis Results and Discussion Converting Cumulative Hydrogen to Copper Corrosion Rate Corrosion Measurements in Pure Water Cells Corrosion Measurements in Highly Saline Brine Cells Overheating Period Pure Water Experiments Overheating Brine Experiments Overheating Solubility of Copper Oxides Solubility of Copper Oxides in Pure Water Solubility of Copper Oxides in Chloride Rich Brines Conclusions and Recommendations References Appendices Appendix A: Experiment Parameters and Details A1: Experiment Parameters A2: Cell Timelines Appendix B: Experimental Procedures B1: Wire Preparation Procedure B2: Cell Preparation Procedure for Pure Water Experiments B3: Solution Preparation Procedure for Brine Experiments B4: Hydrogen Collection Procedure B5: Correcting for Sensor Drift B6: Unit Conversion Calculator Appendix C: Sample Calculations C1: Monolayers of Copper Consumed C2: Depth of Copper Dissolved from Oxides v

6 List of Tables Table 1: Ranges of groundwater composition at repository depths between 500 and 1000 m (King, 2007)... 8 Table 2: Cell identification for hydrogen monitored pure water experiments Table 3: Cell identification for hydrogen monitored brine solutions Table 4: Dissolved copper concentration from assayed solutions, [Cu]total, compared to the calculated copper based on hydrogen measurements, [Cu]corr. *BDL-below detection limit of mm vi

7 List of Figures Figure 1: Example of a Used Fuel Container (UFC) design (Keech, 2015) Figure 2: The evolution of temperature and redox conditions in a deep geological repository (King, 2007) Figure 3: Pourbaix diagram of copper in pure water at 100 C, [Cu]TOT = 10-8 mol/kg (Puigdomenech & Taxén, 2000) Figure 4: Pourbaix diagram for copper in 5 molal [Cl - ] at 100 C, [Cu] =10-4 molal (Beverskog & Pettersson, 2002) Figure 5: Surface profile of stripped copper sample demonstrating surface roughening due to under-deposit corrosion (King et al., 2001) Figure 6: Schematic of electrochemical cell setup for measuring hydrogen evolution from copper wires Figure 7: Sample reading from hydrogen probe with calibrations Figure 8: Cell D1 Cumulative hydrogen and calculated corrosion rate of bare copper wires immersed in deoxygenated pure water at 75 C, following testing at 30 and 50 C.27 Figure 9: Cell D2- Cumulative hydrogen and calculated corrosion rate of bare copper wires immersed in deoxygenated pure water at 75 C Figure 10: Cell O1 - Cumulative hydrogen and calculated corrosion rate of pre-oxidized copper wires immersed in deoxygenated pure water at 75 C, following testing at 30 and 50 C Figure 11: Corrosion rates, as a function of ph and temperature, of all saline brine experiments to date vii

8 Figure 12: Compilation of cumulative hydrogen produced in all pure water experiments, including overheating period. ( ) indicates refreshing of electrolyte Figure 13: Cumulative hydrogen produced in saline brines for experiments in 50 C waterbath at time of overheating Figure 14: Cumulative hydrogen produced in saline brines for experiments in 75 C waterbath at time of overheating viii

9 1 Introduction Rapid innovation and industry expansion has placed a global demand on greener electricity generation methods. Nuclear power provides an abundant supply of energy, with little greenhouse gas emissions and relatively small amounts of manageable waste. It saves the environment from millions of tonnes of carbon dioxide per year which would be emitted from the burning of fossil fuels (World Nuclear Association, 2015). In comparison to other energy sources, the nuclear industry is unique in that it is the only one that takes full responsibility for the management and disposal of its produced waste. For these reasons, nuclear power has been a major source of energy in Canada over the past decades, contributing approximately 15% of the country s electricity generation (World Nuclear Association, 2015). Canada s line of nuclear power reactors are known as CANDU (Canada deuterium uranium), which use heavy water (deuterium oxide) as a moderator and coolant, and natural uranium as the fuel. The Nuclear Waste Management Organization (NWMO) was established in 2002 to explore approaches for managing Canada s nuclear waste. 1

10 2 Background Information 2.1 Used Nuclear Fuel in Canada CANDU reactors are fuelled by fuel bundles containing natural uranium pellets. Each bundle is 0.5 m in length, 0.1 m in diameter, and holds 28 to 37 Zircaloy fuel tubes in a cylindrical array. Once the fissile components within the uranium have undergone fission to generate electricity, the remaining reaction by-products in the used fuel bundles are considered waste. The waste, or used nuclear fuel, remains radioactive for thousands of years, and must be contained and carefully managed. To date, Canada has produced over 2 million used nuclear fuel bundles, enough to fill six hockey rinks to the top of the boards. This waste is currently stored in interim surface facilities at nuclear reactor sites. The facilities require ongoing care and maintenance and are not ideal for prolonged storage of nuclear waste, hence nuclear industries around the world have taken research initiatives in developing a proper long-term solution for the disposal of nuclear waste (Nuclear Waste Managment Organization, 2012, 2013). 2.2 Canada s Strategy for the Long-term Storage of Nuclear Waste The NWMO is developing a method known as Adaptive Phased Management (APM) for the long-term storage of nuclear waste. This management plan focuses on centralized containment and isolation of used nuclear fuel in a deep geological repository (DGR). The method is consistent with waste disposal programs in other countries such as France, Finland, the United Kingdom, and Sweden. 2

11 The DGR will be implemented m underground in a suitable rock formation, either crystalline or sedimentary, and will consist of a network of waste storage rooms. A number of surface facilities will also be established to support the design, construction and operation of the repository. The surface handling facilities will receive used fuel bundles from interim storage sites, and will package them into durable containers, termed used fuel containers (UFCs), before transferring them underground (Nuclear Waste Managment Organization, 2012). 2.3 Used Fuel Container (UFC) The used fuel containers must be durable and corrosion resistant, in order to withstand the repository environment over thousands of years. The most recent UFC model is shown in Figure 1; it consists of a steel vessel 100 mm in width, which encloses the used fuel bundles and provides mechanical support, and is coated with a 3 mm copper layer which serves as a corrosion barrier (Scully & Edwards, 2013). The UFC s will be placed in respective boreholes drilled within the floor of the repository storage rooms, upon which excavations will be sealed with a thick layer (>30 cm) of bentonite clay. The bentonite seal acts as a buffer, resisting physical and chemical changes which may occur in the repository environment, such as glacial events and migration of corrosive species within groundwater (Nuclear Waste Managment Organization, 2012). 3

12 Copper Coating Lid Used Fuel Bundles Steel Vessel Figure 1: Example of a Used Fuel Container (UFC) design (Keech, 2015). 2.4 Effect of the Repository Environment on the UFC The environment to which the used fuel container is exposed to will vary over the evolution of the repository and will affect the corrosion stability of the outer copper coating. This evolution can be classified by changes in temperature, redox conditions and the degree of saturation with groundwater (King, 2007) Temperature The temperature evolution within the repository can be divided into two phases; an initial warming phase followed by a long-term cooling phase. The warming phase is initiated after the repository has been closed, during which the temperature is expected to increase due to radiogenic heat from the used nuclear fuel. The maximum surface temperature the container will attain depends on its arrangement and orientation to surrounding containers, as well as the properties of the surrounding bentonite clay and the cooling rate of the waste, which depends on its average burnup and time spent in interim storage. The UFC and repository layout is 4

13 designed in such a way to allow a maximum container surface temperature of 100 C (King, 2007), though the actual temperature is not expected to exceed 80 C. The maximum temperature is expected to be attained approximately 30 years after placement of the containers; it is then that the containers surface is most vulnerable to corrosion (King, 2007). After this time the repository will enter a long-term cooling phase, with the temperature falling to near-ambient after years. The temperature evolution is demonstrated in Figure Redox Conditions Redox conditions refer to the balance of the reducing and oxidizing reactions occurring at the surface of the copper canister and the adjacent repository environment. As shown in Figure 2, the evolution of the redox conditions can be divided into an initial oxidizing period, followed by a long-term anoxic period (King, 2007). Upon closure of the repository, trapped air will create an oxidizing environment, in which oxygen will take up electrons from copper, resulting in corrosion of the UFC surface. As oxygen is consumed the environment will transition into an anoxic state, reducing corrosion (King, 2008). The degree to which copper corrodes in this anaerobic environment is the primary focus of this thesis. 5

14 Figure 2: The evolution of temperature and redox conditions in a deep geological repository (King, 2007) Degree of Saturation within the Repository The degree of saturation refers to the time it takes for the sedimentary rock to take up groundwater and equilibrate with the bentonite clay. There are expected to be three saturation stages over the evolution of the repository: 1. Initial Dry-out Phase Immediately following the placement of the containers and sealing of the repository, pore water within the bentonite clay is driven away due to a thermal gradient created from radiogenic heating. The environment is oxic and there is very little oxygen consumption or corrosion (King, 2007). 2. Transition Phase Over time the bentonite will start to gradually saturate and swell with groundwater. Corrosion of copper will begin due to the formation of a thin water layer, containing dissolved oxygen, 6

15 on the surface of the UFC. The corrosion reaction will gradually consume the oxygen, leading to an anoxic environment (King, 2007, 2008). 3. Saturated Phase The bentonite becomes completely saturated with groundwater and develops its maximum swelling pressure. Oxygen is completely consumed, following onset of anoxic conditions in which corrosion rate decreases. Full saturation is expected to occur tens to thousands of years following repository closure. In general, copper corrosion is more prevalent during the early, warm, oxidizing and unsaturated period in the repository environment. Corrosion is expected to decrease and nearly cease when the environment transitions into a cool, non-oxidizing and saturated state (King, 2007) Effect of Salinity in Groundwater Table 1 summarizes groundwater compositions in potential host rock formations at repository depths of m. For both the Canadian Shield crystalline rock and the Ordovician sedimentary rock, groundwater salinity increases with depth, is essentially oxygen free, and shows no evidence of reduced sulfur species (King, 2007). Further, both sets of groundwater are chloride based solutions and have near-neutral ph. However groundwater of the Ordovician rock tends to be more saline than that of the Canadian Shield at equivalent depths. For example at a 500 m depth, sedimentary deposits exhibit groundwater salinities of approximately mg/l total dissolved solids (TDS), compared to just 2000 mg/l in crystalline deposits (King, 2008). This may be troublesome, as high concentrations of chloride 7

16 negatively impact the corrosion stability of the copper coated UFC; more on this in Section Table 1: Ranges of groundwater composition at repository depths between 500 to 1000 m, TDS= total dissolved solids (King, 2007). Canadian Shield Crystalline Rock Components Eh (mv) ph TDS (g/l) Na (g/l) K (g/l) Ca (g/l) Mg(g/L) Cl (g/l) SO4 (g/l) Minimum Maximum Ordovician Sedimentary Rock Components Eh (mv) ph TDS (g/l) Na (g/l) Ca (g/l) Cl (g/l) SO4 (g/l) Shale -225 Slightly acidic Limestone Copper Considerations For centuries copper has been used in underground applications due to its high corrosion resistance and environmental durability. Many of the copper pipes used to transport water underground in Egypt nearly 5000 years ago are still in existence (Myers & Cohen, 1984). Recovered natural copper deposits immersed in clay-rich mudstone showed practically no corrosion, despite having formed nearly 200 million years ago (Assembly of First Nations, n.d.). In addition, it is relatively abundant and low in cost compared to other noble metals. For these reasons copper has been chosen as the coating material for the UFCs, but it still requires thorough investigation in high salinity environments, as these are known to weaken its corrosion resistance. Recent efforts by the NWMO have focused on developing a reasonable corrosion allowance, on the millimetre scale, to account for various corrosion mechanisms which may affect the copper layer. The corrosion allowance is based on penetrating processes occurring from 8

17 uniform corrosion, under-deposit corrosion and microbial influenced corrosion; however this is at low salinities (below 100 g/l salinity) and assuming corrosion in oxygen free water is negligible (Scully & Edwards, 2013). In the highly saline groundwater of the Ordovician rock formations, some anoxic corrosion allowance may be necessary (Scully & Edwards, 2013). Research at the University of Toronto calculated the contribution of anoxic corrosion to be 2-4 mm over 1 million years, as measured in chloride rich water (350 g/l salinity). These measurements are under continuing investigation and will be discussed in Section Possible Corrosion Mechanisms In order to arrive at a suitable penetration allowance, several corrosion types were considered: uniform, localized, microbial, and stress corrosion cracking. The following gives a general briefing of copper corrosion and reviews each mechanism General Corrosion Principles A metal corrodes when it releases electrons via an oxidation reaction: Cu(s) Cu e Equation 1 These free electrons must be taken up by an oxidant such as O2, via a reduction reaction: or 1 O 2 2(aq) + 2 e + 2H + H 2 O(l) Equation 2 1 O 2 2(aq) + 2 e + H 2 O(l) 2OH Equation 3 If an electron acceptor is not present corrosion does not occur. The main oxidants under investigation in repository groundwater are: dissolved oxygen, hydrogen ions (or water), and sulfate (Puigdomenech & Taxén, 2000). 9

18 Further, in order for corrosion to occur the oxidation-reduction reaction must be thermodynamically possible, that is, there must be a spontaneous release of free energy (- G). For example, in some circumstances hydrogen ions present in pure water may act as electron acceptors: H + + e 1 H 2 2(aq) Equation 4 If Equation 4 is able to proceed spontaneously in the system, then H + may be an adequate electron acceptor. This is the case for the corrosion of iron metal in pure water, as the oxidation-reduction reaction is thermodynamically favorable. However with copper, the reaction: Cu(s) + H + Cu H 2 2(aq) Equation 5 is not energetically favourable and should theoretically not occur (Puigdomenech & Taxén, 2000). However some experimental studies have determined the corrosion of copper in pure water does occur, leading the NWMO to investigate this possibility Uniform Corrosion Uniform corrosion in oxygen free environments has been considered in pure water, chloride solutions and hydrogen sulfide solutions; these are described below Corrosion of Copper in Anoxic Pure Water A Pourbaix (EH/pH) diagram of copper in pure water at 100 C is presented in Figure 3. It demonstrations that copper should not corrode in pure water since the metal oxidation lines are far above the hydrogen equilibrium line, meaning H + cannot serve as an electron acceptor. 10

19 Figure 3: Pourbaix diagram of copper in pure water at 100 C, [Cu] TOT = 10-8 mol/kg. Created by (Puigdomenech & Taxén, 2000). However the thermodynamic constraints of the corrosion of copper in anoxic pure water were questioned by (Hultquist, 1986), who claimed his experiments demonstrated the evolution of hydrogen in a copper-pure water system. He assumed copper corrodes via: 2 Cu(s) + H 2 O(l) Cu 2 O(s) + H 2 (g) Equation 6 but later specified that the first corrosion product formed is copper hydroxide (Hultquist, Chuah, & Tan, 1989): Cu(s) + H 2 O(l) CuOH(s) H 2(g) Equation 7 which could then be converted to copper oxide: 2CuOH(s) Cu 2 O(s) + H 2 O(l) Equation 8 It is known that Reaction 7 may occur, however is limited to only the first atomic layer of copper. The corrosion model proposed by (Hultquist et al., 2015) suggests that corrosion could exceed the outermost layer of copper, due to hydroxide and hydrogen penetration into the bulk 11

20 of the metal through grain boundaries. The hydrogen partial pressure required to counteract the dissociation of copper was found to be approximately 1 mbar at 73 C (Hultquist et al., 2009), and a corrosion rate of approximately 0.1 µm/year at room temperature was concluded (Hultquist et al., 2013). Some authors have stated that corrosion of copper by water is indeed possible at very low hydrogen pressures, approximately mbar, and Cu + concentrations (Macdonald & Sharifi- Asl, 2011; Scully & Edwards, 2013). However other studies, such as those done by (Eriksen, Ndalamba, & Grenthe, 1989) and (Simpson & Schenk, 1987) were unsuccessful in demonstrating this. The contradiction in literature led to further experimentation on the copper-pure water system at the University of Toronto. These results will be discussed in this thesis Corrosion of Copper in Anoxic Water with High Chloride Concentrations Chloride ions (Cl - ) in Canadian groundwater may influence the corrosion of copper by forming stable Cu-Cl species during hydrogen evolution. The overall reaction is (Scully & Edwards, 2013): Cu(s) + ncl + H 2 O(l) CuCl n + OH + 1 H 2 2(g) Equation 9 More specifically, the following are most likely to occur at repository salinities: Cu(s) + 2Cl + H 2 O(l) CuCl 2 + OH H 2(g) Equation 10 Cu(s) + 3Cl + H 2 O(l) CuCl OH H 2(g) Equation 11 With increasing Cl - concentration successively higher Cu-Cl complexes become increasingly more stable. Pourbaix diagrams on the copper-chlorine-water system, modeled by (Beverskog & Puigdomenech, 1998), demonstrated that: 12

21 1. At [Cl - ] = 0.2 molal, CuCl2 - species predominate at all temperatures and corrosion occurs at 100 C. 2. At [Cl - ]= 1.5 molal, CuCl3 2- species predominate at 5-25 C and 100 C, while CuCl2 - corrosion products predominate at C, and corrosion occurs at all temperatures. 3. At [Cl - ]=5 molal, CuCl3 2- species predominate and copper corrodes at all temperatures (Beverskog & Pettersson, 2002). Thus, in high concentrations and especially at elevated temperature, chloride is detrimental to the immunity of copper. This is particularly a concern in the highly saline groundwater of the Ordovician sedimentary rock, where Cl - concentrations can reach maximum values of nearly 5.6 molal (Table 1) at elevated temperatures. The Pourbaix diagram modeling this worst case, can be seen in Figure 4; note the hydrogen line crosses into the passivating region, indicating corrosion indeed occurs in the presence of hydrogen gas evolution. 13

22 Figure 4: Pourbaix diagram for copper in 5 molal [Cl - ] at 75 C, [Cu] =1 mmolal and 1µmolal (Senior, Newman, Abghari, & Keech, 2013). Generated using the software Medusa, written by Puigdomenech Corrosion of Copper in Anoxic Water with Sulfides Sulfate reducing bacteria may exist in the repository environment. Sulfide forms very stable corrosion products with copper, such as Cu2S, which can reduce the reduction potential of copper metal to such an extent that the hydrogen ions present in water become adequate electron acceptors and cause copper to oxidize (Puigdomenech & Taxén, 2000): 2Cu(s) + HS + H + Cu 2 S(s) + H 2 (g) Equation 12 Although sulfate reducing bacteria have not been found in host DGR groundwater, as a precautionary measure, the NWMO is considering their effect at low concentrations, 10-5 M (Scully & Edwards, 2013). 14

23 2.6.3 Localized Corrosion Localized attack in an otherwise resistant surface results in pitting corrosion. Metals which form passive films are especially susceptible to pitting, due to local breakdown of the film at isolated sites (Jones, 1996). In a repository environment pitting corrosion of copper was considered during both the aerobic and anaerobic phases (Scully & Edwards, 2013). During the aerobic phase, upon closure of the repository, traces of O2 may be present for the next 100 years, and pitting induced by Cl - anions must be considered. However this mechanism is unlikely to occur due to the limited supply of oxygen. During the anaerobic phase, once oxygen and passive films have been fully consumed, the bare copper surface will be exposed to the environment. Traditional pitting is unlikely to occur, however a study done by (King, Litke, & Ryan, 1992) found that non-uniform general corrosion is possible. They exposed copper samples to groundwater-saturated buffer material and examined their stripped surface profiles after nearly 2 years; an example profile is shown in Figure 5. They observed that the entire surface exhibited roughening and had corroded to different extents. Therefore this is not an example of pitting in the usual sense of permanently separated anode and cathode sites, but rather a milder form of under-deposit corrosion causing surface roughening (King, Ahonen, Taxén, Vuorinen, & Werme, 2002; King, 2007). 15

24 Figure 5: Surface profile of stripped copper coupon following exposure to groundwater-saturated buffer material at 50 C for 733 days (King et al., 2002) Microbial Corrosion Copper can be susceptible to corrosion in the presence of microbial metabolites, particularly sulfate reducing bacteria. However at repository depths oxygen and nutrient levels are low and salinity is high, reducing chances of bacterial survival. Further, protection of the thick bentonite layer is expected to essentially eliminate microbial activity at the copper surface Stress Corrosion Cracking (SCC) Normally ductile metals subjected to tensile stress in specific environments may undergo SCC, forming brittle fractures and propagating cracks. SCC is highly chemically specific, in that certain metals will only undergo it if placed in environments with particular chemical species (Jones, 1996). 16

25 SCC of copper results due to rupture of a formed oxide film, which can be induced by ammonia, nitrite, or acetate, none of which are significantly present in Canadian groundwater. Hence SCC during the aerobic phase of the repository is unlikely to occur. 17

26 3 Research Objectives The aim of this research is to support the NWMO in gaining a better understanding of the corrosion behaviour of copper under simulated repository environments. To obtain this, the following objectives were outlined: 1. Determine whether corrosion of copper in pure anoxic water occurs by monitoring hydrogen evolution. 2. Determine the effect chlorides have on the corrosion rate of copper by monitoring hydrogen evolution in anoxic highly saline solutions, across various ph and temperatures relevant to the NWMO. 3. Examine the influence surface oxides have on the systems in Objectives 1 and 2. 18

27 4 Experimental Procedures This section outlines the materials, equipment and techniques used in this study to examine the corrosion behaviour of copper. 4.1 Hydrogen Evolution over Freely Corroding Copper In order to simulate the UFC-repository interface on a laboratory scale, copper wires were loaded into glass electrochemical cells (manufactured by Scientific Glass Design). For Objective 1, deionized anoxic water was used as the electrolyte, whereas chloride rich anoxic brine was used for Objective 2. Once filled, the cells were sealed and placed in temperature controlled waterbaths. As the corrosion of copper generates hydrogen gas, the sealed cells were allowed to accumulate H2 in their headspace. This hydrogen was routinely purged out and collected into a tedlar gas bag. The contents of the bag were subsequently analyzed with a hydrogen monitor and based on the quantity of hydrogen, the corrosion rate of copper was calculated. A schematic of the cell setup is shown below. 19

28 Figure 6: Schematic of electrochemical cell setup for measuring hydrogen evolution from copper wires Hydrogen Monitored Pure Water Cells To study the system in Objective 1 four simultaneous experiments were performed in pure water, modeling the Cu-H2O system under freely corroding conditions in electrochemical cells. The experiments were initiated at 30 C or 75 C, and were either deoxygenated or oxygenated. The deoxygenated cells contained oxygen free water from the initiation of the experiment, whereas oxygenated cells included a brief period of exposure to oxygenated water before being replenished with deoxygenated water for the remainder of the experiment. The purpose of the oxygenated cells was to mimic the real life transition of the repository environment from an aerobic state, when the UFC is initially placed, to an anaerobic state, when the repository has been sealed and oxygen has been consumed. 20

29 The four pure water cells have been monitored for approximately 2 years to date and are identified as follows: Table 2: Cell identification for hydrogen monitored pure water experiments. Cell Name Initial Condition Initial Temperature ( C) D1 Deoxygenated 30 D2 Deoxygenated 75 O1 Oxygenated 30 O2 Oxygenated Hydrogen Monitored High Salinity Brine Cells To study the system in Objective 2, the copper corrosion in deoxygenated chloride rich brines was investigated through several experiments of various ph, 3-9, and temperatures, C. ACS grade NaCl of 99.5% purity was used as the source of chloride. A detailed list of prior and completed experiments (ID numbers 6 to 23), containing experimental parameters and calculated corrosion, can be found in the thesis of (Abghari, 2013). There are four cells that were started during this research program that are still under investigation; these are listed in Table 3. For experiments 25 onwards the experimental procedure was slightly altered; the wires were enclosed in the cell for a minimum of one week after which the overhead gas was analyzed, if H2 was present the electrolyte was exchanged with an identical fresh electrolyte. The reason for this was to remove dissolved copper resulting from initial dissolution of the air-formed oxide film. Table 3: Cell identification for hydrogen monitored brine solutions. Cell ID Number Initial ph Initial Temperature ( C) [NaCl] (mol/kg)

30 4.1.3 Preparation of Copper Wires Copper wire, 0.25 mm in diameter and 99.9 % in purity (supplied by Goodfellow Metals, UK), was cut into 5 cm length pieces and assembled into bundles of 100. For each experiment 26 bundles, or 2600 wires, were used to maintain a surface area of 1024 cm 2. The high surface area was chosen to ensure detectable hydrogen production. The wires were cleaned via; immersion in acetone in an ultrasonic bath, following rinsing with deionized water (18 MΩ.cm), and immersion in 1 M nitric acid for two minutes of ultrasonication. The wires were subsequently rinsed repeatedly with deionized water, and finally ultrasonicated in ethanol. They were left to dry in a desiccator overnight before use Electrolyte and Cell Preparation The electrolyte and cell preparation procedure for pure water experiments was as follows: 1. Deionized water (18 MΩ.cm) was boiled for 15 minutes in an Erlenmeyer flask while being sparged with high purity nitrogen gas (below 10 ppb O2) for the deoxygenated tests, or pure oxygen (99.9%) gas for the oxygenated tests. 2. Prepared copper wires were loaded into an experimental cell. 3. The cell was subsequently purged with the test gas (nitrogen or oxygen) for 15 min. 4. The anoxic or oxic water was transferred from the gas saturated Erlenmeyer flask to the experimental cell through a syringe, until a volume of 150 ml was reached. 5. The cell headspace was further purged with test gas to ensure no contamination with air before the cell was sealed. For deoxygenated highly saline brine experiments the following steps were modified: ml of 0.25, 2.5, or 5 molal NaCl solution was prepared in a mixing cell using nitrogen (below 10 ppb O2) sparged anoxic water from the Erlenmeyer flask. The ph 22

31 of the solution was maintained using a buffer (acetate, disodium phosphate or citrate) and adjusted with HCl or NaOH. 5. The solution in the mixing cell was sparged with nitrogen for an hour before being transferred, via deoxygenated tubing, into the experimental cell. The cell headspace was further purged with nitrogen to ensure no air contamination before sealing. 6. The cells were placed in one of three waterbaths set at 30, 50, and 75 C Replenishing the Electrolyte Following the cell preparation outlined above, the pure water experiments underwent a preliminary treatment phase: 1. All four cells were placed in the 75 C waterbath for 1-2 weeks. The purpose of this was to dissolve air-formed oxides as rapidly as possible. 2. Following this period the cells headspaces were purged with nitrogen, the expelled gas being analyzed for hydrogen, none or trace quantities were detected. 3. For the oxygenated experiments, the cell headspaces were further purged with nitrogen for 5 minutes, in order to remove residual oxygen from the gas phase. Any dissolved oxygen in solution was allowed to be naturally consumed, as this occurs rapidly at elevated temperatures. 4. For deoxygenated experiments, the fluids within the cells were purged out using nitrogen flow, and fresh deoxygenated water was injected into the cells (in the same manner as described in Steps 4-5 of Section 4.1.4). The purpose of this was to remove any copper in solution resulting from dissolved oxide. 5. The cells were set to their experimental temperature; 30 or 75 C as described in Table 2. Similarly, for the highly saline brine experiments, in order to assess and minimize the effect of air-formed oxide on the hydrogen measurements, the following was performed for Experiments 25-29: 23

32 1. Following Step 6 in Section above, the cells were maintained at their corresponding temperatures for one week. 2. The cells headspaces were purged with nitrogen, the expelled gas being analyzed for hydrogen. 3. The electrolyte was purged from the cell using nitrogen flow and fresh electrolyte was injected into the cell (in the same manner as described in Steps 4-5 of Section 4.1.4). 4. The cells were placed back in their corresponding waterbaths Hydrogen Analysis As mentioned, sealed cells were left in a freely corroding state to accumulate hydrogen. This hydrogen was routinely evacuated with nitrogen flow and collected into a tedlar gas chromatography bag as seen in Figure 6. The sampling periods ranged between 2 weeks to 1 month, in order to determine whether sampling frequency influenced the equilibration of the system, however this did not prove to be a factor as longer intervals would be required. The hydrogen was analysed using a Hydrosteel 6000 hydrogen monitor (manufactured by IonScience, UK), which measures the hydrogen in gas using a solid-state electrochemical sensor (proton pump). The output of the sensor is given in unites of pl/cm 2 /s, because it is usually used to monitor hydrogen effusion from metals. The hydrogen sensor response was found to vary up to 5% due to sensor temperature, which itself is a function of ambient temperature and prior use. To correct for this two calibrations were performed; one immediately before the start of a sample, and one immediately after. The calibrant used was certified gas, nominally ppm H2 in nitrogen. An example of a typical analysis is presented in Figure 7. Following the first check against the first calibration, the sensor was left to return to a reading of zero before analysis of the sample. After the sample, 24

33 the sensor was again left to return to zero before the second calibration was performed. The second calibration corrects for the drift of the hydrogen sensor over the sample analysis period. Integration of the sample reading curve yields the total hydrogen volume collected from the tedlar bag, which can be converted to moles through the ideal gas law. Figure 7: Sample reading from hydrogen probe with calibrations. 25

34 5 Results and Discussion 5.1 Converting Cumulative Hydrogen to Copper Corrosion Rate By utilizing the assumption that copper corrosion generates hydrogen via the mechanism: 2Cu(s) + 2H + 2Cu + + H 2 (g) Equation 13 and that corrosion is uniform, the corrosion rate of copper can be calculated: Corrosion Rate ( nm year ) = dn H2 dt cm cm Equation 14 where; 2 is the Cu to H2 stoichiometric factor, 7.11 cm 3 is the molar volume of copper, 1024 cm 2 is the wire surface area, 10 7 converts from cm to nm, and converts from days to years. The differential term is determined by plotting the cumulative moles of hydrogen, determined by the method described in Section above, against the experimental duration. A curve fit is then applied to these data points, typically linear or an exponential association, and differentiated to give dn H2 dt Corrosion Measurements in Pure Water Cells This section includes the results from the pure water experiments after 350 days, presented in graphs of cumulative hydrogen versus experimental duration and their calculated corrosion rates. 26

35 Corrosion of Bare Copper in Anoxic Pure Water The cumulative hydrogen generation and calculated corrosion rates at 75 C for cells D1 and D2 are presented in Figure 8 and Figure 9, respectively. For the first 129 days cell D1 showed no hydrogen production at 30 C and 50 C, this timeframe is hence excluded from the plot. The cell was subsequently transferred to the 75 C waterbath and immediately started producing H2, continuing over a duration of 218 days. The hydrogen production shows a slight non-linear trend, corresponding to a corrosion rate below 2 nm/year. Cell D2 demonstrated consistent and linear hydrogen generation from its initiation at 75 C and over 347 days, corresponding to a nearly constant corrosion rate of 0.83 nm/year. Near the end of this monitoring period it was observed that the initially shiny red-orange wires were discolouring and turning a medium brown color, possibly indicating formation of Cu2O or CuOH on the surface, or, microscopic roughening resulting in absorption of light. 8.0x Cumulative Hydrogen (moles) 6.0x x x Cumulative Hydrogen Corrosion Rate Experimental Duration at 75 C (days) Corrosion Rate (nm/year) Figure 8: Cell D1 Cumulative hydrogen and calculated corrosion rate of bare copper wires immersed in deoxygenated pure water at 75 C, following testing at 30 and 50 C. 27

36 8.0x Cumulative Hydrogen (moles) 6.0x x x Cumulative Hydrogen Corrosion Rate Corrosion Rate (nm/year) Experimental Duration at 75 C (days) Figure 9: Cell D2- Cumulative hydrogen and calculated corrosion rate of bare copper wires immersed in deoxygenated pure water at 75 C Corrosion of Oxide Covered Copper in Anoxic Water Following the pre-treatment of cells O1 and O2 to oxygenated water at 75 C for 1 week, the wires turned black within hours, indicating the rapid formation of CuO on the surface. During this time no hydrogen was observed due to oxygen consumption. Following this oxygenated period the cells headspaces were purged with nitrogen to remove any residual oxygen from the gas phase (the solution was left to consume dissolved O2), and cell O1 was set to 30 C, while O2 was set to 75 C. Similarly to cells D1 and D2, cell O1 produced no hydrogen at 30 C and 50 C, however showed immediate hydrogen generation when transferred to 75 C, corresponding to a corrosion rate below 1 nm/year, as seen in Figure 10. Oddly, cell O2 produced no hydrogen over 350 days at 75 C; the reason for this is not understood. After extended duration at 75 C 28

37 it was observed that the black wires turned orange-brown, possibly due to the transition of CuO to Cu2O and CuOH, but the exact composition of the corrosion product must be investigated in future work. 2.0x Cumulative Hydrogen (moles) 1.5x x x Cumulative Hydrogen Corrosion Rate Experimental Duration at 75 C (days) Corrosion rate (nm/year) Figure 10: Cell O1 - Cumulative hydrogen and calculated corrosion rate of initially-oxidized copper wires immersed in deoxygenated pure water at 75 C, following testing at 30 and 50 C. The deoxygenated cells produced over double the amount of total hydrogen compared to the oxygenated cells, though both were on the µmoles scale. This is most likely due to the protection of the oxide layer, probably a mix of CuO and Cu2O, formed during the initial oxic stage of the oxygenated experiments. In general, all pure water cells exhibited slow copper corrosion rates below 1 nm/year, over a time scale of 350 days. After 350 days, there was a series of waterbath drying events which caused a significant increase in corrosion rate. This will be discussed in Section

38 Monolayers of Copper Consumed The cumulative hydrogen from the plots above can also be used to calculate the monolayers of copper that have reacted on the wires. This calculation is given in Appendix B. Recall the reaction from Equation 7: Cu(s) + H 2 O(l) CuOH(ads) + 1 H 2 2(g) This is theoretically a surface phenomenon, meaning that H2 production should cease when only the outermost monolayer of copper is consumed by hydroxyl ions. However if the number of monolayers exceeds one, this would imply a new mechanism in which the corrosion is a penetrating process (Senior, Newman, Abghari, & Keech, 2013). One explanation for this scenario could be that the OH groups adsorbed on the copper surface can move into the bulk of the metal through grain boundaries, providing free Cu surface for continuous CuOHads complexation, and hence continuous hydrogen evolution (Belonoshko & Rosengren, 2012). Cells D1 and D2 each produced 5.9 x 10-6 moles of hydrogen over the testing period, corresponding to 3.9 monolayers of copper consumed. Cell O1 produced 2 x 10-6 moles of hydrogen, corresponding to 1.3 monolayers of copper. This implies that the corrosion is indeed a penetrating process, and not a short term surface phenomenon Corrosion Measurements in Highly Saline Brine Cells The thesis of (Abghari, 2013) can be referred to for details on prior experiments in concentrated chloride solutions. Figure 11 is a 3D compilation of all stable corrosion rates recorded to date; note experiment 29 is excluded because it did not generate a stable corrosion rate due to 30

39 overheating (more on this in the following section). It can be seen that corrosion rates are highest at 75 C, while much lower at 30 and 50 C. The ph does not appear to have as great an influence as temperature. Further, while the Pourbaix diagrams created by (Beverskog & Pettersson, 1998, 2002) predicted that increasing NaCl concentration would result in higher corrosion rates (see Section ), this was not necessarily observed. Rather, at 75 C and ph 5, the highest corrosion rate observed was at 0.25 molal NaCl (Experiment 27), as opposed to 2.5 and 5 molal solutions of equivalent ph and temperature (Experiments 28 and 25, respectively). This may be because an excess of Cl - ions can accumulate on the copper surface, ultimately reducing corrosion. In general, corrosion rates in saline brines did not exceed 10 nm/year. Corrosion Rate (nm/year) 10 8 [NaCl] (mol/kg) ph Temperature ( C) Figure 11: Corrosion rates, as a function of ph and temperature, of all saline brine experiments to date. 31

40 5.2 Overheating Period There was a series of waterbath drying events after the stable corrosion rate period, where the water in the waterbath evaporated, leaving the cells in direct contact with the heating element, however the volumes of the electrolytes within the cells were not affected, due to the tight seal of the cell. This caused an unknown but significant rise in temperature, which resulted in an increased rate of hydrogen production in all of the ongoing experiments. When temperature control was restored, by refilling the baths with room temperature water and allowing them to reach their set temperatures, the hydrogen production subsequently ceased. This may have occurred, in an additional 2-3 years perhaps, had the experiments continued to be at constant temperatures and were allowed to equilibrate, reaching a saturation limit of Cu + in solution. However with the waterbath drying events, rapid heating and cooling of the glass cells (when the bath was refilled) may have created a point of ingress for oxygen. Oxygen within the system could stop hydrogen production by the following manners: 1. In the case of the brine experiments, Cu is oxidized by O2, which can raise the quantity of Cu + in solution to an extent were the copper reduction potential is sufficiently above H2 evolution. 2. In addition, for the pure water experiments, it is also possible that O2 can hinder H2 production by forming a sufficiently thick oxide layer on the wires. If the former is true, then hydrogen should be restored once the solution in the cell has been exchanged with fresh electrolyte. If the latter is true, then hydrogen production would not regenerate even after the solution has been refreshed. 32

41 Since the hydrogen production is dependent on temperature, the corrosion rates during the overheating period could not be calculated as the temperature was unknown. The corrosion rates of these experiments during the controlled temperature interval were discussed in Section Pure Water Experiments Overheating Figure 12 compiles all pure water experiments, and presents the cumulative hydrogen that was generated from the start of the experiments and throughout the overheating period. The steady corrosion rates of experiments D1, D2 and O1 at 75 C, over the first 347 days (before overheating), were calculated in Section Recall that cells D1 and O1, initially started at 30 C, only started producing hydrogen when transferred to 75 C on day 129, as seen in Figure 12. Cell O2, which was initiated at 75 C, showed no hydrogen until overheating. During the overheating period, it can be seen that all experiments indeed experienced a sudden jump in hydrogen production, as they all shared the same 75 C waterbath. Even cell O2, which had not produced any hydrogen from its initiation at 75 C, started to during overheating. The jump is more notable for the deoxygenated cells as opposed to the oxygenated. A key factor here is the surface condition of the copper at the start of the experiments; the initial oxygen in the oxygenated cells contributed to the formation of an effective oxide layer. All experiments subsequently stopped producing hydrogen. To understand the reason for this, the electrolyte in each cell was replaced with fresh deoxygenated deionized water (by method described in Section 4.1.5), marked by the symbol ( ) in Figure 12, and was assayed for copper. The results of the assayed solutions are presented in Section 5.3. Generally, the copper concentration was either below the detection limit, 0.05 µg/ml, or marginally over. Further, after refreshing the 33

42 electrolyte, the cells did not regenerate hydrogen. A thick oxide layer may be hindering H2, more on this in Section x10-5 Overheating D2 Cumulative Hydrogen (moles) 1.6x x x x D1 O1 O Experimental Duration (days) Figure 12: Compilation of cumulative hydrogen produced in all pure water experiments, after preliminary treatment. ( ) indicates refreshing of electrolyte Brine Experiments Overheating Figure 13 compiles brine experiments 25 and 28, and presents the cumulative hydrogen that was generated from the start of the experiments and throughout the overheating period. As mentioned, these experiments were initiated at 30 or 50 C, though hydrogen was not detected at these temperatures. They were subsequently transferred to the 75 C waterbath and immediately started producing hydrogen. The steady corrosion rates at 75 C were below 8 nm/year. The cells were placed back in the 50 C waterbath on day 303, to test whether 34

43 hydrogen production would again terminate, and indeed this was observed. Day 400 was when this bath dried out, resulting in increased hydrogen evolution. Once temperature control at 50 C was restored, it was increased to 75 C to test whether hydrogen would evolve, however as in the pure water experiments, it had completely ceased. The solutions in the cells were refreshed and the cells were brought back to 50 C. The results of the assayed solutions are presented in Section 5.3. Unlike the pure water results, significant copper was detected in solution, up to 262 µg/ml. After a few weeks, the cells started producing hydrogen again. This did not occur immediately after flushing as oxygen that may have been reintroduced into the system needed to be consumed. What is of great interest is that hydrogen was now evolving at 50 C; the reason behind this will be investigated in future work. 6.0x10-5 Overheating 50 C 75 C 50 C Cumulative hydrogen (moles) 5.0x x x x x C 50 C 75 C 50 C Exp. 28 Exp Experimental duration (days) Figure 13: Cumulative hydrogen produced in brines for experiments 25 and 28. Figure 14 compiles the overheating results for experiments 27 and 29. Experiment 27 produced no hydrogen at 30 and 50 C, but maintained a constant corrosion rate at 75 C until the bath 35

44 went dry. As Experiment 29 was started at 75 C on day 376, just prior to overheating, it did not get an opportunity to generate a stable corrosion rate. Similarly to Experiments 25 and 28, after refreshing the solution Experiment 29 started producing hydrogen. This implies that the hydrogen production ceased after overheating due to copper accumulation within the solution, likely caused by oxygen ingress. Experiment 27 was not refreshed as it is under current observation. 1.4x x10-4 Overheating 75 C Cumulative Hydrogen (moles) 1.0x x x x x C 50 C 75 C Exp. 27 Exp Experimental Duration (days) Figure 14: Cumulative hydrogen produced in brines for experiments 27 and Solubility of Copper Oxides Following the overheating period samples of electrolyte were analyzed for dissolved copper. This was performed at ANALEST, the analytical facility at the University of Toronto. The copper concentration was assayed via Inductively Coupled Plasma Atomic Emission Spectrometry (ICP AES), with a detection limit of 0.05 µg/ml. This dissolved copper 36

45 concentration, denoted [Cu]total, can be compared to the calculated copper based on hydrogen measurements, denoted [Cu]corr. Any discrepancy between these two values must come from dissolved oxides, formed during wire preparation or oxygen ingress Solubility of Copper Oxides in Pure Water As discussed in Section , copper species in pure water can result from the corrosion reaction via Equation 6, or Equation 7 and Equation 8, as well as from oxides (CuO and Cu2O) formed in air during wire preparation, or during any oxygen ingress. The copper concentrations of the assayed solutions, [Cu]total, are compared to the amounts calculated through the hydrogen measurements, [Cu]corr, below: Table 4: Dissolved copper concentration from assayed pure water solutions, [Cu] total, compared to the calculated copper based on hydrogen measurements, [Cu] corr. *BDL=below detection limit of mm. Experiment [Cu] total (mm) [Cu] corr (mm) Total Monolayers of Cu Consumed O1 initial final N/A D1 initial BDL 0 final BDL D2 initial BDL 0 final Cell O2 is still under investigation and was not assayed. For the deoxygenated experiments, D1 and D2, an attempt was made to measure the monolayers of copper consumed by initial air-formed oxides; by expelling and analyzing the electrolyte after 14 days from the start of the experiments. However this could not be done since the initial assayed solutions presented no detectable copper. All of the final assayed solutions, taken after the overheating event, presented [Cu]total values orders of magnitude below [Cu]corr, meaning that the corrosion products, possibly CuOH and 37

46 Cu2O, have very low solubility in pure water and remain on the wires. A study done by (Palmer & Bénézeth, 2008) investigated the solubility of copper oxides in pure water, and determined the solubility limit of Cu2O to be approximately mm at neutral ph and C. This agrees with the assayed solution which were below the detection limit of mm. Assuming [Cu]total is negligible, the number of monolayers of copper consumed was determined through the total hydrogen evolved, as seen in Table 4. The monolayers consumed throughout the duration of the experiments are all greater than unity, ranging between 2 to12, implying that the corrosion process is indeed an ongoing penetrating mechanism Solubility of Copper Oxides in Chloride Rich Brines Copper species in chloride rich brines can originate from either of the following mechanisms: From dissolved oxides, [Cu]oxide, either initially formed in air: CuO(s) + Cu(s) + 6Cl + 2H + 2(CuCl 3 ) 2 + H 2 O(l) Equation 15 Cu 2 O(s) + 6Cl + 2H + 2(CuCl 3 ) 2 + H 2 O(l) Equation 16 or during the experiment because of oxygen ingress: 4Cu(s) + O 2 (g) + 4H Cl 4(CuCl 3 ) 2 + 2H 2 O(l) Equation 17 and from corrosion with hydrogen production, [Cu]corr: Cu(s) + 3Cl + H + (CuCl 3 ) H 2 2(g) Equation 18 Note that Equation 15 to Equation 16 delay the onset of reaction Equation 18 until they have reached completion. 38

47 The analyses of Experiments is presented below. Note that [Cu]oxide is simply determined by subtracting [Cu]corr from [Cu]total. Table 5: Dissolved copper concentration from assayed saline solutions, [Cu] total, compared to the calculated copper based on hydrogen measurements, [Cu] corr, and dissolved oxides, [Cu] oxide. Depth of copper [NaCl] Temperature [Cu]total [Cu]corr [Cu]oxide Experiment dissolved from (mol/kg) ( C) (mm) (mm) (mm) [Cu]oxide (nm) 25 initial final overheat initial final overheat initial final 75-overheat initial final 75-overheat Primarily it can be seen that these experiments contain significant amounts of copper in solution relative to the pure water experiments. This is expected considering the tendency of the system to form stable Cu-Cl species in solution. However the discrepancy between [Cu]total and [Cu]corr in the final solutions is a surprising observation. The purpose of exchanging the electrolyte after one week was to rid the system of any air-formed oxides, and indeed evidence of this is present in the initial solutions as [Cu]corr is zero (or nearly), meaning the total dissolved copper originates entirely from [Cu]oxide. However dissolved oxide is still prevalent in the final solutions; here [Cu]total does not originate completely from Equation 18 as would be anticipated, but also from [Cu]oxide. This could imply that either one week was insufficient in dissolving all the air-formed oxide, or more likely, there was oxygen ingress after overheating. In general, these results indicate that the source of copper in the salt experiments comes from both dissolved oxides and corrosion. In some instances the contribution of dissolved oxides was several times more than corrosion. Further, accumulation of copper species in the 39

48 overheating experiments, likely due to oxygen ingress, ultimately stopped the corrosion reaction. In a repository, the bentonite clay and copper canister can be considered a closed system, as the clay barrier limits both inward diffusion of oxygen and outward diffusion of dissolved corrosion products (King, Lilja, & Vähänen, 2013). Thus corrosion of the copper canisters is expected to stop once the system has been saturated with dissolved copper and reached equilibrum. This would have likely been demonstrated had the experiments been under constant temperature control over an additional time duration. 40

49 6 Conclusions and Recommendations The purpose of this research was to support the NWMO in investigating the stability of copper, as it will be used as the coating for used fuel containers. After long-term experimentation on the corrosion behaviour of copper in simulated repository environments, the following can be concluded: 1. The corrosion of copper in pure deoxygenated water was demonstrated at rates below 1 nm/year at 75 C, assuming uniform corrosion, while no corrosion was detected at lower temperatures. Calculations on the monolayers of copper consumed proved that the corrosion mechanism was a penetrating process, and not a short lived surface phenomenon. Further, the corrosion rates were accelerated during an overheating period of unknown temperature. Once controlled temperature was restored, the corrosion rates for all the experiments ceased. Very little copper was found in the analyzed solutions, indicating the corrosion products, possibly CuOH and Cu2O, have a low solubility and stay on the copper s surface. Future work will involve characterizing the surface products on the wires, after corrosion in pure water over extended temperature control. This will be done through XPS and with the use of an Ultra High Vacuum Sample Transfer Device, which is scheduled to arrive at the University of Toronto in July The corrosion rates of copper in deoxygenated chloride rich solutions did not exceed 10 nm/year at 75 C, assuming uniform corrosion, while no corrosion was detected 41

50 initially at 30 and 50 C. The ph does not appear to influence the corrosion rate to the extent of temperature. Further, the overheating period accelerated the corrosion rates and likely introduced significant oxygen into the cells, which resulted in the accumulation of copper species within the systems, ultimately stopping corrosion. The corrosion of copper in the repository is expected to terminate once a dissolved copper solubility limit is reached, as the clay barrier will limit inward and outward diffusion. After replenishing the experiments with new electrolyte after the overheating period, some experiments started generating corrosion rates at 50 C. Why corrosion was occurring now and not when initiated at this lower temperature is not well understood and must be further investigated in future work. It is possible that undissolved oxides may have been hindering corrosion at the start of the experiment, but further investigation on this is needed. 42

51 7 References Abghari, B. (2013). Corrosion of Copper in Concentrated Aqueous Chloride Under Anaerobic Conditions (Master's Thesis). University of Toronto. Retreived from Assembly of First Nations. (n.d.). Nuclear Waste Technical Backgrounder Series Part Two: Used Nuclear Fuel and Long-Term Storage. Retreived from Belonoshko, A. B., & Rosengren, A. (2012). A Possible Mechanism of Copper Corrosion in Anoxic Water. Philosophical Magazine, 92(36), Beverskog, B., & Puigdomenech, I. (1998). Pourbaix Diagrams for the System Copper-Chlorine at C (Report No. 98:19). Nykoping, Sweden: Swedish Nuclear Power Inspectorate. Beverskog, & Pettersson, S. (2002). Pourbaix Diagrams for Copper in 5 m Chloride Solution (Report No. 02:03). Nykoping, Sweden: Swedish Nuclear Power Inspectorate. Eriksen, T. E., Ndalamba, P., & Grenthe, I. (1989). Short Communication: On the Corrosion of Copper in Pure Water. Corrosion Science, 29(10), Hultquist, G. (1986). Hydrogen Evolution in Corrosion of Copper in Pure Water. Corrosion Science, 26(2), Hultquist, G., Chuah, G. K., & Tan, K. L. (1989). Comments on Hydrogen Evolution from the Corrosion of Pure Copper. Corrosion Science, 29(11), Hultquist, G., Graham, M. J., Kodra, O., Moisa, S., Liu, R., Bexell, U., & Smialek, J. (2013). Corrosion of Copper in Distilled Water without Molecular Oxygen and the Detection of Produced Hydrogen (Report No. 2013:07). Stockholm, Sweden: Swedish Radiation Safety Authority. Hultquist, G., Graham, M. J., Kodra, O., Moisa, S., Liu, R., Bexell, U., & Smialek, J. L. (2015). Corrosion of Copper in Distilled Water Without O 2 and the Detection of Produced Hydrogen. Corrosion Science, 95, Hultquist, G., Szakálos, P., Graham, M. J., Belonoshko, A. B., Sproule, G. I., Gråsjö, L., Dorogokupets, P., Danilov, B., AAstrup, T., Wikmark, G., Chuah, G.K., Eriksson, J.C., Rosengren, A. (2009). Water Corrodes Copper. Catalysis Letters, (132), Jones, D. A. (1996). Principles and Prevention of Corrosion (2nd ed.). New Jersey: Prentice-Hall Inc. Keech, P. G. (2015). An Overview of the NWMO Used Fuel Container [Internal Powerpoint Slides]. Toronto, Canada: Nuclear Waste Managment Organization. 43

52 King, F. (2007). Status of the Understanding of Used Fuel Container Corrosion Processes- Summary of Current Knowledge and Gap Analysis (Report No. TR ). Toronto, Canada: Nuclear Waste Managment Organization. King, F. (2008). Theory Manual for the Copper Corrosion Model for Uniform Corrosion (Report No. TR ). Toronto, Canada: Nuclear Waste Managment Organization. King, F., Ahonen, L., Taxén, C., Vuorinen, U., & Werme, L. O. (2002). Copper Corrosion under Expected Conditions in a Deep Geological Repository (Report No ). Eurajoki, Finland: Posiva. King, F., Lilja, C., & Vähänen, M. (2013). Progress in the Understanding of the long-term Corrosion Behaviour of Copper Canisters. Journal of Nuclear Materials, 438(1-3), King, F., Litke, C., & Ryan, S. (1992). A Mechanistic Study of the Uniform Corrosion of Copper in Compacted Na-Montmorillonite/Sand Mixtures. Corrosion Science, 33(12), Macdonald, D. D., & Sharifi-Asl, S. (2011). Is Copper Immune to Corrosion When in Contact With Water and Aqueous Solutions? (Report No. 2011:09). Swedish Radiation Safety Authority. Myers, J. R., & Cohen, A. (1984). Conditions Contributing To Underground Copper Corrosion. American Water Works Association Journal, 76(8), Nuclear Waste Managment Organization. (2012). Description of Canada s Repository for Used Nuclear Fuel and Centre of Expertise. Toronto, Canada. Retreived from Nuclear Waste Managment Organization. (2013). Implementing Adaptive Phased Management 2013 to Toronto, Canada. Retrieved from _oct23.pdf Palmer, D. A., & Bénézeth, P. (2008). Solubility of Copper Oxides in Water and Steam. Proceedings of the 14th International Conference on the Properties of Water and Steam (pp ). Oak Ridge, USA: Oak Ridge National Laboratory. Puigdomenech, I., & Taxén, C. (2000). Thermodynamic Data for Copper Implications for the Corrosion of Copper Under Repository Conditions (Report No. TR-00-13). Stockholm, Sweden: Swedish Corrosion Instituite. Scully, J. R., & Edwards, M. (2013). Review of the NWMO Copper Corrosion Allowance (Report No. TR ). Toronto, Canada: Nuclear Waste Managment Organization. Senior, N. A., Newman, R. C., Abghari, B., & Keech, P. G. (2013). Influence of ph, Temperature, and Salinity Upon Anoxic Copper Corrosion. Proceedings of the 16th International Conference on Environmental Degredation of Materials in Nuclear Power Systems-Water Reactors. Toronto, Canada: University of Toronto. 44

53 Simpson, J. P., & Schenk, R. (1987). Short Communication: Hydrogen Evolution from Corrosion of Pure Copper. Corrosion Science, 27(12), World Nuclear Association. (2015). Nuclear Basics. Retrieved April 23, 2015, from 45

54 Appendices Appendix A: Experiment Parameters and Details A1: Experiment Parameters Pure Water Experiments Experiment Initial Condition Preliminary Phase at 75 C Temperatures before Overheating ( C) D1 Deoxygenated Electrolyte replaced after 14 days D2 Deoxygenated Electrolyte replaced after 14 days 75 O1 Oxygenated Headspace purged after 7 days O2 Oxygenated Headspace purged after 7 days 75 Brine Experiments Experiment [NaCL] molal Buffer Used ph Temperatures before Overheating ( C) 25 5 Citrate Citrate Citrate Citrate A2: Cell Timelines The following timelines summarize the key events that occurred during the lifetime of each experiment. Brine Experiments Cell 25 Date (Duration) Comments (day 1) Cell started at 30 C (day 14) No H 2 detected; solution flushed from cell [ANALEST Certificate #: ANAL (2013/04/16)]. Replaced with new solution (day 38) No H 2 detected; Cell transferred to 50 C waterbath (day 48) H 2 detected sporadically on one measurement (day 56) to No H 2 detected (day 139) 46

55 (day 139) Cell transferred to 75 C waterbath; from now onward H 2 continuously detected (day 317) Cell placed back in 50 C; from now onward no H 2 detected (day 413) to Overheating period; waterbath evaporated once or twice, causing (day 532) unknown increase in temperature and H 2 production (day 532) After temperature was restored to 50 C; No H 2 detected from here onward (day 765) Cell placed in 75 C waterbath; No H 2 detected from here onward (835) Solution in cell flushed [ANALEST Certificate #: ANAL (2013/04/16)]. Replaced with new solution and placed cell in 50 C waterbath Cell 27 Date (Duration) Comments (day 1) Cell started at 30 C (day 10) Minimal H 2 detected; solution flushed from cell [ANALEST Certificate #: ANAL (2013/04/16)]. Replaced with new solution (day 35) No H 2 detected; Cell transferred to 50 C waterbath (day 136) No H 2 production for the last 100 days. Cell transferred to 75 C waterbath, from here onward H 2 continuously detected (day 410) to Overheating period; waterbath evaporated once or twice, causing (day 529) unknown increase in temperature and H 2 production (day 529) After temperature was restored to 75 C; decreasing H 2 detected from here onward (day 763) Only one port leading to the cell can be open; the other one is jammed. Limited operation is now possible (flushing of head space is OK, but flushing solution would introduce O 2) Cell 28 Date (Duration) Comments (day 1) Cell started at 50 ºC (day 6) Sporadic H 2 detected; No data logged into the excel sheet (day 7) Solution flushed from cell [ANALEST Certificate #: ANAL (2015/05/01)]. Replaced with new solution (day 12) No H 2 detected (day 26) Minimal H 2 detected. Cell transferred to 75 ºC waterbath; from now onward H 2 continuously detected (day 204) Cell placed back in 50 C; from now onward no H 2 detected (day 300) to (day 419) Overheating period; waterbath evaporated once or twice, causing unknown increase in temperature and H 2 production (day 419) After temperature was restored to 50 C; no H 2 detected from here onward (day 652) Cell placed in 75 C waterbath; no H 2 detected from here onward (day 670) Solution in cell flushed [ANALEST Certificate #: ANAL (2015/05/01)]. Replaced with new solution and cell placed in 50 C waterbath (day 721) H 2 detected at 50 C from here onward. 47

56 Cell 29 Date (Duration) Comments (day 1) Cell started at 75 ºC (day 26) High level of H2 detected but file number not recorded (day 34) Solution in cell flushed [ANALEST Certificate #: ANAL (2015/05/01)]. Replaced with new solution (day 49) Cell s headspace was purged with O 2 for 15 min (day 61) High level of H2 detected; meaning O 2 was consumed rapidly (day 90) to (day 209) Overheating period; waterbath evaporated once or twice, causing unknown increase in temperature and H 2 production (day 209) After temperature was restored to 75 C; decreasing H 2 detected from here onward (day 512) Solution in cell flushed [ANALEST Certificate #: ANAL (2013/04/16)]. Replaced with new solution. Pure Water Cells Cell D1 Date (Duration) Comments (day 1) Preliminary phase; Cell started with deoxygenated water in 75 C waterbath for 14 days in order to dissolve air-formed oxide (day14) Cell purged with N 2 and H 2 detected. Solution in cell flushed [ANALEST Certificate #: ANAL (2013/04/16)]. Replaced fresh deoxygenated water. Cell placed in 30 C waterbath (day 63) No H 2 detected so far so cell placed in 50 C waterbath (day 143) Only minimal H 2 detected, cell placed in 75 C; from now onwards H 2 continuously detected (day 389) to Overheating period; waterbath evaporated once or twice, causing (day 536) unknown increase in temperature and H 2 production (day 536) After temperature was restored to 75 C; decreasing H 2 detected from here onward (day 775) Solution in cell flushed [ANALEST Certificate #: ANAL (2015/05/01)]. Replaced with new deoxygenated water. From here onwards no H 2 was detected at 75 C. Cell D2 Date (Duration) Comments (day 1) Preliminary phase; Cell started with deoxygenated water in 75 C waterbath for 14 days in order to dissolve air-formed oxide (day14) Cell purged with N 2 and slight H 2 detected. Solution in cell flushed [ANALEST Certificate #: ANAL (2013/04/16)]. Replaced fresh deoxygenated water. Cell remains in 75 C waterbath (day 21) H 2 detected from here onwards 48

57 (day 389) to Overheating period; waterbath evaporated once or twice, causing (day 536) unknown increase in temperature and H 2 production (day 536) After temperature was restored to 75 C; decreasing H 2 detected from here onward (day 839) Solution in cell flushed [ANALEST Certificate #: ANAL (2013/04/16)]. Replaced with fresh deoxygenated water solution. No H 2 detected from ere onward. Cell O1 Date (Duration) Comments (day 1) Preliminary phase; Cell started with oxygenated water in 75 C waterbath for 7 days, wires turned black within hours (day 7) Cell s headspace purged with N 2, expelled gas analyzed for H 2, none detected. Headspace purged with N 2 for 5 min, solution was NOT replaced. Cell was left to cool in air for an hour then placed in 30 C (day 60) After no H 2 detection, cell placed was place in 50 C waterbath (day 140) After no H 2 detection, cell was placed in 75 C waterbath; from here onwards H 2 was continuously detected (day 386) to Overheating period; waterbath evaporated once or twice, causing (day 533) unknown increase in temperature and H 2 production (day 533) After temperature was restored to 75 C; decreasing H 2 detected from here onward (day 828) Solution in cell flushed [ANALEST Certificate #: ANAL (2015/05/01)]. Replaced with new deoxygenated water. From here onwards no H 2 was detected at 75 C. Cell O2 Date (Duration) Comments Preliminary phase; Cell started with oxygenated water in 75 C waterbath for 7 days, wires turned black within hours (day 7) Cell s headspace purged with N 2, expelled gas analyzed for H 2, none detected. Headspace purged with N 2 for 5 min, solution was NOT replaced. Cell remained in 75 C (day 18) H 2 detected from here onwards (day 386) to Overheating period; waterbath evaporated once or twice, causing (day 533) unknown increase in temperature and H 2 production (day 533) After temperature was restored to 75 C; decreasing H 2 detected from here onward. 49

58 Appendix B: Experimental Procedures B1: Wire Preparation Procedure Revision History Author Date Procedure ID 0 NS 12/10/2011 DS257-exp-pro-wireprep 1 NS 07/01/ Overview The following procedure details the necessary steps to accurately and reproducibly manufacture the wire bundles required for the hydrogen generation experiments. 2. Materials Copper wire of 0.25 mm diameter and 99.5% purity, supplied by Goodfellow, is to be used. 3. Manufacture of Wire Bundles The objective is to produce wire bundles, each containing 100 wires of 50 mm length. 26 bundles are required per experiment. 3.1 Equipment Requirements Wire, 5 mm ID PTFE tubing (20 sections ~2 cm in length and 52 rings ~0.5 mm length), elastic bands, scissors. 3.2 Health and Safety Labcoat, eye protection and gloves to be worn at all times. 3.3 Procedure 1. Insert the two screws in the indicated holes on the preparation desk in WB129 (110 cm apart). 2. Tying the copper wire onto one screw, loop the wire across the screws a total of 25 full times (i.e. 50 lengths of wire). 3. Tie the wires together using elastic bands. 4. Ensure that there are 20 PTFE tube sections (>2 cm length) to hand. 5. Using the scissors marked 'wire' and the affixed 1 meter ruler as guide, hold the wires and cut at the 1 m mark. 6. Slip a PTFE tube section over the cut ends onto the wire. 7. Using the 1 meter rule as the guide, cut the 50 mm wire length. 8. Repeat steps 6 and 7 until no spooled wire is left. 9. Two sections of wire are required to produce a single bundle. Combine the wires manually and loosely bind using two PTFE rings. 10. Store the 26 wire bundles in a sealed container marked with the experiment number. 11. Remove the screws from the desk and store. 4. Surface Preparation You will require gloves, acetone, fresh deionised water and ~100 ml 1 M nitric acid Mark a weighing boat with the experiment number and weigh, recording the result on the experimental form. 2. Zero the balance. 50

59 3. Taking the binding rings off the bundles, transfer the wire bundles onto the weighing boat and record the result to 3 decimal places. 4. Transfer the wires into a clean glass beaker and immerse with acetone. 5. Ultrasound for 30 minutes, recording start and finish times, and swilling the wires periodically. 6. Drain the acetone into the waste acetone bottle. 7. Wash the wires with deionised water (18.2 MΩ conductivity) and drain. 8. Immerse the wires with fresh deionised water and ultrasound for three minutes (always based upon a clock, not the ultrasound timer). Drain. Repeat three times. 9. Ultrasonicate the wires in the 1 M nitric acid for two minutes, swilling continously. 10. Drain the acid into the waste nitric acid bottle (stored under the fume hood). 11. Rapidly soak with de-ionised water three times (no ultrasound). 12. Immerse the wires in ethanol, ultrasound for 5 minutes, drain. 13. Drain the ethanol and spread the wires on tissue paper, on large weighing boats. 14. Place the boats/wires in the designated desiccator. 15. In WB231, Purge the desiccator with nitrogen for 10 minutes, then vacuum the desiccator to evaporate the ethanol (vacuum pump is located in WB231). Leave for a minimum of one night. Wires can be prepared in advance and stored under vacuum in the dessicator. 16. Experimental details to be recorded in the Experimental Index file. B2: Cell Preparation Procedure for Pure Water Experiments Revision History Author Date Procedure ID 0 NS DS257-sensor-pro-cellprep-v1 1 NS Overview There are two fundamental cell types required by project DS257 the equilibrium cell, used to repeat literature work, and the corrosion cell, where the headspace is periodically purged and analyzed using the hydrogen sensor. The following procedure details the necessary steps to set up the corrosion cell. There are four corrosion cells to be used, covering two variables: 1. Each cell will be at a fixed temperature, either 30 or 75 C. 2. Cells will either start in anoxic fashion, or will have a preliminary aerobic phase. Both aerobic and anoxic tests have a preliminary phase the anoxic tests, to remove any copper cations that are present from the dissolution of air-formed oxide, and the aerobic tests, specifically to form a preliminary compact oxide film. 1.1 Anoxic tests The experimental apparatus is prepared in the following manner: 1. Wires are prepared as per procedure DS257-exp-pro-wireprep. 2. The wires are loaded into the cell. 3. The cell is purged of air, using high purity nitrogen. 4. The de-oxygenated de-ionized water (procedure DS257-exp-pro-deoxygen) is injected. 5. The cell is left to stand for one week at 75 C, whereupon the headspace is tested. 51

60 6. Should hydrogen be present, the cell is comprehensively purged of gas and water, and fresh de-oxygenated water injected. This is to remove any initial copper cation concentration, formed from the dissolution of copper oxide. 7. The cell is set at the experimental temperature. 1.2 Aerobic tests 1. Wires are prepared as per procedure DS257-exp-pro-wireprep. 2. The wires are loaded into the cell. 3. The cell is purged of air, using oxygen (purity of 4.7). 4. Oxygenated de-ionized water, produced following an identical procedure to the deoxygenated water, but using pure oxygen in place of nitrogen, is injected into the cell. 5. The cell is maintained at 75 C for one week. 6. The cell is then purged of oxygen using high purity nitrogen, and a hydrogen reading taken. 7. The cell is then set at the experimental temperature. Notes; both oxic and anoxic cells have a one week preparatory stage, where the cell is held at 75 C. 2. Materials Sensor test cell Copper wire of 0.25 mm diameter and 99.9% purity, supplied by Goodfellow, is to be used, and prepared following procedure DS257-exp-pro-wireprep. 200 ml of deoxygenated deionized water (procedure DS257-exp-pro-deoxygenation). 60 ml syringe. BIP Nitrogen cylinder. 3. Anoxic Cell Method 17. Load a batch of wires into a cell. 18. With both glass stop-cocks open, connect the cell to the nitrogen cylinder, and purge for 10 minutes; use a bubble trap to ensure no back-flow of air. 19. Seal the cell at the 3-way valve and disconnect the nitrogen. 20. Draw 60 ml of deoxygenated water into the syringe, and connect to the 3-way valve (Figure 1). 21. Inject 10 ml of water into the valve (to purge valve of air and fill with deoxygenated water), to waste from the upper connection (used to connect the nitrogen cylinder). Rotate the valve, and inject 50 ml into cell. Repeat, to inject 100 ml water into the cell. Inject a little nitrogen to empty cell glass tubing of water, and seal at the glass stop-cock. 22. Seal the other glass stop-cock and disconnect the bubble-trap. 23. Preliminary phase only: Leave for one week at 75 C, then test for hydrogen. If hydrogen is present, purge the contents of the cell in an anoxic fashion, discarding the electrolyte (there is no value in testing it). Re-prepare the cell by following steps 2. to 6. holding the cell upside down to expel the water. (Collect in mixing cell). 24. The cell is now running. To analyze, follow procedure DS257-cor-pro-analysis. 25. Upon conclusion of the experiment, follow procedure DS257-exp-pro-cellfinal. 4. Oxic cell Method 1. Load a batch of wires into a cell. 2. With both glass stop-cocks open, connect the cell to the oxygen cylinder, and purge for 10 minutes; use a bubble trap to ensure no back-flow of air. 3. Seal the cell at the 3-way valve and disconnect the oxygen. 4. Draw 60 ml of oxygenated water into the syringe, and connect to the 3-way valve (Figure 1). 52

61 5. Inject 10 ml of water into the valve (to purge valve of air and fill with oxygenated water), to waste from the upper connection (used to connect the oxygen cylinder). Rotate the valve, and inject 50 ml into cell. Repeat, to inject 100 ml water into the cell. Inject a little oxygen to empty cell glass tubing of water, and seal at the glass stop-cock. 6. Seal the other glass stop-cock and disconnect the bubble-trap. 7. Leave for one week at 75 C. 8. Transfer to the experimental temperature. 9. The cell is now running. To analyze the headspace, follow procedure DS257-cor-proanalysis. 10. Upon conclusion of the experiment, follow procedure DS257-exp-pro-cellfinal. B3: Solution Preparation Procedure for Brine Experiments Revision History Author Date Procedure ID 0 NS 8/03/ Overview This procedure details the steps required to prepare the stock solutions required for the brine experiments. The overall procedure is as follows: The NaCl is weighed into a mixing cell, to which a fixed volume of deionized water is added. Hydrochloric acid solution is added to provide the required ph (which is monitored concurrently), and finally, further deionised water is added to make a known volume of solution. The solution is then de-oxygenated through purging with high purity nitrogen gas, the final ph recorded, and the brine transferred to the experimental cells. 2. Solutions The experimental solution is to contain 2.5 or 5 molal chloride (added as NaCl), at various ph's between 0 and Hydrochloric acid solution The HCl solution is made by diluting 3 ml of concentrated HCl solution to 100 ml with deionised water. 2.2 Brine solution The quantity of NaCl required for a solution with molality n can be determined using the calculator provided in 'DS25-TH-lit-NaClmolal.xls'. For 5 m solutions, the quantity required is g, and for 2.5 m solutions, g, for 80 ml of solution (note, this is not the volume of water added, but the final volume of the solution). Weigh the analar grade NaCl into the mixing cell base and add the stirrer bar. Calibrate the ph probe against ph 2, 4 and 7 standards. Fit into the mixing cell lid. Transfer the mixing cell to a magnetic stirrer. Using the syringe marked 'water', inject 60 ml of deionised water into the cell. Start stirring the brine. 53

62 Using the syringe marked 'acid', inject a known quantity to lower the ph to a suitable value. As with the water, purge the mixing cell valve with the solution (1 ml in this case), and inject the 8 ml into the cell. Using a similar method, add 1 M NaOH and 0.1 M NaOH, as appropriate, to reach the desired ph. Record the volumes added. Add de-oxygenated water to bring the total volume to 80 ml, as described in the table below: Molality NaCl (g) H2O (ml) Connect the mixing cell to the nitrogen cylinder to purge the brine for one hour. Seal the cell and record the final ph (unstirred). B4: Hydrogen Collection Procedure Revision History Author Date Procedure ID 0 NS 25/04/2012 DS25-PHE-pro-H2collection-v1 1. Overview This procedure deals with the methodology for removing the accumulated hydrogen from a test cell prior to analysis. The process consists of two steps; 1. Sample bag connection: An empty tedlar bag is connected to the cell. 2. Sample transfer: Nitrogen is injected into the cell, directly from the cylinder, transferring accumulated hydrogen from the cell into the sample bag. The purpose underlying the use of a tedlar bag for the collection of the accumulated cell hydrogen is that it prevents the condensation of water within the hydrogen sensor. 2. Cell Design An overview of the cell design is presented in the figure below. The three-way valve is used as the outlet, to collect the sample gas. The metering valve is used as the inlet. 54

63 3. Bag connection 3.1 Sample Bag Preparation 1. The 'sample' bag is connected to the coiled silicone tubing, which in turn is connected to the 'inlet' (i.e. central position) of the cell outlet 3-way valve. 2. A 250 ml beaker is then filled with cold water, into which the coiled tubing is lowered. This is to prevent condensation within the sample tedlar bag. 3. The syringe labelled 'gas', is connected to the upper port of the 3-way valve, using a short section of silicone tubing. 4. The 3-way valve is directed upwards to connect the syringe and 'sample' bag. 5. The 'sample' bag (with an open valve), and the connected tubing are then evacuated using the syringe. 6. Once complete, the valve is rotated to face the cell, discharging the excess pressure of the cell into the sample bag. 3.2 Connection of the Nitrogen Line to the Cell 8. The nitrogen cylinder (5.0 grade, 2 ppm O 2) is fitted with a length of narrow-bore vinyl pipe with a 1/4" Swagelok fitting on the end. This fitting is to be connected, very loosely at first, onto the cell metering valve, which remains closed until step Open the cylinder, increase the pressure in the regulator to 5 psig and open the regulator needle valve, flowing nitrogen through the line and out of the Swagelok fitting, for a minimum of 30 seconds. This is to purge the pipe of any atmospheric oxygen. 10. Tighten the Swagelok valve. The gas will stop flowing once the line pressure stabilizes at 5 psig. 4. Sample Collection 1. Gradually open the cell metering valve to flush the cell with nitrogen. Once the 'sample' bag contains ~500 ml of gas, close the metering valve. 2. Isolate the cell from the 'sample' bag by rotating the 3-way valve 90. Close the valve on the tedlar bag also. 3. Disconnect the 'sample' bag and place aside for analysis. 4. Seal the nitrogen cylinder (cylinder valve, regulator and needle valve) and disconnect the pipe from the cell metering valve. 5. Record the time and any deviations from the procedure and observations in the laboratory book. 55

64 B5: Correcting for Sensor Drift Revision History Author Date Procedure ID 0 NS 25/4/2012 DS25-TH-pro-correctSensDrift-v0 1. Overview This document describes the calculation steps used to determine the corrosion rate of carbon steel under anaerobic conditions, through the production of molecular hydrogen (H 2). As described in DS25-PHE-pro-H2collection, the accumulated hydrogen from within a cell is purged using high purity nitrogen, and collected in a gas-impermeable (tedlar) bag. This gas is left to cool, to prevent condensation within the hydrogen sensor. The hydrogen probe has been demonstrated to display variability in terms of both initial offset and drift, and these must be corrected for in order to accurately determine the accumulated hydrogen. The analysis of the accumulated gas is preceded by a calibration period (against certified gas), to determine the offset. Similarly, a second calibration period is conducted to determine the sensor drift. A typical curve is presented in Figure 1. The three phases are recorded in a single file. Figure 15: Sampling phases; calibration, sample and calibration. 1.1 Step 1 The probe sensitivity is firstly corrected through determination of a multiplication factor; the probe is zeroed using atmospheric air and then used to sample certified gas (nominally 100 ppm H 2). As demonstrated in procedure DS21-TH-pro-unitconv, the certified gas, measured in ppm, can be converted directly to pl/cm 2 /s by dividing by For example, 100 ppm hydrogen should result in a reading of 307 pl/cm 2 /s on the probe, assuming that the sample gas is under atmospheric pressure. The ratio between the actual reading and the theoretical value is the offset, and every 56

65 recorded value is adjusted accordingly. The probe also displays drift, as demonstrated by the second calibration peak (Figure 2). Figure 2: Hydrogen sensor drift. The probe drift is corrected for by comparing the initial and final stable calibration values and using a linear gradient as a fractional correction factor for all recorded data. The hydrogen probe outputs the sensor temperature, which clearly increases with usage. No clear relationship between sensitivity and sensor temperature was found. Figure 16: Linear data adjustment. 57

66 2. Assumptions Inlet gas is at atmospheric pressure; elevated pressure increases the gas flow through the pump, increasing the hydrogen reading. To prevent this, the inelastic tedlar bags used for these experiments are only partially filled. Tests on internal pressure demonstrated negligible changes in pressure when the bag was partially filled. B6: Unit Conversion Calculator Revision History Author Date Procedure ID 0 NS 25/4/2012 DS25-TH-pro-unitConv-v0 1. Overview The purpose of this document is to detail the method to convert the hydrogen concentration, as recorded by the hydrogen sensor in pl/cm 2 /s, to other units; ppm and moles/second. 2. Conversion to ppm 1. The hydrogen sensor is operated in 'Low Temperature' mode; and ordinarily uses a gas collecting plate of 14.4 cm diameter (162.9 cm 2 ). Multiplication by this factor removes the cm -2 units, i.e [H 2 reading] (pl/s). 2. The probe operates at a flow rate of 0.5 ml/s. Division of the hydrogen concentration in pl/s yields: [H 2 reading] (pl/ml). 3. Division by 1000 yields pl/µl, i.e. ppm. [H 2] (ppm) = [H 2] (pl/cm 2 /s) However, the probe is internally calibrated and it is necessary to adjust the output against certified gas using a multiplication factor. 58

67 Appendix C: Sample Calculations C1: Monolayers of Copper Consumed Calculation parameters: Wire surface area: 1024 cm 2 Cu cell parameter: pm (fcc) Figure C1: Lattice cell for copper Table C1: The copper density along the indicated plane, and the hydrogen gas generation corresponding to a monolayer of that plane. Plane Copper atomic density Evolved H 2 per monolayer (atoms/cm 2 ) (moles/cm 2 ) (moles) (100) (110) (111) All calculation were made using the (111) plane. C2: Depth of Copper Dissolved from Oxides Calculation parameters: Electrolyte volume: 150 ml Molar volume of copper: 7.11 cm 3 /mol Dissolved copper oxide in solution: [Cu]oxide mm Conversation factor: 10 Depth of dissolved copper from oxides (nm) = [Cu] oxide mm 7.11 cm3 mol 1024cm