Corrosion Questions in English

(B) The process of rusting involves the oxidation of iron and the reduction of oxygen.
$1$. At the anode: $2Fe(s) \to 2Fe^{2+}(aq) + 4e^-$.
$2$. At the cathode: $O_2(g) + 4H^+(aq) + 4e^- \to 2H_2O(l)$.
$3$. The $Fe^{2+}$ ions are further oxidised by atmospheric oxygen to $Fe^{3+}$ ions,which form hydrated ferric oxide $(Fe_2O_3 \cdot xH_2O)$,known as rust.
$4$. Option $B$ is incorrect because metallic iron is oxidised,not reduced,and it forms $Fe^{2+}$ or $Fe^{3+}$ ions,not $Fe^{2-}$ ions.

(C) . Magnesium acts as a sacrificial anode because it is more reactive than iron. It undergoes oxidation instead of iron,thereby protecting the iron hull of the ship from corrosion caused by water and salt.

(A) Galvanization or galvanizing is the process of applying a protective zinc coating to steel or iron to prevent rusting.
The most common method is hot-dip galvanizing,in which the parts are submerged in a bath of molten hot zinc.

(B) Galvanization is the process of applying a protective zinc coating to iron or steel to prevent rusting.
$Zn$ is used for this purpose because it acts as a sacrificial anode,protecting the iron even if the coating is scratched,due to its more negative standard reduction potential $(E^o = -0.76 \ V)$ compared to iron $(E^o = -0.44 \ V)$.

(B) . Magnesium is more electropositive and reactive than iron. When attached to iron pipes,magnesium acts as a sacrificial anode and is more readily oxidized (converted into positive ions) than iron,thereby protecting the iron pipes from corrosion.

(B) The corrosion resistance of aluminium is due to the formation of a thin,non-porous layer of $Al_2O_3$ (aluminium oxide) on its surface when the metal is exposed to air. This layer prevents further oxidation of the underlying metal,a process known as protective passivation.

(C) Aluminium forms a protective oxide layer on its surface,which prevents further oxidation,whereas the oxide layer formed on iron is porous and does not provide protection.

(B) $Al$ (Aluminium) reacts with atmospheric oxygen to form a thin,stable,and non-porous layer of aluminium oxide $(Al_2O_3)$ on its surface.
This oxide layer acts as a protective barrier that prevents further oxidation and corrosion of the underlying metal.

(B) Corrosion is a process where a metal is oxidized by the loss of electrons to oxygen and facilitated by water and other electrolytes. This process involves the formation of an electrochemical cell on the surface of the metal,making it an electrochemical phenomenon.

(D) The rusting of iron is an electrochemical process.
In the presence of moisture and oxygen,the surface of iron acts as a small electrochemical cell.
The reaction at the anode is the oxidation of iron: $2Fe(s) \rightarrow 2Fe^{2+} + 4e^-$.
The reaction at the cathode involves the reduction of oxygen in the presence of $H^{+}$ ions: $O_2(g) + 4H^{+}(aq) + 4e^- \rightarrow 2H_2O(l)$.
Thus,the presence of $H^{+}$ ions catalyses the cathodic reaction,thereby accelerating the rusting process.

(A) At the anode,iron is oxidised: $2Fe(s) \to 2Fe^{2+}(aq) + 4e^-$.
At the cathode,dissolved oxygen is reduced in the presence of water: $O_2(g) + 2H_2O(l) + 4e^- \to 4OH^-(aq)$.
The overall reaction is $2Fe(s) + O_2(g) + 2H_2O(l) \to 2Fe(OH)_2(s)$.
$Fe(OH)_2$ is further oxidised by atmospheric oxygen to $Fe_2O_3 \cdot xH_2O$ (rust).

(A) In the electrochemical theory of aqueous corrosion,the metal surface acts as a small electrochemical cell.
At the anode,oxidation of iron occurs: $Fe_{(s)} \to Fe^{2+}_{(aq)} + 2e^-$.
At the cathode,the electrons released at the anode are consumed by the reduction of oxygen in the presence of $H^+$ ions:
$O_{2(g)} + 4H^+_{(aq)} + 4e^- \to 2H_2O_{(l)}$.

(D) . Iron.
Iron undergoes oxidation in the presence of moisture and oxygen to form hydrated iron$(III)$ oxide,commonly known as rust.
The chemical reaction is: $4Fe(s) + 3O_2(g) + 2xH_2O(l) \to 2Fe_2O_3 \cdot xH_2O(s)$.

(D) Rust is a hydrated iron$(III)$ oxide,which is chemically represented as $Fe_2O_3 \cdot xH_2O$.
It is formed by the oxidation of iron in the presence of moisture and oxygen.
Among the given options,$Fe_2O_3 + Fe(OH)_3$ represents the composition of rust as a mixture of hydrated oxides and hydroxides.

(A) Galvanization is the process of applying a protective $zinc$ coating to iron or steel to prevent rusting.
This is typically achieved through a hot-dip process where the iron is submerged in molten $zinc$.
$Zinc$ acts as a sacrificial anode,protecting the iron from oxidation even if the coating is scratched.

(A) Galvanisation is the process of applying a protective $Zn$ coating to steel or iron to prevent rusting.
It involves the deposition of a layer of zinc on the surface of iron or steel.
The zinc acts as a sacrificial anode,meaning it corrodes preferentially to the iron,thereby protecting the underlying metal from oxidation.
Therefore,the correct option is $A$.

(B) For the rusting of iron,both oxygen and water are necessary. $Fe$ does not rust in dry air. In moist air,both oxygen and water vapor are present,which leads to the formation of hydrated iron$(III)$ oxide,$Fe_2O_3 \cdot xH_2O$,commonly known as rust.

(A) The correct option is $(A)$.
This is an example of cathodic protection or sacrificial protection.
$Mg$ has a more negative standard reduction potential than $Fe$,meaning it is more reactive and acts as an anode.
$Mg$ loses electrons more readily than $Fe$ $(Mg \rightarrow Mg^{2+} + 2e^-)$,thereby preventing the oxidation of iron.

(C) Chrome plating provides a protective layer over the steel surface,which acts as a physical barrier against corrosive agents like oxygen and moisture.
This prevents the oxidation of iron,thereby protecting the steel from corrosion.
It is a form of surface coating rather than cathodic or anodic protection.

(D) The most convenient method to protect the bottom of a ship made of iron is by connecting it with a more reactive metal block,such as $Mg$ (magnesium).
This is known as cathodic protection or sacrificial protection.
Since $Mg$ is more reactive than $Fe$ (iron),it acts as the anode and undergoes oxidation,while the iron acts as the cathode and is protected from corrosion.

(C) The process of coating iron or steel with a thin layer of zinc to prevent corrosion is known as $Galvanization$.

(D) Galvanisation is the process of applying a protective $Zn$ coating to iron or steel to prevent rusting.
Since $Zn$ is more reactive than $Fe$,it acts as a sacrificial anode and protects the iron even if the coating is scratched.
Therefore,$Zn$ plating is the most durable method for protecting iron against corrosion.

(B) Magnesium acts as a sacrificial anode because it is more reactive than iron. It undergoes oxidation instead of iron,thereby protecting the iron hull of the ship from corrosion caused by water and salt.

(A) The overall corrosion reaction is: $Fe(s) + 2H^+ + \frac{1}{2}O_2 \to Fe^{2+} + H_2O(l)$.
The standard cell potential is calculated as: $E^o_{cell} = E^o_{cathode} - E^o_{anode}$.
$E^o_{cell} = 1.23 \ V - (-0.44 \ V) = 1.67 \ V$.
The Gibbs free energy change is given by: $\Delta G^o = -nFE^o_{cell}$.
Here,$n = 2$ (number of electrons transferred),$F = 96500 \ C/mol$.
$\Delta G^o = -2 \times 96500 \times 1.67 \ J/mol$.
$\Delta G^o = -322310 \ J/mol = -322.31 \ kJ/mol$.
Rounding to the nearest integer,the value is $-322 \ kJ$.

(D) The most effective method to protect the iron hull of a ship from corrosion is cathodic protection.
In this process,a more reactive metal like magnesium $(Mg)$ is attached to the iron hull.
Since $Mg$ is more electropositive than iron,it acts as a sacrificial anode and gets oxidized instead of iron,thereby preventing the iron from rusting.

(A) The rusting of iron is an electrochemical process.
In this process,iron $(Fe)$ acts as an anode and undergoes oxidation by losing electrons: $Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-$.
Therefore,the rusting of iron is primarily due to the loss of electrons by iron atoms.

(B) Galvanization is the process of applying a protective zinc coating to iron or steel to prevent rusting. The iron is coated with a layer of zinc $(Zn)$ metal to protect it from corrosion.

(B) Galvanization is the process of applying a protective $Zn$ coating to steel or iron to prevent rusting.
The most common method is hot-dip galvanizing,in which the parts are submerged in a bath of molten $Zn$.

(C) The rusting of iron is an electrochemical process.
At the anode,iron is oxidized to ferrous ions: $2Fe(s) \rightarrow 2Fe^{2+}(aq) + 4e^-$.
At the cathode,oxygen is reduced in the presence of $H^+$ ions: $O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)$.
The $Fe^{2+}$ ions further react with oxygen to form $Fe^{3+}$ ions,which form hydrated ferric oxide,$Fe_2O_3 \cdot xH_2O$.
During this process,$Fe(OH)_2$ is formed initially,which is then further oxidized to $Fe(OH)_3$ and finally dehydrated to form rust $(Fe_2O_3 \cdot xH_2O)$.

(C) In the corrosion of iron,the anodic reaction is $Fe \rightarrow Fe^{+2} + 2e^-$ and the cathodic reaction is $O_2 + 4H^+ + 4e^- \rightarrow 2H_2O$.
The standard cell potential is calculated as:
$E^{\circ}_{cell} = E^{\circ}_{cathode} - E^{\circ}_{anode}$
Given $E^{\circ}_{cathode} = 1.23 \, V$ and $E^{\circ}_{anode} = -0.44 \, V$.
$E^{\circ}_{cell} = 1.23 \, V - (-0.44 \, V)$
$E^{\circ}_{cell} = 1.23 + 0.44 = 1.67 \, V$

(A) Rusting of iron is an electrochemical process.
To prevent rusting,iron is connected to a more electropositive metal (like $Zn$ or $Mg$).
Since the more electropositive metal acts as an anode and undergoes oxidation,the iron object acts as a cathode and is protected from corrosion.
This method is known as cathodic protection.

(C) Corrosion of iron is an electrochemical process that requires both oxygen and moisture.
It can be prevented or minimized by creating an impermeable barrier (like paint or galvanization) on the surface of the iron to prevent contact with air and moisture.
Option $A$ is incorrect because iron should be in contact with a metal having a lower reduction potential (like $Zn$) for cathodic protection.
Option $B$ is incorrect as iron does not corrode significantly in oxygen-free water.
Option $D$ is incorrect because salt water acts as an electrolyte,increasing the conductivity and rate of the electrochemical reaction,not because the potential of iron changes.

(C) Rust is a hydrated iron$(III)$ oxide,which is formed by the oxidation of iron in the presence of oxygen and moisture.
The chemical formula for rust is represented as $Fe_2O_3 \cdot xH_2O$.

(A) Corrosive sublimate is the common name for mercury$(II)$ chloride,which has the chemical formula $HgCl_2$.
It is a white crystalline solid and is highly toxic.
In contrast,$Hg_2Cl_2$ is known as calomel.

(B) The process of coating iron or steel with a thin layer of zinc to prevent corrosion is known as galvanization. $Zn$ is more reactive than $Fe$ and acts as a sacrificial anode,protecting the iron even if the coating is scratched. Therefore,the correct answer is $Zn$.

(D) Corrosion is an electrochemical process in which metals are oxidized by the atmosphere,leading to their degradation. The factors affecting the rate of corrosion include:
$(i)$ The position of the metal in the electrochemical series; more reactive metals corrode faster.
$(ii)$ The presence of impurities in the metal,which creates local galvanic cells.
$(iii)$ The presence of electrolytes or gases like $CO_2$ in water,which increases the conductivity and acidity of the medium,accelerating the process.
Therefore,all the given factors affect corrosion.

(D) The rusting of iron is an electrochemical process that occurs in the presence of moisture and oxygen.
It is significantly accelerated or catalyzed by the presence of $H^{+}$ ions,which increase the rate of the cathodic reaction.

(A) In galvanization,iron is coated with zinc. Since the standard oxidation potential of zinc is higher than that of iron,zinc acts as a sacrificial anode and protects iron from corrosion.
Similarly,when a magnesium bar is connected to an iron pipe,magnesium (which has a higher oxidation potential than iron) acts as a sacrificial anode and undergoes oxidation,thereby protecting the iron pipe from corrosion.
This process is known as cathodic protection.

(A) During the corrosion of iron,$Fe$ is oxidized to $Fe_2O_3 \cdot xH_2O$ (hydrated ferric oxide).

(D) Copper lies below hydrogen in the electrochemical series,meaning its standard reduction potential is higher than that of hydrogen ($E^o_{Cu^{2+}/Cu} = +0.34 \ V$ vs $E^o_{H^+/H_2} = 0.00 \ V$).
Therefore,copper cannot displace $H_2$ from dilute acids,and the reaction is non-spontaneous.
Since the reaction is non-spontaneous,the free energy change $(\Delta G)$ for this process is positive.
Thus,the Assertion is incorrect because copper does not corrode in acidic solutions,but the Reason is correct.

(N/A) In the process of corrosion,due to the presence of air and moisture,oxidation takes place at a particular spot of an object made of iron. That spot behaves as the anode. The reaction at the anode is given by:
$Fe_{(s)} \longrightarrow Fe^{2+}_{(aq)} + 2e^-$
Electrons released at the anodic spot move through the metallic object and go to another spot of the object.
There,in the presence of $H^{+}$ ions,the electrons reduce oxygen. This spot behaves as the cathode. These $H^{+}$ ions come either from $H_2CO_3$,which are formed due to the dissolution of carbon dioxide from air into water,or from the dissolution of other acidic oxides from the atmosphere in water.
The reaction at the cathode is given by:
$O_{2(g)} + 4H^{+}_{(aq)} + 4e^- \longrightarrow 2H_2O_{(l)}$
The overall reaction is:
$2Fe_{(s)} + O_{2(g)} + 4H^{+}_{(aq)} \longrightarrow 2Fe^{2+}_{(aq)} + 2H_2O_{(l)}$
Also,ferrous ions are further oxidized by atmospheric oxygen to ferric ions. These ferric ions combine with moisture,present in the surroundings,to form hydrated ferric oxide $(Fe_2O_3 \cdot xH_2O)$,which is rust.
Hence,the rusting of iron is envisaged as the setting up of an electrochemical cell.

(N/A) Corrosion is the process where the surface of metallic objects is slowly coated with oxides or other salts of the metal. In this process,a metal is oxidized by the loss of electrons to oxygen,leading to the formation of metal oxides.
Characterization of metal corrosion:
$(i)$ The rusting of iron.
$(ii)$ The tarnishing of silver.
$(iii)$ The development of a green coating on copper and bronze.
Damage due to corrosion:
Corrosion causes enormous economic loss every year. It results in significant structural damage to buildings,bridges,ships,and all objects made of metals,especially those made of iron.

(N/A) Corrosion of iron is commonly known as rusting. It occurs in the presence of water and air.
Chemistry of corrosion: The chemistry of corrosion is complex,but it can be considered essentially as an electrochemical phenomenon.
Anode: At a particular spot on an iron object,oxidation takes place,and that spot behaves as an anode. The reaction is:
Oxidation: $2 Fe_{(s)} \rightarrow 2 Fe^{2+}_{(aq)} + 4 e^{-}$
$E^{\circ}_{(Fe^{2+}|Fe)} = -0.44 \ V$
Cathode: Electrons released at the anodic spot move through the metal to another spot,where they reduce oxygen in the presence of $H^{+}$ ions. These $H^{+}$ ions are formed due to the dissolution of carbon dioxide from the air into water,forming carbonic acid $(H_{2}CO_{3})$,or from other acidic oxides in the atmosphere. This spot behaves as a cathode with the reaction:
Reduction: $O_{2_{(g)}} + 4 H^{+}_{(aq)} + 4 e^{-} \rightarrow 2 H_{2}O_{(l)}$
$E^{\circ}_{(H^{+}|O_{2}|H_{2}O)} = 1.23 \ V$
Overall reaction: The overall reaction is the sum of the anodic and cathodic reactions:
$2 Fe_{(s)} + O_{2_{(g)}} + 4 H^{+}_{(aq)} \rightarrow 2 Fe^{2+}_{(aq)} + 2 H_{2}O_{(l)}$
The $Fe^{2+}$ ions are further oxidized by atmospheric oxygen to form hydrated ferric oxide,$Fe_{2}O_{3} \cdot xH_{2}O$,which is rust.

(N/A) Prevention of corrosion is of prime importance. The damage caused by corrosion is as follows:
$I$. Metal wastage leads to significant economic losses.
$II$. Corrosion of structural components like bridges can lead to structural failure and accidents.
$III$. Corrosion in machinery reduces efficiency,causes malfunctions,and eventually leads to complete failure.
Different methods to prevent metal corrosion:
$I$. Coating the surface with paint or protective films.
$II$. Applying chemical inhibitors like bisphenol.
$III$. Coating the surface with other metals ($Sn, Zn$,etc.) that are either inert or act as sacrificial anodes. An electrochemical method involves attaching a more reactive metal (like $Mg$ or $Zn$) which corrodes preferentially,protecting the base metal.
$IV$. Electroplating with inert metals to protect the surface. For example,$Ag$ plating on $Cu$ or $Au$ plating on $Ag$.

(A) The corrosion of metals involves the oxidation of the metal surface when exposed to atmospheric conditions.
$1$. For $Fe$ (Iron),it reacts with oxygen and moisture to form hydrated iron$(III)$ oxide,$Fe_2O_3 \cdot xH_2O$,which is a reddish-brown substance known as rust.
$2$. For $Cu$ (Copper),it reacts with atmospheric $CO_2$ and moisture to form a basic copper carbonate,$CuCO_3 \cdot Cu(OH)_2$,which appears as a green coating.
$3$. For $Ag$ (Silver),it reacts with atmospheric sulfur compounds (like $H_2S$) to form silver sulfide,$Ag_2S$,which appears as a black coating.

(A) In the electrochemical process of iron corrosion,the surface of the iron acts as the anode,where iron is oxidized: $Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-$.
The dissolved oxygen in the water layer on the surface acts as the cathode,where oxygen is reduced: $O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)$.

(N/A) The corrosion of iron is an electrochemical process occurring in the presence of moisture and oxygen.
$(a)$ At the anode,iron undergoes oxidation:
$2Fe(s) \rightarrow 2Fe^{2+}(aq) + 4e^-$
$(b)$ At the cathode,oxygen undergoes reduction in the presence of $H^+$ ions:
$O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)$
$(c)$ The overall reaction is obtained by adding the anodic and cathodic reactions:
$2Fe(s) + O_2(g) + 4H^+(aq) \rightarrow 2Fe^{2+}(aq) + 2H_2O(l)$

(A) During the corrosion of $Fe$ (iron),the metal surface acts as an electrochemical cell.
At the anodic site,$Fe$ undergoes oxidation: $Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-$.
The electrons released at the anodic site move through the metal to the cathodic site.
At the cathodic site,these electrons are consumed by the reduction of oxygen in the presence of $H^+$ ions: $O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)$.
Therefore,the electron movement is from the anodic site to the cathodic site.

(A) During the corrosion of iron $(Fe)$,the metal acts as an anode and undergoes oxidation: $2Fe(s) \rightarrow 2Fe^{2+}(aq) + 4e^-$.
The electrons released at the anode move through the metal and are consumed at the cathodic region.
In the presence of atmospheric oxygen and moisture (which provides $H^+$ ions),the oxygen is reduced: $O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)$.
Thus,$O_2$ is the species that gets reduced.

(N/A) The chemical formula of rust is $Fe_2O_3 \cdot xH_2O$,where $x$ represents a variable number of water molecules.