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Terms in this set (33)
a reaction that happens on the surface of a metal when the metal reacts with oxygen and water (either liquid or vapour) from the air or water around it. Corrosion reactions usually happen slowly.
- Corrosion can be defined as the deterioration of materials by chemical processes. In a sense, corrosion can be viewed as the spontaneous return of metals to their ore.
= iron corrosion
iron + oxygen + water --> hydrated iron oxide (rust)
red-brown layer which forms on the surface of iron or steel [only iron and steel* rust]; it is soft and crumbly and flakes off easily
*Steel = hard alloy of iron mixed with carbon and other elements
Rusting is a problem because . . .
rust crumbles easily, exposing the underlying iron/steel and leaving it to rust again. The cycle repeats, degrading the iron until the structure is destroyed.
- Rusting destroys about 20% of the world's iron and steel every year.
How to prevent rusting
1. Barrier protection: coating with a layer (barrier) that keeps out air and water
2. Metal plating, either with a non-reactive metal or metal that doesn't corrode easily (e.g. gold or chromium), or with a more reactive metal (e.g. zinc or magnesium) which corrodes instead of the iron/steel and forms a coating that stops water and oxygen from reaching the iron/steel underneath
3. Sacrificial protection
4. Alloying: mixing iron/steel with other metals to form an alloy which is resistant to corrosion
5. Cathodic protection
Barrier protection (examples)
* paint - for bridges and cars
* plastic coating - for garden furniture and wire netting
* greasing/oiling - for tools and moving parts of machinery
Metal plating (examples)
chrome (electroplated chromium) plating on automotive trim; gold-plated electrical connectors (to provide corrosion-resistant electrically conductive layer on the copper connectors)
Using a more reactive metal to protect a less reactive metal from corrosion - the more reactive metal corrodes instead. Galvanisation is a form of sacrificial protection. Another way is to attach bars of the more reactive metal (e.g. zinc or magnesium) to the iron/steel object (especially one submerged in water).
Sacrificial protection (examples)
* Metal plating (with more reactive metal)
* Attaching zinc/magnesium blocks to the hull of a ship
* Connecting zinc/magnesium blocks with conducting cables to an underwater iron pipe/pipeline.
Coating iron with zinc by dipping it into molten zinc. The zinc will react with oxygen in the air to form zinc oxide, which is a hard substance without cracks or pores and thus prevents oxygen and water from getting at the iron underneath. If the zinc coated surface is scratched a little, it will just generate more zinc oxide.
* Example: galvanized corrugated iron roofing.
Stainless steel (alloy example)
mixture of steel with chromium (minimum 10.5% chromium content), nickel and other elements. T he chromium combines with oxygen in the air to form a tightly-packed layer of chrome-containing oxide, which protects the steel underneath.
Electrochemical (galvanic) corrosion
When metals react with oxygen during corrosion, an electrochemical process takes place where the metal atoms lose electrons to form rust (the metal oxide). In particular, a current is generated internally by "galvanic action", and electrons flows through the electrically conductive metal from the part serving as the "anode" to the part serving as the "cathode".
At the anode, the iron dissolution releases electrons and the Fe2+ ions produced are oxidized by oxygen in the air to form rust (hydrated iron oxide). (The electrons released at the anode travel through the metal to the cathode, where oxygen is reduced from the air and the electrons combine with oxygen and water to form hydroxyls (OH-)
4 elements of electrochemical corrosion
1. Anode - The electrode where galvanic reaction(s) generate electrons and positive ions are formed. Corrosion occurs at the anode. (Note that this is the negative electrode since electrons are left behind by the oxidation action that occurs).
2. Cathode - The electrode that receives electrons and where negative ions are formed. The cathode is protected from corrosion. (Note that this is the positive electrode since chemical reactions are taking electrons away here).
3. Electrolyte-The conductor through which current is carried. In the case of corrosion, water serves as the electrolyte transporting ions to and from the metal.
4. Return Current Path - The metallic pathway connecting the anode to the cathode. It is often the underlying metal substrate.
Preventing electrochemical corrosion
Removing any one of the 4 elements will stop the current flow and galvanic corrosion will not occur.
Cathodic protection (CP) is a technique to control the corrosion of a metal surface by making it work as a cathode of an electrochemical cell, which receives rather than gives up electrons - thus preventing corrosion (which involves iron losing electrons to form rust). This is achieved either by placing the protected metal in contact with another more reactive metal (sacrificial anode method) or by connecting an external direct current power supply to the metal object (impressed current method).
Cathodic protection (uses)
to protect steel, water or fuel pipelines and storage tanks, steel pier piles, ships, offshore oil platforms and onshore oil well casings.
Sacrificial anode method
Involves putting the iron/steel object in contact with a more reactive metal such as zinc or magnesium (with more negative electrochemical potential). This produces a galvanic cell in which the reactive metal works as an anode and provides a flow of electrons to the iron/steel object, which then becomes the cathode. The cathode is protected and the anode progressively gets destroyed, and is hence called a sacrificial anode.
Impressed current method
Involves connecting an external direct current power supply from an inert anode to the metal being protected and maintaining a continual negative charge (flow of electrons) through the metal so the entire surface becomes a cathode (receiving, rather than becoming the source of, electrons) and does not corrode.
- This type of cathodic protection is commonly used to protect oil pipelines and other buried structures such as buried propane storage tanks.
Corrosion resistant metals
Many metals (e.g. aluminium, tin, magnesium, chromium) react quickly with oxygen from the air and thus corrode - however, unlike iron, they are resistant to corrosion as they are protected by a thin coating of metal oxide that forms on the surface of the metal and acts as an impenetrable barrier that prevents further destruction.
Metals that do not corrode
gold, platinum, silver (in clean air)
Copper reacts with carbon dioxide from the air to form a layer of copper carbonate (the green patina), which covers the surface of the copper and protects the copper underneath.
- Note: Copper also reacts slowly with oxygen from the air to form a thin brown-black layer of copper oxide, which protects the copper underneath from further corrosion.
Reactivity Series (Metals)
*Carbon & Hydrogen: non-metals added for comparison
[14 metals + C + H]
- Metals at the top of the reactivity series (e.g. Group 1 metals like potassium, sodium and lithium) corrode easily as they react quickly with oxygen in the air to form an oxide layer on the surface of the metal.
- Metals at the bottom of the reactivity series (e.g. gold, platinum) do not corrode.
Experiment to determine conditions of rusting: Method
Set up 6 labelled test tubes with iron wool/iron nails as follows (to compare the relative rusting that occurs):
1. Wool/nail in 5 cm3 water in open tube [air + water]
2. Wool/nail in sealed tube with anhydrous calcium chloride (drying agent) [air + no water]
3. Wool/nail in sealed tube with boiled water and layer of oil on top [no air + water]
4. Wool/nail in vegetable oil in sealed tube [no air + no water]
5. Wool/nail in 5 cm3 boiled water with nitrogen in sealed tube [nitrogen + no air + water]
6. Wool/nail in 5 cm3 boiled water with oxygen in sealed tube [oxygen + no air + water]
Experiment to determine conditions of rusting: Experiment Notes
*Boiling water forces out (removes) dissolved air
*Layer of oil prevents any further air from dissolving in the boiled water
*Flushing tube with oxygen/nitrogen flushes out air
Experiment to determine conditions of rusting: Observations
- No rusting in Tubes 4 and 5
- Slight rusting in Tubes 2 and 3 (possibly because some water/air still present)
- Rusting in Tubes 1 and 6, with 6 showing the most rust - indicating that the iron is reacting with oxygen from the air to form rust
Experiment to show air (oxygen) used during rusting: Method
1. Wet iron wool is put into the bottom of a test tube, which is inverted and placed in a beaker of water.
2. Measure the length* of the column of air in the test tube.
3. Leave for 1 week and then measure the new length of the column of air.
*Can also use to calculate volume of air in tube where V =πr2 x l
Experiment to show air (oxygen) used during rusting:
The length of the column of air is less after one week. This is because the rusting of wool removes oxygen from the air (reacts with the wool to form rust), drawing water up the tube.
- By observing the change in the column/volume of air in the tube, the percentage of air which has been removed can be found. This should be about 20% - approximately the percentage of oxygen in the air.
[If the percentage of air removed is significantly less, may leave the tube for another week to see if more air is used up]
Experiment to show sacrificial corrosion: Method
Set up 6 test tubes with iron nails immersed in liquified rust indicator gel, as follows:
1. Iron nail only (control)
2. Iron nail with copper
3. Iron nail with magnesium
4. Iron nail with zinc
5. Iron nail with lead
6. Iron nail with tin
*small piece of non-iron metal wrapped tightly around iron nail using pliers
Experiment to show sacrificial corrosion: Observations
Relative rusting of iron nail (in ascending order): nail with zinc (no rust); nail with magnesium (slight rust); nail with tin (some rust); nail with copper (very rusty on non-covered part); control nail (very rusty at the top); nail with lead (very rusty)
- The more reactive the metal wrapping it (e.g. magnesium and zinc), the less the iron nail rusts: the more reactive metal gives sacrificial protection and is preferentially oxidized way leaving the less reactive metal (the iron nail) intact.
- On the other hand, the rusting of iron is accelerated by less reactive metals (such as lead and tin): the non-covered parts of the nail rust preferentially to the surrounding pieces of lead/tin.
Why avoid buying bashed steel tin cans from a supermarket?
The tin plating provides a barrier between the steel can and water/air (tin oxide film forms which resists further corrosion), but once the tin layer is scratched, the steel underneath is exposed and will rust preferentially.
Why do iron/steel objects not rust much in dry desert regions?
Because water, as well as oxygen (from the air), are needed for iron/steel to rust.
Why do iron/steel objects near the coast rust faster?
Because sea spray of salt water accelerates the rusting chemistry.
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