Discussion Examples for Corrosion Experiments | Redox Reactions, Metal Ionization Tendency, and Corrosion Protection

What Is a Corrosion Experiment?

A corrosion experiment is an experiment that investigates how metals change under the influence of water, oxygen, acids, salts, other metals, and other environmental factors.
Changes such as discoloration of the metal surface, formation of rust, changes in mass, gas evolution, and changes in solution color are observed to discuss how easily metals are oxidized and under what conditions corrosion can be prevented.
Corrosion is an important phenomenon related to the durability and safety of metallic materials.

In many cases, corrosion occurs through the simultaneous progress of an oxidation reaction in which a metal loses electrons and a reduction reaction in which oxygen or hydrogen ions accept electrons.
In other words, corrosion is a redox reaction and may proceed through the formation of small battery-like reaction sites on the metal surface.
Therefore, corrosion becomes easier to understand when considered from an electrochemical perspective.

Main Items to Include in the Results

In the results of a corrosion experiment, organize the type of metal used, type of solution, pH, presence or absence of chloride ions, presence or absence of oxygen, immersion time, temperature, changes in the metal surface, presence or absence of rust or precipitate, mass changes, potential differences, comparison conditions, and other information.
Because corrosion is greatly affected by environmental conditions, it is important to record the conditions specifically.

Main Items to Include in the Results

• Type of metal sample
• Pretreatment of the metal surface
• Type of solution
• Solution concentration
• pH conditions
• Presence or absence of chloride ions
• Presence or absence of oxygen
• Temperature
• Immersion time
• Discoloration of the metal surface
• Formation of rust or precipitate
• Presence or absence of gas evolution
• Mass change
• Presence or absence of contact between dissimilar metals
• Presence or absence of corrosion-protection treatment
• Degree of corrosion progression
• Causes of error and points for improvement

Reference Experimental Values and Calculation Examples for Corrosion Experiments

Here, metals are immersed in electrolyte solutions, and mass changes, surface observations, potential differences, and corrosion-protection effects are compared to organize redox reactions, the ionization tendency of metals, and the ease with which corrosion proceeds.

Corrosion is an electrochemical phenomenon that includes the reaction in which a metal is oxidized and becomes metal ions.
The way corrosion progresses changes greatly depending on the type of metal, ions in the solution, oxygen, chloride ions, contact with dissimilar metals, and the presence or absence of corrosion-protection treatment.

Reference Experimental Conditions

Ionization Tendency of Metals and How to View Corrosion

In general, metals with greater ionization tendencies more readily release electrons and become positive ions and are therefore considered more easily oxidized.
However, actual corrosion is also affected by oxide films, passive films, solution conditions, oxygen supply, and other factors.

Mass Changes When Metals Are Immersed in NaCl Aqueous Solution

The following shows an example in which each metal was immersed in a 3.5 mass% NaCl aqueous solution for 24 hours and the masses before and after immersion were measured.
The larger the mass loss, the more the metal is considered to have dissolved and the more corrosion progressed.

Example Calculation of Mass-Loss Percentage

The mass-loss percentage expresses as a percentage how much mass decreased relative to the mass before immersion.

Mass-loss percentage (%) = Mass loss ÷ Mass before immersion × 100

For example, for iron, the mass before immersion was 6.482 g and the mass loss was 0.051 g, so it can be calculated as follows.

Mass-loss percentage = 0.051 ÷ 6.482 × 100 = 0.79%

In this reference example, the mass-loss percentage of iron was relatively large, indicating that corrosion readily progressed in the NaCl aqueous solution.

Example Calculation of Corrosion Rate

The corrosion rate per unit area and unit time can be calculated from the mass loss.
Here, the corrosion rate is simply expressed using the following equation.

Corrosion rate (mg/cm²·day) = Mass loss (mg) ÷ Specimen area (cm²) ÷ Immersion time (day)

The mass loss of iron is 0.051 g, or 51 mg.
If the specimen area is 10.0 cm² and the immersion time is 1 day, the corrosion rate is calculated as follows.

Corrosion rate = 51 mg ÷ 10.0 cm² ÷ 1 day = 5.1 mg/cm²·day

The larger this value, the more metal was lost over the same area and the same period.

Comparison of Corrosion Rates Based on Mass Loss

Comparison of Iron Corrosion Under Different Solution Conditions

Corrosion is greatly affected not only by the type of metal but also by the solution conditions.
Here, examples are shown for iron specimens immersed for 24 hours in different solutions.

The mass loss of iron was larger in the NaCl aqueous solution than in pure water.
In addition, corrosion progressed even more readily under acidic conditions, while corrosion was suppressed under conditions where a corrosion inhibitor was added.

Example Calculation of Inhibition Efficiency by a Corrosion Inhibitor

The effect of a corrosion inhibitor can be determined by comparing the corrosion rate without the inhibitor with the corrosion rate in the presence of the inhibitor.

Inhibition efficiency (%) = (Corrosion rate without inhibitor − Corrosion rate with inhibitor) ÷ Corrosion rate without inhibitor × 100

If the corrosion rate of iron in NaCl aqueous solution is 5.1 mg/cm²·day and the corrosion rate in the presence of the corrosion inhibitor is 1.4 mg/cm²·day, the inhibition efficiency is calculated as follows.

Inhibition efficiency = (5.1 − 1.4) ÷ 5.1 × 100 = 72.5%

Therefore, the addition of the corrosion inhibitor was considered to suppress the corrosion of iron by approximately 72.5%.

Comparison of Corrosion Caused by Contact Between Dissimilar Metals

When different metals are electrically connected and placed in an electrolyte solution, galvanic corrosion may occur because of the potential difference.
The metal with the greater ionization tendency and more active potential tends to become the anode and its corrosion is accelerated.

Example Observation of Displacement Reactions with Metal Ions

One method of confirming the ionization tendency of metals is to place a piece of metal in an aqueous solution containing ions of another metal.
A metal with a greater ionization tendency reduces the metal ions in the solution while itself being oxidized and dissolving.

Example of How to Write the Results

When each metal was immersed in a 3.5 mass% NaCl aqueous solution for 24 hours, the mass loss of iron was 0.051 g and the corrosion rate was 5.1 mg/cm²·day, which were the largest values among the metals measured in this experiment.
In contrast, the mass loss of copper was 0.005 g and the corrosion rate was 0.5 mg/cm²·day, indicating relatively little corrosion.

When the iron specimens were immersed in different solutions, the corrosion rate was 0.8 mg/cm²·day in pure water, 5.1 mg/cm²·day in 3.5 mass% NaCl aqueous solution, and 7.4 mg/cm²·day in an acidic aqueous solution.
These results confirmed that chloride ions and acidic conditions promote corrosion of iron.

In addition, when a corrosion inhibitor was added to the NaCl aqueous solution, the corrosion rate of iron decreased to 1.4 mg/cm²·day.
The inhibition efficiency compared with the condition without the corrosion inhibitor was 72.5%, so the addition of the corrosion inhibitor was considered to have greatly suppressed corrosion.

Points for Connecting the Results to the Discussion

In a discussion of a corrosion experiment, it is important to relate not only mass changes and appearance but also redox reactions, ionization tendency, solution conditions, and corrosion-protection effects.

• Can a larger mass loss be explained as indicating greater progress of corrosion?
• Does the ionization tendency of the metals correspond to the actual corrosion results?
• Can examples such as aluminum, which has a large ionization tendency but is protected by an oxide film, be explained?
• Can the reason corrosion progressed more readily in NaCl aqueous solution and acidic aqueous solution be explained?
• When dissimilar metals are in contact, which metal becomes the anode and is more susceptible to corrosion?
• Can the concept of sacrificial protection, in which zinc protects iron, be explained?
• Can corrosion suppression by a corrosion inhibitor be explained using corrosion rate and inhibition efficiency?

Example Discussion

In this experiment, the corrosion behavior of metals was compared in NaCl aqueous solution.
In iron, reddish-brown rust was clearly observed, the mass loss was 0.051 g, and the corrosion rate was 5.1 mg/cm²·day.
In contrast, the mass loss of copper was as small as 0.005 g, and only slight changes were observed on the surface.
This was considered to result from iron being more easily oxidized than copper and therefore more susceptible to corrosion in an aqueous solution containing chloride ions.

Zinc also showed a mass loss, but when iron and zinc were brought into contact, there was little rust on the iron and white corrosion products were observed on the zinc side.
This was considered to result from zinc having a greater ionization tendency than iron and being preferentially oxidized as the anode.
The phenomenon in which zinc corrodes itself and thereby suppresses corrosion of iron can be explained as sacrificial protection.

In addition, when iron and copper were brought into contact, corrosion of iron progressed more readily than when iron was used alone.
This was because a potential difference arose between iron and copper, causing iron, the more active metal, to act as the anode and copper, the more noble metal, to act as the cathode.
As a result, the oxidation reaction of iron was promoted and corrosion was considered to have progressed near the contact area.

Regarding the solution conditions, the corrosion rate of iron was larger in NaCl aqueous solution and acidic aqueous solution than in pure water.
In NaCl aqueous solution, chloride ions may have readily destroyed the surface film and promoted localized corrosion.
In addition, under acidic conditions, hydrogen ions readily participate in the reduction reaction, making the oxidation reaction of the metal more likely to proceed.

Under conditions where a corrosion inhibitor was added, the corrosion rate of iron decreased from 5.1 mg/cm²·day to 1.4 mg/cm²·day, and the inhibition efficiency was 72.5%.
This was considered to result from the corrosion inhibitor adsorbing onto the metal surface and forming a protective film or suppressing the anodic or cathodic reaction.
Therefore, to suppress corrosion, it is important to appropriately select the combination of metals, solution conditions, and corrosion-protection treatment.

Summary

In corrosion experiments, the ease with which corrosion progresses can be specifically evaluated by comparing metal mass changes, surface observations, contact between dissimilar metals, solution conditions, and the presence or absence of corrosion inhibitors.

In this reference example, corrosion of iron progressed greatly in NaCl aqueous solution, while corrosion of copper was relatively small.
In addition, zinc acted as sacrificial protection for iron, and the addition of a corrosion inhibitor reduced the corrosion rate of iron.
In a report, it is useful to discuss redox reactions, ionization tendency, galvanic corrosion, and corrosion-protection effects in relation to one another.

Corrosion Is a Redox Reaction

In metal corrosion, an oxidation reaction occurs in which metal atoms lose electrons and become metal ions.
For example, in iron, Fe loses electrons and becomes Fe2+.
These electrons are consumed in reduction reactions involving oxygen, hydrogen ions, or other species.
In other words, metal corrosion is a phenomenon in which oxidation and reduction reactions occur simultaneously.

For corrosion to proceed, there must be a location where the metal releases electrons and a location where a reaction that accepts those electrons occurs.
If there are nonuniform areas on the metal surface, local anodic and cathodic regions may form and the reaction may proceed like a small battery.
This can be explained using the concept of a local cell.

Oxidation of a metal: M → Mn+ + ne-

Oxidation of iron: Fe → Fe2+ + 2e-

Corrosion of Iron and Formation of Rust

During corrosion of iron, Fe is oxidized to Fe2+ and subsequently undergoes further oxidation and hydroxide formation to become hydrated iron oxides known as rust.
Rust is not a single substance but is treated as a complex mixture of oxides and hydroxides including Fe2O3·nH2O and FeOOH.

Iron rust generally appears reddish brown, and corrosion may progress from the surface into the interior.
Unlike a dense and protective film such as the oxide film on aluminum, iron rust is porous and allows water and oxygen to pass through, so it does not easily stop corrosion completely.
Therefore, corrosion-protection treatment is important for iron.

Fe → Fe2+ + 2e-

O2 + 2H2O + 4e- → 4OH-

Fe2+ + 2OH- → Fe(OH)2

Effects of Water and Oxygen

Water and oxygen are strongly involved in the corrosion of many metals, including iron.
Water assists ion movement and provides an environment in which electrochemical reactions can proceed.
Oxygen acts as a reactant in the reduction reaction and accepts electrons released by the metal.
Therefore, corrosion progresses readily under conditions where both water and oxygen are present.

Corrosion progresses less readily in dry air because little moisture is present.
In addition, in water from which oxygen has been removed, the oxygen-reduction reaction is less likely to occur, so corrosion may be suppressed.
When comparing experimental conditions, it is important to discuss that the coexistence of water and oxygen, rather than water or oxygen alone, is important for corrosion.

Effect of Electrolytes

Corrosion progresses more readily in environments containing electrolytes such as sodium chloride aqueous solution.
When an electrolyte is present, the conductivity of the solution increases and the movement of electrons and ions associated with reactions on the metal surface becomes easier.
This allows the reactions at the anodic and cathodic regions to continue more readily.

Iron and steel corrode readily in seawater because water, oxygen, and salts are all present.
If corrosion progresses more readily in salt water than in pure water, one possible cause is that the electrolyte makes corrosion current flow more readily.
The presence or absence of electrolytes greatly affects the corrosion rate.

Effect of Chloride Ions

Chloride ions, Cl-, are representative ions that promote metal corrosion.
Cl- may destroy protective films on metal surfaces and promote localized corrosion.
In particular, in metals such as stainless steel and aluminum that are protected by surface oxide films, chloride ions may destroy the films and cause pitting corrosion.

Corrosion progresses readily in salt water and seawater not only because their conductivity is increased by electrolytes but also because Cl- destabilizes protective films.
If corrosion is large only under experimental conditions containing chloride ions, the effect of Cl- should be discussed.

Effect of Acidic Conditions

Under acidic conditions, a large amount of H+ is present, so corrosion of metals may be promoted.
The metal is oxidized and releases electrons, and H+ accepts those electrons to produce H2.
Reactions in which metals such as zinc and iron dissolve in acids and evolve gas can be explained by this concept.

Under acidic conditions, protective oxide or hydroxide films may also dissolve more readily.
Therefore, even metals whose corrosion is suppressed by protective surface films may corrode more readily in acidic solutions.
If bubbles are observed in an experiment, they can be discussed as corrosion reactions accompanied by hydrogen evolution.

Zn → Zn2+ + 2e-

2H+ + 2e- → H2

Zn + 2H+ → Zn2+ + H2

Ionization Tendency of Metals and Corrosion

The ionization tendency of metals is a concept that represents the order in which metals readily lose electrons and become positive ions.
Metals with greater ionization tendencies tend to be more easily oxidized and more susceptible to corrosion.
For example, zinc and iron ionize more readily than copper and silver and may corrode more readily depending on the conditions.

However, actual susceptibility to corrosion is not completely determined by ionization tendency alone.
Some metals, such as aluminum, have a large ionization tendency but form a dense oxide film on the surface that suppresses corrosion.
Therefore, both ionization tendency and surface films must be considered.

Concept of Local Cells

Metal surfaces are not completely uniform and contain differences such as scratches, dirt, crystal structure, and oxygen concentration.
Therefore, even on the same metal surface, one area may act as an anode while another area acts as a cathode.
Such a state resembling a small battery is called a local cell.

When a local cell forms, the metal is oxidized and dissolves at the anodic region, while reactions such as oxygen reduction occur at the cathodic region.
As a result, localized corrosion may occur in which only part of the metal surface corrodes.
The reason scratches and surface nonuniformity promote corrosion can be explained by the formation of local cells.

Galvanic Corrosion Caused by Contact Between Dissimilar Metals

When different types of metals are in contact in an electrolyte solution, a battery may be formed because of the potential difference and corrosion of one of the metals may be accelerated.
This is called galvanic corrosion or dissimilar-metal contact corrosion.
The metal with the greater ionization tendency tends to become the anode and is preferentially oxidized.

For example, when iron and copper are brought into contact and placed in salt water, iron may become the anode and become more susceptible to corrosion.
On the other hand, when zinc is brought into contact with iron, zinc is more easily oxidized than iron, so the zinc corrodes first and the iron may be protected.
This leads to the concept of sacrificial protection.

Concept of Sacrificial Protection

Sacrificial protection is a method in which a metal that ionizes more readily than the metal to be protected is brought into contact with it, allowing that metal to corrode first and thereby protecting the target metal.
A representative example is the use of zinc to protect iron.
Zinc ionizes more readily than iron, so zinc becomes the anode and is oxidized, while iron is protected as the cathode.

Galvanized steel is iron plated with zinc, and even if the surface is scratched, zinc may corrode sacrificially and protect the iron.
This property is a corrosion-protection method that applies the concepts of the ionization tendency of metals and contact between dissimilar metals.

Zn → Zn2+ + 2e-

Difference Between Tin Plating and Zinc Plating

In corrosion protection of iron, the difference between zinc plating and tin plating is important.
Zinc ionizes more readily than iron, so even if the iron is scratched, the zinc may corrode sacrificially and protect the iron.
In contrast, tin ionizes less readily than iron, so it can protect the iron while the surface remains completely covered, but if the coating is scratched and iron is exposed, the iron becomes more likely to corrode preferentially.

This difference can be explained by the relationship between the ionization tendencies of metals and galvanic corrosion caused by contact between dissimilar metals.
Corrosion protection that merely covers the surface behaves differently when damaged from corrosion protection that also acts sacrificially.
When discussing plating for corrosion protection in a report, attention should be paid to the type of plating metal.

Corrosion Protection by Oxide Films

Aluminum and stainless steel suppress corrosion by forming dense oxide films on their surfaces.
Aluminum is a metal with a large ionization tendency, but a thin Al2O3 oxide film forms on its surface, making corrosion less likely to progress into the interior.
Protection by such a film may be called passivation.

However, if the oxide film is destroyed by chloride ions, strong acids, or strong bases, corrosion may progress.
If aluminum normally shows little change but corrodes in a specific solution during an experiment, destruction of the oxide film should be considered.

Corrosion Protection by Painting

Painting is a corrosion-protection method in which the metal surface is covered with a resin or paint to block contact with water, oxygen, and electrolytes.
Because water, oxygen, and ion movement are necessary for corrosion to proceed, corrosion can be suppressed by keeping these away from the metal surface.
Painting is commonly used to protect steel materials from corrosion.

However, if scratches or pinholes are present in the coating, water and oxygen can enter through those areas and localized corrosion may progress.
The effectiveness of corrosion protection by painting greatly depends on adhesion, coating thickness, and the presence or absence of defects.
Comparing corrosion at coated and scratched areas in an experiment makes it possible to discuss the role of the protective coating.

Corrosion Protection by Plating

Corrosion protection by plating includes methods that cover the metal surface to prevent contact with the environment and methods that act as sacrificial protection.
In cases such as tin plating, where water and oxygen are blocked by covering the metal surface, it is important that the plating layer remain undamaged.
If the coating is scratched, the underlying metal may be exposed and corrosion may progress.

On the other hand, in zinc plating, zinc ionizes more readily than iron, so even when scratches are present, zinc may be oxidized first and protect the iron.
The corrosion-protection effect of plating changes depending on the type of plating metal, film thickness, adhesion, and presence or absence of scratches.

Factors Affecting Corrosion Rate

Corrosion rate changes depending on the type of metal, solution pH, chloride-ion concentration, oxygen concentration, temperature, surface condition, and presence or absence of contact with dissimilar metals.
In general, when the temperature increases, reaction rates increase and corrosion may progress more readily.
In addition, scratches or dirt on the surface make local cells more likely to form.

The higher the conductivity of the solution, the more easily corrosion current flows and the more readily corrosion progresses.
Therefore, corrosion may be greater in salt water than in pure water and in strongly acidic solutions than in weakly acidic solutions.
When comparing corrosion rates, it is important to change and consider one condition at a time.

Evaluation of Corrosion by Mass Change

The progress of corrosion may be evaluated from changes in the mass of a metal sample.
When the metal dissolves into the solution, its mass decreases.
On the other hand, if rust or oxides remain attached to the surface, the apparent mass may increase because oxygen and water have been added.
Therefore, care is required when interpreting mass changes.

Measuring the decrease in the metal itself after removing corrosion products has a different meaning from measuring the mass including the corrosion products.
In a report, it should be clearly stated what the measured mass represents.
If surface rust is not sufficiently removed, errors arise in the evaluation of corrosion amount.

Causes of Error in Corrosion Experiments

Causes of error in corrosion experiments include differences in pretreatment of metal surfaces, presence or absence of scratches and dirt, differences in sample area, errors in solution concentration, changes in pH, temperature differences, differences in immersion time, differences in oxygen supply, solution stirring conditions, mass-measurement errors, and insufficient removal of corrosion products.
Because corrosion is sensitive to surface condition, even small differences among samples can greatly affect the results.

Causes of apparently large corrosion include scratches on the surface, high chloride-ion concentration, acidic conditions, contact between dissimilar metals, and increased temperature.
Causes of apparently small corrosion include formation of oxide films, insufficient dissolved oxygen, insufficient immersion time, and corrosion products acting like protective films.
Organizing the causes according to the experimental conditions makes the discussion more specific.

When the Results Can Be Considered Good

A corrosion experiment can be considered to have produced good results when clear differences in the degree of corrosion are observed depending on the type of metal and solution conditions and those differences can be explained by redox reactions, ionization tendency, the effect of electrolytes, and the presence or absence of corrosion-protection treatment.
For example, if corrosion progresses in salt water and is small under dry conditions, the results clearly demonstrate the effects of water and electrolytes.

In addition, if corrosion of iron is suppressed when it is in contact with zinc, it can be explained as sacrificial protection, while if corrosion of iron progresses when it is in contact with copper, it can be explained as galvanic corrosion caused by contact between dissimilar metals.
If the observations correspond to theory, the experimental results can be considered reasonable.

Example Discussions When the Experiment Did Not Go Well

When a corrosion experiment does not go well, possible causes should be considered from results such as difficulty observing differences in corrosion, a metal different from that expected corroding, unnatural mass changes, variation in surface changes, or only small differences from the control experiment.
Organizing the causes according to the metal surface, solution conditions, immersion time, oxygen supply, measurement method, and sample area makes the discussion easier.

How to Write Points for Improvement

In a discussion of a corrosion experiment, writing not only the causes of error but also points for improvement makes the report easier to organize.
Points for improvement can be organized according to pretreatment of the metal samples, solution conditions, immersion conditions, observation and measurement methods, and corrosion-protection conditions.

Improvements to Metal Samples

• Polish the metal surfaces under the same conditions
• Remove oil and contamination
• Keep the sample areas consistent
• Check for the presence or absence of scratches
• Keep electrical-contact and contact conditions constant
• Do not touch the samples with bare hands

Improvements to Solution and Environmental Conditions

• Prepare the solution concentrations accurately
• Measure or adjust the pH
• Keep the chloride-ion concentration consistent
• Keep the temperature constant
• Keep the immersion time exactly the same
• Keep the oxygen-supply conditions constant

Improvements to Observation and Measurement

• Take photographs before and after corrosion
• Record the color and degree of rust using consistent levels
• Standardize drying conditions before weighing
• Standardize the method for removing corrosion products
• Perform multiple measurements and calculate the average
• Prepare a control experiment

Difference Between a Superficial Discussion and a Good Discussion

In a discussion of a corrosion experiment, simply writing that “it rusted” or “it changed color” results in a superficial discussion.
A good discussion relates redox reactions, the ionization tendency of metals, the roles of water, oxygen, and electrolytes, local cells, and corrosion-protection methods.

Examples of Expressions That Can Be Used in Reports

The following expressions can be used when writing the results and discussion of corrosion experiments.
Adjust the necessary parts according to your own experimental results.

• Corrosion is a reaction in which a metal loses electrons and is oxidized.
• In corrosion, the oxidation reaction of the metal and reduction reactions of oxygen or hydrogen ions proceed simultaneously.
• Water assists ion movement and provides an environment in which corrosion reactions can proceed.
• Oxygen participates in a reduction reaction that accepts electrons and is a factor that promotes corrosion.
• When an electrolyte is present, the conductivity of the solution increases and corrosion becomes more likely to proceed.
• Chloride ions may destroy protective films and promote localized corrosion.
• Metals with greater ionization tendencies tend to be more easily oxidized.
• When dissimilar metals are in contact, the potential difference may promote corrosion of one of the metals.
• In sacrificial protection, a metal that ionizes more readily than the metal to be protected is corroded first.
• Painting and plating suppress corrosion by preventing contact with water and oxygen.

Points to Check When Discussing Corrosion Experiments

Checking the following points before writing the report makes the discussion easier to write.

• Is corrosion explained as a redox reaction?
• Is it stated that the metal is oxidized and becomes metal ions?
• Are reduction reactions involving oxygen or H+ considered?
• Are the effects of water and oxygen explained?
• Are the effects of electrolytes and chloride ions considered?
• Is the ionization tendency of metals related to corrosion susceptibility?
• Is the concept of local cells used?
• Is galvanic corrosion caused by contact between dissimilar metals explained?
• Is the mechanism of sacrificial protection explained?
• Are corrosion-protection methods using oxide films, painting, and plating considered?
• Is the meaning of mass change interpreted correctly?
• Do the points for improvement correspond to the causes of error?

Summary

A corrosion experiment is an experiment that investigates the phenomenon in which a metal reacts with its surrounding environment and becomes oxidized.
Corrosion proceeds through a combination of an oxidation reaction in which the metal loses electrons and reduction reactions in which oxygen or hydrogen ions accept electrons.
Water, oxygen, electrolytes, chloride ions, and acidic conditions are important factors that promote corrosion.

The susceptibility of metals to corrosion is related to their ionization tendencies.
However, in practice, the presence or absence of oxide films, surface condition, contact with dissimilar metals, pH, chloride-ion concentration, and other factors also have an effect.
Sacrificial protection using zinc, painting, plating, and oxide films are representative methods for preventing corrosion.

In a report, rather than simply writing that “it rusted” or “it corroded,” organize and discuss redox reactions, the ionization tendency of metals, local cells, galvanic corrosion caused by contact between dissimilar metals, sacrificial protection, the effects of water, oxygen, and electrolytes, causes of error, and points for improvement.
Corrosion experiments are important experiments for understanding the relationship between the properties of metallic materials and electrochemical reactions.