Electrochemistry Experiment Discussion Examples | Redox Potentials and the Reactivity of Metal Ions

What Is an Electrochemistry Experiment?

An electrochemistry experiment investigates the relationship between redox reactions and electrical energy.
In a redox reaction, one substance loses electrons and another substance receives electrons.
When this electron transfer is observed through electrodes and an external circuit, it can be measured as a potential difference or electric current.

For example, when a metal is immersed in a metal-ion solution, the metal may dissolve and become ions, or metal ions may receive electrons and be deposited as metal.
The ease with which such reactions occur is affected by the type of metal, the concentration of metal ions, the redox potential, the condition of the electrode surface, and other factors.

In a report, the observed potential and changes in the metal are explained in relation to oxidation reactions, reduction reactions, and electrode potentials.

Basics of Oxidation and Reduction

Oxidation is a reaction in which a substance loses electrons.
Reduction is a reaction in which a substance receives electrons.
In electrochemistry, the electrode at which oxidation occurs is sometimes called the anode, while the electrode at which reduction occurs is called the cathode.

Oxidation: loss of electrons
Reduction: gain of electrons

The reaction in which a metal becomes a metal ion is oxidation.
The reaction in which a metal ion is deposited as a metal is reduction.
Which reaction occurs is determined by the oxidizing agent, reducing agent, and electrode potentials involved.

Results to Observe in an Electrochemistry Experiment

In electrochemistry experiments, not only the potential value but also changes in the electrode surface, changes in solution color, metal deposition, bubble formation, and changes in current are observed.
These results provide clues for discussing which electrode underwent oxidation or reduction.

Main Items to Include in the Results

• Types of metal electrodes used
• Types and concentrations of metal-ion solutions used
• Measured potential difference or electromotive force
• Direction of electrode connections
• Changes in the electrode surfaces
• Presence or absence of metal deposition
• Presence or absence of dissolution of the metal electrodes
• Changes in solution color
• Presence or absence of bubble formation
• Changes in measured values over time
• Comparison with theoretical values or standard electrode potentials
• Sources of error

Reference Experimental Values for Electrochemistry and Examples of Analysis of Redox Potentials and Metal-Ion Reactivity

Here, reference experimental values are organized for discussing reactions between metals and metal ions, redox potentials, metal deposition and dissolution, and the electromotive force of cells.
Standard electrode potentials, the Nernst equation, changes in potential caused by concentration, the order of metal reactivity, and measurement errors are organized in a form that can be readily used in reports.

Electrochemistry deals with redox reactions involving the transfer of electrons.
A reaction in which a metal loses electrons and becomes metal ions is oxidation, while a reaction in which metal ions receive electrons and become metal is reduction.
Because different metals differ in how readily they lose electrons and how readily their metal ions receive electrons, differences arise in metal reactivity and cell electromotive force.

Reference Experimental Conditions

Basics of Oxidation and Reduction

When a metal piece is placed in a metal-ion solution, the reaction proceeds when the combination consists of a metal that is readily oxidized and metal ions that are readily reduced.

Reference Values of Standard Electrode Potentials

The standard electrode potential E° is an indicator of how readily reduction occurs under standard-state conditions.
The larger the value, the more readily the ion is considered to receive electrons and be reduced.

Metals with smaller standard electrode potentials tend, as elemental metals, to release electrons and become cations more readily.
Therefore, Mg and Zn are considered more readily oxidized than Cu and Ag.

Examples of Reactions Between Metal Pieces and Metal-Ion Solutions

The following are examples of observations when a metal piece is placed in a solution containing ions of another metal.
When a reaction proceeds, the metal piece may dissolve and the metal ions in solution may be deposited as metal.

When Zn is placed in CuSO₄ solution, Zn dissolves as Zn²⁺ and Cu²⁺ is deposited as Cu.
This is because Zn is more readily oxidized than Cu.

Order of Metal Reactivity

The order of metal reactivity can be estimated from displacement reactions between metal pieces and metal ions.

Highly reactive metals readily lose electrons and become cations and therefore readily reduce other metal ions.
In this reference example, the ease of oxidation is considered to decrease in the order Mg, Zn, Fe, Cu, and Ag.

Example Measurement of the Electromotive Force of a Daniell Cell

In a Daniell cell using Zn and Cu electrodes, Zn is oxidized and Cu²⁺ is reduced.

Negative electrode: Zn → Zn²⁺ + 2e⁻

Positive electrode: Cu²⁺ + 2e⁻ → Cu

Overall: Zn + Cu²⁺ → Zn²⁺ + Cu

The standard electromotive force is calculated by subtracting the standard electrode potential of the negative electrode from that of the positive electrode.

E°_cell = E°_cathode − E°_anode

Example Calculation of the Electromotive Force of a Daniell Cell

If the standard electrode potential of Cu²⁺/Cu is +0.34 V and that of Zn²⁺/Zn is −0.76 V,

E°_cell = 0.34 − (−0.76) = 1.10 V

Therefore, the electromotive force of the Daniell cell under standard-state conditions is calculated to be 1.10 V.
If the measured value is 1.06 V, it is slightly smaller than the theoretical value, but the direction of the cell reaction is considered to agree with theory.

Electromotive Force and Spontaneity of Reactions

If the electromotive force is positive, the redox reaction in that direction can be judged to proceed readily and spontaneously.

Effect of Concentration According to the Nernst Equation

Electrode potential is also affected by metal-ion concentration.
At 25 °C, it can be expressed as follows.

E = E° + (0.0592/n) log[oxidized form]/[reduced form]

For the metal electrode Mⁿ⁺ + ne⁻ → M, if the activity of metal M is approximated as 1,

E = E° + (0.0592/n) log[Mⁿ⁺]

As the metal-ion concentration decreases, the reduction potential decreases because the amount of ions available for reduction becomes smaller.

Example of Changes in Electromotive Force Caused by Concentration Differences

In a Zn-Cu cell, changing the Cu²⁺ or Zn²⁺ concentration changes the electromotive force.

The electromotive force of a cell is affected not only by standard electrode potentials but also by the concentrations of oxidized and reduced species.

Example Calculation of Metal Deposition

In electrolysis and displacement reactions, the amount of metal deposited can be considered from the amount of electrons transferred.
Here, the reduction of Cu²⁺ and deposition of Cu are used as an example.

Cu²⁺ + 2e⁻ → Cu

If the actual amount deposited is smaller than the theoretical value, current efficiency, detachment from the electrode surface, losses during washing and drying, and side reactions should be considered.

When the Measured Electromotive Force Is Smaller Than the Theoretical Value

Even if the measured value is lower than the theoretical value, it should not immediately be concluded that the reaction equation is incorrect.
The electrode surface, internal resistance, concentration, and contact condition should be checked.

Example of Changes in Electromotive Force With Measurement Time

If a cell remains connected, the electromotive force may gradually change as the reaction proceeds and the electrode surfaces change.

When comparing electromotive forces, it is important to standardize the timing of measurements.

Examples of Predicting Redox Reactions

Standard electrode potentials can be used to predict whether one metal can reduce ions of another metal.

In the direction in which the difference in E° is positive, the redox reaction is considered to proceed spontaneously more readily.

Example of How to Write the Results

When a Zn piece was placed in CuSO₄ solution, a reddish-brown solid was deposited on the Zn surface and the blue color of the solution became lighter.
This occurred because Zn was oxidized to Zn²⁺ and Cu²⁺ was reduced to Cu.
The standard electrode potentials are +0.34 V for Cu²⁺/Cu and −0.76 V for Zn²⁺/Zn, so Cu²⁺ is more readily reduced.

The standard electromotive force of the Zn-Cu cell was calculated as E°_cell = 0.34 − (−0.76) = 1.10 V.
The measured electromotive force was 1.06 V, which was slightly smaller than the theoretical value.
This difference may have been caused by an oxide film on the electrode surface, resistance of the salt bridge and wires, deviations in solution concentration, polarization, and similar factors.

Comparison of the reactions between metal pieces and metal-ion solutions showed that Zn and Fe reduced Cu²⁺ and caused Cu to be deposited, whereas Cu did not reduce Zn²⁺.
From these results, Zn and Fe are considered more readily oxidized and more reactive than Cu.

Points for Connecting the Results to the Discussion

In an electrochemistry discussion, it is important to relate the observed metal deposition and dissolution to oxidation, reduction, standard electrode potentials, electromotive force, and concentration changes.

• Can you explain which metal was oxidized and which metal ion was reduced?
• Can you discuss metal deposition and changes in solution color in relation to electron transfer?
• Can you compare the ease of reduction from the relative standard electrode potentials?
• Can you correctly determine the positive and negative electrodes of the cell?
• Can you calculate the electromotive force using E°_cell = E°_cathode − E°_anode?
• Can you explain that a reaction tends to proceed spontaneously when the electromotive force is positive?
• Can you use the Nernst equation to explain that metal-ion concentration affects electrode potential?
• Can you discuss internal resistance, electrode surfaces, concentration deviations, and polarization as reasons the measured value may be lower than the theoretical value?
• Can you explain that electromotive force may change depending on measurement time and contact conditions?

Example Discussion

In this experiment, reactions between metal pieces and metal-ion solutions and the electromotive force of cells using metal electrodes were investigated.
When Zn was placed in CuSO₄ solution, reddish-brown Cu was deposited on the Zn surface.
This occurred because Zn lost electrons and became Zn²⁺, while Cu²⁺ received electrons and became Cu.
Therefore, in this reaction Zn was oxidized and Cu²⁺ was reduced.

Comparing the standard electrode potentials, Cu²⁺/Cu is +0.34 V and Zn²⁺/Zn is −0.76 V.
Cu²⁺ more readily receives electrons and is reduced, while Zn more readily loses electrons and is oxidized.
Because of this potential difference, the reaction Zn + Cu²⁺ → Zn²⁺ + Cu is considered to proceed spontaneously.

In the Zn-Cu cell, the Zn electrode functioned as the negative electrode and the Cu electrode as the positive electrode.
From E°_cell = E°_cathode − E°_anode, the standard electromotive force was calculated as 0.34 − (−0.76) = 1.10 V.
The measured value was 1.06 V, which was slightly smaller than the theoretical value.
This difference may have resulted from an oxide film or contamination on the electrode surfaces, internal resistance of the salt bridge, poor wire contact, and deviations in solution concentration.

In addition, changing the metal-ion concentration changed the electromotive force.
This is because electrode potential depends not only on the standard electrode potential but also on the concentrations of oxidized and reduced species.
According to the Nernst equation, when the Cu²⁺ concentration decreases, the potential of the Cu electrode decreases and the overall electromotive force of the Zn-Cu cell also decreases.
Therefore, when comparing cell electromotive forces, not only the types of metals but also the solution concentrations must be kept constant.

In some cases, the electromotive force decreased as the measurement time became longer.
This may have occurred because the progression of the cell reaction changed the condition of the electrode surfaces and the ion concentrations and caused polarization.
To accurately compare electromotive forces, it is important to prepare the electrode surfaces, confirm the contact of the salt bridge and wires, and perform measurements at the same timing.

Summary

In electrochemistry, metal deposition, dissolution of metal pieces, and cell electromotive force can be explained by considering redox reactions as electron-transfer processes.
Metal ions with larger standard electrode potentials are more readily reduced, while metals with smaller standard electrode potentials tend to be more readily oxidized as elemental metals.

This reference example covered oxidation and reduction, standard electrode potentials, reactions between metal pieces and metal ions, the order of metal reactivity, Daniell cells, electromotive-force calculations, the Nernst equation, concentration differences, amounts of metal deposition, differences between measured and theoretical values, and changes with measurement time.
In a report, it is useful to relate observations to redox potentials and explain why a reaction proceeds and why a measured value differs from the theoretical value.

What Is a Redox Potential?

A redox potential is an indicator of how readily a redox couple receives electrons or loses electrons.
In general, the higher the reduction potential, the more readily the chemical species receives electrons and is reduced.
Conversely, a metal with a low reduction potential tends to lose electrons and be oxidized more readily.

For a combination of a metal and metal ions, comparing electrode potentials makes it possible to consider which metal is more readily oxidized and which metal ions are more readily reduced.

How to Think About Standard Electrode Potentials

A standard electrode potential is the electrode potential of a redox couple under standard-state conditions.
Under standard-state conditions, factors such as ion concentration in solution and gas pressure are standardized.
Because potentials measured experimentally are often obtained under conditions different from the standard state, they may not completely agree with standard electrode potentials.

Even so, comparison of standard electrode potentials allows prediction of the direction in which a reaction is more likely to proceed.
In a cell reaction, reduction tends to occur on the side with the higher standard reduction potential, while oxidation tends to occur on the side with the lower value.

How to Interpret the Reactivity of Metal Ions

The reactivity of metal ions can be considered in terms of how readily they receive electrons.
Metal ions that are readily reduced may receive electrons from another metal and be deposited as metal.
For example, when a metal is immersed in a solution containing ions of another metal, the formation of a deposit on the metal surface may indicate that the metal ions were reduced.

On the other hand, if the metal dissolves into the solution, that metal is considered to have been oxidized.
In this way, considering metal deposition together with dissolution of the metal electrode makes it easier to explain the redox reaction.

Relationship With the Ionization Tendency of Metals

The ionization tendency of a metal refers to the tendency of the metal to lose electrons and become a cation.
Metals with a greater ionization tendency are more readily oxidized and tend to enter solution as metal ions.
On the other hand, ions of metals with a smaller ionization tendency may be more readily reduced and deposited as metal.

In electrochemistry experiments, the direction of metal displacement reactions and cell reactions can be explained by relating ionization tendency to redox potentials.

Discussion of Metal Displacement Reactions

A metal displacement reaction is a reaction in which one metal reduces ions of another metal while itself being oxidized.
When a metal plate is placed in a metal-ion solution, this may be observed as another metal being deposited on the surface of the metal plate.

This reaction occurs more readily with a combination of a metal that is readily oxidized and metal ions that are readily reduced.
In a report, the deposited metal, changes in solution color, and changes in the surface of the metal plate are used as the basis for discussion.

Discussion of the Electromotive Force of a Cell

In a cell, oxidation and reduction occur separately at two electrodes, and the difference in their potentials is observed as the electromotive force.
The electromotive force can be considered as the difference between the cathode and anode electrode potentials.

Electromotive force = Cathode potential − Anode potential

The theoretical electromotive force can be calculated from standard electrode potentials, but measured values may deviate because of concentration, temperature, internal resistance, the salt bridge, the condition of the electrode surfaces, and other factors.

The Nernst Equation and the Effect of Concentration

Electrode potential changes with the concentration of metal ions.
Standard electrode potential is the value under standard conditions, but because concentrations in an experiment may differ from standard conditions, the actual potential also changes.
The equation expressing this relationship is the Nernst equation.

Even when detailed calculations are not performed, it can be explained that “when the metal-ion concentration changes, the electrode potential changes, and the measured electromotive force also changes.”

Discussion of a Concentration Cell

A concentration cell is a cell in which a potential difference arises because the concentrations of the solutions on the two sides differ, even though the same metal and the same metal ions are used.
When the metal-ion concentrations differ, the electrode potentials also differ, producing an electromotive force.

This experiment demonstrates that a potential difference can arise solely from a concentration difference even when the same type of metal is used.
It is therefore important for understanding that electrode potential depends on concentration.

Effect of the Electrode Surface Condition

In electrochemistry experiments, the condition of the electrode surface strongly affects measurement results.
If an oxide film, dirt, oil, or deposits are present on the electrode surface, electron transfer may be hindered or the actual surface area participating in the reaction may change.
As a result, the potential may not stabilize, may deviate from the theoretical value, or the reaction may become slower.

Role of the Salt Bridge and Associated Errors

In cell experiments, two half-cells may be connected by a salt bridge.
The salt bridge maintains charge balance in the solutions while preventing direct mixing.
If the salt bridge does not function properly, current may not flow readily or the potential may become unstable.

Poor contact of the salt bridge, bubbles, inappropriate concentration, and drying can cause deviations or instability in measured values.

Effect of Internal Resistance

Real cells have internal resistance.
The solution, salt bridge, electrode surfaces, and connections act as resistances, and values lower than the theoretical electromotive force may be measured.
In particular, when measurement is performed while current flows, a voltage drop occurs because of internal resistance.

Effect of Liquid Junction Potential

Where different solutions are in contact, a liquid junction potential may arise because of differences in the migration rates of ions.
A salt bridge can reduce this effect, but cannot completely eliminate it.
Liquid junction potential becomes a source of error in precise potential measurements.

When the Potential Changes Over Time

If the potential changes over time, possible causes include changes in the electrode surfaces, changes in ion concentrations, formation of oxide films, metal deposition, and insufficient mixing of the solution.
In electrochemistry experiments, the value immediately after measurement begins may differ from the value after stabilization.

Discussion of Metal Deposition

When metal is deposited on an electrode surface, the metal ions in solution are considered to have received electrons and been reduced.
The color and shape of the deposited metal are affected by the type of metal, deposition rate, current density, and condition of the electrode surface.

If deposition is uneven, contamination of the electrode surface or local differences in current density may have affected the result.

Discussion When a Metal Electrode Dissolves

When a metal electrode dissolves, the metal is considered to have lost electrons and become metal ions.
This is an oxidation reaction.
Changes such as clouding of the electrode surface, loss of mass, and changes in solution color may be observed.

Discussion of Changes in Solution Color

Some metal ions have characteristic colors.
When electrochemical reactions change the concentration or oxidation state of metal ions, the color of the solution may change.
For example, when a colored metal ion is reduced and deposited as metal, its concentration decreases and the solution color may become lighter.

Discussion When Bubbles Are Generated

If bubbles are generated during an electrochemistry experiment, electrolysis of water, generation of hydrogen or oxygen, or reaction with an acid may be involved.
In addition to reactions of metal ions, reactions involving water, hydrogen ions, or hydroxide ions may occur as electrode reactions.

Discussion of Overpotential

In actual electrode reactions, a potential larger than the theoretically required potential may be necessary.
This difference is called overpotential.
Overpotential changes depending on the electrode material, electrode surface, type of reaction, current density, and other factors.

In particular, the effect of overpotential is large in gas-evolution reactions, and the actual reaction order may not always be completely explained by theoretical values alone.

Redox Potential and Reaction Spontaneity

Whether a redox reaction proceeds spontaneously can be considered from the potential difference of the entire cell.
If the electromotive force of the cell is positive, the redox reaction in that direction is considered to proceed spontaneously more readily.
Conversely, an external voltage must be applied to cause a reaction that does not proceed spontaneously.

Causes of Differences Between Theoretical and Measured Values

In electrochemistry experiments, theoretical and measured values often do not agree.
Possible causes include solution concentrations differing from standard conditions, temperature differing from the standard state, contamination of electrode surfaces, poor salt-bridge contact, internal resistance, liquid junction potentials, and measurement-instrument errors.

When the Result Can Be Considered Good

A good result in an electrochemistry experiment is indicated when the measured potential does not greatly contradict the theoretical or literature value and when the observed electrode reactions agree with the direction predicted from the redox potentials.
In addition, if the potential is stable and changes in the electrode surfaces and solution correspond to the reaction equations, the result can be considered reliable.

Example Discussion When the Experiment Did Not Go Well

When an electrochemistry experiment does not go well, possible causes can be considered from results such as an unstable potential, a large deviation from the theoretical value, absence of metal deposition, unexpected changes in the electrode, or generation of many bubbles.
It is easier to organize the discussion by separately considering concentration, electrode surfaces, salt bridges, internal resistance, connections, measurement time, and side reactions.

How to Write Points for Improvement

In a discussion of an electrochemistry experiment, including points for improvement as well as sources of error makes the report easier to organize.
Improvements are easier to write when divided into solution preparation, electrode treatment, measurement procedures, and apparatus conditions.

Improvements to Solution Preparation

• Prepare metal-ion concentrations accurately
• Measure at concentrations close to standard conditions
• Mix the solution sufficiently
• Keep the temperature as constant as possible
• Avoid contamination by impurities or other metal ions

Improvements to Electrode Treatment

• Clean the electrode surfaces
• Remove oxide films and contamination
• Keep the electrode surface areas consistent
• Immerse the electrodes sufficiently in the solution
• Replace or clean electrodes with deposited material when necessary

Improvements to Measurement Procedures

• Connect the salt bridge correctly
• Avoid bubbles and poor contact
• Read the value only after the potential has stabilized
• Do not reverse the polarity
• Check the range and connections of the measuring instrument
• Perform multiple measurements to confirm reproducibility

Difference Between a Superficial Discussion and a Good Discussion

In a discussion of an electrochemistry experiment, simply writing that “a potential appeared” or “metal was deposited” results in a superficial discussion.
A persuasive discussion can be produced by connecting redox potentials, electron transfer, metal-ion reactivity, electrode surfaces, and concentration conditions.

Examples of Expressions That Can Be Used in Reports

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

• The measured potential difference is considered to originate from the difference in electrode potentials between the two redox couples.
• Metal ions with higher reduction potentials more readily receive electrons and are more readily reduced.
• Metals with lower reduction potentials more readily lose electrons and are more readily oxidized.
• Because a deposit was observed on the metal surface, the metal ions in solution are considered to have been reduced.
• If a metal electrode dissolved, that metal may have been oxidized and become metal ions.
• Possible reasons the measured value differed from the value predicted from standard electrode potentials include concentration conditions and internal resistance.
• According to the Nernst equation, electrode potential depends on metal-ion concentration.
• An oxide film or contamination on the electrode surface may hinder electron transfer and make the measured value unstable.
• Poor salt-bridge contact may hinder ion movement and affect the measured electromotive force.
• In electrochemical reactions, side reactions such as gas evolution must also be considered in addition to the target reaction.

Points to Check When Discussing an Electrochemistry Experiment

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

• Have you explained which substance was oxidized?
• Have you explained which substance was reduced?
• Have you related electrode potentials and standard electrode potentials to reactivity?
• Have you explained metal-ion reactivity in terms of how readily electrons are received?
• Have you related metal deposition and electrode dissolution to redox reactions?
• Have you compared the theoretical and measured electromotive forces?
• Have you considered the effect of concentration conditions?
• Have you considered the electrode-surface condition as a source of error?
• Have you considered the effects of the salt bridge, internal resistance, and liquid junction potential?
• Have you discussed the possibility of side reactions such as bubble formation?
• Have you checked whether the measured value was read after it stabilized?
• Do the points for improvement correspond to the sources of error?

Summary

In electrochemistry experiments, electron transfer caused by redox reactions is observed as a potential difference, current, and changes in electrode surfaces.
By comparing redox potentials, it is possible to consider which metals are more readily oxidized and which metal ions are more readily reduced.

The reactivity of metal ions is related to their ease of reduction and to the ionization tendency of metals.
Metal ions with higher reduction potentials more readily receive electrons and may be reduced and deposited as metal.
On the other hand, metals with lower reduction potentials more readily lose electrons and are considered more likely to dissolve into solution as metal ions.

In a report, do not simply write that “a potential appeared” or “metal was deposited.”
Discuss the results by relating oxidation reactions, reduction reactions, standard electrode potentials, the Nernst equation, concentration conditions, and the condition of the electrode surfaces.
If deviations between theoretical and measured values can also be explained in terms of concentration, the salt bridge, internal resistance, liquid junction potential, and electrode-surface contamination, a persuasive electrochemistry laboratory discussion can be produced.