OPTION 9.6 – SHIPWRECKS,
CORROSION AND CONSERVATION
1. The chemical composition of the ocean implies its potential role as an electrolyte.
1.2.1 – Identify the origins of the minerals in oceans as: – leaching by rainwater from terrestrial
environments, – hydrothermal vents in mid-ocean ridges.
Mineral ions in oceans come from two main sources: the leaching by rainwater from terrestrial
environments and hydrothermal vents in mid-ocean ridges. Rainwater washes over and leaches
mineral ions from rocks and soils in the process of weathering. These ions, including Ca 2+, Mg2+
and HCO32–, are carried by the water-cycle into the sea. When seawater enters hydrothermal
vents in mid-ocean ridges, or when it cools mineral-rich magma, it also leaches out mineral ions.
1.2.2 – Outline the role of electron transfer in oxidation-reduction reactions.
1.2.3 – Identify that oxidation-reduction reactions can occur when ions are free to move in liquid
electrolytes.
Oxidation is the loss of electrons, reduction is the gain. In oxidation-reduction reactions, the
electron moves from the anode (the site of oxidation) to the cathode (the site of reduction).
Without this electron transfer, there will be no reaction.
Redox reactions can occur when ions are free to move in liquid electrolytes. An example of this is
the displacement of copper by zinc in a copper solution. Zinc metal comes in contact with copper
ions and u get electron transfer. Zinc loses electrons and becomes cations while copper cations
gain electrons. If the ions cannot move freely, for example two compartments separated by a very
fine semipermeable membrane that Cu2+ cannot penetrate, then the redox reaction will not
occur.1
1.2.4 – Describe the work of Galvani, Volta, Davy and Faraday in increasing understanding of electron
transfer reactions.
1.3.1 – Process information from secondary sources to outline and analyse the impact of the work of
Galvani, Volta, Davy and Faraday in understanding electron transfer reactions.
Luigi Galvani was an Italian physician in the eighteenth century. In the 1780s he undertook a
series of investigation into the twitching of frog leg muscles with a static electricity generator. He
also obtained twitching by pressing a brass hook into the frog’s spinal cord and hanging the hook
on an iron railing. He concluded that animal tissue contained an “electric fluid”, through which a
force, “animal electricity”, could acted upon the tissue. This was the first example in our
understanding of electrolytes in electron transfer reactions.
Alessandro Volta was a contemporary of Galvani, who later refuted his theory of “animal
electricity”. He believed that the twitching of the frog legs were due to the two different metal
pieces holding the legs. In 1800, Volta made his most famous discovery, the voltaic pile, a stack
of silver and zinc discs separated by felt pads soaking in brine (salt solution). This discovery led
to our understanding of using electron transfer reactions to generate electricity.
Humphry Davy was inspired by work of other scientists who managed to electrolyse and
decompose water into hydrogen and water using a voltaic pile. This led to his electrolysis
experiments where he decomposed potassium hydroxide (potash) and molten sodium hydroxide
1
Credit where credit is due: this very good explanation was suggested by phenol. All kudos is owed to him.
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(soda) into previously undiscovered elements, potassium and sodium. These investigations led to
our understanding of electron transfer reactions as the basis of electrolysis.
Michael Faraday was Davy’s laboratory assistant, and as a result continued Davy’s work on
electrolysis. He was the first person to define the laws of electrolysis and quantify electrolysis. He
also introduced the terminology used today such as electrolyte, cation and anion. His discoveries
led a better understanding of electron transfer reactions in electrolysis.
2. Ships have been made of metals or alloys of metals.
2.2.1 – Account for the difference in corrosion of active and passivating metals.
The difference between corrosion of active and passivating metals is the oxide layer formed: in
active metals, such as iron, the oxide layer is permeable to oxygen and water, so that fresh metal
beneath will be attacked by corrosion. Passivating metals (lead, zinc, aluminium, chromium) do
not corrode readily because initial corrosion produces non-porous oxides, which are impervious
against oxygen and water, and thus against further corrosion.
2.2.2 – Identify iron and steel as the main metals used in ships.
2.2.3 – Identify the composition of steel and explain how the percentage composition of steel can determine
its properties.
2.3.3 – Gather and process information from secondary sources to compare the composition, properties and
uses of a range of steels.
Iron and mild steel are the main metals used in ships. Pure iron is relatively soft, and carbon
atoms in steel interrupt the lattice structure, preventing them from sliding easily, increasing
strength. Mild steel (iron plus 0.2% – 0.5% carbon) is flexible, malleable and easily beaten into
bars; because of this, it is used in cars and ships in areas that need to absorb damage (eg.
crumple-zones) or shock.
Structural steel (iron plus 0.5%–0.8% carbon) is stronger and less malleable/flexible than mild
steel; because of this, it is used in load-bearing applications such as girders, beams, bridges and
railway tracks. Tool steel (iron plus 0.8%–1.5% carbon) is the strongest of the three and has the
highest melting-point; because of this, it is used in tool-parts, such as saws and drills bits.
2.2.4 – Describe the conditions under which rusting of iron occurs and explain the process of rusting.
2.3.2 – Use available evidence to analyse and explain the conditions under which rusting occurs.
When iron metal is in contact with oxygen AND water (sometimes in the form of moisture), the
iron is oxidised in Fe2+ ions. The oxygen and water are reduced to hydroxide:
Oxidation:
Reduction:
Fe(s)
→
½O2(g) +
Fe2+(aq) +
H2O(l) +
2e–
2e–
2OH–
→
E° = + 0.44V
E° = + 0.40V
Since the total potential of both half-equations is positive, the reaction is spontaneous. The Fe2+
ions and the OH– ions form a green precipitate, Fe(OH)2. Fe(OH)2 will oxide further in the
presence of oxygen and water to form Fe(OH)3:
4Fe(s)
+
3O2(g)
+
6H2O(l)
→
4Fe(OH)3(s)
Fe(OH)3 will then dehydrate to form rust, Fe2O3·2H2O. Therefore the overall equation for the
formation of rust is:
4Fe(s)
+
3O2(g)
+
4H2O(l)
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→
Fe2O3·2H2O(s)
Any piece of iron (or steel) will be more likely
to oxidise in some places and not in others.
It will oxidise where there are fine cracks,
thin, damaged, twisted or bent sections. This
is because here the iron atoms are further
apart, the lattice is disrupted (by carbon
atoms) and the electrons flow more easily
and are therefore more easily lost. This
section of iron is said to be anodic because
oxidation occurs here.
The electrons released from the oxidation of iron flow away to another area which will act as the
cathode. Here, the water and oxygen are reduced to hydroxide and the iron is prevented from
corroding at the cathode.
The hydroxide irons and the iron ions move through the water and meet to form rust between the
cathode and the anode.
Acidic conditions accelerate corrosion because the cathodic reaction becomes the reduction of
oxygen and hydrogen ions, which has a higher reduction potential than (and thus happens in
preference to) that of oxygen and water:
Fe(s)
→
½O2(l) +
Oxidation:
Reduction:
Fe2+(aq) +
2H+(aq) +
2e–
2e–
→
H2O(l)
E° = + 0.44V
E° = +1.23V
Rules to work out where anode and cathode are:
 The place where oxygen can most easily reach becomes the cathode.
 If oxygen can reach all places evenly, the place where iron is disrupted is the anode.
2.3.1 – Identify data, select equipment, plan and perform a first-hand investigation to compare the rate of
corrosion of iron and identified form of steel.
One iron nail and one galvanised iron nail of the same size were initially weighed on an electronic
balance and then left in 50mL beakers with 20mL of water in each at SLC. Every day any signs of
corrosion were recorded and the nails were dried and reweighed. A graph of the masses of each
nail was drawn up after five days and mass losses were compared.
3. Electrolytic cells involve oxidation-reduction reactions.
3.2.1 – Describe, using half-equations, what happens at the anode and cathode during electrolysis of
selected aqueous solutions.
An electrolytic cell is one where electrical energy is applied to cause a redox reaction to happen.
Unlike galvanic cells, in electrolytic cells:




The cathode (reduction electrode) is negatively charged
The anode (oxidation electrode) is positively charged.
Anions (negative) carry charge towards the anode.
Cations (positive) carry charge towards the cathode.
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Examples of electrolytic cells include:
CuCl2(aq):
Cu2+(aq)
H2O(l)
+
→
2e–
→
Cu (goes towards cathode)
½O2(g) + 2H+ + 2e– (bubbling at anode)
CuSO4(aq):
Cu2+(aq)
H2O(l)
+
→
2e–
→
Cu (goes towards cathode)
½O2(g) + 2H+ + 2e– (bubbling at anode)
CuSO4(aq) with copper anode:
Oxidation: Cu2+(aq) + 2e– → Cu E° = – 0.34V
Reduction: Cu → Cu2+(aq) + 2e– E° = + 0.34V
(Copper is deposited on the cathode)
AgNO3(aq) with silver anode:
Oxidation: Ag(s) → Ag+(aq) + e– E°= – 0.80V
Reduction: Ag+(aq) + e– → Ag(s) E°= + 0.80V
(Silver is deposited on the cathode)
3.2.2 – Describe factors that affect an electrolysis reaction: – effect of concentration, – nature of electrolyte,
– nature of electrodes.
Factors which affect the rate of electrolysis are concentration, nature of electrolyte and the nature
of electrodes:
The higher the concentration of the electrolyte, the greater the rate of electrolysis; the higher
concentration means more ions can migrate through the solution at a faster rate.
The nature of the electrolyte (that is its state) is important; most electrolytes are aqueous, since
dissolved ions can readily move within solution. Solid electrolytes (of toothpaste consistency)
allow for a slower reaction since ion migration is slower. Liquid electrolytes are rarely used
because a very high temperature is needed to maintain them, but they react quickly due to the
high concentration of ions.
The nature of the electrode is important in determining the rate of reaction; a reactive metal will
more readily react in electrolysis than an inert one. A high surface area, or more precisely, a high
contact surface area with the electrolyte (eg. a mesh-type electrode) will mean a faster reaction
rate.
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3.3.1 – Plan and perform a first-hand investigation and gather first data to identify the factors that affect
the rate of an electrolysis reaction.
Oxidation: I–(aq) → ½ I2(aq) + e– E° = – 0.77V
Reduction: H2O(l) → ½H2(g) + OH–(aq) E° = – 0.83V
We use the amount of bubbling of hydrogen gas as
the indicator of the rate of electrolysis. The
equipment was set up as the diagram. The rate of
bubbling was measured on a scale of 1 to 5.
To show how concentration affected the rate, we
changed the potassium iodide solution to 2M.
To show how the nature of electrolyte affected the
rate, we changed to using molten potassium iodide.
To show how the nature of electrodes affected the
rate, we used copper electrodes; then we used
cylindrical carbon electrodes instead of rectangular.
4. Iron and steel corrode quickly in a
marine environment and must be protected.
4.2.1 – Identify the ways in which a metal hull may be protected including: – corrosion resistant metals, –
development of surface alloys, – new paints.
The ways in which a metal hull of a ship can be protected include:
Corrosion resistant metals:
the use of passivating metals (such as chromium) by
themselves or alloyed to steel means that the hull will have
corrosion-resistant properties.
Surface alloys:
currently in development stages, surface alloys are made
using a laser, which fires particles of passivating metals onto a
steel hull, giving it cathodic protection.
New paints:
with polymer technology or with passivating metal particles
form a physical impervious layer which protects the hull
against corrosion.
4.2.2 – Predict the metal which corrodes when two metals form an electrochemical cell using a list of
standard potentials.
Since corrosion is really the oxidation of a metal into its ionic form, when two metals of different
oxidation (electron-losing) potentials form an electrochemical cell, the one with the higher
oxidation potential will become the anode and corrode. The other will become the cathode, and
will be protected from corrosion.
4.2.3 – Outline the process of cathodic protection, describing examples of its use in both marine and wet
terrestrial environments.
4.2.4 – Describe the process of cathodic protection in selected examples in terms of the oxidation/reduction
chemistry involved.
4.3.4 – Gather and process information to identify applications of cathodic protection, and use evidence to
identify the reasons for their use and the chemistry involved.
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The basis of cathodic protection is to render the whole piece of iron that needs to be protected as
the cathode, so that it does not oxidise. In marine and terrestrial environments, forms of cathodic
protection such as galvanising, sacrificial anodes and applied voltage are used.
When zinc is electroplated onto iron in the process of galvanising, the iron is protected from
corrosion. This is because:
Oxidation:
Fe(s)
Zn(s)
→
→
Fe2+(aq) +
Zn2+(aq) +
2e–
2e–
E° = + 0.44V
E° = + 0.76V
Since the oxidation potential of zinc is greater than that of iron, it will oxidise in preference (ie
become the anode) and render the iron as the cathode. Moreover, zinc, being a passivating metal,
will form an impervious layer of zinc oxide around the iron, giving further protection against
corrosion. Galvanised iron is used in marine environments such as galvanised nails and
terrestrial environments such as roofs and cars.
A similar principle is used in the process of using sacrificial anodes. A metal with greater
oxidation potential than iron is used so that it will oxidise in preference to iron. Zinc is usually
chosen over other metals (of greater oxidation potential) to decrease the rate of reaction and
lengthen the life of the anode. It is used in the hulls and propellers of ships and underground steel
piping (a moist terrestrial environment).
In the process of applied voltage, iron is connected to low voltage circuit, so that electrons are
pushed into it constantly, rendering it cathodic. This form of protection is usually used in
conjunction with the sacrificial anodes, in large ships, metal girders and some cars.
4.3.1 – Identify data, gather and process information from first-hand or secondary sources to trace
historical developments in the choices of materials used in the construction of ocean-going vessels with a
focus on the metal used.








By 3000BC, some societies were constructing wooden ships. Wood was the most
common shipbuilding material until the nineteenth century.
Metals were used in some early ships, such as Viking longboats which had iron and
bronze fittings.
Around 1500AD the development of iron nails made it possible to connect wooden
planks to frames and bulkheads. This made the hull stronger and less flexible, but the
nails were susceptible to rapid corrosion.
By the 1800s composite ships were built using wooden planks over iron frames. The
first all iron ship was the British Vulcan, a passenger barge, launched in 1818. By
1870 more than 90% of the ships produced in the UK were iron.
Although they needed constant maintenance due to corrosion, iron ships had
numerous advantages over wooden ones: stronger, thus safer, more economical,
easier to repair, could be built larger, carry more cargo, travelled faster, was not
susceptible to fire from cannonball explosions.
By the late 1800s shipbuilders began to use steel alloys, which meant lighter and
stronger ships.
In the early 1900s the invention of electric welding meant faster and better
construction of steel ships.
Other developments in the twentieth-century included the progressive improvement
in steel alloys, incorporating aluminium, chromium, titanium, zinc and nickel. Modern
steels are lighter, stronger and more corrosion-resistant than before.
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4.3.2 – Identify data, choose equipment, plan and perform a first-hand investigation to compare 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.
A variety of metals and alloys were gathered, including equal-sized strips of mild steel, structural
steel, zinc, tool steel and copper. These were placed into separate test tubes of 10mL of
seawater. The signs of corrosion were recorded and measured daily for 5 days (on a scale of 1 to
5), looking for evidence of oxidation of the sample. Those with the least amount of corrosion were
identified to be the most suitable for marine vessels.
4.3.3 – 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.
Equal-sized nails were setup with the following protections: galvanised, painted steel nail, waxed
steel nail and steel nail with no protection. They were placed in a test tube with 10mL of water,
each half-exposed at SLC. All other variables were kept constant. The signs of corrosion were
recorded and measured daily for 5 days (on a scale of 1 to 5), looking for evidence of corrosion in
each sample. The most effective protection(s) showed the least evidence of corrosion.
5. When a ship sinks, the rate of decay and corrosion
may be dependent on the final depth of the wreck.
5.2.1 – Outline the effect of: – temperature, – pressure on the solubility of gases and salts.
5.2.2 – Identify that gases are normally dissolved in the oceans and compare their concentrations in the
oceans and their concentrations in the atmosphere.
5.2.3 – Compare and explain the solubility of selected gases at increasing depths in the oceans.
Generally, salts increase in solubility with increasing temperature. Gases decrease, on the other
hand, in solubility with increasing temperature. Gases also increase in solubility with increasing
pressure.
The concentration of O2 is high in the atmosphere, and low in the ocean. The concentration of
CO2 is relatively low in the atmosphere, and can be high in the ocean around reefs and living
things which respire CO2.
The sea surface is usually highly oxygenated due to wave action and photosynthesis of
phythoplankton. As the depth increases, oxygen levels are used up and drop because of the
respiration of marine organisms. Between 500m to 1000m, oxygen minima develop. The oxygen
level rises at even greater depths due to deep cold ocean current bringing in dense oxygenated
water from the polar seas.
5.2.4 – Predict the effect of low temperatures at great depths on the rate of corrosion of a metal.
The predicted effect of low temperatures at great depths on the rate of corrosion is that the
corrosion would be slow, for two reasons: all chemical reactions (including redox) slow down at
low temperatures, and at great depths the oxygen concentration would not be very high.
5.3.1 – Perform a first-hand investigation to compare and describe the rate of corrosions in different: –
oxygen concentrations, – temperatures, – salt concentrations.
Differing O2: three equal-sized samples of iron were placed and submerged in three test tubes
with the following contents: one was 10mL ordinary tap-water, another was 10mL tap-water with a
layer of oil on top to lower any surface O2, third was 10mL boiled tap-water with a layer of oil.
Each day the evidence of corrosion were observed and recorded on a scale of 1 to 5.
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Differing temperatures: three equal-sized samples of iron were placed and half-submerged in
5mL tap-water; one was left at SLC, another in the incubator at 60°C, the third left in the fridge at
7°C. Each day the samples were observed and any evidence of corrosion was recorded on a
scale of 1 to 5.
Differing salt concentrations: three equal-sized samples of iron were placed and halfsubmerged in test tubes with the following contents: 5mL distilled water, 5mL of 0.5M NaCl(aq) and
5mL of 2M NaCl(aq). Evidence of corrosion was observed and recorded daily on a scale of 1 to 5.
5.3.2 – 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.
A metal wreck has sunken onto the ocean floor at 1000m. Therefore:

A low dissolved O2 concentration would mean a less corrosion: O 2 is required
for corrosion to occur and low levels mean a slower reaction and less corrosion.
A low temperature (1 to 4°C) would mean less corrosion: all chemical reactions
(including redox) slow down at low temperatures, meaning slower reaction and
less corrosion.
The lack of light will mean the lack of living organisms at that depth. Organisms
respire CO2 which forms an equilibrium with water to become carbonic acid.
The lack of acidic environments would also slow down corrosion: the reduction
potential of O2 and H+ is higher than that of O2 and H2O, therefore faster
corrosion would occur if the environment was acidic.


6. Predictions of slow corrosions at
great depths were apparently incorrect.
6.2.1 – Explain that ship wrecks at great depths are corroded by electrochemical reactions and by anaerobic
bacteria.
6.2.2 – Describe the action of sulfate-reducing bacteria around deep wrecks.
6.2.3 – Explain that acidic environments accelerate corrosion in non-passivating metals.
Shipwrecks of great depths, such as the Titanic, were found to be corroded beyond predictions.
This was mainly due to two reasons: high acidic levels around wrecks due to the respiration of
carbon dioxide from encrusting marine organisms (eg. barnacles) and the sulfate-reduction action
of certain bacteria.
Certain types of anaerobic bacteria rely on the reduction of the sulfate ion rather than oxygen to
generate energy from respiration. Sulfate reacts with hydrogen ions to form hydrogen sulphide
and water:
SO42–(aq)
+
10H+(aq)
+
8e–
→
H2S(aq)
+
4H2O(l)
Hydrogen sulfide is a weak acid, and forms an equilibrium:
H2S(aq)

H+(aq)
+
HS–(aq)
Hydrogen sulfide produced may interact with lead or silver to form black deposits of lead or silver
sulfide.
This type of bacteria flourishes in deep-ocean environments due to the lack of oxygen, and
therefore the lack of competition.
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Acidic environments promote corrosion because:
Reduction:
Reduction:
½O2(l) +
½O2(g) +
2H+(aq) +
H2O(l) +
2e–
2e–
→
→
H2O(l)
2OH–
E° = +1.23V
E° = + 0.40V
As can be seen, the reducing potential of oxygen and hydrogen ions are greater than that of
oxygen and water. Therefore the reduction of oxygen and hydrogen ions happen in preference,
and, because of the higher potential, at a faster rate. Therefore the rate of corrosion is
accelerated in acidic environments.
6.3.1 – Perform a first-hand investigation to compare and describe the rate of corrosion of metals in
different acidic and neutral solutions.
Three equal-sized samples of iron were placed and submerged in three test tubes with the
following contents: one was 10mL distilled water, another was 10mL 0.5M hydrochloric acid, the
third was 10mL 1M hydrochloric acid. Each day the evidence of corrosion were observed and
recorded on a scale of 1 to 5.
7. Salvage, conservation and restoration of objects from wrecks
requires careful planning and understanding of the behaviour of chemicals.
7.2.1 – Explain that artefacts from long submerged wrecks will be saturated with dissolved chlorides and
sulfates.
7.2.2 – Describe the processes that occur when a saturated solution evaporates and relate this to potential
damage to drying artefacts.
When a saturated solution evaporates, any salt crystallises. In the case of wooden artefacts, salt
water has entered into the cells of the wood; if they are simply evaporated, any salt will crystallise
within the cell and potentially cause damage to the artefact.
7.2.3 – Identify the use of electrolysis as a means of removing salt.
Anode: 2Cl–
 Cl2(g) + 2e–
4OH–  2H2O(l) + O2(g) + 4e–
2H2O(l)  4H+ + O2(g) + 4e–
Cathode: 2H2O(l) + 2e–  H2(g) + 2OH–
An alkaline electrolyte, NaOH, is used to prevent
the corrosion of the iron as iron is passivated in
solutions with a pH >8. The chlorine ions are slowly
extracted from the corroded metal and oxidised on
the anode. Hydrogen gas formed at the cathode
helps loosen the corrosion and other concretions.
7.2.4 – Identify the use of electrolysis as a means of cleaning and stabilising iron, copper and lead artefacts.
7.2.5 – 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.
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When an artefact is extracted, such as iron, crystallisation of salt and acidic oxidation must be
avoided:
Acidic Oxidation:
4FeCl2(s) + 4H2O(l) + O2(g) → 2Fe2O3(s) + 8HCl(aq)
Therefore, the artefact must be first stabilised by storing it in an inhibitive solution such as 2 to 5%
sodium hydroxide; this ensures no further damage to the artefact by acidic oxidation.
Crystallisation of the saturated solution is then avoided by the process of desalination, where the
object is submerged in fresh water; the salt within wooden artefacts is diluted and leached out by
osmosis. A 5% solution of polyethylene glycol (PEG) is added to the water to add support to
fragile wooden artefacts. An example of this process is a wooden ship at the Maritime Museum,
which, because it could not be submerged, was sprayed daily for ten years to desalinise it.
The artefact is then treated to remove concretions, or deposits, of CaCO 3. The concretions are Xrayed to find their size and position; they are removed very carefully by chisel and hammer, or by
dental drills. A thin of concretion is left on which will be removed by hydrogen gas during
electrolysis.
Electrolysis is used to reverse the process of corrosion: the corroded object is made into the
cathode, the anode is usually platinum and the electrolyte a solution of sodium hydroxide. For
example, a corroded silver fork can be restored as follows:
Anode: 4OH– → 2H2O(l) + O2(g) + 4e–
2H2O(l) → 4H+ + O2(g) + 4e–
Cathode:
2H2O(l) + 2e– → H2(g) + 2OH–
Ag+(aq) + e– → Ag(s)
Any silver ions will be reduced back onto the
cathode as silver metal and the hydrogen gas will
remove remaining concretions. Often this process
is done with a mesh anode, wrapped around the
cathode.
The artefact is then preserved for display usually by a combination of protections: the artefact
would most often be varnished to prevent oxygen and water re-entering; it would be kept in a
display cabinet which regulates an isolated environment and a voltage applied to it for cathodic
protection.
7.3.1 – Perform an investigation and gather information from secondary sources to compare conservation
and restoration techniques applied in two Australia maritime archaeological projects.
The Vernon 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 (which could cause unnecessary damage) and also 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
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paint. The timber stocks were saturated with a zinc napthenate solution (which retards the growth
of organisms such as mould).
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).
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).
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