IN THIS FILE:
Archaeological material from marine sites presents some of the most difficult problems that are confronted by the conservator. The techniques of preserving this material have been carefully studied, and various techniques have been developed. The following sections describe the most applicable and useful approaches and conservation techniques that can be used to process and document metals recovered from shipwrecks and other underwater sites.
During most of the history of metallurgy only a few metals have been used. The metals of antiquity (iron, tin, copper, lead, silver, and gold) are those which were recognized and intentionally used with consistent regularity to manufacture tools, weapons, ornaments, hardware, and other paraphernalia. Each of these metals was used individually and in combination with the others, or with zinc or tin, to form more serviceable alloys, such as bronze, brass, and pewter.
From the moment of manufacture, the various metals and their alloys, except for gold, react with their environment and begin a corrosion process that converts them to more stable compounds. Before competent conservation techniques can be applied to a metal artifact, it is essential that the conservator be aware of the corrosion products that result from exposure to different environments. The nature of the corrosion products determines the technique and procedures that can be used effectively.
The corrosion of metals can be discussed in terms of terrestrial environments with temperate, tropical, and desert subdivisions, as well as aquatic environments, with salt and fresh water subdivisions. A more simplified approach is to look at the presence of oxygen and moisture in the environment. In any environment, moisture is a critical variable, and since aquatic environments, especially sea water, are the topic of interest here, dry conditions (where metal corrosion is minimal) are not considered. In sea water, the above variables, along with temperature, pH, and the presence of aggressive anions like chloride, determine the rates and types of corrosion.
The corrosion of iron in sea water will be discussed here; iron conservation treatments are
presented in File 10A and File 10B. The corrosion and preliminary conservation of non-ferrous metals is discussed in File 11; the conservation of cupreous metals is discussed in File 12;
silver and silver alloys are discussed in File 13; lead, tin, and lead alloys in File 14; and gold and gold alloys in File 15.
FERROUS METAL CORROSION
Iron is usually the most common metal recovered from archaeological sites. Due to the variety of conditions and environments within which corrosion can occur and the number and complexity of the corrosion products, iron presents the conservator with the most difficult problems of all the metals of antiquity. Iron corrosion processes are applicable to other metals and make iron a useful introduction to all metallic corrosion. The following information on corrosion processes relies heavily on Potter (1956), Evans (1961),Pourbaix (1966), Hamilton (1976), and Pearson (1987a).
In electrochemical corrosion, a galvanic cell is created when two different metals, or different areas on the same metal, are coupled by means of an electrical or ion-conducting electrolyte. The result is an electrochemical reaction. In essence, electrochemical corrosion is reserved for those processes where a current flows between anodic and cathodic areas situated at different parts of a metallic surface or between two different metals of the same or different material. The electrochemical oxidation of iron results in the formation of ferrous ions as the initial product.
Iron recovered from land sites that has only been exposed to ground and air moisture corrodes essentially by an electrochemical process. The corrosion of iron in sea water proceeds in a similar yet greatly accelerated manner because water generally becomes more corrosive as its salt content increases. Iron corrodes five times faster in sea water than in soil, and 10 times faster in sea water than in air (Cornet 1970:439).
The corrosion of metals submerged in salt water is a complex process. In the case of metal
recovered from shipwrecks, the shipwreck itself has often been explained in terms of a large galvanic cell based
upon the electromotive series of metals (Peterson 1969:30, 1972:244). Stated in general terms, all metals are compared
in an electrochemical cell with a hydrogen electrode, which is given an arbitrary electrode value of 0. Metals
that have a potential more negative than hydrogen in a galvanic cell are said to have a negative electrode potential,
and metals having a potential more positive to hydrogen have a positive electrode potential. By measuring the electromotive
force (EMF) in volts required to balance a galvanic cell formed by a particular metal immersed in a solution of
its salts of normal cation activity and a hydrogen electrode, the metals are arranged according to their relative
chemical activity or electrode potential into an electromotive series of metals (see Table 9.1). The least active
metals are at the top of the series, the most active are at the bottom. The more negative the electrode potential,
the more active the metal, and the greater the tendency for the atoms to lose electrons and form positive ions
which go into solution. When the ions of a metal go into solution, the parent metal always becomes negatively charged,
regardless of its electrode potential sign. When two metals form an electrochemical cell, the metal with the more
negative reduction electrode potential in the electromotive series becomes the anode. It loses electrons and forms
positive ions, which then go into solution. The more positive or noble metal in the cell forms the cathode and
is given cathodic protection, while the anodic metal is preferentially corroded in any resulting electrochemical
|Noble End (Cathodic)|
|Gold (+1 aurous, +3 auric)||
Au+3 + 3e- ⇔ Au
Ag+ + e- ⇔ Ag
|Copper (+1 cuprous)||
Cu++ e- ⇔ Cu
|Copper (+2 cupric)||
Cu+2+ 2e- ⇔ Cu
2H++ 2e-⇔ H2
|Lead (+2 plumbous, +4 plumbic)||
Pb+2 +2e- ⇔ Pb
|Tin (+2 stannous, +4 stannic)||
Sn+2 + 2e- ⇔ Sn
|Iron (+2 ferrous)||
Fe+2 +2e- ⇔ Fe
|Iron (+3 ferric)||
Fe+3 + 3e- ⇔ Fe
Zn+2 + 2e- ⇔ Zn
|Base End (Anodic)|
As a blanket explanation for the corrosion of different metals in a shipwreck in salt water, the large galvanic cell metaphor and the electromotive position of the metals have been overused and are not completely understood. The large mass of different metals associated with a sunken ship in salt water may consist of thousands of independent galvanic cells, each formed between two metals that have different electrode potentials. In order to establish a galvanic cell, the metals must be in very close proximity or in contact with each other. This requirement necessarily limits the metals that can form a galvanic couple to a single encrustation. Even then, variables, such as conductivity of the electrolyte, aggressive ion concentration, and mass transport, may alter or interfere with expected theoretical or laboratory reactions.
Any metallic surface is almost certainly to contain inclusions of more noble metals; it is very rare that a metal is 100 percent 'pure.' For this reason, a metal need not be in contact with a more noble metal to corrode in sea or tap water. An oxide-scale layer on a metal surface will be cathodic to the metal, which will be anodic in the presence of an electrolyte. A metallic couple between the two can form a number of galvanic cells. Electrochemical cells may also form on a chemically homogeneous metal in areas of mechanical stress, such as a dent or a bend, and concentrate the corrosion along this stress line. In shipwreck sites, stress corrosion is a very important factor in the corrosion and deterioration of given iron artifacts. Iron fastenings that are bent during the wrecking and sinking, for example, will corrode preferentially at the point of the bend, leaving a void but with good metal remaining at each end.
Even if a metal is pure, without an oxide layer or area of stress, immersion in a solution, such as sea water, which contains traces of salts of nobler metals can cause the formation of local cells at the metal surface, which effectively corrode the object (Potter 1956:238; Leigh 1973:20). In addition, the effects of different oxygen concentration, temperature, and pH at a metal surface may cause corrosion.
For iron artifacts buried in the ground, pitting is a prominent feature of the corrosion process. This anaerobic environment tends to be chemically reducing and forms soluble ferrous ions, which often diffuse some distance away from the iron surface. When iron is buried in an aerobic soil or exposed on the surface to the air, the ferrous ions initially formed in the corrosion process oxidize to ferric ions, resulting in layers of ferric oxide scale on the metal surface. This ferric oxide scale tends to form layers that may crack and spall due to the differences in the thermal expansion coefficients between the ferrous and ferric corrosion products and the metal. Alternatively, the corrosion products may inhibit additional corrosion by forming a protective film. Air-oxidized artifacts occupy more volume than the original metal and usually have obvious layers of ferric oxide scale. If salts, such as sodium chloride, are present in the environment, a very conductive solution is formed, and electrochemical corrosion is accelerated.
The electrochemical corrosion of metals has been detailed in a number of sources. In the case of iron, it has been shown that in any electrochemical cell where iron establishes a metallic couple in salt water with a more noble metal, such as copper or silver, or even with another piece of iron or a different part of the same iron object, the anodic and cathodic reactions are the same (see Potter 1956:236-237; Evans 1963:28).
At the surface of the more noble metal (the cathode), the following reaction occurs:
2H2O + 2e → H2 + 2(OH)-
The hydroxides combine with the sodium ion in the solution to form sodium hydroxide as the cathodic product:
Na+ + OH- → NaOH
At the anode, the reaction is the production of ferrous ions:
Fe+ - 2e → F+2
which, in turn combine with chloride in the salt water to form ferrous chloride as the anodic product:
Fe+2 + 2CI- → FeCl2
On exposure to air, or solutions containing dissolved oxygen, the ferrous chloride oxidizes to ferric chloride and ferric oxide. Ferrous chloride and ferric chloride and are freely soluble and may yield ferrous hydroxide when they combine with the cathodic product sodium hydroxide:
FeCl2 + 2NaOH → Fe(OH)2 + 2NaCl
In solutions containing dissolved oxygen, a secondary reaction oxidizes the ferrous hydroxide to a ferric state. In the presence of hydroxyl ions in a neutral or slightly alkaline solution, this hydrated ferric hydroxide (any form of ferric oxide with internal water, i.e., common rust) is precipitated on or around the electrodes of the cell. The sequence of reaction at an iron anode in the presence of oxygen as stated by Potter (1956:236) is:
|ferrous ion||Fe - 2e → Fe+2|
|ferrous hydroxide||Fe+2+ 2OH- → Fe(OH)2|
|hydrated ferric hydroxide (red-brown rust)||4Fe(OH)2 + O2 → 2H2O + 2Fe2O3 · H2O|
The primary anodic reaction of electrochemical corrosion of iron is the production of ferrous ions. The secondary stage, the oxidation of the ferrous ion compounds to a ferric state, is modified in anaerobic environments. Intermediate oxidation products of ferrous hydroxide, such as hydrated magnetite and black magnetite, are formed (Potter 1956:236-237; Evans 1963:28-29, 75):
6Fe(OH)2 + O2 →
4H2O + 2Fe3O4 · H2O
(green hydrated magnetite)
Fe3O4 · H2O → H2O
Depending on the environment, the corrosion products can take on a variety states of division and hydration, as well as a variety of physical forms. It is common to find corroded iron from marine sites with an outer layer of hydrated ferric hydroxide (common rust), which has restricted the supply of oxygen to the ferrous hydroxide briefly formed at the surface of the metal. Laminated corrosion layers consisting of an inner layer of black magnetite, a thin layer of hydrated magnetite, and an outer layer of hydrated ferric hydroxide are formed:
Fe3O4/2Fe3O4 · H2O or 2Fe2O3 · H2O
It is easy to see how two different areas of the same metal object can become anodic and cathodic to form an electrolytic cell. Electrons flow from the anodic area to the cathodic area causing the metal to corrode by forming soluble positive ions at the anode. Millions of these cells over the surface of the metal result in massive oxidation, which continues until an equilibrium state is reached. The corrosion process is halted at the cells when they come into equilibrium but may continue at alternate anodic and cathodic positions on the object until the bulk of the metal is oxidized.
As metals corrode in salt water, there are localized changes in the pH, which upset the equilibrium
between the dissolved calcium carbonate and dissolved carbon dioxide in the sea water (Leigh 1973:205). This results
in insoluble precipitates of calcium carbonate and magnesium hydroxide. These precipitates intermix with sand,
marine life, and corrosion products (especially ferrous hydroxide, ferrous sulfide, and magnetite) to form a hard
dense layer of encrustation or concretion around the metal. The encrustation accumulates on the original metal
surface to form a perfect mold around the object; furthermore, it will actually separate two metal pieces that
were initially touching each other. Such encrustation effectively separates the metals from each other and destroys
the electrochemical cell by cutting off the current flow and/or oxygen supply. It is rare, in fact, to find any two metal objects recovered from a shipwreck that are in direct contact
with each other.
Despite the fact that the corrosion processes are impeded by the formation of encrustation, metal deterioration can continue due to the presence of sulfate-reducing bacteria. These bacteria play a large part in the corrosion of metals, especially iron in salt water. They also adversely affect metals in fresh water, as well as metals buried in the soil under anaerobic conditions (Evans 1963:224; Pearson 1972a:35; Leigh 1973:205). Sulfate-reducing bacterial activity accounts for most of the rapid corrosion of buried iron and steel pipelines in waterlogged clay soils in England (Farrer et al. 1953:80). As much as 60 percent of the corrosion of iron in salt water can be attributed to bacterial action (Pearson 1972a:35).
Sulfate-reducing bacteria, particularly the strains known as Sporovibrio desulphuricans (Pearson 1972a:35) and Desulphovibrio desulphuricans (Farrer et al. 1953:82) are commonly found in salt water, fresh water, and waterlogged soil, where decaying organic material consumes oxygen and creates localized anaerobic environments. Sea water has a large supply of sulfates, and under aerobic conditions, these bacteria use hydrogen to reduce the sulfates (SO4)-2to sulfides (S-2) as a metabolic by-product according to the reaction:
H2SO4 + 8H → H2S + 4H2O
In this process, the hydrogen that accumulates on the iron as a cathodic product polarizes the cathode in an oxygen-free environment. The polarization of the cathode ordinarily halts the electrochemical corrosion process. However, the utilization of hydrogen in the metabolism of the bacteria depolarizes the cathodic areas of the cell, allowing corrosion to continue unabated. In addition, the hydrogen sulfide formed as a metabolic by-product reacts not only with iron but with all of the metals of antiquity (except gold) and accelerates the corrosion process. The hydrogen sulfide reacts with the ferrous ion from the anodic areas to produce ferrous sulfide and ferrous hydroxide, two major corrosion compounds of iron associated with objects recovered from the sea (Leigh 1973:205). On iron, the corrosion process (Pearson 1972a:34-35) proceeds as follows:
Fe+2 + H2S → FeS + 2H+
3Fe+2 + 6OH 3Fe(OH)2
4Fe + H2SO4 + 2H2O → FeS + 3Fe(OH)2
The life cycle of sulfate-reducing bacteria stimulates both the cathodic and anodic reactions of the electrochemical corrosion process. In some cases, however, the precipitation of a continuous film or layer of iron sulfide may stifle rather than stimulate the anodic reaction (Evans 1963:225). Without the presence of sulfate-reducing bacteria, the corrosion of iron and other metals in anaerobic environments would be inhibited.
One observation that has become apparent is that the presence of wood in direct association
with most metals has an adverse effect on them. Apparently, this results from the fact that as wood decays it consumes
oxygen, thus creating an anaerobic environment that stimulates the establishment of sulfate-reducing bacteria.
The wood also provides nourishment for the bacteria. This corrosion reaction is most evident on iron, silver, and
lead in direct contact with wood.
There are comparatively few differences between the corrosion processes of iron and those of mild steel, wrought iron, and numerous low alloy steels (Evans 1963:93). Even cast iron oxidizes by the same processes, including the action of sulfate-reducing bacteria. In addition, when cast iron is submerged in salt water, it undergoes a corrosion process called graphitization (Patoharju 1964:316, 1973:3; Pearson 1972a:10). In this reaction the salt water conducts a current between the anodic pearlite and the cathodic graphite flakes in the iron to form a galvanic cell. The pearlite corrodes, leaving a porous framework of graphite filled with the iron corrosion products discussed earlier. This graphite framework can maintain the original form of the object with little outward change in appearance, but with a considerable loss of density and mechanical strength. The process can proceed until the bulk of the metallic iron has corroded within the graphite framework. Ultimately the graphite framework will become incapable of supporting the object by itself, resulting in the deformation of the object.
The corrosion processes of iron in aquatic environments are generally known, but the reactions are complex and subject to many unpredictable variables. The majority of the resulting corrosion products, however, are predictable to a considerable degree of accuracy. This knowledge, while not replacing analytical tests for validation on particular objects, is usually sufficient to determine what conservation alternatives are warranted for any given artifact if its history is known.
The most commonly encountered iron corrosion products are:
|FeCl2||ferrous chloride, anhydrous|
|FeCl2 · H2O||ferrous chloride, hydrated|
|Fe3O4 or FeO ·Fe2O3||ferro-ferrous oxide (magnetite)|
|2Fe3O4 · H2O||magnetite, hydrated|
|2Fe2O3 · 3H2O||ferric hydroxide (common rust)|
|FeCl3||ferric chloride, anhydrous|
|FeCl3 · x H2O||ferric chloride, hydrated|
On metal objects recovered from shipwrecks, the most prevalent iron corrosion products are ferrous sulfide, magnetite, ferrous hydroxide, and iron chlorides. Many iron objects will completely convert to ferrous sulfide, leaving only a loose slush within a natural mold of encrustation. Other iron objects will completely mineralize to a massive oxide, magnetite, but retain their structural integrity and surface detail; some will be completely degraded to a loose granular oxide. In each case, iron sulfides are present in varying degrees. Also, iron chlorides are always a component part of any of the iron corrosion products. The main difference between exposed and buried iron concretions is the prevalence of magnetite in buried encrustations and the prevalence of ferro-hydroxide and ferric oxide in encrustations exposed to open sea water (North and MacLeod 1987:78), especially in high-energy zones, such as reefs.
Once iron has been removed from a marine environment, the corrosion process will continue, and even accelerate, unless certain precautions are taken. It is essential that iron artifacts be properly stored in an inhibitive solution to prevent further corrosion. If the iron in an encrustation is exposed to the air or to an uninhibited solution, the ferrous compounds can oxidize to a ferric state, which will occupy a greater volume and will scale off the surface. This process can disfigure an object and eventually destroy it. Without exception, it is much better to conserve iron artifacts with ferrous corrosion products rather than one with ferric corrosion products. Every precaution should be taken to prevent the ferrous corrosion products from oxidizing to ferric products through proper storage and treatment.
The greatest damage to marine iron after recovery is caused by the iron chlorides. The formation of ferrous chloride has already been shown by the reaction: Fe+2 + 2Cl- FeCl2, which, in turn, oxidizes to ferric chloride and ferric oxide in the general reaction: 6FeCl2 + 3O2 2FeCl3 + 2Fe2O3. Both of these reactions are gross oversimplifications, but the reaction proposed by Eriksen and Thegel (1966:90) -- Fe + 2NaCl + 2H2O FeCl2 + 2NaOH + H2 -- is not thermodynamically feasible. Regardless of the exact equation, both the ferrous chloride and ferric chloride combine with water to form hydrates: FeCl2 x H2O and FeCl3 x H2O, where x is normally 2, 4, or 6. It is these hydrated chlorides that cause the trouble. Upon exposure to moisture and oxygen, they hydrolyze to form ferric oxide or ferric hydroxide and hydrochloric acid. The hydrochloric acid in turn oxidizes the remaining non-corroded metal to ferrous chloride and hydrogen, or ferric chloride and water. In a simplified form, some or all of the following reactions may continue until no metal remains:
Fe - 2e →Fe+2
Fe+2 + 2Cl- → FeCl2
4FeCl2+ 4H2O + O2 → 2Fe2O3 + 8HCl
4FeCl2 +7H2O + O2 → 2Fe2O3 · 3H2O + 8HCl
2FeCl3 + 3H2O → Fe2O3 + 6HCl
4FeCl3 + 9H2O → 2Fe2O3 · 3H2O + 12HCl
Fe0 + 2HCl → FeCl2 + H2
4Fe0 + 3O2 + 12HCl → 4FeCl3 +
Of the above corrosion products, it is possible by electrolytic reduction to reduce the ferrous compounds Fe(OH)2, FeCl2, FeS and the ferrous oxide portion of Fe3O4, which probably exists as FeO · Fe2O3. It is not possible to reduce the ferric compounds in an aqueous solution. This problem is discussed further under the Electrode Potential heading of the section on Iron Conservation.
The preceding discussions on metal corrosion are necessarily brief and primarily refer to those corrosion products most commonly found on metals recovered from salt water. An early paper on the major corrosion products of the different metals can be found in Gettens (1963,1964). The most current and detailed discussion regarding the corrosion of marine iron can be found in North and MacLeod (1987).
PRELIMINARY STEPS: DOCUMENTATION, STORAGE,
AND MECHANICAL CLEANING
The Conservation of Antiquities and Works of Art (Plenderleith and Werner 1971) is an early, general reference for the conservation of archaeological material from any environment. While it contains only a few direct references to the objects recovered by marine archaeologists, many basic conservation techniques are discussed in a general manner. It is regrettable that much of the extant literature on the conservation of iron and other metals from salt water has either been neglected, oversimplified, or is misleading as to alternative procedures, cost, time involved, and problems encountered. (See, for example, Peterson 1964, 1969, 1972; Townsend 1964, 1972; Eriksen and Thegel 1966; Marx 1971; Wilkes 1971.) Presently, the most comprehensive survey of the techniques of conserving iron from marine sites, as well as other material, is Conservation of Marine Archaeological Objects (Pearson 1987a).
Sea-recovered metals present the most difficult problems the conservator is likely to encounter, but all of the conservation procedures used on these metals can be applied to metals recovered from other environments. The absence of marine encrustation and excessive chloride contamination considerably reduces the length of time required to process and stabilize objects from non-marine environments.
Regardless of the conservation technique utilized, it is essential to understand that no treatment is sufficient unto itself. It is but a part of a series of conservation processes designed to assure that a lasting preservation will be achieved. The duty of the conservation laboratory is to take the metal specimen as received and deliver a stabilized object. This involves a number of sequential and alterative steps. The major steps in the conservation of metals recovered from marine environments are:
1. Preliminary Steps
a. Initial documentation
b. Storage prior to treatment
c. Mechanical cleaning
d. Preliminary evaluation
3. Final Steps
a. Rinse after treatment
e. Periodic inspection
When one is responsible for conserving material resulting from an archaeological excavation, the basic approach to conservation should be that once an encrustation or any artifact has been delivered to the laboratory for treatment the conservator must (1) preserve and stabilize the specimen as well as possible, (2) recover useful archaeological information, and (3) acquire data for conservation research. These are possible only if extensive records are maintained, including detailed description, radiographs, black-and-white photographs, color slides, and notes on the preservation procedures used. Since all photographic negatives and prints are kept as a permanent record, they should undergo archival processing and be stored in a cool, dry, dark cabinet for maximum protection. Digital images are also recommended. All records should be well organized in a well-designed and readily accessible data base.
Proper records will include all the pertinent archaeological
data, identifications, descriptions, and the complete conservation procedure for each artifact. Considerable archaeological
data exists in the form of associations and provenience of artifacts within each encrustation. This information
is recoverable only by in situ observation and recording by the conservator. In other words, the conservator is
in a unique position to supply valuable archaeological data necessary to reconstruct details of the past. The conservation
data record the treatment history of every specimen, thereby accumulating valuable research records on the evaluation
of particular conservation techniques. If any specimen needs re-treatment in the future, the card provides valuable
information on why the original treatment failed and how to reverse the process.
STORAGE OF IRON PRIOR TO TREATMENT
The focus of the following section is upon metals recovered from sea water but is equally applicable to even non-metal artifacts recovered from other environments. The major alternative storage solutions are discussed.
Artifacts from the sea are usually encrusted together; they may even form large masses weighing well in excess of a ton. A single encrustation may contain objects of a variety of materials, including metal, wood, bone, and fiber. In order to prevent further corrosion, disintegration, or collapse, these materials must be kept wet throughout the period between recovery and treatment. It is necessary, therefore, to select a solution in which all materials can be safely stored. Since iron artifacts are likely to be the most common objects found in encrustation, solutions which provide good protection for iron but do not adversely affect other metals and materials should be selected. During storage, the encrustation should be left intact. The encrustation forms an excellent protective coating, which retards corrosion, prevents the chemical conversion of corrosion products already present, protects the artifacts from additional deterioration, and preserves artifact associations until they can be properly documented. Once processing has begun, and the different materials are removed, individual artifacts can be placed in a more desirable storage solution pending further conservation.
Iron recovered from a marine environment should be stored in an inhibitive aqueous solution. An inhibitive solution is any solution containing a substance that diminishes or prevents the corrosion of metals. Alkaline inhibitive solutions or inhibitive solutions containing oxidizing agents are commonly used in conservation.
The most common alkaline inhibitors used in conservation are sodium hydroxide, sodium carbonate, and sodium sesquicarbonate. Solutions containing these alkalies will prevent the corrosion of iron in oxygenated water as long as they are in concentrations sufficient to maintain a pH which passivates the iron (i.e., makes it chemically inactive) through the formation of an oxide film on the metal. In general, iron can be passivated in a chloride-free solution with a pH above 8. (See Hamilton1976:21-25 for a more thorough discussion of storage environments.) In inhibitive solutions with a pH of less than 8, the presence of oxygen will increase the rate of deterioration; the corrosion will be localized and the attack will be even more intense than if no inhibitor had been used (Evans 1963:151). Passivation of iron is difficult or impossible at a pH below 8, relatively easy at a pH above 8, and very easy between pH 10 and 12 (Pourbaix 1966:312). Iron will corrode by hypoferrate formation in solutions above pH 13 that are free from oxidizing agents. Thus, if iron is stored in an alkaline solution with the pH maintained between 10 and 13, the iron will remain passified and will not corrode. Figure 10B.7
A 5 percent sodium carbonate (pH 11.5) or a 5 percent sodium sesquicarbonate (pH 9.7) storage solution is sufficient for most iron objects if chlorides are not abundant. At high chloride concentrations, prolonged storage in either of these two solutions is not advisable unless additional alkali is added or the solution is changed often. Because the pH of these solutions is borderline to the corrosion domain, they are not recommended for long-term storage of iron objects from a marine environment. They can be used only for short-term field use or other temporary laboratory storage.
If long-term storage is required, an inhibitive solution containing oxidizing agents can be used, but the difficulties of properly disposing oxidizing agents prevent their more widespread use. From an environmental viewpoint, it is safer to take the extra time required to monitor the storage of iron in alkaline solutions.
Various chromate compounds, such as potassium chromate, potassium dichromate, and sodium chromate make effective storage mixtures. They are more reliable than alkaline inhibitors as long as the concentration and pH are maintained at safe levels. Chromate solutions prevent corrosion by forming a very thin passivating film of ferric oxide and chromic oxide on the surface of the metal (Pearson 1972a:14). This oxidizing solution creates an environment where a much wide range of pHs and electrode potentials of the surface of the metal is in the passivation range of the metal (Pourbaix 1966:74). It must be emphasized, however, that the chromate solution must be alkaline. The natural alkalinity of chromate (i.e., pH 9.1 to 9.3) is an important factor in passivating iron. Dichromates (pH <7) are more acidic than chromates and will not passivate iron unless an alkali is added. The addition of alkali, e.g., NaOH, converts dichromates to chromates and establishes the natural pH of chromate.
Like the alkaline inhibitors, the protection offered by a chromate solution will be suppressed if the pH is too low. Should this occur, the chromate can stimulate an intense localized attack on the surface of the metal and create pits covered with membranous blisters of iron hydroxide. The iron hydroxide inhibits contact between the chromate and the iron surface and allows the anodic reaction Fe Fe+2 to take place (Evans 1963:141; Kranz 1969:20). It is necessary that the pH of the chromate solution be maintained in a range of 9.0 to 9.5; if the pH of the solution falls below 9, corrosion of the object will be worse than if no inhibitors were used. The concentration of chloride in chromate solutions is not as critical a factor as it is in alkaline inhibitors, as long as there is free hexavalent chrome (Cr+6) present in the solution (Worth Carlin 1975, personal communication). For this reason, chromate solutions are particularly suited for the storage of iron from chloride-contaminated environments.
The pH of the chromate solution must be checked regularly, as some of the chromate in the solution is reduced and the solution may become acidic. If this occurs, additional alkali must be added to convert the dichromates to chromate and re-establish the natural pH range of 9.0 to 9.5.
Chromate solutions have the serious disadvantage of being highly toxic if ingested and inhaled. Chromates are strong irritants, and some are highly flammable if they come in contact with organic material. Chromate solutions should not be discharged into city sewage lines or natural drainages, since they kill beneficial bacteria. Most cities have regulations concerning the disposal of solutions containing hexavalent chrome, and a conservation laboratory must, of course, comply with them. If proper disposal cannot be arranged, chromate solutions should not be used.
There are several methods of treating a chromate solution for disposal. One method is described in Pearson (1972a:62): the chromate solution is acidified with concentrated sulfuric acid to a pH of 4. Sodium metabisulfite is added until the solution turns bright green. This reduces the hexavalent chrome to trivalent chrome. The solution is then neutralized with a 40 percent sodium hydroxide solution to precipitate out chromium hydroxide. The chromium hydroxide is allowed to settle as a sludge in the bottom of the vat. The solution is drained into the sewer lines, and the chromium hydroxide, which is insoluble, is disposed in a chemical dump.
Excellent long-term storage results have been achieved with
a 0.1N solution of potassium dichromate (K2Cr2O7) with sodium hydroxide. Many wrought iron artifacts recovered from
16th-century Spanish ships have been stored in this solution for more than three years with no apparent corrosion.
The pH may have to be adjusted to and maintained at 9.0 to 9.5 by the addition of sodium hydroxide as the chrome
is reduced. Sodium chromate also provides good results and has the added benefit of being cheaper than potassium
9.1. Stability conditions of pure iron in a
Use of De-ionized
Most conservation literature recommends that all storage solutions be prepared with distilled or de-ionized water. The exception to this rule occurs when the material to be conserved contains more chloride than is present in the local water supply. Tap water should be used for all storage solutions and electrolytes until the chloride level of the solution is less than that of the tap water. Using rain water in solutions as an interim step between tap water and de-ionized water will considerably reduce conservation costs. Following this procedure will result in significant financial savings when a large number of chloride-contaminated materials are to be processed. Unaltered de-ionized or distilled water should never be used as a storage solution for metal artifacts. They are generally slightly acidic and hence highly corrosive. Furthermore, when de-ionized or distilled water is used for rinsing or as a bath for detecting chlorides, immersion of the object should be kept to a minimum and adequate precautions taken.
Upon delivery to the conservation laboratory, marine archaeological material is typically covered with a dense and often thick encrustation. Removing the artifacts from this tough mass is analogous to removing objects from inside concrete blocks. Since most of the objects are hidden from view, radiographs are indispensable for determining the context of an encrustation and for serving as a guide in extracting the artifacts it contains. In order to x-ray large encrustations, large industrial X-ray machines, such as a 260 KVP water-cooled Picker Industrial X ray, must be used. Small- to medium-sized encrustation, up to approximately 1 x 1 m, can be x-rayed intact, and most of the constituent artifacts identified. It is often desirable to make a tracing or overlay from the individual X-ray plates. On the overlay, all the discernible artifacts are traced. Any specimens not detectable on the radiographs can be drawn, and their correct provenience can be ocated as they are encountered in the encrustation. Useful notations, such as catalogue numbers, condition of specimens, etc., can easily be added to the overlay.
For removing the encrustation, mechanical cleaning is the only feasible alternative. To accomplish this, a variety of hammers and chisels are indispensable, especially on the very large pieces. By hammering and chiseling along cleavage lines, the encrustation can be detached from large objects with little or no damage to the artifacts. For the extraction of smaller specimens, however, pneumatic tools are more efficient and less destructive. Chicago Pneumatic Weld Flux chisels are particularly serviceable for removing large amounts of encrustation and for extracting large, less fragile articles. Smaller, more precisely controlled Chicago Pneumatic Air Scribes, with their more delicate scribes and various chisels, which can be custom-made in the laboratory, are ideal for removing the encrustation from small, fragile pieces and for getting into restricted places. The pneumatic air scribe is much more durable than any comparable electric scribes or vibrotools. Combined use of the two types of pneumatic tools, the chisel and the scribe, is often necessary. They can, for example, be most effective in freeing movable parts, such as loaded breech chambers, iron lifting rings, and swivels on cannons.
When in doubt as to how much to mechanically remove, it is much safer to leave a thin layer of encrustation on the surface of the metal where the original surface cannot be determined. When the object is placed in electrolysis, the evolution of hydrogen bubbles at the surface of the metal will loosen the encrustation and mechanically remove any left on the metal. While electrolysis can be used to remove encrusted artifacts, the whole process of conservation is considerable speeded up if the encrustation removal is aided by mechanically removing by hand as much of the encrustation as possible.
The encrusted bores of cannons present a special problem. Tube drills are ideal, but each caliber gun requires a differently sized tube drill. Since these are quite expensive, they may not be practical for many laboratories. A suggested procedure is to use a hammer or chisel to remove as much as possible of the encrustation from the muzzle of the bore. A sandblaster can be then used to cut through the encrustation. Periodically, a steel rod ground to a chisel-like point should be used to roughen the surface of the encrustation so that the sand will work more efficiently. This technique may sound brutal, but very little, if any, harm is done to the cannon. The sandblaster will actually cut a hole through the encrustation without touching the metal. That is, a layer of encrustation is usually left on the surface of the bore. If necessary, additional encrustation can be removed with iron rods. Once a hole is made through the cannon bore, the cannon can be set up in an electrolytic bath with a center auxiliary anode. Hydrogen evolution in the bore will remove the remaining encrustation. This works very well on wrought-iron breech-loading hooped-barrel guns because cleaning is facilitated by the cannon tube being open at both ends.
The use of acids is generally a slow, ineffective process. Acids, especially hydrochloric acid, attack the metal oxide as readily as the encrustation and so are too damaging to consider. Even if successful, chemical techniques pose problems for recording associations and measurements and for making casts from molds of completely oxidized artifacts.
It has been known for a long time that when encrusted iron
artifacts are placed in electrolysis, that the encrustation is loosened off the surface of the metal. This technique,
termed 'deganguing' by the French (Montlucon 1986, 1987) can be quite effective on singularly encrusted objects,
but can be destructive when used on encrustation containing possible molds of corroded artifacts that can be cast
in epoxy. Additionally, some information in the encrustations is in the form of associations that must be visually
documented by the conservators. 'Deganguing' is an effective tool, that has been in use, to varying degrees for
years, but it should be used judiciously when working with complex, heavily encrusted objects that may contain
numerous other artifacts.
PRELIMINARY ARTIFACT EVALUATION
After each artifact is removed from an encrustation, it must be rinsed, carefully examined, and its condition evaluated to determine the most appropriate conservation treatment. It is useful to classify the metal specimens into one of three categories analogous to those suggested by Western (1972:83). These are based upon weight/size ratio, close visual inspection, testing the surface with a magnet, probing the corrosion layers with a dental pick, and occasionally using X rays. The categories include:
Only after these decisions have been made should treatment begin.