Difference between revisions of "Rust" - New World Encyclopedia
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{{otheruses}} | {{otheruses}} | ||
| − | [[Image:Removing rust with sand.jpg|thumb|250px|A | + | [[Image:Removing rust with sand.jpg|thumb|250px|A blacksmith removes rust with sand prior to welding.]] |
[[Image:Rusted floorboards.jpg|thumb|250px|Rust damage in automobiles can create hidden dangers.]] | [[Image:Rusted floorboards.jpg|thumb|250px|Rust damage in automobiles can create hidden dangers.]] | ||
| − | [[image:Rust03102006.JPG|thumb|250px| | + | [[image:Rust03102006.JPG|thumb|250px|Rusting can completely eat away iron. Note the galvanization on the unrusted portions.]] |
| − | [[Image:Rust_screw.jpg|thumb|250px|right| | + | [[Image:Rust_screw.jpg|thumb|250px|right|A rusted bolt.]] |
{{steels}} | {{steels}} | ||
| − | '''Rust''' is the | + | '''Rust''' is the material formed when [[iron]] or its [[alloy]]s corrode in the presence of [[oxygen]] and [[water]]. It is a mixture of [[iron oxide]]*s and [[hydroxide]]*s. In today's world, iron is commonly used in the alloy known as [[steel]]. Thus, rusting usually refers to the corrosion of steel. |
| − | + | == Process of rusting == | |
| − | [[ | + | Iron is found naturally as [[iron oxide]]* in the [[ore]] [[hematite]]*, and metallic [[iron]] tends to return to a similar state when exposed to air and water. Energy is given off when rust forms. |
| − | When steel contacts water, an [[electrochemistry|electrochemical]] process | + | The process of rusting of iron can be summarized as three basic stages: |
| + | * The oxidation of iron metal to form iron(II) ions; | ||
| + | * the formation of [[hydroxide]] ions; and | ||
| + | * their reaction together, with the addition of oxygen, to create rust. | ||
| + | |||
| + | When steel contacts water, an [[electrochemistry|electrochemical]] process begins. On the surface of the metal, [[iron]] (Fe) atoms release [[electron]]s (e<sup>−</sup>) to form iron(II) ions (Fe<sup>2+</sup>). This process, called the oxidation of iron, may be represented as follows: | ||
:Fe → Fe<sup>2+</sup> + 2e<sup>−</sup> | :Fe → Fe<sup>2+</sup> + 2e<sup>−</sup> | ||
| − | The [[electron]]s released travel to the edges of the water droplet, where there is plenty of dissolved oxygen. They [[reduction (chemistry)|reduce]] the oxygen and water to hydroxide ions: | + | The [[electron]]s released travel to the edges of the water droplet, where there is plenty of dissolved oxygen. They [[reduction (chemistry)|reduce]] (combine with) the oxygen and water to form hydroxide (OH<sup>−</sup>) ions: |
:4e<sup>−</sup> + O<sub>2</sub> + 2H<sub>2</sub>O → 4OH<sup>−</sup> | :4e<sup>−</sup> + O<sub>2</sub> + 2H<sub>2</sub>O → 4OH<sup>−</sup> | ||
| − | The [[hydroxide]] ions react with the iron(II) ions | + | The [[hydroxide]]* ions react with the iron(II) ions to form iron(II) hydroxide (Fe(OH)<sub>2</sub>). The hydroxide in turn reacts with more dissolved oxygen to form hydrated iron(III) oxide (Fe<sub>2</sub>O<sub>3</sub>.''x''H<sub>2</sub>O). The general form of the reactions may be written as follows: |
:Fe<sup>2+</sup> + 2OH<sup>−</sup> → Fe(OH)<sub>2</sub> | :Fe<sup>2+</sup> + 2OH<sup>−</sup> → Fe(OH)<sub>2</sub> | ||
:4Fe(OH)<sub>2</sub> + O<sub>2</sub> → 2(Fe<sub>2</sub>O<sub>3</sub>.''x''H<sub>2</sub>O) + 2H<sub>2</sub>O | :4Fe(OH)<sub>2</sub> + O<sub>2</sub> → 2(Fe<sub>2</sub>O<sub>3</sub>.''x''H<sub>2</sub>O) + 2H<sub>2</sub>O | ||
| − | + | Hydrated iron oxide is permeable to air and water, allowing the underlying metal to continue to corrode even after a surface layer of rust has formed. Over time, the entire iron mass may convert to rust and disintegrate. | |
| + | |||
| + | Corrosion tends to progress faster in seawater than freshwater because seawater—which contains higher concentrations of ions from various salts (especially [[sodium chloride]]*)—conducts electricity more readily. Rusting is also accelerated in the presence of [[acids]], but inhibited by [[alkali]]s, through [[corrosion|passivation]]. Rust can often be removed through [[electrolysis]], however the base metal object cannot be restored by this method. | ||
== Rust prevention == | == Rust prevention == | ||
| − | |||
| − | [[Galvanization]] consists of coating metal with a thin layer of another such metal. | + | By contrast, the corrosion of [[aluminum]] involves the formation of [[aluminum oxide]]*, which coats the surface of the metal and protects it from further oxidation. This process is known as ''[[passivation]]*''. Likewise, stainless steel resists rusting by forming a [[passivation layer]]* of [[chromium(III) oxide]]*. Passivation also occurs with [[magnesium]], [[copper]], and [[zinc]]. |
| + | |||
| + | [[Galvanization]] consists of coating metal with a thin layer of another such metal. Typically, zinc is applied by either [[hot-dip galvanizing]] or [[electroplating]]. Zinc is traditionally used because it is cheap, easy to refine and adheres well to steel. Zinc also provides [[cathodic protection]] to metal that itself is unplated, but close enough that any water touching bare iron is also in contact with some zinc. The zinc layer acts as a galvanic anode rusting in preference. Galvanization often fails at seams, holes and joints, where the coating is pierced. | ||
More modern coatings add aluminium to the coating as ''zinc-alume'', aluminium will migrate to cover scratches and thus provide protection for longer. These rely on the aluminium and zinc oxides protecting the once-scratched surface rather than oxiding as a [[sacrificial anode]]. | More modern coatings add aluminium to the coating as ''zinc-alume'', aluminium will migrate to cover scratches and thus provide protection for longer. These rely on the aluminium and zinc oxides protecting the once-scratched surface rather than oxiding as a [[sacrificial anode]]. | ||
Revision as of 14:30, 10 November 2006
- For other uses, see Rust (disambiguation).
| Iron alloy phases |
|---|
|
Austenite (γ-iron; hard) |
| Types of steel |
|
Carbon steel (≤2.1% carbon; low alloy) |
| Other iron-based materials |
|
Cast iron (>2.1% carbon) |
Rust is the material formed when iron or its alloys corrode in the presence of oxygen and water. It is a mixture of iron oxides and hydroxides. In today's world, iron is commonly used in the alloy known as steel. Thus, rusting usually refers to the corrosion of steel.
Process of rusting
Iron is found naturally as iron oxide in the ore hematite, and metallic iron tends to return to a similar state when exposed to air and water. Energy is given off when rust forms.
The process of rusting of iron can be summarized as three basic stages:
- The oxidation of iron metal to form iron(II) ions;
- the formation of hydroxide ions; and
- their reaction together, with the addition of oxygen, to create rust.
When steel contacts water, an electrochemical process begins. On the surface of the metal, iron (Fe) atoms release electrons (e−) to form iron(II) ions (Fe2+). This process, called the oxidation of iron, may be represented as follows:
- Fe → Fe2+ + 2e−
The electrons released travel to the edges of the water droplet, where there is plenty of dissolved oxygen. They reduce (combine with) the oxygen and water to form hydroxide (OH−) ions:
- 4e− + O2 + 2H2O → 4OH−
The hydroxide ions react with the iron(II) ions to form iron(II) hydroxide (Fe(OH)2). The hydroxide in turn reacts with more dissolved oxygen to form hydrated iron(III) oxide (Fe2O3.xH2O). The general form of the reactions may be written as follows:
- Fe2+ + 2OH− → Fe(OH)2
- 4Fe(OH)2 + O2 → 2(Fe2O3.xH2O) + 2H2O
Hydrated iron oxide is permeable to air and water, allowing the underlying metal to continue to corrode even after a surface layer of rust has formed. Over time, the entire iron mass may convert to rust and disintegrate.
Corrosion tends to progress faster in seawater than freshwater because seawater—which contains higher concentrations of ions from various salts (especially sodium chloride)—conducts electricity more readily. Rusting is also accelerated in the presence of acids, but inhibited by alkalis, through passivation. Rust can often be removed through electrolysis, however the base metal object cannot be restored by this method.
Rust prevention
By contrast, the corrosion of aluminum involves the formation of aluminum oxide, which coats the surface of the metal and protects it from further oxidation. This process is known as passivation. Likewise, stainless steel resists rusting by forming a passivation layer of chromium(III) oxide. Passivation also occurs with magnesium, copper, and zinc.
Galvanization consists of coating metal with a thin layer of another such metal. Typically, zinc is applied by either hot-dip galvanizing or electroplating. Zinc is traditionally used because it is cheap, easy to refine and adheres well to steel. Zinc also provides cathodic protection to metal that itself is unplated, but close enough that any water touching bare iron is also in contact with some zinc. The zinc layer acts as a galvanic anode rusting in preference. Galvanization often fails at seams, holes and joints, where the coating is pierced. More modern coatings add aluminium to the coating as zinc-alume, aluminium will migrate to cover scratches and thus provide protection for longer. These rely on the aluminium and zinc oxides protecting the once-scratched surface rather than oxiding as a sacrificial anode.
There are several other methods available to control corrosion and prevent the formation of rust, colloquially termed rustproofing. Cathodic protection makes the iron a cathode in a battery formed whenever water contacts the iron and also a sacrificial anode made from something with a more negative electrode potential, commonly zinc or magnesium. The electrode itself doesn't react in water, but only to provide electrons to prevent the iron rusting.
Bluing is a technique that can provide limited resistance to rusting for small steel items, such as firearms; for it to be successful, water-displacing oil must always be rubbed onto the blued steel.
Corrosion control can be done using a coating to isolate the metal from the environment, such as paint. Large structures with enclosed box sections, such as ships and modern automobiles, often have a wax-based product (technically a slushing oil) injected into these sections. This may contain rust inhibiting chemicals as well as forming a barrier. Covering steel with concrete provides protection to steel by the high pH environment at the steel-concrete interface. However, if concrete covered steel does corrode, the rust formed can cause the concrete to spall and fall apart. This creates structural problems.
To prevent rust corrosion on automobiles, they should be kept cleaned and waxed. The underbody should be sprayed to make sure it is free of dirt and debris that could trap moisture. After a car is washed, it is best to let it sit in the sun for a few hours to let it air dry. In winter, or in salty conditions, cars should be washed more regularly as salt (sodium chloride) can accelerate the rusting process.
See also
ReferencesISBN links support NWE through referral fees
- Jones, Denny (1996). Principles and Prevention of Corrosion, 2nd edition, Upper Saddle River, New Jersey: Prentice Hall. ISBN 0-13-359993-0.
- Working Safely with Corrosive Chemicals
External links
- A Primer on Rust, thorough rust prevention and removal information
- NACE International - Professional society for corrosion engineers
- corrosion-doctors.org - Site dedicated to corrosion of all types
- [1] - European Federation of Corrosion
- [2] - Corrosion Properties of Stainless Steel
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