chemistry - shipwrecks

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 Earths crust slowly move apart releasing hot magma

located below the waters 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 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.

His experiment turned out to

be wrong. However, his

experiment sparked

tremendous interest in the

study of electricity.

- Volta 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.

Established the idea that

chemical reactions could

produce electricity. His

battery Volta became successful and as the best

way to make electricity.


- Davy 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 battery made

electricity from

chemical reactions.

Discovered many new chemical elements by

electrolysis.

Developed the technique of

electrolysis.

- Faraday 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.

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%


- 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 Typical use

- Iron Pure element Soft & malleable Corrodes very

slowly

No common use in

modern times

- Cast Iron 3-4% Carbon Hard, but brittle Corrodes rapidly

Cast engine blocks.

Decorative iron

lacework - Structural Steel About 0.5%

Carbon

Very strong, but

can be shaped

when hot.

Corrodes rapidly

Reinforcing bars in

concrete

Girders in

buildings and

bridges

Railway lines

- Mild Steel Less than 0.2% carbon

Fairly strong, but

malleable &

ductile.

Corrodes rapidly.

Car bodies

Pipes

Roofing steel

Nuts & bolts

Ship hulls

Chains

- Stainless Steel Up to 20% chromium and

nickel

Hard & shiny

Resistant to

corrosion

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.


Fe 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).


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- H2 (g) +OH-(aq) - 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) 1/2O2 (g) + 2H + 2e -


- 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+ + 2e- V= -1.23

And

2I- I2 (aq) + 2e- V= -0.62 - Meanwhile, at the cathode, the competition is between potassium and water.

K + e- K (s) V= -2.94

H2O (l) + e H2 (g) +OH- 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


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) + 2e- E= -1.36V

2H20 (l) O2(g) +4H+ +4e- E= -1.23V

2Br-(aq) Br2 (l) + 2e- E= -1.06V

2I- (aq) I2 (s) + 2e- E= -0.54V

The relative tendency of these species to be oxidised at the anode in

electrolysis reactions is I- > Br- > H2O > Cl-

- Nature 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


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 trawler-

sized 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.


- 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) + 2e- E= +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


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


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.

- Its 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


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, non-

polar 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.

Compare and explain the solubility of selected gases at increasing depths in

the oceans.

- With increasing ocean depths, temperature decreases, and pressure increases.

Comparisons of common gases in Air & SeaWater

Gas Conc. In Atmosphere

(%)

Max. % dissolved in

surface waters

Oxygen 21% 0.9%

Nitrogen 78% 1.5%

Carbon Dioxide 0.035% 5.4%


- 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.


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- + 9OH-

- Every 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.


CO2(g)+3H2O(l) CO2(aq)+3H2O(l) H2CO3(aq)+2H2O(l) CO3-

2(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- 4OH- E= +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.


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.


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 Method Examples of use

Impregnating polymers

into wood

Wooden object is soaked in

natural plant resins until the

polymer penetrates and sets

in the object.

Used to restore strength and

stabilise wooden objects.

Heat treatment of iron

artefacts

Object is heated in an

electrical furnace in an

atmosphere of hydrogen. Iron

chlorides and oxides are

reduced to iron.

Helps to restore the structure

of badly corroded objects.

Acid baths Soak objects in a dilute acid bath.

Cleaning copper artefacts:

dissolve corrosion layers in

dilute citric acid.

Wax coatings/lacquers Apply a thin coating of grease, wax or lacquer appearance unchanged.

Preserve objects from further

corrosion for display.

Desalination Remove salt from the surface layer

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 Cooks 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).


- The cannons were dried for 48 hours at 120C. 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).

chemistry - shipwrecks

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