shipwrecks summary complete

8/12/2019 Shipwrecks Summary COMPLETE

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2/17 Jananee Sundarakumar | Chemistry | Shipwrecks |HSC Course

Descr ibe the work of Galvani , Vol ta, Davy an d Faraday in increasing understanding of electron

transfer & Process informat ion from secondary sources to out l ine and analyse the impact of the

wo rk of Galvani, Volta, Davy and Faraday in unders tanding electron trans fer reaction s

Luigi Galvani:

o An Italian physician in the 1780s he undertook a series of investigation into the twitching of

frog leg muscles with a static electricity generator. He concluded that animal tissue contained

an electric fluid, through which a force, animal electricity, could acted upon the tissueo He also obtained twitching by pressing a brass hook into the frogs spinal cord and hanging

the hook on an iron railing.

o Galvani, studying the connection between life and electricity through investigations involving

connecting a nerve and muscle of a frog with two dissimilar metals, concluded that electricity

was generated by the animal itself.

o Impact:

This was the first example in our understanding of electrolytes in electron transfer

reactions. In this experiment, Galvani created a galvanic cell (in which an electron

transfer reaction occurs), although his interpretation of his observations was incorrect

and he did not understand the electron transfer process.

His foundational work provoked Volta to research further into electrochemistry.

Volta

o Was an existing at the same time as Galvani, proposed that, in Galvanis experiment,

electricity was generated by the

o Contact between metals in the moist electrolyte of the frogs bodily fluids thus refuting his

theory of animal electricity.

o He (incorrectly) believed that the twitching of the frog legs were due to the two different metal

pieces holding the legs. In 1800, Volta supported this hypothesis by the invention of the

Voltaic cell, a stack of silver and zinc discs separated by felt cardboard pads soaking in brine

(salt solution). This relies on an electron transfer reaction to operate.

o Impact:

Although Volta did not realise that an electron transfer reaction, as opposed to contact

between metals, was what was creating electricity, his work was significant in that it

facilitated Davys experimentation with electrolysis (another form of electron transfer

reaction) and stimulated further research into electrochemistry (such as the

construction of a useable galvanic cell; the first source of direct electrical current).

Humphry Davy:

o Explained correctly that chemical reactions resulted in creation of electricity in the Voltaic cell,

and not between two different metals.

o He also recognised that the reaction bringing about this current, where decomposition

reactions (is the separation of a chemical compound into elements or simpler compounds), he

used the understanding from this to create electrolysis apparatus and to decompose water

into hydrogen and oxygen (further proving water is compound not element).

o He developed improved versions of Volta's pile to decompose many substances. Heelectrolysed molten salts to isolate metals such as sodium, potassium, and calcium.

o Impact:

Through his experiments, he concluded that it was electrostatic forces that held

together the elements of compounds, precursory to ionic theory developed by

Arrhenius and the understanding that it was the movement of electrons that caused

current and oxidation-reduction reactions that generated electricity

He used electrolysis to decompose compounds, isolate elements, and show in

chemical reactions there was a definite connection between reaction and electricity.

These investigations led to our understanding of electron transfer reactions

as the basis of electrolysis.

Michael Faraday:o He preformed studies to relate the amount of substance produce to the quantity of electricity

that passed. Thus developed two major laws of electrolysis based on experimental evidence.

The first states: the amount of product of electrolysis is proportional to the quantity of

electrical charge that has flowed through the circuit. (Q=It)



The second states: the masses of substances produced are proportional to the molar

masses of each substance

o Impact:

Allowed the electron transfer reaction of electrolysis to be understood quantitatively

His discoveries supported the suggestion made later that electricity consisted of small

particles (electrons), contributing to the understanding of particle theory and electron

transfer that we have today.

He was the first person to define the laws of electrolysis and quantify electrolysis. Healso 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

Acco unt for the di f ferences in corros ion of act ive and passivat ing metals

Corrosion: the degradation (through oxidation) of a metal, so that it loses strengths and becomes

unable to fulfil its intended purpose.

o The most common type of corrosion is known as rusting.

A passivating metal is a reactive metal that readily forms an unreactive, impervious surface coating

(usually an oxide), by the reaction with oxygen or water. It protects the metal from further reaction (i.e.

corrosion).

o An example of an passivating metal is aluminum, it forms a layer of aluminum oxide

4Al (s)+ 3O2(g) 2Al2O3(s)

Prevents oxygen going through the rest of the aluminum metal, hence no further

corrosion.

Other metals in include Zinc (Zn) {the passivation of Zinc is called galvanising},

Chromium (Cr) which forms Cr2O3, and stainless steel (which also contains chromium

to passivate through).

Note: some non-metals such as copper also form a protective layer of oxide (or

carbonate) but these are not termed passivating, as they are not reactive metals. An active metalis a reactive metal, such as magnesium or sodium, that does not passivate (i.e. form a

protective coating against further reaction). Most are located in the first two columns of the periodic

table.

Ident i fy iron and steel as the main m etals used in ships

Since 1860s ocean-going ships have been made from steel and iron.

They both have the advantages of being hard, high mechanical strength, and can be welded.

Ident i fy the com posi t ion of steel and expla in how the percentage comp osi t ion of steel can determine

i ts proper t ies & Gather and process informat ion from secondary sources to compare the

comp osi t ion, proper t ies and uses a range of steels

Alloysare partial or complete solid solution of one or more elements.

In general, alloys are harder than pure metals, as the presence of differently sized atoms interrupts the

orderly structure of the metal lattice, preventing atoms in the lattice from sliding over each other.

In general there are 2 different types of alloys

o Homogeneous: meaning the presence of different atoms is evenly spread out

o Heterogeneous: meaning the presence of different atoms is not evenly spread out

Steelis an alloy of iron, carbon (no more than 2%), and often varying amounts of other metals.

An increase in any of the element metals leads to different types of steel. The choice of steel for a

specific purpose depends on the properties necessary for that purpose as well as the environment the

alloy will be exposed to.
http://www.mondofacto.com/facts/dictionary?tablehttp://www.mondofacto.com/facts/dictionary?table



Element Property effects

Aluminium Steel becomes easily shaped, and forms passivating layer

Nickel High corrosion resistant

Tungsten High strength at high temperatures

Vanadium Wear resistance

Composition, properties and uses of a range of steels

o NOTE: the rest of the composition is made up of IRON

Descr ibe the condi t ions und er which rust ing of iron occ urs and expla in the process of rust ing &

Use avai lable evidence to analyse and expla in the con di t ions und er which ru st ing occu rs

Rusting: the most common form of corrosion that involves redox reactions

o It produces red hydrated iron oxide (Fe2O3.H2O)

Conditions for rusting:

o Iron: the metal to be oxidised.

o

Water: acts as an electrolyte to connect the anodic and cathodic sites, and is also reduced(discussed below).

o Both oxygen and water are necessary: O2 and H2O are needed because together they are

reduced to form hydroxide ions.



o Impurity sites in the iron: these are needed to act as cathodes, where the reduction (gain of

electrons) of O2can occur.

This is clearly seen through the slow process of rusting in pure iron, it doesnt have

good reducing surface

Conditions that affect rusting:

o Salt water:

Salt water accelerates rusting because it is an electrolyte and therefore more

conductive than fresh watero Concentration of oxygen:

Oxygen concentration pushes equilibrium to the right therefore increasing the amount

of hydroxide produced, increases the chances for the formation of iron oxide

Reduction: O2 (g) + 2H2O (l)+ 4e

4OH

(aq)

Furthermore the oxygen is needed to form rust, i.e. it receives electrons from

iron oxide (gets reduced). So an increase in oxygen concentrations increases

all the above rates.

o Mechanical stress:

When iron is under stress, example at bends, sharpened joins or edges, the stress

disrupts the metallic lattice orderly crystal structure of iron, hence becomes distorted

and this makes it easier for individual Fe atoms to break away from the crystal as Fe2+

ions.

o Acidic electrolyte:

In the process of rusting, when iron is oxidised, water is reduced. The reaction is:

o O2(g)+ 2H2O (l)+ 4e- 4OH

-(aq)

By Le Chateliers principle, as [OH-] becomes small (i.e. the electrolyte becomes

more acidic), the equilibrium will favour the forward reaction and hence the corrosion

of iron.

o Temperature:

As with all chemical reactions, corrosion occurs more rapidly at higher temperatures,

as high temps. increase rates of collision

o

Contact with other metals: If iron is in contact with metals that are weaker reductants, such as tin or copper, it

rusts more readily as less active metals act as cathodic sites.

This process is called galvanic corrosion.

Process of rusting:

When iron metal is in contact with oxygen AND water (sometimes in the form of moisture), at anodic

sites, iron atoms oxidise (i.e. lose electrons) to form Fe2+

ions.o Oxidation: Fe (s) Fe

2+(aq) + 2e

E = + 0.44V

The electrons from this half-reaction flow to the cathodic sites, usually an impurity such as carbon (i.e.

where reduction (gain of electrons) can take place)



o There, the electrons reduce the oxygen dissolved in the thin film of moisture on the irons

surface OR in the water touching the iron (if the iron is submerged)

o Reduction: O2 (g) + 2H2O (l)+ 4e

4OH

(aq) E = + 0.40V

o Note: in order for this galvanic cell to continue to operate, there has to be a migration of ions

through the moisture layer from one location to the other. This migration of ions is to preserve

the electrical neutrality (for the flowing of electrons that are needed for reduction/oxidation

reactions). This is why salt solutions are ideal, as they provide an ample supply.

Simplified rusting formula:

o 4Fe (s) + 3O2 (g)+ 2H2O (l) 2(Fe2O3.H2O)(s)

3. Electrolytic cells involve oxidation-reduction reactions

Descr ibe, using hal f -equat ions, what happens at th e anode and cathod e dur ing electro lys is of

selected aqueous solut ions

Galvanic Cell Electrolytic Cell

Chemical Electrical energy Electrical Chemical energy

Produces electrical current Requires an input of electrical current

Spontaneous reaction Reaction is forced

Anode is NEGATIVE () Anode is POSITIVE (+)Cathode is POSITIVE (+) Cathode is NEGATIVE ()

E0is positive (electricity produced) E0is negative (electricity required)

Anode= oxidation, cathode = reduction Anode= oxidation, cathode = reduction

Electron flow= oxidation reduction Electron flow= oxidation reduction

Electrolysis of concentrated NaCl solution

o Diagram:

o Equations:

Oxidation: 2Cl-(aq)Cl2(g)+ 2e

-(bubbles of Cl2formed)

Reduction: H2O(l)+ e-1/2H2+ OH

-(bubbles of H2formed)

Redox: 2H2O(l)+ 2Cl-(aq)H2(g) + 2OH

-(aq) + Cl2(g)

Electrolysis of molten NaCl

o Diagram:


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o Similarly, if the Cl- (aq) concentration is decreased, the E will decrease even further, hence

favouring the water oxidation.

Nature of electrolyte:

o Products vary according to the nature of electrolytes.

o For example in molten electrolytes, there exists only 2 substances (the anions and cations);

hence the yield is a product at each electrode. For example, the electrolysis of molten NaCl,

others include magnesium chloride (MgCl2) which yields magnesium and chlorine gas.

o However, electrolytes that consist of aqueous solutions may result in water being reacted atthe electrodes in preference to the other alternatives (that are wanted).

o For example, an aqueous solution of magnesium chloride would yield hydrogen gas at the

cathode rather than magnesium. This occurs because the E value for the reduction of water

to hydrogen, has a LOWER numerical value that that for the reduction of magnesium ions

(Mg2+

).

o And again at the cathode (though depending on the concentration) the oxidation of water to

oxygen has a lower numerical value (-1.23V) than chloride ions into chlorine gas (-1.26V), so

it will preferably oxidise.

o This happens when there are two species present that have similar tendencies to be oxidised

or reduced; that is their standard electrode potentials are similar.

Nature of electrodes

o Reactions that are out using inert electrodes, will simply act as conductors, or a site where

oxidation and reduction can take place, IN PREFERENCE to the electrolyte.

o For example, in the electrolytic refining of copper, blister copper is used as the anode and the

cathode is pure copper. The blister copper electrode oxidises, while layers of pure copper

build up around the inert cathode.

o At the anode the reaction is Cu(s) Cu2+

(aq) + 2e-, however if another inert electrode is

added (i.e. there are now 2 inert electrodes), the anode reaction simply becomes 2H2O (l)

4H+(aq) + O2(g) + 4e

-, this is because there is no other solid wishing to be oxidised

except water (remember copper sulfate exist as ions in the solution).

4. Iron and steel corrode quickly in a marine environment and must be protected

Ident i fy the ways in which a metal hul l may be protected including corrosion resistant m etals,

developm ent of sur face al loys, new p aints

Corrosion resistant metals:

o A common way to protect a metal is simply using one that does not corrode in the first place,

that is use a substitute material.

o This is done in buildings and parts on the ship, where aluminium has nearly completely

replaced all steel in windows, and door frames.

Aluminium can form a passivating layer, hence protecting the iron.

o Other examples include, stainless steel (that does not corrode) has replaced carbon steel,

where it is able to take in high heat (heat favours any chemical reaction and induces rusting). Another reason is that they can passivate (form chromium (III) oxide which doesnt

allow oxygen to react with bare metal), that is to form a protective layer of oxide which

prevents rusting.

Development of surface alloy:

o One method is through the use of stainless steel-like properties on the surface of a metal hull.

It would be able to resist rusting, as the bare metal under it is not exposed.

This is done by bombarding a hull with ions of chromium and nickel

The metal ions are formed in a high temperature gaseous discharge (plasma) and are

directed onto the surface of ordinary steel where they become embedded as atoms

and so form a surface alloy

This stainless steel surface resists corrosion. New paints:

o Painting a steel object protects it from oxygen and water and hence from corrosion.



o This is an inexpensive and potentially effective method. However, the paint needs regular

maintenance (which is a nuisance for submerged steel objects) and can pollute the

environment by leaching out chemicals.

o Polymer-based paints, such as Rustmaster pro,are a fairly recent development.

When applied to steel, a polymer layer impervious to water and oxygen is formed.

Additives in the paint also react with the steel to form an intermediate layer of highly

insoluble ionic layer pyroaurite.

This ionic layer is tightly bonded to the steel surface and extends well into the polymerlayer. It prevents the movement of ions hence rusting.

This paint is effective, does not require maintenance as frequently, and does not

release chemicals into the environment.

o Phosphate coatings are used on steel parts for corrosion resistance, or as a foundation for

subsequent coatings or painting.

It serves as a conversion coating in which a dilute solution of phosphoric acid and

phosphate salts is applied via spraying or immersion, chemically reacts with the

surface of the part being coated to form a layer of insoluble, crystalline phosphates

i.e. it forms iron (III) phosphate, which is insoluble and adheres to the iron tightly,

forming a protective layer

Predict the metal which corrodes wh en two m etals form an electroch emical cel l using a l is t of

standard potent ia ls

In a galvanic cell the metal that corrodes (oxidises), is the one that has the highest oxidation potential

(this can be obtained by flipping the reduction potential table values).

This is because, the one with higher oxidation potential is the one that wants to OXIDISE MOST.

Out l ine the process of cathodic p rotect ion, descr ib ing examples of i ts use in bo th m ar ine and w et

terrestr ial environments

Cathodic protection(also known as galvanic protection):

o the process of protecting a metal, by making a metal the cathode (reduction, ie gains

electrons) of an galvanic cell (electrochemical cell) in order to prevent its oxidation.

o Marine and terrestrial environments, forms of cathodic protection such as galvanised steel,

sacrificial anodes and applied voltage are used.

Descr ibe the process of cathodic protect ion in s elected examples in terms of the oxidat ion/reduct ion

chemistry involved

Sacrifical electrodes (more specifically anodes):

o A sacrificial electrode is usually a bloke of zinc or magnesium (or anything that will preferably

oxidise in preference to iron). This is done by the galvanic action, the more reactive metal, in

this case (Mg or Zn) oxidise sacrifices preferentially to the iron, giving up electrons to the iron

and so preventing iron oxidising to form Fe

+2

. Fe(s) Fe

2+(aq) + 2e

E = + 0.44V

Zn(s) Zn2+

(aq) + 2e

E = + 0.76V

o Since the oxidation potential of zinc is greater than that of iron, it will oxidise in preference (i.e.

become the anode) and render the iron as the cathode

o If any Fe+2

ions did form, they would be reduce, that is be converted back to iron solid (Fe (s))

by the electrons given up by Zn or Mg.

Zn (s) + Fe2+

(aq) Zn2+

(aq) + Fe2+

(aq)

Note: since the action is galvanic, it is essentially a galvanic cell.

However galvanic cells require an electrolyte (i.e. solution containing free ions

that make it electrically conductive), so too does sacrificial electrodes.

Hence the medium should be sea water/moist earth, such that the build ofMg

+2 ions is balanced, and electrons can flow without resistance from the

solid block of Mg to the iron.



o Also Mg should not be used in moist ground, as it can react with water violently and

deteriorate it much quicker.

Applied voltages (also known as impressed current/voltage):

o Another form of cathodic protection is to use an inert (i.e. non-sacrificial) anode such as

platinum-coated titanium. A suitable voltage is applied between it and the metal wanted to be

protected, rendering it cathodic (i.e. hull of ship, or underground pipes).

o This voltage forces electrons into potentially active sites on the steel hull and so prevents

oxidation.

Fe2+

(aq) + 2e

Fe (s) E = + 0.44V

o However because there may be extra electron flow from the voltage source, water (as

moisture) may also react

Anode at : 2H2O (l)+ 2e- H2(g)+ 2OH

-(aq) ORO2 (aq)+ 2H2O (l)+ 4e

-

4OH-

Cathode: 2H2O (l) 4H+

(aq) + O2 (g)+ 4e-

Ident i fy data, gather and process informat ion from f i rst -hand or secondary sources to trace

histor ica l developm ents in the choice of m ater ia ls used in the construc t ion of o cean -going vessels

with a focus on the m etals used

Wood was the most common shipbuilding material until the nineteenth century.

Metals were used in some early ships:

o Viking longboats which had iron and bronze fittings to cover oak planks.

o Greeks and Romans using iron to replace the weaker copper.

Around 1500AD Lead and then copper sheeting was fixed to the hull of ships to prevent biofouling by

marine organisms, iron nails made it possible to connect.

o This made the hull stronger and less flexible, but the nails were susceptible to rapid corrosion.o Copper nails were used in their place.

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

o 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.

o Iron was the preferred metal because it was readily available by that time, and because of its

hardness, and strength.

o

However, the process of making steel was unsophisticated and difficult to control, resulting inhigh amounts of carbon (0.2%), sulfur (0.1-0.2%), and phosphorus (0.1%) and a very brittle

steel, especially at low temperatures.



In the early 1900s the production of steel improved after World War II, when manufacturers were able

to remove sulfur and phosphorus from steel, and the invention of electric welding meant faster and

better construction of steel ships.

o Companies experimented with the addition of elements such as manganese,

o Ships became more durable and tougher, but much less brittle.

Other developments in the twentieth-century included the progressive improvement in steel alloys,

incorporating aluminium, chromium, titanium, zinc and nickel.

o Modern steels are lighter, stronger and more corrosion-resistant than before. Other developments at this time:

o Bronze (alloys of copper and tin) was used instead of steel as propellers (because they resist

corrosion even when exposed to rapidly moving, highly aerated water).

o Aluminium, being resistant to corrosion and extremely light, was used in the superstructures

and fittings in ships.

o Stainless steel, being extremely resistant to corrosion but very expensive, was used in fittings

and equipment of ships.

The metal predominantly used in ships is still steel, primarily because of the abundance and low cost

of iron, and the tensile strength, and malleability of steel.

o Improved methods of protecting steel from corrosion are constantly being developed.

Gather and proc ess inform at ion to ident i fy appl icat ions of c athodic protect ion, and use avai lable

evidence to ident i fy the reason for their use and th e chemistry invo lved

Uses and applications:

o Sacrificial anodes are often used for accessible steel structures, for example:

Ships hulls, rudders and propellers

Water tanks (above ground).

o Impressed current systems are used for:

Underground steel structures (because sacrificial anodes would not be able to be

replaced) such as pipelines and fuel storage tanks.

Ship hulls, rudders and propellers.

Reasons for use:o Both effectively prevent the corrosion of iron.

o Both are easily maintained; maintenance is simply replacing corroded sacrificial anodes.

o Both are relatively inexpensive corrosion prevention methods, although impressed current

systems can be expensive to set up.

o Impressed current systems release no pollutants (whereas sacrificial anodes release metal

ions).

o Impressed current systems allow the consistent prevention of corrosion of buried steel

structures.

5. When a ship sinks, the rate of decay may be dependent on the f inal depth of the wreck

Out l ine the ef fect of temperature and pressure on th e solubi l i ty of gases and sal ts

Gases Salts

Increase in temperature Decrease solubility Generally increases

Increase in pressure Increase solubility Nothing

And vis versa

Ident i fy that gases are normal ly dissolved in the oceans and com pare their concentrat ions in the

oceans to their concentrat ions in the atmosphere

Gases are normally dissolved in the oceans

Gas Concentration in the Concentration in ocean (at surface)



atmosphere

Oxygen 21%

1.5 x 10-5

mol/L

Up to 8.5ml/L

2.8 x 10-4 mol/L

Nitrogen 79%

3.3 x 10-2

mol/L

Very littlelow solubilty

5.6 x 10-4

mol/L

Carbon

dioxide

0.04%

1.5 x 10-5

mol/L

Up to 46ml/L

1.4 x 10-5

mol/L

Compare and expla in the solubi l i ty of s elected gases at increasing depths in the oceans

Carbon dioxide

o Carbon dioxide has a higher concentration in sea water (16%) than its concentration in the

atmosphere would lead one to expect. This is due to the fact that carbon dioxide can react

with water to form carbonic acid.

o Carbon dioxide is consumed near the surface of the ocean in the photosynthesis of plants.

Further down, it is still produced by the processes of respiration and decomposition, but there

are few plants to consume it, so its concentration is much higher. It also has a very high

solubility under the conditions present in very deep sea water (low temperature, high

pressure).

Nitrogen

o Nitrogen makes up approximately 48% of the gas dissolved in water at the oceans surface,

compared to 78% of the gas in the atmosphere (v/v). Even though nitrogen is not very soluble

in water, it has a high concentration in water due to its high partial pressure. Because there

are few processes that remove or create nitrogen in the ocean, its concentration does not

change heavily with depth.

Oxygen

o Near surface:

As well as being diffused from the atmosphere, oxygen is produced in water by the

photosynthetic action of phytoplankton. Since they require light to photosynthesise, they

are only found in the top 100 metres of the ocean.

o 100-1000m below surface:

Oxygen is still consumed in the process of respiration, but is not as readily replenished by

contact with the atmosphere and photosynthesis.

o Below 1000 metres:

Deep ocean currents bring some oxygen-rich water from the surface.

Oxygen makes up approximately 36% of dissolved gases at the surface, compared to 21%

of the atmosphere (v/v). It has a greater concentration in the ocean because it is more

soluble than nitrogen, and because it is produced by phytoplankton.

Predict the ef fect of low temperatures at great depths on the rate of corrosion of a metal& Use

available evidence to predic t the rate of corrosio n of a metal wreck at great depths in the oceans and

give reasons for the predict ion m ade

The predicted effect of low temperatures at great depths on the rate of corrosion is that the corrosion

would be slow, for two reasons:

o All chemical reactions (including equilibrium reactions) slow down at low temperatures,

meaning slower reaction and less corrosion.

Note: yes the 'exothermic' reaction is favoured in the equilibrium reaction, and would be

the reaction that is created, however it is at an unbelievable slow rate.

o Another reason is, which is probably the most important; lack of light, this will mean the lack of

living organisms at that depth due to photosynthesis to produce oxygen:

A low dissolved O2 concentration would mean a less corrosion: O2 is required for

corrosion to occur and low levels mean a slower reaction and less corrosion.

6. Predict ions of slow corrosion at great depths were apparent ly incorrect

Expla in that ship wrecks at great depths are corroded by electrochemical react ions and by

anaerobic bacteria & Descr ibe the act ion of sul fate reducing bacteria around deep wrecks



Before the shipwrecked Titanic was discovered, it was predicted that the ship would corrode very

slowly because of the low temperature, low salt concentration, and low concentration of dissolved

oxygen. However, the ship was found to be corroded extensively due to a number of factors.

For shipwrecks in deep ocean water, some of the corrosion is caused by the normal electrochemical

process of iron oxidising and reacting to form rust.

However most of the corroision is by anaerobic (requiring no oxygen) sulfate-reducing bacteria. They

respire (obtain energy) by reducing sulfate to sulfur. An example, is the bacteria Desulfovibrio family.

o SO42- (aq) S (s) S-2(aq) For these reactions to occur they need electrons (as the above are all reducing reactions), this comes

from organic compounds such as carbohydrates (i.e. CH2O) that oxidise and release electrons.

o CH2O (aq)+ H2O (l) CO2(g)+ 4H++ 4e

-

Hence the bacteria allow the corrosion process to occur even in the absence of oxygen.They tend to

populate shipwrecks because they provide an excellent source of nutrients, and allow the corrosion

process to occur even in the absence of oxygen.

o The SUM reduction reaction is:

SO42-

(aq)+ 5H2O (l)+ 8e- HS

-(aq)+ 9OH

-(aq)

SO42-

(aq)+ 10H+ (aq)+ 8e- H2S (aq)+ 4H2O (l)

o The oxidation reaction is:

4Fe (s) 4Fe2+

(aq) + 8e (it has been multiplied by 4 so we can balance electrons)

o Net ionic equation:

4Fe (s)+ SO42-

(aq)+ 5H2O (l) 4Fe2+

(aq) + HS-(aq)+ 9OH

-(aq)

o The Fe2+

then react with HS-and OH

-(aq) to form insoluble FeS and Fe(OH)2

4Fe2+

(aq)+ HS-(aq)+ 7OH

-(aq) FeS (s)+ 3Fe(OH)2 (s)+ H2O (l)

o Upon adding the above 2 reactions, and cancelling one H2O on each side, the overall reaction

is:

4Fe (s)+ SO42-

(aq)+ 4H2O (l) FeS (s)+ 3Fe(OH)2 (s)+ 2OH-(aq)

This form to form insoluble black iron (II) sulfide and iron (II) hydroxide along the steel (if there is

oxygen this hydroxide will become converted to rust).

These compounds combine to form finger growths, that are reddish-brown known as resticles.Rusticles also contain many other compounds such as calcium carbonate (CaCO3).

Expla in that acid ic environments accelerate corrosion in non -passivat ing m etals

Now in rusting, the normal reduction reaction, at the cathode is:

o O2(g)+ 2H2O(l)+ 4e

4OH

E = + 0.40V

When working with acidic solutions, the half reaction is usually put in terms of H+, rather then OH

.

This is done by adding 4H+ to each side of the equation above, then 'convert' 4OH

+ 4H

+

4H2O(l). Then cancel 2 H2O from each side to obtain:

o O2(l) + 4H+

(aq) + 4e H2O(l) E = +1.23V

Both equations represent the same thing, oxygen reducing into water; all that changes is how the

acidity of the equation is represented.

Now there are 2 reasons for why acidity accelerates corrosion.

o The half reaction above has a standard potential of 0.4V; this is the value at [OH

] = 1 mol/L,

as the acidity increases (i.e. the concentration of OHdecreases), the following occurs. For

example a pH of 7, [OH] = 1.0 x 10

-7mol/L, the non-standard electrode potential is 0.81V, at

the pH = 4. At pH 3, [OH

] = 1.0 x 10-11

mol/L, it is 0.99V. Thus the fact that the electrode

potential increases means that the half reaction has a greater tendency to occur as acidity

increases.

o By Le Chateliers principle, looking at equation 2 above (doesn't really matter), as the

electrolyte becomes more acidic, [OH] becomes small, hence the equilibrium will favour the

forward reaction and hence the corrosion of iron.

7. Salvage, conservation and restoration of objects from wrecks requires careful planning an

understanding of the behaviour of chemicals



Expla in that ar tefacts from long su bmerged wrecks wi l l be saturated with dissolved chlor ides and

sulfates

Aretefactsare objects from past human societies, they are valued as they provide insight into history.

When long-submerged objects have been recovered, they are generally in very poor condition.

o Wooden artefacts will usually have all their cellulose eaten away, and there pores will

impregenated with salts and other ions such as chlorides and sulfates occupy the pores of an

artefact (microscopic in the case of metals) and saturate the object.o Since most artefacts, are usually metals, which are cations, they will ALSO attract anions in

their reactions by precipitation.

Descr ibe the processes that occur when a saturated solu t ion evaporated and relate th is to the

potent ia l damage to dryin g ar tefacts

When a solution of a solid such as sodium chloride (NaCl) is evaporated, the solution becomes more

concentrated, as there is less water to dilute the solution. The water evaporates and hence the salt

remains.

Eventually it becomes saturated with salts, the water is removed, and the solid starts to crstyallise until

the solid structure forms.

This formation of solid salt crystals throughout porous objects causes an increase in pressure, thus

cracking and damage by distorting the shape of the object, or even reaction that can further accelerate

corrosion.

o 4FeCl2 (aq)+ 10H2O (l) + O2 (g) 4Fe(OH)3 (s)+ 8HCl (aq)

o 4FeCl2(s)+ 4H2O(l)+ O2(g) 2Fe2O3(s)+ 8HCl(aq)

Ident i fy the use of electro lys is as a m eans of removing sal t

Artefacts that have long been submerged, will often carry a huge amount of salt content. One method

of removing these salts is to leach them out of the object.

The process involves putting the object in fresh distilled water for a period of weeks -months,

depending on how high the concentration of salt is.

The action involves osmosis, where molecules move from an higher of higher concentration (the

object) to an area of lower solute concenctration (the water bath). Water is regularly changed toaccelerate this process.

Generally, dilute sodium hydroxide (NaOH) is used then distilled water, as it speeds the removal of

chloride (though it doesnt not matter). It does this because the hydroxide ion (OH-) can replace

chloride (Cl-) in iron, which exists as Fe(OH)2, this will be removed by further means.

MOST chloride is SOLUBLE and present as iron (ii) chloride (FeCl2) or iron (iii) chloride (FeCl3) or

copper (ii) chloride (CuCl2).

HOWEVER, There are some salts are extremely hard to remove by these conventional means, they

are insoluble hydroxy chloride salts. If the object is iron they include Fe(OH)Cl, or if copper Cu(OH)Cl.

These chlorides react with water and oxygen to form rust and hydrochloric acid which even further

expediting the corrosion process:

o Fe(OH)Cl (s)+ H2O (l) Fe(OH)2 (s)+ HCl (aq)o 4Fe(OH)Cl (s)+ 6H2O (l)+ O2 (g) 4Fe(OH)3 (s)+ HCl (aq)

o Similarly for copper compounds.

A more efficent way of removing chloride from such objects, is through electrolysis.

Ident i fy the us e of electro lys is as a m eans o f cleaning and stabi l is ing iron, copper and lead artefacts

Conservation:Cleaning, stabilisation and preservation of an artefact to prevent further deterioration,

involving as little interventive treatment as possible.

Restoration:Returning an artefact to a condition as similar to its original as possible.

The metal object is made the cathode (such that it reduces, gains electrons), and through this the

salts (or even if theres any iron (ii) hydroxide that is wanting to form rust) can accept electrons, then

following occurs:

o Fe(OH)2 (s)+ 2e-

Fe (s)+ 2OH-(aq)

o Fe(OH)Cl (s)+ 2e-

Fe (s)+ OH-(aq)+ Cl

-(aq)

The above reactions can basically be represented as Fe2+

(aq)+ 2e-

Fe (s)



This means that the salt (or iron(ii) hydroxide) is converted to solid iron, which builds up

on the object as that is the cathode (site of electron acceptance), and the rest of the ions

flow into the solution.

This effectively means, salt is removed, and the needed part of that salt is used to repair

the worn object.

Discuss the range of c hemical procedures wh ich can be u sed to clean, preserve and s tabi l ise

ar tefacts from w recks and, where possib le, provide an example of the use of each pro cedure

Artefacts from ships include anything that was used on that ship, it can include wooden objects, lead ,

copper, metal etc. However the main things are often restored as:

Silver artefacts

o Generally silver doesn't corrode, until the Desulfovibrio bacteria reduce sulfate to sulfide, the

following reaction can occur:

SO42-

(aq)+ 10H+ (aq)+ 8e- H2S (aq)+ 4H2O (l)

o Then the silver solid reacts with dihydrogen sulfide to form insoluble silver sulfide.

2Ag (s)+ H2S (aq) Ag2S (s)+ H2(g)

o Silver sulfide forms as relatively firm layer on the object and prevents it from corroding, also

this layer further becomes covered by calcium carbonate (also known as concretions). They

form by the following precipitaiton reaction: Ca

2++ CO3

2- CaCO3(aq)

o The layers of concretions (which can become very thick) are removed by dissolving them in

acid.

CaCO3(s)+ 2HCl (aq) CaCl2(aq)+ H2O (l)+ CO2 (g)

After these have been removed, it will leave black objects (as there is black silver sulfide

remaining) , it will be too abrasive to clean the the silver sulfide by acid, as the silver can

also react with acid and damage engravings and embossings, something so fragile as

coins need to have as much detail as can be.

o The formation of silver sulfide can be reversed by making the corroded object the cathode in

an electrolysis cell

The reduction (cathode): Ag2S (s)+ 2e- 2Ag (s)+ S-2(aq) An inert electrode is used as the site of oxidation (such as platnium or stainless steel),

also a solution of sodium hydroxide is chosen as it facilitates (makes easier) the reduction

reaction to occur. Hence the oxidation (anode): 2H2O (l) 4H+

(aq) + O2 (g)+ 4e-

In this way, the silver sulfide is converted back to silver which embeds back to the object,

with little damage to the markings on the object. If further stabilisation is required, clear

laquer is applied (though not necessary).

Note: silver chloride (AgCl) salt , or any other types of insoluble salts , rarely forms on

silver.

Iron

o Iron artefacts are usually covered in concretions of calcium carbonate or coral, which is

usually removed mechanically carefully by a hammer (MRI magnetic resonance image canbe used to display this layer).

o The piece is saturated with soluble chlorides and sulfates, and covered in insoluble iron

compounds such as hydroxides, oxides, and hydroxychlorides.

o Desalination, which is the removal of salts from an object, is done through the washing of the

object for extended periods of time.

Usually dilute (~0.5mol/L of sodium hydroxide, or carbonate or bicarbonate).

The reason for this, is it includes the fact that it makes the solution more alkaline hence

inhibiting further corrosion, as it is favoured in acidic environment

However the main reason is the fact that they accelerate the removal of chloride ions this

is done by the reaction Fe(OH)Cl (s)+ X-(aq) Fe(OH)X (s)+ Cl

-(aq)

The X-

can be OH-

, HCO3-

or CO32-

, when this form with iron, it can be removed beelectrolysis, at least stopping the damaging of chloride.

o Electrolysis can be performed to remove soluble (the ones that are left) and insolubes

chlorides compounds.



The metal object is made the cathode (such that it reduces, gains electrons), and through

this the salts (or even if theres any iron (ii) hydroxide that is wanting to form rust) can

accept electrons,

Then following occurs:

FeCl2 (aq)+ 2e-

Fe (s)+ 2Cl-(aq)

FeSO4 (aq) + 2e-

Fe (s)+ SO4-2

(aq)

Fe(OH)2 (s)+ 2e-

Fe (s)+ 2OH-(aq)

Fe(OH)Cl (s)+ 2e-

Fe (s)+ OH-(aq)+ Cl

-(aq)

The above reactions can basically be represented as Fe2+

(aq)+ 2e-

Fe (s)

This means that the salt (or iron(ii) hydroxide) is converted to solid iron, which builds up

on the object as that is the cathode (site of electron acceptance), and the rest of the ions

flow into the solution.

This effectively means, salt is removed, and the needed part of that salt is used to repair

the worn object.

At the cathode, when the voltage is abit higher then normal the following reaction occurs:

2H2O (l)+ 2e-

2OH-(aq)+ H2 (g)

The following reaction causes the production of hydrogen bubbles at the cathode, which

helps loosen any remaining concretions. Copper

o Copper artefacts do not corrode badly because copper is unreactive and poisonous to marine

organisms. However, they do form surface layers of CuCl2and Cu(OH)Cl.

The surface of the artefact is cleaned by chemical stripping; it is placed in a 5-10% citric

acid and 1-2% th iourea (a corrosion inhibitor) mixture to dissolve surface deposits.

Chlorides can be removed by soaking in sodium hydrogen carbonate or pure water.

Electrolysis can be performed with sodium carbonate electrolyte (NaOH damages alloys

of copper).

o Then following occurs at the cathode (reduction):

CuCl2 (aq)+ 2e-

Cu (s)+ 2Cl-(aq)

Cu(OH)Cl (s) + 2e-

Cu (s)+ OH-

(aq)+ Cl-

(aq)o Then following occurs at the anode (oxidation):

2H2O (l)+ 2e-

2OH-(aq)+ H2 (g)

The following reaction causes the production of hydrogen bubbles at the cathode, which

helps loosen any remaining concretions.

o The artefact is washed and coated with crystalline wax or clear acrylic lacquer.

Lead

o Lead artifacts have few concretions because lead is poisonous to marine organisms (but still

can form). Lead oxidises slowly to form insoluble: PbS, 2PbCO3.Pb(OH)2, PbCl2, or Pb(OH)Cl

(You can check this with solubility rules).

o Usually lead artifacts include cannon balls, weights, and lead sheeting.

PbS and CaCO3concretions are removed with 10% hydrochloric acid: PbS (s)+ H

+(aq) Pb

2+(aq)+ H2S (aq)

Other insoluble lead compounds are removed with 5% EDTA solution in an ammonia-

ammonium buffer; EDTA forms complexes with lead ions, allowing the insoluble

compounds to dissolve.

Electrolysis can be used to remove hydroxychlorides with sodium carbonate electrolyte.

o Then following occurs at the cathode (reduction):

PbCl2 (aq)+ 2e-

Pb (s)+ 2Cl-(aq)

Pb(OH)Cl (s) + 2e-

Cu (s)+ OH-(aq)+ Cl

-(aq)

o Then following occurs at the anode (oxidation):

2H2O (l)+ 2e-

2OH-(aq)+ H2 (g)

The lead is finally washed with H2SO4and distilled water, and is then dried thoroughly and

coated with crystalline wax.

Wood

o When wood is immersed in sea water, there is rapid decay, wood worms known as toredo

worms and various bacteria, fungi attack the wood, quickly eating its useful cellulose and


17/17

leaving the lignin behind (which is what maintains its shape). Special circumstances can

however slow down the process:

If the wreck or object is buried in silt or mud, which excludes the worms and aerobic

bacteria, this reduces the rate.

If the water is very cold, unsalted, and not in great water current movement (ie near the

ocean floor reaches 40C, and 0.4% of salt such as the Baltic sea, the water movement

over the wreck restricts the inflow of fresh oxygen for bacteria).

o When long submerged wooden objects, they are very fragile (due to missing cellulose), andthe salt is dangerously damaging if it is left to dry, at it allows to shrink, wrap and split.

o Wooden artifacts are first cleaned of silt and mud by spraying them with cold water for an

extended period of time and gently brushing them.

o The water impregnating the wood is replaced by poly(ethylene glycol), which is a inert wax (or

oil as well), a polymer soluble in water with the structure (-CH2-CH2-O-)n- note: different

values of 'n' vary the properties such as viscosity (sticky) and melting points, hence why it can

be called a wax or oil.

o This fills all the cavities (i.e. missing cellulose) in the wood and stabilizes it by restoring some

of its structural strength and stopping further degradation.

o Final the object is dried to remove residual water and coated with higher melting point

poly(ethylene glycol) wax.

Perform an invest igation and gather informat ion from s econdary sources to com pare conservation

and restoration techniques appl ied in two Austra l ia mar i t ime archaeological projects

Endeavourcannons Vernon anchors

- Cannon was stored in mixture of seawater and

10% formalin (to kill bacteria).

- Coral was removed mechanically with

hammers

- Cannon was placed in a 2% NaOH solution (to

prevent acidic corrosion). It was electrolysed

with a current density of 10A/m2, with cannon

as cathode and mild steel as anode. After the

baths, the cannons were washed with fresh

water at each refreshing of the electrolyte (The

electrolyte was replaced after two weeks, then

weekly until chloride concentration stabilised at

20 ppm.)

- Cannon was placed in distilled water with

chromate ions (8.7 x 10-3

mol/L) to remove

chlorides and hydroxides and form a protective

layer of iron oxide and chromic oxide.

- Cannon was dried for two days at 120oC, and

then immersed in crystalline wax (as a further

protective layer) for five days, at 135oC to

ensure maximum penetration of wax.

- Wax was cooled to just above melting point.

- Surface corrosion and remains of protective

paint were removed by blasting (smoothing

rough surface) with copper slag.

- Iron was treated with zinc epoxy paint.

- Timber components were saturated with zinc

naphenate solution to prevent growth of

organisms

- 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.

- The anchors are supported by an aluminium

mesh (cathodic protection).

shipwrecks summary complete

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