Iron Rust

9 Key excerpts on "Iron Rust"

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  eBook - PDF

  Steel Surfaces

  A Guide to Alloys, Finishes, Fabrication, and Maintenance in Architecture and Art

  • L. William Zahner(Author)
  • 2020(Publication Date)
  • Wiley

    (Publisher)

  WHAT IS RUST? Iron wants to return to its natural state of oxides and hydroxides, which are essentially different mineral forms. For art and architectural steel assemblies, corrosion is in the form of a general surface change that occurs when the steel object comes in contact with a moist environment. The reaction is defined as follows: 4Fe + 6H 2 O + 3O 2 → 4Fe (OH) 3 This is the common corrosion of iron. The formula describes how iron present on the surface, along with water and oxygen react to form Fe (OH) 3 , also known as iron (III) hydroxide or ferric hydroxide. The oxygen in the equation comes from the atmosphere and from dissolved oxygen in the moisture on the surface. Oxygen readily dissolves in water, particularly natural waters. Ferric hydroxide, Fe (OH) 3 , the product in the formula, is insoluble and comes out of solution and deposits on the steel or flakes off. The ferric hydroxide is a reddish-brown colored rust. Often forming with this oxide is FeO (OH), or ferric oxyhydroxide. It is soluble and forms on steel when standing water is on the surface. When this occurs, it is considered hydrated, FeO (OH) • nH 2 O. The color is yellow or orange. Because it is insoluble it will run and drip from the surface as shown in Figure 7.4. As the surface dries out, the reddish color we associate with rusted steel remains. This rust is ferric oxide, Fe 2 O 3 , also known as iron (III) oxide. It is the form of rust that develops over a surface during uniform corrosion. The formula for this is: 2Fe (OH) 3 → Fe 2 O 3 + 3H 2 O Ferric oxide is insoluble and more adherent to the steel surface. When oxygen is in low supply, the ferric oxide can appear black in color. This black stain is very stable and adherent. Figure 7.5 shows the black form of this oxide, formed from heavy chloride attack. 250 Chapter 7 Corrosion Characteristics FIGURE 7.4 Stain from corroding steel set over a copper roof.

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  Corrosion of Metallic Heritage Artefacts

  Investigation, Conservation and Prediction of Long Term Behaviour

  • P Dillmann, G Beranger, P Piccardo, H Matthiessen(Authors)
  • 2014(Publication Date)
  • Woodhead Publishing

    (Publisher)

  2 ] to acquire more relevant and robust prediction of corrosion behaviour of materials. In order to get reliable experimental values of damage rates associated with a well-known wet–dry cycle, old rusted historical artefacts and low-alloy steel samples have been introduced in a climatic chamber for ageing tests. It is essential to determine such values in order to validate the modelling calculations and assumptions. Then, the experimental and calculated values are compared.

  8.2 Atmospheric corrosion of iron

  8.2.1 Description

  The atmospheric corrosion of iron is an electrochemical process occurring in the presence of an electrolyte. The corrosion of iron can be summarized by the following chemical equation:

  4

  Fe + 3O

  2

  + 2H

  2

  O = 4FeOOH

  (8.1)

  (8.1)

  The oxidation of iron by dissolved oxygen in thin electrolyte film leads to the formation of a rust layer composed of various oxides and oxy-hydroxides. By taking into account the fact that corrosion will occur away from atmospheric precipitation, the formation of the electrolyte film will be due to condensation of water from the atmosphere. This condensation depends mainly on the temperature and the relative humidity (RH) of the atmosphere. It is generally considered that corrosion phenomena start when relative humidity is higher than 60% and become more important when RH reaches 80% [3 ]. The variations of temperature and RH (due, for example, to the changeover between night and day) will create a succession of wet (with water condensing on the container walls) and dry cycles. Figure 8.1 shows the measurements of the variations of iron corrosion and oxygen consumption rates as experimentally measured by Stratmann [4 ]. Three different stages can be observed in Fig. 8.1 ; the wetting stage, the wet stage and the drying stage. Each of these stages is characterized by different physico-chemical mechanisms.

  8.1 The wet–dry cycle. Variations of iron (dotted line) and oxygen (solid line) consumption rates as experimentally measured by Stratmann [4

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  Encyclopedia of Iron, Steel, and Their Alloys (Online Version)

  • Rafael Colás, George E. Totten, Rafael Colas(Authors)
  • 2016(Publication Date)
  • CRC Press

    (Publisher)

  The results of these activities as well as the findings for the corrosion of iron and steel obtained so far [1–8,11,25,26] are summarized as: [27] a. In wet environments containing chlorides, iron and steel corrode forming rusts with brownish colors. b. Occasionally, rust surfaces are stained yellow and it appears that “rust fluids” have leaked out, and this new yellow rust is called “flowing rust.” The “flowing rust” indicates that corrosion is not a static process, but that dynamic changes occur during the rusting. c. Rusts develop into a layer structure with repeating dense and coarse layers. d. The chloride ions are concentrated at layer boundaries, most notably at the rust/substrate interface. e. The application of wet pH-test paper to rust surfaces exposed by removing the surface layers clearly indi- cates the presence of acidic (pH <3) and alkaline (pH >10) regions. f. Microscopically, there are numerous cracks along the layer planes as well as in the direction vertical to the layer planes, and there are also pores in the rusts. g. On steels constantly exposed to water, rusts keep growing and finally detach as flakes or even slabs at the substrate/rust interface. h. Rusts formed on heavily corroded steels contain approximately 25% of Fe 3 O 4 , 10% of α-FeOOH, and 5% of β- and γ-FeOOH, respectively, with the balance amorphous components. The dynamic aspects of the corrosion shown here cannot be explained by the Evans model, and Tamura [27] has proposed a model as will be shown next based on a quan- titative evaluation of the chemical reactions pertaining to the corrosion, that is, the calculated compositions of the aquatic environment of the corroding region.

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  Corrosion Science and Technology

  • David E.J. Talbot, James D.R. Talbot(Authors)
  • 2018(Publication Date)
  • CRC Press

    (Publisher)

  Section 3.2.6.2 .

  7.2.3Rusting in Air

  Rusting of bare iron and steel surfaces is generally slower in outside air than in water but is much more variable, ranging from near zero to over 0.1 mm per year and it is less predictable. The factors that influence it differ from location to location and vary with time. They include climate, season and weather, atmospheric pollutants, temperature cycles and the initial conditions and orientations of the iron or steel surfaces. There are so many independent variables that empirical comparisons between the results of limited field tests can be unreliable and the effort to acquire statistically significant information is protracted because rates of rusting can take several years to settle to steady values. More progress can be made by applying scientific principles and logic to interpret accumulated local experience.

  There are three sources of water to sustain the electrochemical processes in air, atmospheric humidity, precipitation and wind- or wave-driven spray. 7.2.3.1Rusting Due to Atmospheric Humidity

  Rust can form on iron even when it is not wet to the touch. It is less in dry than in humid air and more in industrial or marine locations than in rural areas. These observations are rationalized on the basis of Evans* and his colleagues’ classic experiments to clarify the effects of relative humidity and pollutants on rusting. Iron samples do not rust appreciably in pure clean air even when nearly saturated with water vapour, but if as little as 0.01% by volume of sulfur dioxide is present as a pollutant, the metal rusts rapidly when the relative humidity exceeds a threshold value of about 70%. Dust contaminated with certain ionic salts has a similar effect in initiating rusting. Evans proposed a simple but elegant explanation as follows. An aqueous electrolyte is needed to produce rust and although water cannot condense spontaneously from unsaturated air, it can be induced to do so by hygroscopic salts. The role of pollutants is to provide particles of such salts on the metal surface. If the pollutant is sulfur dioxide, as it is in industrial locations, the salt is Fe(II) sulfate, FeSO4

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  Steel Corrosion and Degradation of its Mechanical Properties

  • Chun-Qing Li, Wei Yang(Authors)
  • 2021(Publication Date)
  • CRC Press

    (Publisher)

  The complexation is however much more in details of the chemical process rather than in principle. Since the most commonly exposed environment for steel structures is mainly atmosphere, and that for cast iron is soil (such as underground pipes), this book focusses more on the atmospheric corrosion for steel and corrosion in soil for cast iron and ductile iron. Furthermore, this book is intended for engineers for assessing corrosion-induced damages to steel and steel structures; fundamentals of corrosion science and chemical reactions will not be presented in detail here since they can be easily found in other books, such as Revie and Uhlig (2008) and Marcus (2011). 2.2.1 Electrochemical reactions In essence, steel corrosion is an electrochemical process that occurs when two or more points on the steel surface have a potential difference, and two chemical reactions, i.e., oxidation and reduction, take place simultaneously on the steel surface (Cramer and Covino 2003). The process of steel corrosion involves four basic parts: A metal, which acts as an electrical conductor for anodic and cathodic electronic transfer; An anode, where electrochemical oxidation takes place and electrons are liberated; A cathode, where electrochemical reduction occurs, and electrons transferred from the anode are consumed; A conductive medium, which is the aqueous medium or electrolyte or the local environment the steel is exposed to. The oxidation reaction is also known as anodic reaction, in which electrons are lost, resulting in a non-metallic state. Anode is the point where electrons on the steel surface are lost and corrosion takes place. The oxidation reaction can be expressed as follows (Revie and Uhlig 2008): Fe → Fe 2 + + 2 e − (2.1) The reduction reaction is known as cathodic reaction, which balances the anodic reaction. In this reaction, ions in the electrolyte accept electrons that are released from the electrically connected anode

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  Anticorrosive Nanomaterials

  Future Perspectives

  • Chandrabhan Verma, Chaudhery Mustansar Hussain, Eno Ebenso, Chandrabhan Verma, Chaudhery Mustansar Hussain, Eno Ebenso(Authors)
  • 2022(Publication Date)
  • Royal Society of Chemistry

    (Publisher)

  Chapter 8 Anti-corrosive Applications of Iron, Copper and Titanium Oxides

  RUBY ASLAMa , MOHAMMAD MOBINa , SAMAN ZEHRAa , CHANDRABHAN VERMA,b AND JEENAT ASLAM,c

  a Corrosion Research Laboratory, Department of Applied Chemistry, Faculty of, Engineering and Technology, Aligarh Muslim University, Aligarh 202002, India;

  b Interdisciplinary Research Centre for Advanced Materials, King Fahd University of Petroleum and Minerals, Dhahran 31261, Saudi Arabia;

  c Department of Chemistry, College of Science, Taibah University Yanbu-30799, Al-Madina, Saudi Arabia[email protected] ; [email protected]

  8.1 Introduction

  Corrosion is an irreversible deterioration of materials and an expensive phenomenon being one of the significant issues occurring on Earth. There are numerous factors for a material's corrosion, such as operational characteristics, environmental situation, non-equilibrium phases, the ineffectiveness of protective films, or categories of materials used as a protective film. Moreover, the mechanism of an electrochemical reaction involved in the corrosion procedure, diverse atmospheres with operational circumstances, namely temperature, pressure, or humidity, affect the rate of corrosion to a large extent depending upon the materials involved.

  1 ,2

  Inhibitors, protective coatings, cathodic/anodic protection, electroplating, painting, and lubricating techniques can all be used to combat metal degradation. Chemists have proposed alternate corrosion inhibitors to gradually phase out harmful hexavalent chromates’ usage in response to the requirement.

  Metal elements can combine to generate a wide range of oxide compounds. These can have a wide range of structural geometries and an electrical structure that can be metallic, semiconducting, or insulating in nature. Metal oxides are a fascinating group of materials that are now being studied extensively. They have unusual mechanical stress tolerance, excellent optical transparency, extraordinary carrier mobilities, and magnetic characteristics controlled by particle size and shape.3 Oxides are employed in constructing microelectronic circuits, sensors, piezoelectric devices, fuel cells, corrosion-resistant coatings, and catalysts in technological applications. Oxide systems are considered superb protective films due to their significant stability and ability to serve as diffusion barriers for various reacting ionic species.

  4 ,5

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  Corrosion and Surface Chemistry of Metals

  • Dieter Landolt(Author)
  • 2007(Publication Date)
  • EPFL Press

    (Publisher)

  On the other hand, rust layers slow down the rate of corrosion by providing a barrier (although imperfect) between the reactive metal surface and the atmosphere and thus reduces access of oxygen. When a non-corroded steel surface is exposed to the atmosphere, the rate of corrosion is highest. It then diminishes with exposure time as the rust layer builds up. Empirically, the amount of mass corroded as a function of time can be described by the equation: m k t n = (8.34) where m represents the corroded mass, k is a proportionality factor, t is the exposure time, and n is an exponential factor whose value depends on the climatic conditions and the type of steel. Generally, the value of n lies between 0.4 and 0.8. After an initial period of time, often lasting several years, the corrosion rate reaches a steady state value that only depends on the corrosivity of the environment and on the type of steel. The corroded mass then increases linearly with time, meaning that with further build-up of the rust layer its accelerating and inhibiting effects compensate each other. 8.2.3 Reaction mechanisms Reaction of a humid iron surface with oxygen When a polished iron or steel surface is exposed to a humid atmosphere, it becomes promptly covered with a brownish thin film of corrosion products. The reaction is all the more rapid if the surface has been contaminated with ionic species such as sulfates. The reaction mechanisms involved are schematically presented in Figure 8.19. An aqueous film, some micrometer thick, is present on the metal surface where it forms an electrolyte. Atmospheric oxygen diffuses across the liquid and reacts at the metal surface, while iron is oxidized to ferrous ions that dissolve in the electrolyte (Figure 8.19(a)). Fe Fe + 2 2+ → e (8.35)

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  The Iron Oxides

  Structure, Properties, Reactions, Occurrences and Uses

  • Rochelle M. Cornell, Udo Schwertmann(Authors)
  • 2006(Publication Date)
  • Wiley-VCH

    (Publisher)

  Iron objects which are exposed to the atmosphere or are partly immersed in water are often subjected to alternate cycles of wetting and drying (Pourbaix, 1974; Schwit- ter and Böhni, 1980). These cycles may be due to seasonal fluctuations in weather conditions or be the result of tidal movements or of splashing. They cause the corro- 499 18.5 The products of corrosion Fig. 18.4 Schematic represen- tation of a cross section of a corrosion tubercle: A) Surface crust, B) magnetic membrane, C) internal chamber wall, D) fluid interior (Bigham and Tuovinen, 1985, with permis- sion). sion potential of the system to change periodically and this in turn, induces cyclic changes in the composition of the rust (Evans and Taylor, 1972). Such cycles have been simulated in the laboratory and studied using electrochemical and magnetic techniques combined with Mössbauer spectroscopy (Stratmann and Hoffmann, 1989; Marco et al., 1989). Stratmann and Hoffmann (1989) found that a dry, cor- roded iron surface had a corrosion potential of ca. +0.2 V which upon wetting, gra- dually shifted to –0.4 V owing to retarded diffusion of oxygen from the air to the me- tal. The reactive component of the rust, lepidocrocite, was thus reduced via an inter- mediate (probably green rust) to magnetite with simultaneous corrosion of the me- tal. During the drying cycle, the oxygen diffused back through the pores in the oxide layer and the magnetite was oxidized to maghemite. If, however, reduction went only as far as the intermediate state, this phase was oxidized to lepidocrocite. During the wetting/drying cycles, the morphology of the oxide particles changed and this broke up the rust and prevented its adhesion to the underlying metal. At potentials lower than –0.5 V, any goethite in the rust was partly reduced, but usually, the poten- tial drop in the wetting/drying cycles was only sufficient to reduce the thermodyna- mically less stable Fe III oxides.

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  Corrosion Chemistry

  • Volkan Cicek, Bayan Al-Numan(Authors)
  • 2011(Publication Date)
  • Wiley-Scrivener

    (Publisher)

  Corrosion of steel is known by engineers as the result of electrochemical reaction when different potentials are developed by electrically connected metal parts in contact with a solution containing free ions. The so-called electrode potential is dependent on the particular metal and the nature of the solution. Comparative values of electrode potentials may be measured against a standard electrode-electrolyte system. For example, if hydrogen is considered of zero V electrode potential, then lead, iron, zinc and aluminum potentials are 0.13, 0.44, 0.75 and 1.66 V, respectively. When a metal is placed in an electrolyte and different electrode potentials are generated, current flows through the system, causing attack on the more anodic metal (i.e., the metal with the more negative electrode potential). The cathodic metal (i.e., the metal with the more positive electrode potential) remains unattacked. The reaction on the cathodic metal may be deposition of metal, liberation of hydrogen or formation of OH- hydroxyl ions.

  Corrosion may also occur without the presence of different electrode potentials if there is an applied electrical current due to the pickup of stray electrical currents from electrical conductors and equipment or the incidence of induced electrical currents.

  14.2 Steel Structures 14.2.1 Corrosive Environments Steel structures may be exposed to a variety of corrosive elements:

  1. Water, moisture and humidity

  2. Salt-laden air and rain

  3. Chemicals from the atmosphere, splashes or spills

  14.2.2 The Corrosion Process in Steel Structures A clear understanding of the corrosion process is essential to understand the steps to inhibit corrosion with protective coatings.

  Oxygen combines with iron, the major element in steel, to form rust. This electrochemical process returns the iron metal to the state that it existed in nature-iron oxide. The most common form of iron oxide or iron ore found in nature is hematite (Fe2 03 ), which is equivalent to what we call rust. Iron in iron ore is separated from the oxide to yield usable forms of iron, steel and various other alloys through rigorous electrochemical reduction processes.

  The process of combining iron and oxygen, called oxidation, is accompanied by the production of a measurable quantity of electrical current, which is why this is called an electrochemical reaction. For the reaction to proceed, an anode, a cathode and an electrolyte must be present. This is termed a corrosion cell. In a corrosion cell, the anode is the negative electrode where corrosion occurs (oxidation), the cathode is the positive electrode end and the electrolyte is the medium through which an electrical current flows.

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