Choosing an option

Two major considerations for teachers In selecting the one option (of the five available) for their students to study, are first their students' likely interest in the topic and secondly their ability to cope with the amount and complexity of the material. Other considerations are likely to be how well the material integrates with and reinforces core material and how straight-forward and/or predictable are likely examination questions on the material. The availability of equipment for laboratory work is another consideration.

Judging by the number of pages I needed to cover the material in each of the three options treated in Conquering Chemistry, the Shipwrecks, Corrosion and Conservation option appears to be the shortest, and the material does to a significant extent reinforce and extend core material, though some of it can be difficult to understand for weaker students. In some ways the material in Industrial chemistry may be more straight forward, but apart from chemical equilibrium it does not flow as directly from core material. There appears to be much more material in Forensic chemistry and much of it is quite complex for students whose background in organic chemistry is just the core material. While a good option for students with a strong biological interest, it could be daunting for weaker students.

With the current heavy emphasis on factual recall in exam questions, the quantity of material to be assimilated is probably a key consideration in selecting an option. 

In writing Conquering Chemistry I thought that the other two options, Biochemistry of movement and Chemistry of art, contained too much material of too diverse a nature for the seven weeks available for the study of the option. Of course these options will particularly appeal to certain groups of students – those strongly interested in sport and dance and those with a strong artistic bent – and strong motivation can easily overcome barriers that may seem daunting to others.

Some teaching points

HSC exam questions in the Options section of the paper differ somewhat from questions in the core section in that . 

Because of this absence of calculations and problems from the options sections of the HSC exam paper, there will be no Further Exercises sections on the Options pages (though a few exercises will be interspersed through the text). 
   

All of the above are with inert electrodes – electrodes that do not undergo chemical change during the electrolysis. Common cases in which an electrode reacts are electrolysis of aqueous solutions of copper and silver salts using a copper or silver anode respectively (Cu ® Cu²⁺ and Ag ® Ag⁺).

Most electrolyses are not particularly sensitive to the concentration of the ions present, the exception being chloride where concentrated solutions produce chlorine gas  while dilute ones form hydroxide (page 350).

Correlation with electrode potentials
What happens during electrolysis of aqueous solutions can be roughly correlated with electrode potentials (pages 410-11). Basically, the higher the electrode potential the more readily the reduction half reaction occurs and the less readily the oxidation (reverse) half reaction occurs. Hence Ag⁺ and Cu²⁺ are more easily reduced than is H⁺, but H⁺ is reduced in preference to Al³⁺ or Na⁺ (E^os for Ag, Cu, H, Al, Na are +0.80, +0.34, 0.00, –1.66 and –2.71 V respectively). Iodide is more easily oxidised than bromide which in turn is more easily oxidised than water; however water is more easily oxidised than is fluoride (E^os for I, Br, Cl and F are +0.54, +1.09, +1.23, and +2.87 V respectively). Water and Cl are close together (1.36 and 1.23 V) so changing the concentration of chloride can change which is oxidised preferentially. (Remember E^os refer to reduction half reactions.) 

Stating things less precisely: at the cathode the species with the algebraically higher E^o gets reduced (H⁺ rather than Mg²⁺) while at the anode the species with the lower E^o gets oxidised (I^– rather than Cl^–)

Exercise

3. Crevice corrosion

Crevice corrosion is corrosion that develops in microscopic gaps between sheets of metal that are held tightly together by physical means, such as by rivets or by nuts and bolts (and washers). It can be more severe than ordinary surface corrosion. When overlapping sheets are immersed in salt water, Figure 1, there will be a thin film of conducting water between the sheets, so normal corrosion will occur. But soon all the oxygen in the gap (crevice) will be used up. Then Fe tends to oxidise to Fe²⁺ in the crevice while oxygen gets reduced to OH^– on the exposed surface. Some of the Fe²⁺ diffuses out of the crevice and meets with the OH^– and so rust builds up around the exposed parts of the joint. If there is chloride in the film of water, it can form insoluble hydroxy chloride with the Fe²⁺ and this makes the water in the crevice slightly acidic. This more acidic environment can accelerate the oxidation of Fe and so accelerate rusting. There will be a build up of rust around the exposed edge of the joint, and in the crevice there will be some normal rust formed (because oxygen does slowly diffuse into it), there will be some deposits of insoluble hydroxy chlorides (see CCHSC page 434), but more importantly there will be considerable pitting of the iron surfaces which can lead to a weakening of the joint.

Figure 1 Crevice corrosion and its effect on rivets

Crevice corrosion is of particular concern when sheets of metal are riveted together, because by eating away the stem of the rivet, it can weaken the rivet and cause it to break. The best way of overcoming crevice corrosion is to avoid such overlapping joints (that is to eliminate crevices). This is why welding (CCHSC page 401) is preferred to riveting.  

4. The Titanic

Interest in the Titanic and its dramatic sinking increased significantly when the  wreck was discovered in 4 km deep water in 1985 and when in the following decade many artefacts and as well part of the hull were recovered from the ocean floor. Locally further interest was aroused by the opening of an extensive display about this ship by the Australian Maritime Museum in the late 1990s. The production of a very popular Hollywood blockbuster movie in 1997 also helped.

Titanic sunk because it hit an iceberg that tore a huge hole in its hull just below the waterline. The reasons commonly given for the hull fracturing are (1) compared with present-day steels, its steel contained more carbon and small but significant amounts of sulfur and these made it more brittle and hence more likely to fracture rather than just bend and dent, and (2) the steel plates were riveted together (rather than being welded as happens today) and had suffered crevice corrosion which had weaken the rivets so that they broke when stressed by the collision; there had been a longer than usual time between launching the hull and completing its fit-out before the maiden voyage and this could have aggravated corrosion.

The steel and riveting in the Titanic were in no way inferior to those used in other ships of that era. It was just bad luck that such a big ship hit such a large iceberg, though perhaps it was travelling faster than was prudent in such conditions; the ship was trying to set a new cross-Atlantic time record.

Finding that the wreck of the Titanic was severely corroded when only limited corrosion had been expected led to the discovery of bacterial corrosion (CCHSC page 430). 

Concentration and solubility of gases in sea water
The solubility of a gas in water is the maximum concentration of that gas in water when the water is in contact with the gas. The solubility of a gas is proportional to the pressure of the gas above the solution and it increases as the temperature of the solution decreases (CCHSC pages 426-7). For oxygen gas
          S = kP 
where S is solubility, P is the pressure of oxygen gas above the solution and k is the proportionality constant. k = 2.0 x 10^–5 mol L^–1 kPa^–1 so for an oxygen pressure of 100 kPa the solubility is 2.0 x 10^–3 mol L^–1.
However the concentration of a gas in water can be significantly less than the solubility if the solution is not in continuous contact with the gas. This is the situation that applies in deep sea water. The concentration of oxygen is much less than the solubility.

Exercise

5. Metals for shipbuilding

Twice in recent years (2004 and 2007) the HSC exam paper has had extended-response questions based upon developments in technology influencing choice of metals for shipbuilding. Actually over the past one hundred years or so there have not been any dramatic changes in the metals used for shipbuilding – only incremental changes such as improved steels. In recent years there has been a trend towards aluminium and its alloys for smaller vessels such as ferries and catamarans and leisure craft, but large ocean-going cargo and passenger ships still have steel hulls. The account in CCHSC pages 401-4 should be adequate for all HSC purposes.
   

6. Metals in contact

The key point to remember here is that when two metals are in contact, the one with the algebraically lower (smaller) electrode potential will oxidise preferentially, so magnesium will oxidise instead of iron, iron will oxidise instead of tin.

The common situations where corrosion of metals in contact arises are
· copper water pipes connected to steel pipes
· tin cans (steel coated with tin)
· galvanising
· cathodic protection of ship hulls and underground pipelines (using sacrificial anodes).  

Note however that the two metals must be in physical contact (or at least connected by a conducting wire) for this preferential oxidation to occur. If instead the two metals are placed in the one electrolyte solution (as in the 2005 HSC exam paper), then each corrodes or reacts independently of the other.

In that 2005 exam question students were asked to predict what would happen if pieces of magnesium and iron not in physical contact were placed in the one electrolyte solution that contained both Mg²⁺ and Fe²⁺ ions. The Mg metal would react with the Fe²⁺ ions
Mg(s) + Fe²⁺(aq)  ®  Mg²⁺(aq) + Fe(s)
The magnesium would become covered with a blackish deposit of iron as the magnesium itself slowly dissolved. 
In addition the piece of iron would rust – develop a brownish coating of hydrated iron(III) oxide:
Fe(s)  ®  Fe²⁺(aq) + 2e^–
O₂(aq) + 2H₂O(l) + 4e^–  ®  4OH^–(aq) 
Fe²⁺(aq) + 2OH^–(aq)  ®  Fe(OH)₂(s)   ...  ®  Fe₂O₃.xH₂O(s)

An excellent example of corrosion of metals in contact was the Statue of Liberty in New York, USA (see below).

Exercise

7. Acidic and alkaline forms of equations

To convert an alkaline form to an acidic form add H⁺s to both sides and proceed similarly.

Electrolysis of water
Such conversions are often required in writing equations for so-called electrolysis of water experiments. 

For the electrolysis of dilute sulfuric acid (hydrolysis of water) we would use acidic forms:
At the cathode: 2H⁺ + 2e^– ® H₂
At the anode: 2H₂O ® O₂ + 4H⁺ + 4e^–
Overall (double the first and add to the second): 4H⁺ + 2H₂O ® O₂ + 4H⁺ + 2H₂
Cancelling out 4H⁺ gives: 2H₂O ® 2H₂ + O₂

For electrolysis of sodium hydroxide solution (also hydrolysis of water) we would use alkaline forms:
At the cathode: 2H₂O + 2e^– ® H₂ + 2OH^–
At the anode: 4OH^– ® O₂ + 2H₂O + 4e^–
Overall: 2H₂O ® 2H₂ + O₂ (double the 1st, add to 2nd, cancel out 2H₂O + 4OH^–)

For electrolysis of near neutral solutions of salts such as sodium nitrate, magnesium sulfate or potassium fluoride which are also electrolysis of water we could use either the acidic pair or the alkaline pair (but don't use one of each!). 

Oxidation by O₂ in corrosion contexts
In corrosion contexts oxidation is often by oxygen gas either dissolved in water or from the air. Again we can write the reduction half equation in either acidic or alkaline forms:
O₂ + 4H⁺ + 4e^– ® 2H₂O
O₂ + 2H₂O + 4e^– ® 4OH^–

8. Different E^os for different forms of similar half reactions

From a table of standard electrode potentials we find that the two forms of what is essentially the one process have different E^os: for example

The reason lies in the 'standard' part of the name. The standard electrode potential is the value when all solutes in the reaction (equation) are present at concentrations of 1.00 mol/L (and gases present at partial pressures of 1.00 atmosphere). So the value for equation (i) is for [H⁺] = 1.00 mol/L (and so with [OH^–]  = 1.0 x 10^–14 mol/L).  The E^o value for equation (ii) is for [OH^–] = 1.0 mol/L and so [H⁺] = 1.0 x 10^–14 mol/L).

We can calculate a (non-standard) electrode potential for both of these half reactions for [H⁺] = [OH^–] = 1.0 x 10^–7 mol/L (that is pH = 7.0; neutral solution). It is 0.82 V for both of them.

Getting the same value for both should be no surprise: after all both equations simply represent the reduction of oxygen gas in water. Or looked at from the other direction, they both represent the oxidation of water to oxygen gas.

Similarly for

the first value refers to [H⁺] = 1.00 mol/L while for the second [OH^–] = 1.0 mol/L. Again we can calculate a value for pH = 7.0

Again the same value should be no surprise because both half equations just represent the reduction of water to hydrogen gas.

(For the benefit of teachers: these calculations are using the Nernst equation which you would have met in first year university chemistry. Check it out in any standard first year university textbook.)

Remember, it does not matter whether we write the equations as above or with one electron only: H⁺ + e^– ® ½H₂ . The E^o value is the same (see CCHSC page 59).

9. Australian maritime archaeological projects

The syllabus requires students to gather information from secondary sources to compare conservation and restoration techniques applied in two Australian maritime archaeological projects.  The best known projects are:

Cannons from Cook's Endeavour (CCHSC page 438)
The six cannons that Cook jettisoned were located by scientists from the Philadelphia Institute (USA) in 1969. They recovered some of them; an Australian group recovered the rest in 1970 along with about 90 iron ingots (used as ballast in the ship). The cannons were restored by what was then called the Australian Defence Science Services of the (Commonwealth) Department of Supply. Of the cannons, one went to the Philadelphia Institute, one to the Greenwich Maritime Museum (in Great Britain), one to the National Museum of New Zealand; of the three that stayed in Australia, one is at Kurnell, one at the Cooktown Museum and the other at the National Maritime Museum in Sydney.

11. Websites

Unfortunately the addresses of all three websites listed on CCHSC page 441 have changed. They now are:
http://www.conservationsolution.com/projects/about-projects  (Click on Archaeological Elements and Antiquities or Objects and select items about Titanic)
http://www.museum.wa.gov.au/collections/maritime/march/Batavia/batavia.html
http://www.mq.qld.gov.au/exhibitions/pandoragallery.asp
  

1. The Statue of Liberty (metals in contact)

The Statue of Liberty is a 46 m high statue of a woman holding aloft a flaming torch. It stands on a small island near the entrance to New York harbour. Liberty was a gift from France to commemorate the France-US alliance during the American war of independence. It was built in France and transported in pieces to New York where it was assembled and mounted on a large pedestal. It was finally completed in 1886.

The statue was made with an internal steel support structure (skeleton) with a copper 'skin' that took on the characteristic bluish-green colour of air-oxidised copper – a mixture that includes copper hydroxide and carbonate called a patina. The patina is not only attractive in appearance but also it protects the underlying copper from further corrosion because it forms an impervious layer.

The people who made the statue knew about corrosion of metals in contact and tried to avoid this corrosion problem by using strips of asbestos coated with shellac to insulate the steel supporting bars from the copper skin. However over time (about 100 years), particularly with movement of the copper skin relative to the steel supports (due to copper's greater thermal expansion than steel), this insulation was worn away and severe corrosion set in. The steel supports rusted with the supporting tips of many of them breaking off so that the copper skin was able to flap about in the wind and tear and become damaged. 

Finally in 1984-86 the statue underwent a major restoration. The rusted steel bars were replaced with stainless steel ones (to resist corrosion) and as added protection teflon sleeves were used to insulate the copper skin from the stainless steel supports. These sleeves also allowed the copper skin to move more easily with expansion. Apart from repairing some minor damage, the copper skin with its lovely patina was preserved.

Today a greater threat to the statue is acid rain. It slowly dissolves the protective patina (H⁺ + Cu(OH)₂ and CuCO₃) and exposes the thin copper skin to further corrosion. 
   

2. Acid sulfate soils (and acid mine drainage)

On page 430 it is explained how certain bacteria can reduce sulfate (to sulfide) and so bring about corrosion of iron at ocean depths too great for there to be sufficient dissolved oxygen to corrode the iron. This may bring to mind an environmental problem that often occurs in Australia, namely acid sulfate soils (which is not a very informative name!).

Some soils when they are disturbed – dug up for farming or excavated for sub-divisions and buildings – produce acid which drains into waterways and so has a detrimental effect upon aquatic life. The problem here is oxidation of sulfide to sulfate – in some ways the opposite to the bacterial oxidation of shipwrecks.

The soils that produce acidic run-offs are ones that contain iron pyrites, FeS₂. This unexpected formula arises because the two S atoms are covalently bonded together to form the S₂^2– ion: ^–S–S^–. The peroxide anion O₂^2– in Na₂O₂ or BaO₂ is similar: compare with the structure of hydrogen peroxide H–O–O–H .

On exposure to air (oxygen) iron pyrites oxidises to Fe(II) sulfate and sulfuric acid:
2FeS₂(s) + 7O₂(g) + 2H₂O(l) ® 2Fe²⁺(aq) + 4SO₄^2–(aq) + 4H⁺(aq)
(the products could be written 2FeSO₄ + 2H₂SO₄)
This overall equation is made up of the two half equations:
FeS₂ + 8H₂O ® Fe²⁺ + 2SO₄^2– + 16H⁺ + 14e^–
O₂ + 4H⁺ + 4e^– ® 2H₂O

The Fe²⁺ is then oxidised to Fe³⁺ (at the low pH involved, generally by bacteria such as Thiobaccillus ferrooxidans) and precipitates out as a yellow-brown gelatinous Fe(OH)₃. This leads to further acid production:
4Fe²⁺(aq) + O₂(g) + 4H⁺(aq) ® 4Fe³⁺(aq) + 2H₂O(l)
4Fe³⁺(aq) + 12H₂O(l) ® 4Fe(OH)₃(s) + 12H⁺(aq)

Water draining from coal mines is often quite acidic. The cause is this same series of reactions: mining exposes iron pyrites to air and so oxidation occurs. This is called acid mine drainage and is a serious problem for most coal mines.
   

3. Anodisation of aluminium

On page 406 it was stated that aluminium was a passivating metal, meaning that it readily forms a protective oxide layer that prevents further oxidation (corrosion). This protective oxide layer can be made thicker and hence more effective in protecting the underlying metal by a process called anodising.

The aluminium object to be protected is made the anode in an electrolysis cell: the electrolyte is a solution of sulfuric acid and the cathode a piece of unreactive metal (for example, stainless steel or platinum). Under these conditions the process that occurs at the anode is
2Al + 3H₂O ® Al₂O₃ + 6H⁺ + 6e^–
This thickens the oxide layer and so make the aluminium more resistant to corrosion.

If certain dyes (coloured materials) are also present in the electrolyte, they can become incorporated in the oxide layer and so give it a definite colour. In this way aluminium objects can be given distinctive and durable colours. Gold and bronze coloured anodised aluminium objects, such as car and door keys, are fairly common.  Less fashionable now than they were 20 years ago, brightly coloured saucepan lids and cake containers can probably still be seen in your grandmother's kitchen.
   

Answers to selected HSC Exam questions

2007 Question 29(b)

Corrosion is a major problem for vessels that have to operate in a variety of aquatic environments

Analyse how the factors in aquatic environments have impacted on the choice of metals used in the construction of vessels over time.
                                                                                                                                      6 marks

(Transcribed from the 2007 Board of Studies HSC Chemistry examination paper)

2008 Question 30(d)

Several marine archaeological projects exist around Australia. Compare the conservation and restoration techniques used in TWO of these projects with reference to the chemistry applied.
                                                                                                                    7 marks

(Transcribed from the 2008 Board of Studies HSC Chemistry examination paper)

Answers to the 2005 and 2006 Shipwrecks Corrosion and Conservation option questions are given in the complete Answers to HSC exam papers which you can access from the CCHSC home page (Click on it to go there).
   

Answers to exercises