To Study The Effect of Metal Coupling On Rusting of Iron CBSE Class XII Project
What is Rust? Rust is composed of iron oxides. In colloquial usage, the term is applied to red oxides, formed by the reaction of iron and oxygen in the presence of water or air moisture. Other forms of rust exist, like the result of reactions between iron andchloride in an environment deprived of oxygen – rebar used in underwater concretepillars is an example – which generates green rust. Several forms of rust are distinguishable visually and by spectroscopy, and form under different circumstances. Rust consists of hydrated iron(III) oxides Fe2O3·nH2O and iron(III) oxide-hydroxide FeO(OH)·Fe(OH)3. Given sufficient time, oxygen, and water, any iron mass will eventually convert entirely to rust and disintegrate. Surface rust is flaky and friable, and provides no protection to the underlying iron, unlike the formation of patina on copper surfaces. Rusting is the common term for corrosion of iron and its alloys, such as steel. Many other metals undergo equivalent corrosion, but the resulting oxides are not commonly called rust.
Oxidation of Iron Metal When impure (cast) iron is in contact with water, oxygen, or other strong oxidants, or acids, it rusts. If salt is present, for example in seawater or salt spray, the iron tends to rust more quickly, as a result of electrochemical reactions. Iron metal is relatively unaffected by pure water or by dry oxygen. As with other metals, like aluminium, a tightly adhering oxide coating, a passivation layer, protects the bulk iron from further oxidation. The conversion of the passivating ferrous oxide layer to rust results from the combined action of two agents, usually oxygen and water. Other degrading solutions are sulfur dioxide in water and carbon dioxide in water. Under these corrosive conditions, iron hydroxide species are formed. Unlike ferrous oxides, the hydroxides d o not adhere to the bulk metal. As they form and flake off from the surface, fresh iron is exposed, and the corrosion process continues until either all of the iron is consumed or all of the oxygen, water, carbon dioxide, or sulfur dioxide in the system are removed or consumed.
Associated Reactions The rusting of iron is an electrochemical process that begins with the transfer of electrons from iron to oxygen. The iron is the reducing agent (gives up electrons) while the oxygen is the oxidising agent (gains electrons). The rate of corrosion is affected by water and accelerated by electrolytes, as illustrated by the effects of road salt on the corrosion of automobiles. The key reaction is the reduction of oxygen: O2 + 4 e − + 2H2O
4OH−
→
Because it forms hydroxide ions, this process is strongly affected by the presence of acid. Indeed, the corrosion of most metals by oxygen is accelerated at low pH. Providing the electrons for the above reaction is the oxidation of iron that may be described as follows: Fe
Fe2+ + 2 e −
→
The following redox reaction also occurs in the presence of water and is crucial to the formation of rust: 4Fe2+ + O2
4Fe3+ + 2O2−
→
In addition, the following multistep acid-base reactions affect the course of rust formation: Fe2+ + 2H2O
⇌ Fe(OH)2 +
2H+
Fe3+ + 3H2O
⇌ Fe(OH)3 +
3H+
As do the following dehydration equilibria: Fe(OH)2
⇌ FeO
Fe(OH)3
⇌ FeO(OH)
2FeO(OH)
+ H2O + H2O
⇌ Fe2O3 +
H2O
From the above equations, it is also seen that the corrosion products are dictated by the availability of water and oxygen. With limited dissolved oxygen, iron(II)-containing materials are favoured, including FeO and black lodestone or magnetite (Fe3O4). High oxygen concentrations favourferric materials with the nominal formulae Fe(OH) 3-xOx/2. The nature of rust changes with time, reflecting the slow rates of the reactions of solids. Furthermore, these complex processes are affected by the presence of other ions, such as Ca2+, both of which serve as an electrolyte, and thus accelerate rust formation, or combine with the hydroxides and oxides of iron to precipitate a variety of Ca-Fe-O-OH species. Onset of rusting can also be detected in laboratory with the use of Ferroxyl indicator solution. The solution detects both Fe2+ ions and hydroxyl ions. Formation of Fe 2+ ions and hydroxyl ions are indicated by blue and pink patches respectively.
Prevention Because of the widespread use and importance of iron and steel products, the prevention or slowing of rust is the basis of major economic activities in a number of speci alized technologies. A brief overview of methods is presented here; for detailed coverage, see the cross-referenced articles. Rust is permeable to air and water, therefore the interior metallic iron beneath a rust layer continues to corrode. Rust prevention thus requires coatings that preclude rust formation.
Rust-Resistant Alloys Stainless steel forms a passivation layer of chromium(III) oxide. Similar passivation behavior also readily occurs with magnesium, titanium, zinc, zinc oxides,aluminium, polyaniline, and other electroactive conductive polymers. Special "weathering steel" alloys such as Cor-Ten rust at a much slower rate than normal, because the rust adheres to the surface of the metal in a protective layer. Designs using this material must include measures that avoid worst-case exposures, since the material still continues to rust slowly even under near-ideal conditions.
Galvanization Galvanization consists of an application on the object to be protected of a layer of metallic zinc by either hot-dip galvanizing or electroplating. Zinc is traditionally used because it is cheap, adheres well to steel, and provides cathodic protection to the steel surface in case of damage of the zinc layer. In more corrosive environments (such as salt water), cadmium plating is preferred. Galvanization often fails at seams, holes, and joints where there are gaps in the coating. In these cases, the coating still provides some partial cathodic protection to iron, by acting as a galvanic anode and corroding itself instead of the underlying protected metal. The protective zinc layer is consumed by this action, and thus galvanization provides protection only for a limited period of time. More modern coatings add aluminium to the coating as zinc-alume; aluminium will migrate to cover scratches and thus provide protection for a longer period. These approaches rely on the aluminium and zinc oxides re-protecting a once-scratched surface, rather than oxidizing as a sacrificial anodeas in traditional galvanized coatings. In some cases, such as very aggressive environments or long design life, both zinc and a coating are applied to provide enhanced corrosion protection.
Cathodic Protection Cathodic protection is a technique used to inhibit corrosion on buried or immersed structures by supplying an electrical charge that suppresses the electro-chemical reaction. If correctly applied, corrosion can be stopped completely. In its simplest form, it is achieved by attaching a sacrificial anode, thereby making the iron or steel the cathode in the cell formed. The sacrificial anode must be made from something with a more negative electrode potential than the iron or steel, commonly zinc, aluminium, or magnesium. The sacrificial anode will eventually corrode away, ceasing its protective action unless it is replaced in a timely manner. Cathodic protection can also be provided by using a special-purpose electrical device to appropriately induce an electric charge.
Coatings and Painting Rust formation can be controlled with coatings, such as paint, lacquer, or varnish that isolate the iron from the environment. 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. Such treatments usually also contain rust inhibitors. Covering steel with concrete can provide some protection to steel because of the alkaline pH environment at the steel-concrete interface. However rusting of steel in concrete can still be a problem, as expanding rust can fracture or slowly "explode" concrete from within. As a closely related example, iron bars were used to reinforce stonework of the Parthenon in Athens, Greece, but caused extensive damage by rusting, swelling, and shattering the marble components of the building. When only temporary protection is needed for storage or transport, a thin layer of oil, grease, or a special mixture such as Cosmoline can be applied to an iron surface. Such treatments are extensively used when "mothballing" a steel ship, automobile, or other equipment for long-term storage.
Special anti-seize lubricant mixtures are available, and are applied to metallic threads and other precision machined surfaces to protect them from rust. These compounds usually contain grease mixed with copper, zinc, or aluminum powder, and other proprietary ingredients.
Economic Effects Rust is associated with degradation of iron-based tools and structures. As rust has a much higher volume than the originating mass of iron, its build-up can also cause failure by forcing apart adjacent parts — a phenomenon sometimes known as "rust packing". It was the cause of the collapse of the Mianus river bridge in 1983, when the bearings rusted internally and pushed one corner of the road slab off its support. Rust was also an important factor in the Silver Bridge disaster of 1967 in West Virginia, when a steel suspension bridge collapsed in less than a minute, killing 46 drivers and passengers on the bridge at the time.
The Kinzua Bridge in Pennsylvania was blown down by a tornado in 2003, largely because the central base bolts holding the structure to the ground had rusted away, leaving the bridge anchored by gravity alone. Reinforced concrete is also vulnerable to rust damage. Internal pressure caused by expanding corrosion of concrete-covered steel and iron can cause the concrete to spall, creating severe structural problems. It is one of the most common failure modes of reinforced concrete bridges and buildings.
Aim of Experiment To study the effect of metal coupling on the rate of reaction
Materials Required Apparatus- Beakers-15, Iron sheets of 2# size-6, Aluminium rods of 2# size-6, Brass rods of 2# size-6, Zinc sheets of 2# size-6, Measuring cylinders, Chemical Balance, Weight Box. Chemicals- HCl and NaOH
Theory of Experimentation Corrosion is a serious problem of some metals like iron, zinc, aluminium and alloys like brass which are commonly used in day to day life.
•
•
Apart from reducing the life of articles made up of these metals or alloys the chemical substances fo rmed out of corrosion have serious public health problems.
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Replacement of machines or their parts and many other articles in industrial and public dealing lead to huge expenditure.
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Hence, how to reduce or avoid corrosion of articles made up of metals or alloys has been a major subject of study in the field of chemistry and electro-chemistry.
Procedure (i)
Mix 9 ml. of conc. HCl with 241 ml. of water to form 250 ml. of solution.
(ii)
Take this solution in seven different beakers.
(iii)
Mark each beaker serially from 1 to 7.
(iv)
Take the weights of three iron sheets, three aluminium rods, three brass rods and three zinc sheets.
(v)
Now keep iron sheets, aluminium rods, zinc sheets and brass rod in separate beakers.
(vi)
Then take iron + brass, iron + aluminium, iron + zinc, aluminium + zinc and brass + zinc and keep them in different beakers.
(vii) Allow the reactions to occur for 24 hours. (viii)
Note the maximum and minimum temperatures.
(ix)
Now at the end of reaction take out the metals and keep them in sun for some time so that they get dried up quickly .
(x)
Take the weights of each specimen and note the difference.
(xi)
Similarly repeat 1,2,3,4,5,6,7 and 8 steps in a basic solution.
Observation Table
Results & Conclusions 1. The rate of corrosion observed in acidic medium or the mass consumed during the corrosion is in the decreasing order from brass to aluminum. Brass has the highest corrosion rate while aluminium has the least corrosion rate. Brass > Iron > Zinc > Aluminium 2. When coupling of these metals was done each couple showed some difference in their corrosion with respect to each metal kept alone. Iron + Aluminium couple has the highest rate of corrosion while iron +Zinc couple has the lowest rate of corrosion. Rate of corrosion of each couple is in the order of Iron + Aluminium > Brass + Zinc> Iron + Zinc 3. Rate of corrosion in basic medium is in the decreasing order from Brass to Aluminium. The order of rate of corrosion is as below: Brass > Zinc >Iron > Aluminium 4. When these metals were coupled the rate of corrosion was in the decreasing order from: Brass+ Aluminium > Brass + Zinc > Iron + Aluminium 5. Temperature and time of reaction were constant (Temperature was 21° C and time of reaction was 24 hours.)
Significance •
Corrosion is a serious problem of some metals like iron, zinc, aluminium and alloys like brass which are commonly used in day to day life. Apart from reducing the life of articles made up of these metals or alloys the chemical substances formed out of corrosion have serious public health problems. Replacement of machines or their parts and many other articles in industrial and public dealing lead to huge expenditure.
•
Hence, how to reduce or avoid corrosion of articles made up of metals or alloys has been a major subject of study in the field of chemistry and electro-chemistry.
•
The study of the rate of corrosion of different metals or alloys showed gradual decrease in their masses in acidic medium. The decrease is in the order of brass, iron, zinc, aluminium.
Bibliography & Vote of Thanks
