Corrosion Complete Notes PDF

# Corrosion Corrosion is the gradual decaying or eating away of a metal by electrochemical reaction. - It needs an environment known as Corrosion. - In the case of Iron, corrosion is known as rusting of iron. It is a reddish coloured hydrated layer having the formula $Fe_2O_3.3H_2O$ formed at the...

# Corrosion Corrosion is the gradual decaying or eating away of a metal by electrochemical reaction. - It needs an environment known as Corrosion. - In the case of Iron, corrosion is known as rusting of iron. It is a reddish coloured hydrated layer having the formula $Fe_2O_3.3H_2O$ formed at the surface of the Iron. ## Electrochemical Theory of Corrosion Corrosion occurs when a metal is in contact with a conducting solution or is completely immersed in it. This occurs due to the existence of separated anodic and cathodic areas. 1. **At Anodic Areas**, oxidation occurs, and metal is destroyed by either dissolving or forming metal oxide. - $M \longrightarrow M^{2+} + ne^-$ 2. **At Cathodic Areas**, reduction occurs and the dissolved constituent in the conducting medium accepts the electrons to form some ions. - $2H_2O + 2e^- \longrightarrow 2OH^- + H_2 \uparrow$ ## Conditions and Features of Corrosion - Formation of separate anodic and cathodic areas - Presence of conducting medium through which current flows between anodic and cathodic areas. - Corrosion of anodic areas only. - Formation of corrosion products somewhere between anodic and cathodic areas ## Mechanism of Electrochemical Theory of Corrosion Depending upon the nature of the electrolyte, corrosion mechanics are of two types: 1. **Oxygen Absorption Corrosion** - This occurs in the presence of a neutral electrolyte. - Example: Rusting of Iron. - Let's take rusting of Iron as an example. - This type of corrosion takes place in the presence of a neutral electrolyte (example- aqueous solution of NaCl) in the presence of oxygen. - The oxide of Iron formed covers the surface of Iron. -Small cracks on the surface of the oxide layer develop a small anodic area, and the rest part acts as a cathodic area. - At anode, oxidation occurs and metal dissolves with liberation of electrons: $Fe \longrightarrow Fe^{2+} + 2e^-$ - At cathode, this electron flows from anode to cathode through the metal and is accepted by dissolved oxygen. - $ \frac{1}{2}O_2 + H_2O + 2e^- \longrightarrow 2OH^-$ - $Fe^{2+}$ ion from anodic area and $OH^-$ from cathodic area diffused to form ferrous hydroxide ($Fe(OH)_2$). - $Fe^{2+} + 2OH^- \longrightarrow Fe(OH)_2$ (yellowish) - This further oxidizes in the presence of oxygen and it converts into $Fe_2O_3.3H_2O$ which is rust. $Fe_2O_3 .3H_2O$ is black in colour. 2. **Hydrogen Evolution Corrosion** - This occurs in an acidic environment, i.e., in the presence of an acidic electrolyte. Example: Rusting of iron. - At anode, oxidation occurs and Fe metal dissolves to form $Fe^{2+}$ ion with the liberation of electrons. - $Fe \longrightarrow Fe^{2+} + 2e^-$ - At cathode, these electrodes flow through the metal from anode to cathode where $H^+$ ions from acidic solution are eliminated as hydrogen gas. - $2H^+ + 2e^- \longrightarrow H_2(g) \uparrow$ - $Fe + 2H^+ \longrightarrow Fe^{2+} + H_2(g) \uparrow$ ## Control/Prevention of Corrosion The methods used for the protection of metals from corrosion are: - **Proper Designing:** - Design of the material should be such that if corrosion occurs, it is uniform and not localized. - Contact of dissimilar metal in the presence of corrosive solutions (electrolytes) should be avoided. - Anodic area should be as large as possible when two dissimilar metals are in contact. - Sharp corners are poor design as they cause accumulation of solids. - Suitable design should avoid the presence of cracks between adjacent paths of the structure. - Using pure metal and using metal alloys. - **Cathodic Protection:** - To force the metal to behave like a cathode so that corrosion does not occur, and there are two types of cathodic protection: - **Sacrificial Anodic Protection (Galvanic protection):** The metallic structure is connected by a wire to a more anodic metal so that all the corrosion is concentrated at this more active metal, the more active metal itself gets corroded slowly while the parent structure is protected. The active metal used is called Sacrificial Anode. Example: Mg, Zn, Al. - **Impressed current (Direct current) Cathodic Protection (Electrical Protection):** - In this method, current from an external source (DC) is applied in the apposite direction to nullify (neutralize) the corrosion current. This is done to convert corroding metal from anode to cathode. Commonly used anodic materials are graphite, carbon, etc. - **Application of Electrical Protection:** - It is employed in buried (underground) structures including underground pipelines, tanks, transmission lines, towers, and lead up-laid-up ship (marine corrosion). - **Anodic Protection: **- Passivity (passivation) is a process in which a metal develops resistance against the corrosion. This is done by depositing an oxide layer on the surface of the metal. It is applicable to the metal which can be passivated by the deposition of the oxide layer on the surface of the metal structure to be protected. - When Iron rod is dipped in concentrated Nitric acid, the Iron uniformly and rapidly corrodes to form a thin protective Iron hydroxide coating. - The coating prevents the Iron from further corrosion in Nitric acid. - It has been applied in the case of steel or stainless steel and also in case of Ni, Fe, Al, Cr, etc. - **Treatment of metal:** - Alloying the metal, polishing the surface of metal, electroplating, hot dipping, cementation. - **Cementation:** - The metal to be deposited forms an alloy with the surface of the metal on which the deposition is taking place. - **Treatment of medium:** - Maintain pH (acidic = neutral), by adding inhibitors. - **Corrosion Inhibitors:** - Corrosion inhibitors are chemical substances which when added to the electrolytic solution of metal reduces the rate of corrosion. - **Anodic inhibitors (inorganic):** - Examples: Phosphates, Chromates, Silicates, Borate, molybdate, etc. It separates the rate of corrosion acting at the anode and produces insoluble precipitation of anode. They react with the ions of anode and produces insoluble precipitate. The precipitate so formed is absorbed on the metal surface forming a protective layer. - **Cathodic inhibitors (organic):** - Examples: Substituted Urea, Thiourea, Amines, Metal soaps, Bases, etc. In acidic solutions, the main cathodic reactions are $2H^+ + 2e^- \longrightarrow H_2(g) \uparrow$, thus, the corrosion can be controlled by slowing down the diffusion of $H^+$ ions through the cathode. It can be done by using organic inhibitors that get absorbed over the cathodic metal surface and acts as a protective layer. ## Type of Corrosion Corrosion can be of the following types:- 1. **Galvanic Corrosion:** - This type of corrosion takes place when the different (dissimilar) metals are coupled and exposed to atmospheric conditions. The metal with high reduction potential has more tendency to reduce (cathode) and forms a cathode, while the metal having low reduction potential has a tendency to get oxidized and forms an anode. - Example: Zn-Cu Galvanic cell, Zn metal has low reduction potential ($Zn E°=-0.764$) and behaves as an anode where corrosion occurs and Copper behaves as cathode where reduction takes place. 2. **Concentration Cell Corrosion:** - When a metal rod is half dipped in electrolytic solution, the dipped area of the metal is less oxygenated (i.e., lack of oxygen) and behaves as an anode, while the rest portion as cathode (well oxygenated area). - Let's consider a Cu metal rod is half dipped in an electrolytic solution as shown in the figure-$ \frac{1}{2}O_2 + H_2O + 2e^- \longrightarrow 2OH^-$, - **Reaction at cathode:** $ \frac{1}{2}O_2 + H_2O + 2e^- \longrightarrow 2OH^-$ - **Reaction at anode:** $Cu \longrightarrow Cu^{2+} + 2e^-$ - $Cu^{2+} + 2OH^- \longrightarrow Cu(OH)_2$ (yellow rust) 3. **Water Line Corrosion (Oxygen Cell Corrosion):** - Let's take stagnated water in steel tank. The concentration of oxygen above the water surface is greater than under the water line. Here, the metal above the water line, which is oxygen, behaves like cathode and metal surface below the water line is less oxygenated, hence, act as anode. Corrosion acts at anode that occurs at anode below the water line. - **Reaction at cathode:** - $\frac{1}{2}O_2 + H_2O + 2e^- \longrightarrow 2OH^-$ - **Reaction at anode:** - $Fe \longrightarrow Fe^{2+} + 2e^-$ - $Fe^{2+} + 2OH^- \longrightarrow Fe(OH)_2 $(rust) 4. **Pitting Corrosion:** - Let's consider the water drop on the surface of the metal. The metal surface, which is covered by the drop, lacks oxygen and acts an anode, while the metal surface, which is open and in oxygen, acts as cathode. Due to this small anodic and large cathodic area, set up differences of potential (as localized) spots results in the formation of pits and holes in the metal, that is rusting. - **At cathode:** - $\frac{1}{2}O_2 + H_2O + 2e^- \longrightarrow 2OH^-$ - **At anode:**- $Fe\longrightarrow Fe^{2+} + 2e^-$ - $Fe^{2+} + 2OH^- \longrightarrow Fe(OH)_2 $(rust) 5. **Crevice Corrosion:** - When bolts, nuts are attached in a metal, there is a crevice between different metallic objects in contact with the liquid. Crevice area has a lack of oxygen in comparison with the whole metal, thus, crevice area becomes anodic region which exposed areas acts as cathodic cathode. Therefore, corrosion occurs at anodic area (crevice area). This type of corrosion is accelerated by the deposition of dirt; this material restricts the supply of oxygen underneath the covered portion, which becomes anodic area where corrosion occurs. - **At anode:** - $Fe\longrightarrow Fe^{2+} + 2e^-$ - **At cathode:** - $\frac{1}{2}O_2 + H_2O + 2e^- \longrightarrow 2OH^-$ - $Fe^{2+} + 2OH^- \longrightarrow Fe(OH)_2$ (rust) 6. **Intergranular Corrosion:** - This is shown by alloy. It occurs when molten metal is solidified. It starts with random distribution of nuclei in molten mass of each of the grow into grains. Due to this adjustment, adjacent grains do not match, and the mismatched area between the adjacent grains is known as a grain boundary. - **Example:** - During welding of stainless steel (alloy of Fe, Cr, C), the carbide gets precipitated at the grain boundaries, thus repeating the adjacent region of Chromium that becomes anodic area, and it suffers corrosion. This type of corrosion is microscopic attack at plain grain boundries until the grain is completely dissolved. Thus, due to intergranular corrosion, sudden failure of material occurs without any apparent indication of a severe attack. 7. **Stress Corrosion:** - This type of corrosion takes place due to the combined effect of mechanical stress and the corrosive environment on a material. - It is common to alloy, i.e., Zinc and Nickle Brass. - Another example is Caustic embrittlement of mild steel. - When exposed to an alkaline solution at high temperature and stress, it causes failure of material due to caustic embrittlement, which occurs due to: - $Na_2CO_3 + H_2O \longrightarrow 2NaOH + CO_2$ - Due to different concentrations of the electrolyte higher inside cracks and dilute outside, metal dissolved Iron to form Sodium ferroate. ## Treatment of metals - **Alloying the metal:** - In case of Iron, Copper (Cu), Ni, corrosion can be controlled by formation of alloys. - **The strains in the metal:** - It can be removed by heat treatment and annealing as those strains are the focal centers for the corrosion reactions. - **By polishing:** - The surface corrosion of the metal can be decreased due to the removal of oxygen concentration cells. - **Electroplating:** - The coating metal is deposited on the base metal by passing direct current through the electrolyte solution containing a salt of the coating metal. The coating metal is made of an anode and the base metal is the cathode. Example: In steel, tin or zinc plating is applied. Metals like - Au, Ag, Cr, Ni, Cu & Sn, may be electroplated. Noble metals are used for electroplating. - **Hot Dipping:** - The metal to be coated is dipped in the molten bath of the coating metal for sufficient time; then, it is removed along with adhering film. Example: Applying coating of low melting point metals or alloys, i.e., Zn, Pb, Sn. The process providing Zinc coating on the iron is called galvanizing and providing tin coating is called tinning. - **Cementation:** - The metal to be deposited, forms an alloy with the surface of the metal on which the deposition is taking place. The metal to be treated is packed in powdered coating metal in a drum, which can be rotated and is heated at a temp. below the melting point of the more fusible metal. This provides uniform, thin layer (film) on the metal surface.