Immersion Corrosion and Electrochemical Processes Practice Quiz (23 questions)

1. What process is best defined as the destruction of a metal through chemical or electrochemical interaction with its environment?

A. Corrosion caused by environmental chemical reactions

B. Rusting limited to iron-containing materials

C. Friction corrosion caused by repeated surface contact

D. Erosion caused by mechanical surface forces

Answer: (A)

Explanation

Corrosion results from chemical or electrochemical interaction between a metal and its surroundings. Erosion instead results from physical phenomena, while friction corrosion combines physical effects with chemical interaction.

2. Which distinction correctly separates erosion from friction corrosion?

A. Erosion is physical, whereas friction corrosion includes chemical interaction

B. Erosion occurs internally, whereas friction corrosion begins in solutions

C. Erosion is chemical, whereas friction corrosion involves physical wear alone

D. Erosion affects iron, whereas friction corrosion affects non-ferrous metals

Answer: (A)

Explanation

Erosion is the physical destruction of metal, while friction corrosion combines physical effects with chemical interaction. Treating friction corrosion as purely mechanical overlooks its chemical component.

3. Why is corrosion of iron or its alloys specifically called rusting?

A. It occurs through physical abrasion of the iron surface

B. It requires a dry gas that conducts electrical current

C. It produces hydrated iron oxides commonly called rust

D. It describes corrosion found exclusively in non-ferrous metals

Answer: (C)

Explanation

Rusting is corrosion of iron or its alloys that produces rust, composed mainly of hydrated iron oxides. Non-ferrous metals may corrode, but that corrosion is not called rusting.

4. Which environment most clearly supports electrochemical corrosion?

A. A metal exposed to a dry, non-conducting gas

B. A metal reacting with an electrically insulating liquid

C. A metal immersed in an aqueous electrolyte

D. A metal isolated from moisture and ionic substances

Answer: (C)

Explanation

Electrochemical corrosion occurs when a metal contacts an electrolyte-containing environment such as an aqueous solution, soil, or moist air. Dry gases and electrically non-conducting liquids are associated with chemical corrosion instead.

5. What mechanism produces electrochemical corrosion in a metal exposed to an electrolyte?

A. Local short-circuited cells with oxidation and reduction at different sites

B. A dry-gas reaction that proceeds without charge transfer

C. A single uniform reaction with no separation of oxidation and reduction

D. Physical impact that removes metal without an accompanying chemical process

Answer: (A)

Explanation

Electrochemical corrosion is produced by local short-circuited corrosion cells, with oxidation and reduction occurring at different sites. A dry-gas reaction without current flow describes chemical rather than electrochemical corrosion.

6. Which situation represents chemical corrosion?

A. Atmospheric moisture allowing charge transfer through an electrolyte

B. An aqueous electrolyte supporting local oxidation and reduction cells

C. Moist soil connecting separate sites through ionic conduction

D. A dry gas reacting with metal without generating electric current

Answer: (D)

Explanation

Chemical corrosion occurs when dry gases or electrically non-conducting liquids react with a metal without generating or conducting electric current. Moist soil, aqueous electrolytes, and atmospheric moisture can support electrochemical corrosion.

7. What components make up an electrode, or half-cell?

A. A metal surface separated from every ionic conducting solution

B. Two complete electrodes connected through an external circuit

C. An electrolyte that contains no electronic conducting material

D. An electronic conductor and an ionic conductor containing related ions

Answer: (D)

Explanation

An electrode consists of an electronic conductor, such as a metal, and an ionic conductor, such as an aqueous solution containing its ions. A system containing two electrodes is a galvanic cell rather than a single half-cell.

8. What defines a galvanic cell?

A. A dry metal surface reacting without ionic movement in solution

B. A reference device used to measure one electrode's potential

C. A system of two electrodes where charge transfer accompanies surface redox reactions

D. A single electrode containing a metal and its ionic solution

Answer: (C)

Explanation

A galvanic cell contains two electrodes, with ions transferring charge through solutions and oxidation and reduction occurring at electrode surfaces. One metal-solution interface is a half-cell, not a complete galvanic cell.

9. At which electrode are electrons released during a galvanic-cell reaction?

A. The electrolyte, where ions conduct charge

B. The reference electrode, where potential is fixed

C. The anode, where oxidation occurs

D. The cathode, where reduction occurs

Answer: (C)

Explanation

Oxidation occurs at the anode, releasing electrons, whereas reduction occurs at the cathode, accepting electrons. The electrolyte conducts ionic charge but is not the electrode where electron release is defined.

10. Which pair of half-reactions correctly represents metal oxidation at the anode and metal-ion reduction at the cathode?

A. Anode: $M^{z+}+ze^- \rightarrow M$; cathode: $M \rightarrow M^{z+}+ze^-$

B. Anode: $M^{z+} \rightarrow M+ze^-$; cathode: $M+ze^- \rightarrow M^{z+}$

C. Anode: $M \rightarrow M^{z+}+e^-$; cathode: $M^{z+}+e^- \rightarrow M$

D. Anode: $M \rightarrow M^{z+}+ze^-$; cathode: $M^{z+}+ze^- \rightarrow M$

Answer: (D)

Explanation

At the anode, a metal atom is oxidized to an ion while releasing $ze^-$, and at the cathode, the ion gains $ze^-$ to become metal. Reversing the reactions assigns oxidation and reduction to the wrong electrodes.

11. What does the electrochemical series arrange?

A. Metals in increasing order of their electrochemical potentials

B. Metals in decreasing order of their atomic masses

C. Ions in increasing order of their solution concentrations

D. Electrodes in decreasing order of their surface areas

Answer: (A)

Explanation

The electrochemical series ranks metals by increasing electrochemical potential. A lower potential indicates a more active metal, while a higher potential indicates a less active metal.

12. Which conditions define the normal hydrogen electrode?

A. Oxygen at 1013 hPa, hydroxide activity of 1 mol/dm3, and platinum coated with platinum black

B. Hydrogen at 1013 hPa, hydrogen-ion activity of 0.1 mol/dm3, and a copper electrode

C. Hydrogen at 1013 hPa, hydrogen-ion activity of 1 mol/dm3, and platinum coated with platinum black

D. Hydrogen at 1 hPa, hydrogen-ion activity of 1013 mol/dm3, and an uncoated platinum electrode

Answer: (C)

Explanation

The normal hydrogen electrode uses hydrogen at 1013 hPa, hydrogen-ion activity of 1 mol/dm3, and platinum coated with platinum black; its assigned potential is zero. A metal electrode instead has a potential measured relative to a reference electrode.

13. Which expression gives the electrode potential under nonstandard conditions?

A. $E=E^0+\frac{RT}{zF}\ln\frac{a_{\mathrm{ox}}}{a_{\mathrm{red}}}$

B. $E=E^0+\frac{zF}{RT}\ln\frac{a_{\mathrm{red}}}{a_{\mathrm{ox}}}$

C. $E=E^0+\frac{RT}{F}\ln\frac{z a_{\mathrm{ox}}}{a_{\mathrm{red}}}$

D. $E=E^0+\frac{zRT}{F}\ln\frac{a_{\mathrm{ox}}}{a_{\mathrm{red}}}$

Answer: (A)

Explanation

The Nernst equation relates electrode potential to standard potential, temperature, charge number, and the oxidized-to-reduced activity ratio as shown. The alternatives place the charge number or activity ratio incorrectly.

14. When can one metal displace another metal from its salt solution?

A. A lower-potential, more active metal displaces a higher-potential, less active metal

B. Two metals with identical potentials exchange ions spontaneously in solution

C. A metal displaces another metal when their atomic masses are sufficiently different

D. A higher-potential, less active metal displaces a lower-potential, more active metal

Answer: (A)

Explanation

A more active metal has a lower electrochemical potential and can displace a less active metal with a higher potential from its salt solution. The reverse potential relationship does not predict displacement.

15. What reaction occurs when zinc is placed in copper sulfate solution?

A. $Cu+ZnSO_4\rightarrow CuZn+SO_4$

B. $Zn+CuSO_4\rightarrow ZnSO_4+Cu$

C. $Cu+ZnSO_4\rightarrow CuSO_4+Zn$

D. $Zn+CuSO_4\rightarrow ZnCu+SO_4$

Answer: (B)

Explanation

Zinc is more active than copper, so it displaces copper from copper sulfate and forms zinc sulfate. The reverse displacement is not predicted because copper has the higher electrochemical potential.

16. Which metal reaction with an acid illustrates hydrogen displacement by a metal with a negative electrochemical potential?

A. $Cu+2HCl\rightarrow CuCl_2+H_2$

B. $Ca+2HCl\rightarrow CaCl_2+O_2$

C. $Mg+HCl\rightarrow MgCl+H_2O$

D. $Mg+2HCl\rightarrow MgCl_2+H_2$

Answer: (D)

Explanation

Magnesium has a negative electrochemical potential and reacts with hydrochloric acid by displacing hydrogen, producing magnesium chloride and hydrogen gas. The copper equation incorrectly assigns this behavior to copper, while the other equations have incorrect products or balancing.

17. How are anodic and cathodic sites arranged in a local corrosion cell?

A. More active sites are cathodes, less active sites are anodes, and the solution electrically connects them

B. Both types of sites act as anodes, with electrons supplied by dissolved oxygen

C. More active sites are anodes, less active sites are cathodes, and the metal electrically connects them

D. Both types of sites act as cathodes, with ions carrying electrons through the metal

Answer: (C)

Explanation

More active surface sites dissolve anodically, while less active sites support cathodic reduction; the sites are electrically short-circuited through the metal. Reversing the site assignments confuses the direction of local corrosion.

18. Which pair correctly represents the anodic and cathodic reactions during iron corrosion in an acidic solution?

A. Anode: $Fe\rightarrow Fe^{2+}+2e^-$; cathode: $2H^++2e^-\rightarrow H_2$

B. Anode: $Fe^{2+}+2e^-\rightarrow Fe$; cathode: $H_2\rightarrow 2H^++2e^-$

C. Anode: $Fe+2H^+\rightarrow Fe^{2+}+H_2$; cathode: $O_2+2H_2O+4e^-\rightarrow 4OH^-$

D. Anode: $2H^++2e^-\rightarrow H_2$; cathode: $Fe\rightarrow Fe^{2+}+2e^-$

Answer: (A)

Explanation

Iron is oxidized at the anode, releasing electrons, while hydrogen ions are reduced at the cathode to form hydrogen gas. The overall reaction combines these half-reactions as $Fe+2H^+\rightarrow Fe^{2+}+H_2$.

19. What cathodic reaction generally occurs in a neutral, oxygenated environment?

A. $\frac{1}{2}O_2+H_2O+2e^-\rightarrow 2OH^-$

B. $2OH^-\rightarrow H_2O+\frac{1}{2}O_2+2e^-$

C. $\frac{1}{2}O_2+2H^++2e^-\rightarrow H_2O$

D. $O_2+4H^++4e^-\rightarrow 2H_2O$

Answer: (A)

Explanation

In neutral oxygenated conditions, dissolved oxygen is reduced to hydroxide according to the stated cathodic reaction. The other equations describe acidic oxygen reduction or reverse oxidation processes rather than the neutral oxygen reaction.

20. What does the gravimetric method quantitatively determine during corrosion testing?

A. The solution’s change in electrical conductivity

B. The solution’s change in temperature during exposure

C. The specimen’s change in mass during exposure

D. The specimen’s surface roughness after polishing

Answer: (C)

Explanation

Gravimetry quantifies corrosion by measuring how the specimen’s mass changes during testing. Mass loss is associated with corrosion evaluation, whereas mass gain is more commonly applied to high-temperature oxidation studies.

21. Which corrosion condition is least reliably characterized by gravimetry?

A. General surface corrosion across a specimen

B. Overall material loss during immersion

C. Uniform thinning over an exposed surface

D. Localized pitting at isolated surface sites

Answer: (D)

Explanation

Gravimetry is reliable for general or surface corrosion because total mass change reflects broadly distributed attack. Localized pitting, like other concentrated forms of corrosion, can be poorly represented by an overall mass measurement.

22. In the laboratory corrosion setup, what does the steel specimen represent?

A. The tested material exposed to the solution

B. The container holding the corrosive solution

C. The device controlling solution concentration

D. The reference used to calibrate the apparatus

Answer: (A)

Explanation

The steel specimen is the material being tested, while the glass beaker holds the acid, base, or salt solution. Treating the specimen as the container confuses the roles of the two main components.

23. Which combination correctly describes the laboratory apparatus for a corrosion test?

A. Copper electrodes, ceramic vessels, and a solution of unknown concentration

B. Metal powders, graduated cylinders, and a solution heated to boiling

C. Steel specimens, glass beakers, and a solution of specified concentration

D. Polymer coupons, sealed tubes, and distilled water without additives

Answer: (C)

Explanation

The apparatus includes different steel specimens, 100 or 250 ml glass beakers, and an acid, base, or salt solution at a specified concentration. The other combinations introduce materials or conditions not identified as components of this setup.