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HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry 1. The chemical composition of the ocean implies its potential role as an electrolyte. Identify the origins of the minerals in oceans as: - Leaching by rainwater from terrestrial environments: Rainwater leaches - minerals/ions from the soil and rocks that it flows over or seeps through. This water will eventually flow into the ocean via springs and rivers. Hydrothermal vents in mid-ocean ridges: Mid-ocean ridges occur where two of the tectonic plates of the Earth‟s crust slowly move apart releasing hot magma located below the water‟s surface. Liquid magma is released and where it meets seawater, the water becomes super-heated and results in the formation of hydrothermal vents. Outline the role of electron transfer in oxidation-reduction reactions. - REDOX reactions involve the transfer of electrons. - Loss of Electrons = Oxidation - Gain of Electrons= Reduction Identify that oxidation-reduction reactions can occur when ions are free to move in liquid electrolytes. - Electrolytes are substances that dissolve to form ions. An electrolyte solution contains ions that are free to move around, and is the perfect environment for REDOX reactions. Describe the work of Galvani, Volta, Davy and Faraday in increasing understanding of electron transfer reactions. Process information from secondary sources to outline and analyse the impact of the work of Galvani, Volta, Davy and Faraday in understanding electron transfer reaction. Scientist Work Impact - - Galvani Carried out experiments in His experiment turned out to Volta which freshly dissected frogs legs “twitched” when touched by different metal hooks and wires. He believed that electricity came from the frogs as “animal electricity. Suggested that the electricity making the frog twitched was coming from the metal and not the frog. He made huge amounts of electricity from a series of metal plates and paper soaked in salt water. be wrong. However, his experiment sparked tremendous interest in the study of electricity. Established the idea that chemical reactions could produce electricity. His battery „Volta‟ became successful and as the best way to make electricity. HSC- Stage 6 Shipwrecks, Corrosion and Conservation - - Davy Faraday 2 Unit Chemistry Studied the decomposition of chemical compounds by passing electricity through them -> “Electrolysis”. He decomposed water into its elements. He was the first to make samples of sodium and potassium by electrolysis of molten compounds. Established that a He invented equipment for making electricity charge and using it he was able to prove relationship b/w the mass of chemicals reacting and the amount of electricity involved in REDOX reaction. Also experimented with making electricity by magnetic means. battery made electricity from chemical reactions. Discovered many new chemical elements by electrolysis. Developed the technique of electrolysis. He turned electrochemistry into a quantitative science, rather than being merely descriptive. Sparked greater interest in further understanding of acids. Electromagnetism>invention of the electrical motor and generator. 2. Ships have been made of metals or alloys of metals. Account for the differences in corrosion of active and passivating metals. - Generally, the more active a metal is, the faster it will corrode. Passivating metals react with oxygen in the air to form a thin layer of the metal oxide on the surface. Oxide layer is impermeable to air and water and prevents further corrosion. Coating prevents any further reaction between the metal and oxidising agent. These metals tarnish and lose their shiny lustre when exposed to air. This oxide layer usually adheres strongly to the metal but can be removed by abrasion with steel wool or emery paper. - Identify iron and steel as the main metals used in ships. - End of 19th century wood gave way to iron and steel in shipbuilding. - Steel has the advantage of having a good mechanical strength, it is relatively hard and it can be rolled into sheets and pressed into desired shaped. Identify the composition of steel and explain how the percentage composition of steel can determine its properties. Gather and process information from secondary sources to compare the composition, properties and uses of a range of steels. - Steel is iron with small amounts of carbon in it. > 2% HSC- Stage 6 - Shipwrecks, Corrosion and Conservation 2 Unit Chemistry More the carbon, harder the steel, hence more suitable for structural uses, tools, etc. In pure iron the atoms can slide past each other relatively easily when a force is applies (soft). If there are small amounts of different atoms present, their different size disrupts the lattice of atoms and makes it more difficult for adjoining regions to slide past each other. Metal Composition Properties - Iron Pure element Soft & malleable Corrodes very slowly Hard, but brittle Corrodes rapidly - Cast Iron 3-4% Carbon - Structural Steel About 0.5% Carbon Very strong, but can be shaped when hot. Corrodes rapidly - Mild Steel Less than 0.2% carbon Fairly strong, but malleable & ductile. Corrodes rapidly. - Stainless Steel Up to 20% chromium and nickel Hard & shiny Resistant to corrosion Typical use No common use in modern times Cast engine blocks. Decorative iron “lacework” Reinforcing bars in concrete Girders in buildings and bridges Railway lines Car bodies Pipes Roofing steel Nuts & bolts Ship hulls Chains Surgical instruments Cutlery Food handling equipment Kitchen sinks Describe the conditions under which rusting of iron occurs and explain the process of rusting. 1) Oxidation: Iron atoms lose electrons. HSC- Stage 6 Fe Shipwrecks, Corrosion and Conservation → 2 Unit Chemistry Fe2+ + 2e 2) Reduction: Oxygen and water gain electrons O2 + 2H2O + 4e– → 4OH– 3) Electrons move from anode to the cathode through the iron. To form the overall reaction, the oxidation half-reaction must be multiplied by 2, so the electrons will balance. 2Fe(s) + 2H2O (l) + O2 (g) → 2Fe2+ (aq) + 4OH-(aq) 4) These ions will migrate through any moisture present. When they meet each other, insoluble iron (II) hydroxide forms. Fe2+ (aq) + 2OH-(aq) → Fe (OH) 2(s) 5) Iron(II) hydroxide reacts with more oxygen, forming hydrated iron(III) oxide. This is rust. 4Fe(OH)2 + O2 → 2(Fe2O3 . H2O) + 2H2O Identify data, select equipment, and plan and perform a first-hand investigation to compare the rate of corrosion of iron and an identified form of steel. - - - Aim: To compare the rate of corrosion of iron and mild steel. Hypothesis: Mild steel rusts the heaviest. Material: 2 test tubes Pure iron nail & Mild-steel iron nail. Salt solution Test-tube rack Method: Collect two test tubes. Pour same amount of salt solution into both test tubes. Put each nail into a test tube Wait for over a day to clearly observe the difference in the appearance of the nail. Results: Iron shows moderate surface rusting, mild steel shows heavy rusting. Explanations: Iron undergoes rusting at a slow to moderate rate. Mild steel shows accelerated rusting. This presence of carbon promotes rusting. Use available evidence to analyse and explain the conditions under which rusting occurs. - Oxygen is a reactant in the reduction. Water is a reactant in the reduction. An electrolyte environment allows the ions to migrate and react with each other. The electrodes must be connected in order for the electrons to flow from the anode (site of oxidation) to the cathode (site of reduction). HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry 3. Electrolytic cells involve oxidation-reduction reactions. RECALL: - The higher the reduction potential (more positive the value), the harder it is to oxidise the anion. Nitrate, sulfate and fluoride ions are never oxidised, water molecules are oxidised to oxygen instead. The lower the electrode potential (more negative), the harder it is to reduce a metal ion. Silver & copper ions are always reduced. Potassium, sodium, calcium and aluminium are never reduced. At the cathode, there are often two half-reactions competing and often the water reaction wins the competition by having a significantly lower energy/voltage requirement. → - H2O (l) + e- - At the anode, there are also two possible oxidations that occur. Once again, the water often wins because of its lower energy/voltage requirement. - H2O (l) → ½ H2 (g) +OH-(aq) 1/2O2 (g) + 2H + 2e - HSC- Stage 6 - Shipwrecks, Corrosion and Conservation 2 Unit Chemistry The H+ ions and OH- ions will migrate through the solution, find each other and react to form water. So subtracting this from the equation, what really happens is H2O → H2 + 1/2O2 Describe, using half-equations, what happens at the anode and cathode during electrolysis of selected aqueous solutions. - Electrolysis of potassium iodide solution. Voltage applied with inert graphite electrodes. At the anode, the competing oxidations are water and iodide. H2O (l) → 1/2O2 (g) + 2H+ + 2eAnd 2I- - → I2 (aq) + 2e- → H2O (l) + e - V= -0.62 Meanwhile, at the cathode, the competition is between potassium and water. K + e- - V= -1.23 K (s) → ½ H2 (g) +OH- V= -2.94 V= -0.83 The water reaction is highly favoured. It produces a stream of hydrogen gas bubbles, and the hydroxide ions (OH-) cause the phenolphthalein to change from colourless to red as the solution changes from neutral to basic. Whenever several electrode reactions are possible, the one favoured is that with the lowest voltage required (most positive value). Describe factors that affect an electrolysis reaction: - - Effect of concentration: Increased concentration of ions increases the current and consequently the rate of electrolysis. The reaction can also favour the half-equation with least negative potential provided there are high concentrations available. The conc. Of ions in the electrolyte will change as the reaction proceeds, consuming some and producing others. These changes may eventually result in the formation of different products. For example; water and chloride are close to each other on the standard potential table, when a concentrated solution of say, NaCl is the electrolyte, Cl- will oxidise in preference to water, whereas in a dilute solution, water will be oxidised. Nature of electrolyte: Molten Electrolyte: The anion from the salt is oxidised at the anode. The cation is reduced at the cathode. Aqueous Electrolysis: In the electrolysis of aqueous solutions of reactive metals, water is reduced at the cathode in preference to metal ions. Na+(aq) +e- → Na(s) E= -2.71V 2H2O(l) + 2e- → H2(g) + 2OH-(aq) E= -0.83V HSC- Stage 6 - Shipwrecks, Corrosion and Conservation 2 Unit Chemistry In the electrolysis of less active metal salts, such as copper (II) and silver salts, the metal ions are reduced at the cathode in preference to water. Ag+(aq) + e- → Ag(s) E= +0.80V 2H2O(l) + 2e- → H2(g) + 2OH-(aq) E= -0.83V The possible anode reactions: 2Cl- (aq) → Cl2(g) + 2eE= -1.36V 2H20 (l) → O2(g) +4H+ +4eE= -1.23V 2Br-(aq) → Br2 (l) + 2eE= -1.06V 2I- (aq) → I2 (s) + 2eE= -0.54V The relative tendency of these species to be oxidised at the anode in electrolysis reactions is I- > Br- > H2O > ClNature of electrodes: Inert Electrodes: Graphite and platinum are considered inert, as they are rarely involved in electrolysis reactions. They simple provide the surface for electron transfer in the electrode half-reactions. If an inert anode is used, the only oxidation possible is: H2O (l) →1/2O2 (g) + 2H+ + 2e- . However if a copper metal is used as the anode, another reaction becomes possible: Cu(s) → Cu2+ + 2e-. This has a lower voltage requirement than the water oxidation, so this time there is no oxygen production and the copper electrode is eaten away as copper atoms become ions and deposits onto the cathode. Metal electrodes: Reactive metals used for electrodes such as iron, copper or silver as the anode may undergo oxidation in preference to other possible reactions. Because electrode potentials for the oxidation of these metals are relatively positive, particularly compared to the oxidation of water to oxygen gas. Ag (s) →Ag+ (aq) + e- E= -0.80V Cu (s) → Cu 2+ (aq) + 2e- E= -0.34V Fe (s) → Fe 2+ (aq) + 2e- E= +0.44 2H2O (l) → O2 (g) + 4H+ + 4e- E= -1.23 Plan and perform a first-hand investigation and gather first-hand data to identify the factors that affect the rate of an electrolysis reaction. - - Background: One way to determine the rate of electrolysis is to measure the mass of metal deposited on an inert cathode. The greater the rate of electrolysis, the more metal that will deposit. However in this experiment you will determine this rate by considering the amount of I2 (aq) formed. The amount of iodine in solution can be used as an indication of the degree of electrolysis. The more iodide present, the more intense is the blue colour when an iodine solution is added. Aim: To determine the effect of: Applied voltage HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry Concentration of electrolyte Surface area of the electrode Separation of electrodes ….on the effect of electrolysis. - - Method: Prepare the control cell. Prepare 200mL of 0.5mol/L solution of potassium iodide. Place this in a beaker and use two identical clean graphite electrodes. Place the electrodes 5 cm apart and use a voltage of 4 V. Run the electrolysis for exactly 4 minutes. Remove exactly 5 mL of the electrolyte solution. Mix this with a known volume of starch solution in water and observe the intensity of the blue colour of the starch-iodine complex. Varying the voltage. Repeat this experiment keeping all other factors the same but using voltages of 1V, 3V, 5V and 7V. Varying electrolyte concentration. Repeat this experiment keeping all other factors the same but using electrolyte concentrations of 0.25mol/L, 0.75mol/L, 1.0mol/L and 1.5mol/L. Varying electrode surface area. Repeat this experiment keeping all other factors the same but using thinner electrodes, thicker electrodes, and electrodes of different surface areas. Varying electrode separation. Repeat this experiment keeping all other factors the same but using distance between the electrodes of 2 cm, 4 cm, 6 cm and 8 cm. Results: Higher voltages, higher electrolyte concentrations, larger surface areas and electrodes close together increase the rate of electrolysis. Anything that increases electron flow in the circuit will increase the rate of the electrolysis reaction. 4. Iron and steel corrode quickly in a marine environment and must be protected. Identify the ways in which a metal hull may be protected including: - - Corrosion resistant metals: One way to prevent rusting of a steel ship would be to build it from a non-corrosive, passivation metal such as stainless steel, or aluminium. This option is effective, but not economical for large ships. Stainless steel is commonly used for railings and other ship fittings, but is far too expensive to build the hull from. Aluminium is commonly used for small boats and for fishing trawlersized vessels. It is rarely used for large ships due to the cost. Development of surface alloys: A cheaper way to achieve the benefits of stainless steel is to make passivation surface alloys. The steel surface is treated with a laser, which quickly heats the surface, causing it to melt and bombarding the surface with chromium and nickel ions. The ions penetrate the steel surface and form a passivating alloy just a few atoms thick. The bulk of the metal is ordinary, cheap steel with a very thin passivating surface layer, which resists corrosion. HSC- Stage 6 - Shipwrecks, Corrosion and Conservation 2 Unit Chemistry New paints: New polymer-based paints are now available which give protection way beyond what ordinary paints can achieve. These paints form a film of polymer material that is highly effective at preventing oxygen and water penetration. As well, chemicals in the paint react with the metal surface to form an insoluble, ionic layer that totally prevents ion migration, which is necessary for the formation of rust. Predict the metal, which corrodes when two metals form an electrochemical cell using a list of standard potentials. - If two different metals are physically or electrically connected and immersed in an - electrolyte, then some amount of current flows between the two metals. The current is supplied by one of the metals by releasing metal ions to the conductive environment. With two metals in contact, the more active metal forms the anode with the cathode being the less active metal. For iron and copper the statndad electrode potentials are: Fe2+ (aq) + 2e- → Fe (s) E= -0.44V Cu2+ (aq) + 2e- → Cu (s) E= +0.34V This means: Anode: Fe (s) → Fe2+ (aq) + 2eE= +0.44V Cathode: Cu2+ (aq) + 2e- → Cu (s) E= +0.34V At the cathode copper is competing with the dissolved gas: O2(g) + 4H+ +2e- → H2O E= 1.23V Oxygen „wins‟ as it has a more positive electrode potential. Outline the process of cathodic protection, describing the examples of its use in both marine and wet terrestrial environments. - - Wet terrestrial environment: Sacrificial anodes and applied voltages will protect a steel ship in the ocean and be just as effective at protecting a steel pipeline buried in wet ground. As usual, by forcing electrons into the steel with an applied voltage, the steel is forced to become a cathode and cannot oxidize. Marine environment: Sacrificial anodes: If a block of a more reactive metal is attached to a steel structure, it becomes the anode and forces the steel to be the cathode. In the electrolyte environment of the ocean, this is more effective than galvanising and is commonly used to protect the steel hulls of ships. It is important that the anode have electrical contact with the steel so electrons can flow from anode to cathode. At the steel cathode, any Fe2+ ions which may have formed by corrosion are forced to reduce back to iron metal. However, the main cathode reaction is the familiar reduction of oxygen and water. Galvanised steel: It is referred to the steel that has been coated with a thin layer of zinc, or a mixture of zinc and aluminium. The zinc is a passivating metal and forms an impermeable surface coating which effectively seals the surface. Applied voltages: Another way to force the steel to be a cathode is to connect it to the negative terminal of an electricity supply. This forces electrons into any possible corrosion sites in the steel, forcing it to be a cathode and HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry preventing it becoming an anode. The positive side of the electricity supply is simply connected to an inert electrode, which is below water level. The electric circuit is completed by ion migration through the seawater. The steel is forced to be a cathode and cannot corrode. Describe the process of cathodic protection in selected examples in terms of the oxidation/reduction chemistry involved. - - - Consider the following reduction potentials: Fe2+ (aq) + 2e- → Fe (s) E= -0.44V Mn2+ (aq) + 2e- → Mn (s) E= -1.18V 2H2O (l) + O2 (aq) + 4e- → 4OH- (aq) E= +0.34V Manganese has a lower reduction potential than iron and so can be used to protect iron and steel structures. In a galvanic cell, Mn is oxidised in preference to Fe. Mn forms the anode, while Fe becomes the cathode. Now the E for the reduction of dissolved oxygen is greater than that for the reduction of iron and so dissolved oxygen is more likely to be reduced. This will occur on the surface of the iron with the iron acting as no more than the conductor of electrons. Identify data, gather and process information from first-hand or secondary sources to trace historical developments in the choice of materials used in the choice of materials used in the construction of ocean-going vessels with a focus on the metals used. - - - - From ancient times, sea-going ships were always constructed from wood. The wooden planks and beams were fastened together with metal nails and spiked, but no one imagined constructing the ship itself from metal. In the 1600s it was discovered that is thin sheets of copper or lead were nailed onto the hull of a ship it would reduce the growth of seaweed, barnacles, wood-eating worms etc on and in the hull. When the industrial revolution began in England, new machines, engines, railways and bridges were built from cast iron, which had become cheaply available in large quantities. By the 1780s, a few small boats had been built with iron hulls. To the surprise of many, the iron boats not only floated, but also were stronger, lighter and cheaper than wooden equivalents. Over the following 70 years, larger and larger iron-hulled ships were built, until by the 1850s nearly all commercial shipping was being built from iron. One hundred years later, steel is still the majority metal used in shipbuilding. Identify data, choose equipment, plan and perform a first-hand investigation to compare the corrosion rate, in a suitable electrolyte, of a variety of metals including named modern alloys to identify those best suited for use in marine vessels. - Background: By dissolving some phenolphthalein and potassium ferricyanide in salt water you can make ferroxyl indicator. The presence of Fe2+ ions imparts a blue HSC- Stage 6 - - Shipwrecks, Corrosion and Conservation 2 Unit Chemistry colour to the solution. The presence of OH- ions makes the indicator turn pink. You will need 8 identical, clean nails for this experiment. Aim: To examine conditions which prevent or slow down the rate of rustin. Method: Place a nail in a clean test tube and cover with ferroxyl indicator. Allow this to stand for some time without disturbing. Anodic and cathodic regions will show up. This has to do with places around the nail that have been stressed more than others. In separate test tubes, covered with ferroxyl indicator, place the following: An iron nail tightly wound with clean copper wiring. An iron nail tightly wound with clean magnesium ribbon. An iron nail painted with phosphoric acid and allowed to dry first. An iron nail painted with an outdoor paint, An iron nail coated with grease. Two iron nails connected by leads to a 1.5V battery. Results: Plan and perform a first-hand investigation to compare the effectiveness of different protections used to coat a metal such as iron and prevent corrosion. - - Corrosion-resistant metals: This option is effective, but not economical for large ships. Stainless steel is commonly used for railings and other ship fittings, but is far too expensive to build the full from. Surface alloys: It is effective as it is very cheap. The bulk of the metal is ordinary, cheap steel with a very thin passivating surface layer, which resists corrosion. New paints: Not effective as they need regular maintenance and hence proven to be very expensive. Gather and process information to identify applications of cathodic protection, and use available evidence to identify the reasons for their use and the chemistry involved. - It‟s all explained somewhere. Look for it. 5. When a ship sinks, the rate of decay and corrosion may be dependent on the final depth of the wreck. Outline the effect of temperature and pressure on the solubility of gases and salts. Temperature: Salts: As temperature increases the maximum, solubility of ionic compounds also increases. In the colder waters of the ocean depths, you might expect less salt to be dissolved. Gases: As temperature increases the maximum, solubility of most gases decreases. Therefore, you might predict that in the cold ocean HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry depths the amount of dissolved oxygen would be higher than near the surface, due to temperature. Pressure: Salts: Changing the pressure has effect on the solubility of ionic compounds. The deep ocean has the same electrolyte properties as the surface water. Gases: As pressure increases the maximum, solubility of most gases also increases. Therefore, you might predict that in the ocean depths the amount of dissolved oxygen would be much higher than near the surface, due to a very high pressure. However, oxygen can only dissolve at the surface in contact with the atmosphere. The mixing of surface waters into the depths is very slow, and meanwhile oxygen is being consumed by all the ocean creatures respiring. The result is that the concentration of dissolved oxygen in the deep is very slow. Meanwhile, sea animals release carbon dioxide, which dissolves readily, and forms hydrogen ions in solution. Identify that gases are normally dissolved in the oceans and compare their concentrations in the oceans to their concentrations in the atmosphere. - - The most common atmospheric gases (nitrogen and oxygen) are both small, nonpolar molecules. They are only slightly soluble in water, and their concentration in seawater is very low compared to their concentration in the atmosphere. Other gases, such as carbon dioxide are much more soluble because they do not just dissolve, but react with water to form hydrogen carbonate and carbonic acid. Comparisons of common gases in Air & SeaWater Gas Oxygen Nitrogen Carbon Dioxide - Conc. In Atmosphere (%) 21% 78% 0.035% Max. % dissolved in surface waters 0.9% 1.5% 5.4% Compare and explain the solubility of selected gases at increasing depths in the oceans. With increasing ocean depths, temperature decreases, and pressure increases. HSC- Stage 6 - Shipwrecks, Corrosion and Conservation 2 Unit Chemistry Oxygen concentrations are greatest near the surface and at great depths. Nitrogen in ocean waters is within 5% of equilibrium with the atmosphere. Predict the effect of low temperatures at great depths on the rate of corrosion of a metal. - There are very low temperatures, around 4 C. - Very low oxygen levels around 2ppm. - Salinity levels are higher than at the surface. - Seawater is slightly alkaline. - Both low oxygen levels and low temperatures should retard metallic corrosion. - It can be predicted that shipwrecks in very deep ocean water corrode fairly slowly, if at all. Perform a first-hand investigation to compare and describe the rate of corrosion of materials in different: - Aim: To compare and describe the rate of corrosion of materials in different oxygen - - - concentrations, temperature and salt concentrations. Oxygen concentration: Each test tube contains exactly the same metal, salt solution and is stored at the same temperature. Method: Prepare a 150 mL solution of about 3% salt water. Boil 100 mL of this salt water for several minutes to remove oxygen. Completely fill a container with this low-oxygen solution and stopper. Allow it to cool before using. Place three steel nails separately into test tubes. Completely fill the first with the low-oxygen saltwater solution and stopper. Half fill the second test tube with unboiled salt water. Half fill the third tube with tape water. Label the tubes. To each test tube, add 3 drops phenolphthalein. Place the tubes to one side to be monitored over several days. Do not agitate. Result: Higher the oxygen concentration the faster the rate of rusting. Temperatures: Each test tube contains exactly the same metal, gas and liquid. Method: Prepare a 150 mL solution of about 3% salt water. Place three steel nails separately into test tubes. Partly fill them with saltwater solutions. Label the tubes. To each test tube, add 3 drops phenolphthalein. Place one tube upright in a refrigerator, the second in an incubator set at 35% and leave the other to one side on the bench top. Monitor the tubes over several days. Do not agitate. Result: Higher the temperature, the faster the rate of corrosion. Salt concentrations: Each test tube contains exactly the same metal, gas and is stored at the same temperature. HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry Method: Prepare a 150 mL solution of about 3% salt water. Also, prepare a 150 mL solution of about 6% salt water. Place three steel nails separately into each test tubes. Partly fill the first with the 3% salt-water solution. Half fill the second test tube with 6% salt water. Label the tubes. To each test tube, add three drops of phenolphthalein. Place the tubes to one side to be monitored over several days. Do not agitate. Result: Higher the salt concentration the faster the rate of rusting, although the differences may be slight. Use available evidence to predict the rate of corrosion of a metal wreck at great depths in the oceans and give reasons for the prediction made. - Prediction: The rate of corrosion at great depth is extremely slow, if present at all. Reasons: Oxygen levels insufficient for corrosion to occur. And the temperature is very low. 6. Predictions of slow corrosion at great depths were apparently incorrect. Explain that shipwrecks at great depths are corroded by electrochemical reactions and anaerobic bacteria. - - - „Titanic” was found to be in very poor conditions. She was covered in rusticles, hanging like icicles from every steel structure. When analysed, the rust deposits were found to contain a lot of black iron (II) sulphide (FeS), as well as rust (Fe2O3). It turns out that bacteria caused such a reaction. These bacteria are anaerobic, they live without oxygen. Instead of using oxygen to release the energy of food, they rely on the reduction of sulfate. SO42- + 5H2O + 8e- → HS- + 9OHEvery reduction needs an oxidation, and the steel is it! These bacteria live naturally in the bottom mud, but grow rapidly around a wreck because of the increased supply of organic matter. When they grow around the steel of the ship, the oxidation reaction is the familiar: Fe(s) → Fe2+ + 2e- Describe the action of sulfate reducing bacteria around deep wrecks. - The anaerobic bacteria rely on the reduction of sulfate ions, rather than oxygen, to generate energy from respiration. - The reduction of sulfate ions can occur in both neutral and acidic conditions. Explain that acidic environments accelerate corrosion in non-passivating metals. - Concentration of carbon dioxide in the deep ocean is much higher than that of oxygen. When carbon dioxide dissolves, it reacts to form a weak acid. HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry CO2(g)+3H2O(l) CO2(aq)+3H2O(l) H2CO3(aq)+2H2O(l) CO32(aq)+2H3O+(aq) - - This is why the water of the deep ocean is slightly acidic compared to the surface waters, which are neutral or slightly basic. Acidic levels around the wreck due to the many bacteria feeding on the organic matter, and producing extra carbon dioxide and other weak acids. These acidic conditions can accelerate corrosion. In the normal rusting of iron, the cathode reduction is: O2 + 2H2O + 4e- → 4OHE= +0.40V However, in acidic conditions another reduction of dissolved oxygen becomes possible: O2 + 4H+ 4e- → 2H2O E= +1.23V This reaction has higher, more positive electrode potential than the first. For any given oxygen concentration, the „acidic reduction‟ is favoured and runs faster. This promoted the faster anode oxidation of the metal as well, so corrosion is more rapid. Perform a first-hand investigation to compare and describe the rate of corrosion of metals in different acidic and neutral solutions. - Aim: To compare the rate of corrosion of metals in different acidic and neutral solutions. Methods: Place 3 nails into three test tubes and place them on a test tube rack. 1st test tube: Place 3% salt solution + 2 drops water. ( Neutral) 2nd test tube: Place 3% salt solution + 2 drops acid. ( Acidic) 3rd test tube: Place 3% salt solution + 2 drops base. (Basic) Result: The second test tube rusted faster. - - 7. Salvage, conservation and restoration of objects from wrecks requires careful planning and understanding of the behaviour of chemicals. Explain that artefacts from long submerged wrecks will be saturated with dissolved chlorides and sulfates. - Sulfates and chlorides are readily available at great depths of the ocean and present in - reasonably high concentrations. Overtime, the artefact completely saturates with the ions of chloride and sulfate. The ions become impregnated into the structure thus causing a potential damage to the artefact. Describe the processes that occur when a saturated solution evaporates and relate this to the potential damage to drying artefacts. - When a salt solution evaporates, the salts form crystals: If salts become impregnated into the structure of an artefact and it is allowed to dry out, the salts will form crystals. As the crystals expand, they cause pressure within the artefacts structure, weakening the object. Because of this process, ceramic objects and glass may break or crack; leather materials become hardened and strained; metal objects corrode more rapidly. HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry Identify the use of electrolysis as a means of cleaning and stabilising iron, copper and lead artefacts. - Electrolysis can be used to free chloride ions from the insoluble compounds into the - - - - solution – since chloride ions enhance corrosion, they all need to be removed. Initially a low current is used to draw out negative ions (chlorides and sulfates) which become attached to the inert anode. Later a high voltage current is used which generates bubbles of hydrogen gas at the cathode which helps to remove/loosen any remaining concretions. Electrolysis removes deeply embedded chloride ions. Anode: 2Cl– →Cl2 (g) + 2e– 4OH– → 2H2O (l) + O2 (g) + 4e– 2H2O (l) →4H+ + O2 (g) + 4e– Cathode: 2H2O(l) + 2e– →H2(g) + 2OH– Iron artefacts: The first treatment is the mechanical removal of calcareous deposits and removal of seawater ions by connecting the surface to an electrical circuit in a solution of sodium hydroxide. The chloride level in the electrolyte was continually monitored to determine how much has been removed, as chloride ions greatly accelerate the corrosion process. Eventually the cannon was washed in distilled water and dried before being sprayed with a rust inhibitor. Finally, it is coated with wax to protect it from air and water. Copper artefacts: Electrolysis is used to reduce the copper ions back to metallic copper. The process involved passing a current through an electrolytic cell with the artefact as the cathode and a stainless steel anode. A the cathode, the copper is reduced to metallic copper. Hydrogen gas is also formed at the cathode by the reduction of water, with small bubbles forming on the surface of the artefact and helping to loosen any flakes of corroded material. Lead artefacts: Being a metal of low reactivity, lead is fairly resistant to corrosion in a marine environment, but lead artefacts can easily be physically damaged due to the softness of this metal. The artefacts were stored in distilled water and then any calcareous material removed mechanically or by soaking in a weak acid bath. The artefact can then be stripped using disodium salt of ethylenediaminetetraacetic acid (EDTA) or the artefact can be converted back to metallic lead by electrolysis and then coated with a protective wax coating. HSC- Stage 6 Shipwrecks, Corrosion and Conservation 2 Unit Chemistry Discuss the range of chemical procedures, which can be used to clean, preserve and stabilise artefacts from wrecks and, where possible, provide an example of the use of each procedure. Procedure Impregnating polymers into wood Heat treatment of iron artefacts Acid baths Wax coatings/lacquers Desalination Method Examples of use Wooden object is soaked in natural plant resins until the polymer penetrates and sets in the object. Object is heated in an electrical furnace in an atmosphere of hydrogen. Iron chlorides and oxides are reduced to iron. Soak objects in a dilute acid bath. Used to restore strength and stabilise wooden objects. Apply a thin coating of grease, wax or lacquer – appearance unchanged. Remove salt from the surface layer Helps to restore the structure of badly corroded objects. Cleaning copper artefacts: dissolve corrosion layers in dilute citric acid. Preserve objects from further corrosion for display. Iron-based objects. Perform investigation and other information from secondary sources to compare conservation and restoration techniques applied in two Australian maritime archaeological projects. - - - - The Vemon anchors, which are on, display outside the Maritime Museum were conserved and restored in the 1980s. Electrolysis was not applied to the anchors because it required the removal of the timber stocks and because the cast iron was in sufficiently good condition. In 1992, the preservation process removed the outer corrosion and protective paint by blasting with copper slag then polished the surface with garnet. The iron was then treated with zinc epoxy paint. The timber stocks were saturated with a zinc napthenate solution. The Endeavour cannons, part of Captain Cook‟s famous ship were rediscovered in 1969 after being jettisoned when the Endeavour was damaged in the Great Barrier Reef in June 1770. The six cannons were then transported into a salt solution with 10% formalin to kill any bacteria present. Hammers were used to remove hard coral concretions from the cannons. The cannons were then placed in 2% NaOH solution to prevent acidic corrosion; then electrolysis was applied in NaOH baths for many weeks. After the baths, the cannons were washed with fresh water at each refreshing of the electrolyte. After electrolytic treatment, the cannons were washed for 5 months to remove any remaining chloride and hydroxide using distilled water with chromate ions (the chromic oxide formed is a surface protective layer). HSC- Stage 6 - - Shipwrecks, Corrosion and Conservation 2 Unit Chemistry The cannons were dried for 48 hours at 120°C. They were then immersed in molten microcrystalline wax (as a further protective layer) for 5 days to ensure maximum penetration of wax. While the Endeavour cannons required electrolysis because of its badly corroded condition, the Vernon anchors did not. The cannons required many steps of treatment (stabilisation, removal of concretions, electrolysis and surface protection), while the anchors required only two (removal of concretions and surface protection).
