Redox titration is an analytical chemistry experiment used to determine the amount or concentration of a substance in a sample by utilizing a reaction between an oxidizing agent and a reducing agent. Among these methods, titration using potassium permanganate is commonly performed in introductory chemistry and analytical chemistry experiments at universities.
Potassium permanganate titration uses the strong oxidizing power and characteristic reddish-purple color of permanganate ions. Because the endpoint can sometimes be determined without adding a separate indicator, there are discussion points that differ from those of acid-base titration.
This article clearly explains how to interpret redox titration results, endpoint determination in potassium permanganate titration, reaction equations and quantitative relationships, sources of error, and discussion examples that can be used in reports.
Note: This article is a reference intended to assist with discussions of results obtained in chemistry experiments at universities and similar institutions. For actual experimental procedures and safety precautions, always follow the instructions in your university’s laboratory manual and those given by your instructor or TA.
- What Is Redox Titration?
- Characteristics of Potassium Permanganate Titration
- Why Potassium Permanganate Acts as a Self-Indicator
- Why Acidic Conditions Are Important
- Items to Examine in the Results
- Reference Experimental Values and Analysis Examples for Potassium Permanganate Redox Titration
- Reference Experimental Conditions
- Half-Reaction of Permanganate Ions Under Acidic Conditions
- Reaction for the Determination of Fe2+
- Example Titration Data for an Fe2+ Sample
- Example of Calculating the Fe2+ Concentration
- Example of Expressing the Amount of Fe2+ in mg/L
- Example of Standardizing KMnO4 Standard Solution Using Oxalic Acid
- Example of Correcting Fe2+ Concentration Using the Standardized Concentration
- Example of Quantifying an Oxalic Acid Sample
- Example of Endpoint Color Changes
- Example of Insufficient Acidic Conditions
- Example When Heating Is Required
- Example of Blank Correction
- Effect of Over-Titration
- Effect of Air Oxidation of the Reducing Agent
- Changes in the Concentration of KMnO4 Standard Solution During Storage
- Main Sources of Error
- Example of How to Write the Results
- Points to Connect to the Discussion
- Example Discussion Text
- Summary
- How to Think About Reaction Equations and Quantitative Relationships
- Error in Endpoint Determination
- Discussion of Variation in Titration Volumes
- Error Due to the Concentration of the Potassium Permanganate Solution
- Error When Acidic Conditions Are Insufficient
- Error Due to the Oxidation or Reduction State of the Sample
- Error Due to Insufficient Stirring
- Effect of Temperature
- Burette-Reading Error
- How to Write About Excluding a Titration Value
- When the Result Can Be Considered Good
- Example Discussion When the Experiment Did Not Go Well
- How to Write Points for Improvement
- Difference Between a Superficial Discussion and a Good Discussion
- Examples of Expressions That Can Be Used in Reports
- Points to Check When Discussing Redox Titration
- Summary
What Is Redox Titration?
Redox titration is a titration method used to determine the concentration of an unknown sample by utilizing a reaction in which electrons are transferred between an oxidizing agent and a reducing agent. An oxidizing agent is a substance that accepts electrons from another substance, while a reducing agent is a substance that gives electrons to another substance.
Acid-base titration uses a neutralization reaction involving hydrogen ions and hydroxide ions, whereas redox titration considers the amount of reaction based on electron transfer. Therefore, it is important to correctly account for the coefficients in the reaction equation and the number of electrons transferred.
In potassium permanganate titration, permanganate ions act as the oxidizing agent and oxidize reducing substances in the sample. The amount or concentration of the reducing substance can be determined from the titration volume.
Characteristics of Potassium Permanganate Titration
A major characteristic of potassium permanganate titration is that the permanganate ion itself has a strong reddish-purple color. Under acidic conditions, permanganate ions are reduced to manganese(II) ions, which are almost colorless.
As long as reducing substances remain in the sample, the added permanganate ions are immediately reduced and the reddish-purple color disappears. After all of the reducing substances have reacted, a slight excess of permanganate ions causes a pale reddish-purple color to remain. This persistent color is used to determine the endpoint.
| State | Observed Change | Possible Interpretation |
|---|---|---|
| During titration | The reddish-purple color disappears immediately | Reducing substances remain in the sample |
| Near the endpoint | The reddish-purple color becomes slower to disappear | The amount of reducing substance has decreased |
| Endpoint | A pale reddish-purple color remains for a while | Permanganate ions have become slightly excessive |
| Excess addition | The solution becomes dark reddish purple | The endpoint has probably been exceeded |
Why Potassium Permanganate Acts as a Self-Indicator
In potassium permanganate titration, the reddish-purple color of permanganate ions can be used to determine the endpoint. When the reagent used for titration itself serves as an indicator in this way, it can be described as acting as a self-indicator.
Under acidic conditions, permanganate ions are reduced to lightly colored manganese(II) ions. Therefore, while a reducing agent remains in the sample, the reddish-purple color disappears. After the endpoint is exceeded, the color of excess permanganate ions remains, allowing the endpoint to be confirmed as a pale reddish-purple color.
Example Discussion:
In potassium permanganate titration, the permanganate ion itself shows a reddish-purple color, so the endpoint can be determined without adding a separate indicator. While reducing substances remain in the sample, the added permanganate ions are reduced and the color disappears. After all reducing substances have reacted, a slight excess of permanganate ions produces a persistent pale reddish-purple color, so this point can be judged as the endpoint.
Why Acidic Conditions Are Important
In potassium permanganate titration, it is important to carry out the reaction under acidic conditions. Under acidic conditions, permanganate ions are reduced to manganese(II) ions. On the other hand, if the conditions are inappropriate, precipitates such as manganese dioxide may form, making the reaction more complex.
When the reaction conditions change, the number of electrons accepted by the permanganate ion and the products formed also change, which affects the titration calculations. Therefore, in the report, whether the acidic conditions were appropriate and whether any precipitate or turbidity was observed can be discussed.
Example Discussion:
In potassium permanganate titration, it is important to perform the reaction under acidic conditions. If the acidic conditions were insufficient, reactions different from the assumed reaction could occur, such as precipitation of permanganate ions as manganese dioxide. In that case, the quantitative relationship of the reaction would change and errors could occur in the titration volume and concentration calculation.
Items to Examine in the Results
In the results of potassium permanganate titration, organize not only the titration volume but also the endpoint color, variation in the measured values, changes in the sample solution, and the reaction equation used in the titration. As in acid-base titration, it is important to compare multiple titration values and confirm reproducibility.
Main Items to Include in the Results
- Volume of the sample solution
- Concentration of the potassium permanganate solution
- Titration volume for each trial
- Average titration volume
- Color change at the endpoint
- Calculated sample concentration or content
- Variation in titration values
- Presence or absence of precipitate or turbidity
- If a value was excluded, the reason for exclusion
Example of How to Write the Results:
A 10.00 mL sample solution was titrated with an aqueous potassium permanganate solution. The endpoint was taken as the point at which a pale reddish-purple color persisted for approximately 30 seconds. The titration volumes were 12.45 mL for the first trial, 12.32 mL for the second trial, and 12.34 mL for the third trial. Because the first trial may have slightly exceeded the endpoint, the average of the second and third trials, 12.33 mL, was used to determine the sample concentration.
Reference Experimental Values and Analysis Examples for Potassium Permanganate Redox Titration
Here, reference experimental values are organized for discussing titration volume, reaction equations, equivalence relationships, unknown sample concentration, standardization, endpoint determination, and sources of error in redox titration using a potassium permanganate standard solution. Fe2+ determination, oxalic acid determination, standardization of KMnO4 standard solution, the importance of acidic conditions, and the pale red endpoint color are summarized in a form that is easy to use in reports.
The permanganate ion MnO4− is a strong oxidizing agent and is reduced to Mn2+ under acidic conditions. Because KMnO4 itself is purple, a pale reddish-purple color due to excess MnO4− remains at the endpoint, allowing the endpoint to be determined without an indicator.
Reference Experimental Conditions
| Item | Details |
|---|---|
| Type of titration | Redox titration |
| Titrant | Potassium permanganate standard solution KMnO4 |
| Examples of analytes | Fe2+, oxalic acid, hydrogen peroxide, nitrite, etc. |
| Main condition | Measure under acidic conditions |
| Endpoint | Point at which a pale reddish-purple color remains for a while |
| Evaluation items | Titration volume, reaction ratio, unknown sample concentration, standardization, blank correction, sources of error |
Half-Reaction of Permanganate Ions Under Acidic Conditions
Under acidic conditions, permanganate ions are reduced as follows.
MnO4− + 8H+ + 5e− → Mn2+ + 4H2O
| Item | Details | Meaning in Titration |
|---|---|---|
| MnO4− | Oxidizing agent | Accepts electrons from the other substance |
| Mn2+ | Reduction product | Formed under acidic conditions |
| 5e− | Accepts five electrons | 1 mol of MnO4− corresponds to 5 mol of electrons |
| H+ | Provides acidic conditions | If insufficient, the reaction changes |
If the amount of acid is insufficient, the reaction may not proceed ideally and a brown precipitate of MnO2 may form.
Reaction for the Determination of Fe2+
Fe2+ is oxidized to Fe3+.
Fe2+ → Fe3+ + e−
The overall reaction equation with permanganate ions can be expressed as follows.
MnO4− + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O
| Substance | Oxidation / Reduction | Reaction Ratio | Use in Calculation |
|---|---|---|---|
| MnO4− | Reduced | 1 | Determined from the KMnO4 titration volume |
| Fe2+ | Oxidized | 5 | 5 times the molar amount of MnO4− |
| H+ | Reaction condition | 8 | Maintains acidic conditions |
Example Titration Data for an Fe2+ Sample
The following are reference data obtained by titrating 25.00 mL of a sample containing Fe2+ with 0.0200 mol/L KMnO4 standard solution.
| Trial | Sample Volume | KMnO4 Concentration | Titration Volume | Endpoint Assessment |
|---|---|---|---|---|
| 1st | 25.00 mL | 0.0200 mol/L | 15.42 mL | Slightly dark reddish purple |
| 2nd | 25.00 mL | 0.0200 mol/L | 15.36 mL | Good |
| 3rd | 25.00 mL | 0.0200 mol/L | 15.34 mL | Good |
| 4th | 25.00 mL | 0.0200 mol/L | 15.37 mL | Good |
| Average | 25.00 mL | 0.0200 mol/L | 15.36 mL | Average of the 2nd to 4th trials |
This is an example in which the first trial is judged to have been slightly over-titrated while confirming the endpoint color, and the average of the second to fourth trials, 15.36 mL, is used as the representative value.
Example of Calculating the Fe2+ Concentration
Consider the case where the KMnO4 concentration is 0.0200 mol/L, the average titration volume is 15.36 mL, and the sample volume is 25.00 mL.
n(MnO4−) = 0.0200 mol/L × 0.01536 L = 3.072×10−4 mol
Because MnO4− : Fe2+ = 1 : 5,
n(Fe2+) = 3.072×10−4 × 5 = 1.536×10−3 mol
The Fe2+ concentration in 25.00 mL of sample is,
C(Fe2+) = 1.536×10−3 ÷ 0.02500 = 0.06144 mol/L
Therefore, the Fe2+ concentration is determined to be 0.06144 mol/L.
Example of Expressing the Amount of Fe2+ in mg/L
Using 55.85 g/mol as the molar mass of Fe, the Fe2+ concentration of 0.06144 mol/L is converted to mg/L.
Fe2+ amount = 0.06144 mol/L × 55.85 g/mol = 3.431 g/L
3.431 g/L = 3431 mg/L
| Item | Value | Calculation |
|---|---|---|
| Fe2+ concentration | 0.06144 mol/L | Calculated from the titration volume |
| Molar mass of Fe | 55.85 g/mol | Used for conversion |
| Mass concentration | 3.431 g/L | 0.06144 × 55.85 |
| Conversion to mg/L | 3431 mg/L | Multiply g/L by 1000 |
Example of Standardizing KMnO4 Standard Solution Using Oxalic Acid
Because the concentration of KMnO4 standard solution may change during storage, it may be standardized using a substance that is close to a primary standard. Here, an example of standardization using oxalic acid is shown.
2MnO4− + 5C2O42− + 16H+ → 2Mn2+ + 10CO2 + 8H2O
| Item | Value | Calculation / Meaning |
|---|---|---|
| Oxalic acid standard-solution concentration | 0.0500 mol/L | Known concentration |
| Oxalic acid standard-solution volume | 25.00 mL | 0.02500 L |
| Amount of oxalic acid | 1.250×10−3 mol | 0.0500 × 0.02500 |
| Reaction ratio | MnO4− : C2O42− = 2 : 5 | Used to determine the amount of KMnO4 |
| Required amount of MnO4− | 5.00×10−4 mol | 1.250×10−3 × 2/5 |
| KMnO4 titration volume | 25.20 mL | Measured value during standardization |
| Actual KMnO4 concentration | 0.01984 mol/L | 5.00×10−4 ÷ 0.02520 |
More accurate quantitative analysis can be performed by using the concentration of 0.01984 mol/L determined by standardization rather than the nominal concentration of 0.0200 mol/L.
Example of Correcting Fe2+ Concentration Using the Standardized Concentration
Consider the case where the actual KMnO4 concentration is 0.01984 mol/L and the titration volume is 15.36 mL.
| Calculation Condition | KMnO4 Concentration | Fe2+ Concentration | Trend in the Result |
|---|---|---|---|
| Using the nominal concentration | 0.0200 mol/L | 0.06144 mol/L | Slightly overestimated |
| Using the standardized concentration | 0.01984 mol/L | 0.06095 mol/L | After correction |
If the standardized concentration is lower than the nominal value, using the nominal concentration in the calculation causes the sample concentration to be overestimated.
Example of Quantifying an Oxalic Acid Sample
This is an example in which 25.00 mL of an unknown sample containing oxalic acid is titrated with 0.0200 mol/L KMnO4.
| Trial | Sample Volume | KMnO4 Concentration | Titration Volume | Endpoint Assessment |
|---|---|---|---|---|
| 1st | 25.00 mL | 0.0200 mol/L | 10.08 mL | Good |
| 2nd | 25.00 mL | 0.0200 mol/L | 10.12 mL | Good |
| 3rd | 25.00 mL | 0.0200 mol/L | 10.10 mL | Good |
| Average | 25.00 mL | 0.0200 mol/L | 10.10 mL | Representative value |
Because the reaction ratio is MnO4− : C2O42− = 2 : 5, the amount of oxalate ions is 2.5 times the amount of MnO4−.
n(MnO4−) = 0.0200 × 0.01010 = 2.020×10−4 mol
n(C2O42−) = 2.020×10−4 × 5/2 = 5.050×10−4 mol
C(C2O42−) = 5.050×10−4 ÷ 0.02500 = 0.02020 mol/L
Example of Endpoint Color Changes
In KMnO4 titration, the color of the titrant itself is used to determine the endpoint.
| Titration Stage | State in the Solution | Observed Color | Assessment |
|---|---|---|---|
| Before titration | A large amount of reducing agent is present | The KMnO4 color disappears immediately | Before the endpoint |
| During titration | Added MnO4− is consumed by the reaction | A temporary purple color appears and disappears | Reducing agent still remains |
| Near the endpoint | The reducing agent is almost consumed | A pale reddish-purple color begins to remain | Near the endpoint |
| Endpoint | KMnO4 is slightly in excess | A pale reddish-purple color remains for about 30 seconds | Titration complete |
| Over-titration | KMnO4 is in excess | The reddish-purple color is dark | Titration volume is too large |
The endpoint color is not a dark purple but only a slight persistent pale reddish-purple color. Attempting to make the color more clearly visible increases the likelihood of over-titration.
Example of Insufficient Acidic Conditions
Acidic conditions are important in permanganate titration. If the amount of acid is insufficient, MnO4− may not be reduced all the way to Mn2+, and a brown precipitate such as MnO2 may form.
| Condition | Observation | Effect on the Reaction | Direction of Discussion |
|---|---|---|---|
| Sufficiently acidic conditions | The endpoint is pale reddish purple | Reduced to Mn2+ | Ideal titration |
| Slightly insufficient acid | Brown turbidity appears | Possible MnO2 formation | The reaction equation changes |
| Insufficient acid | The endpoint becomes unclear | KMnO4 consumption shifts | Reliability of quantitative results decreases |
| Excessively acidic | Usually few major problems, depending on conditions | Attention to side reactions and sample stability | Check the effect on the sample component |
If a brown precipitate forms, the redox reaction may not be proceeding as expected and the reliability of the titration result decreases.
Example When Heating Is Required
The reaction between oxalic acid and KMnO4 is slow at low temperatures, and the endpoint may be difficult to identify. Moderate heating makes the reaction proceed more readily.
| Temperature Condition | Reaction Progress | Endpoint | Direction of Discussion |
|---|---|---|---|
| Room temperature | Reaction is slow | It takes time for the color to disappear | Over-titration is likely |
| Moderate heating | Reaction becomes faster | Endpoint is relatively clear | Easier to quantify |
| Excessive heating | Attention required for changes in sample or solution | Errors may occur | Keep the temperature condition constant |
If the reaction is slow, the KMnO4 color may not disappear immediately, so care must be taken not to mistake this for the endpoint.
Example of Blank Correction
If the acid, water, or reagents contain substances that consume KMnO4, correction is performed using a blank test.
| Measurement | KMnO4 Titration Volume | Concept of Correction |
|---|---|---|
| Sample titration | 15.36 mL | Consumption by sample + reagents |
| Blank test | 0.06 mL | Consumption derived from acid, water, and reagents |
| Corrected titration volume | 15.30 mL | 15.36 − 0.06 |
Even if the blank-test volume is small, its relative effect may become large for low-concentration samples.
Effect of Over-Titration
If too much KMnO4 is added beyond the endpoint, the amount of reducing agent in the sample is estimated to be larger than it actually is.
| Condition | KMnO4 Titration Volume | Calculated Fe2+ Concentration | Trend in the Result |
|---|---|---|---|
| Appropriate endpoint | 15.36 mL | 0.06144 mol/L | Reference |
| 0.05 mL over-titration | 15.41 mL | 0.06164 mol/L | Slightly high |
| 0.10 mL over-titration | 15.46 mL | 0.06184 mol/L | High |
| 0.20 mL over-titration | 15.56 mL | 0.06224 mol/L | Clear overestimation |
In KMnO4 titration, attempting to make the endpoint color dark makes over-titration more likely, so it is important to use the point at which only a pale reddish-purple color remains as the endpoint.
Effect of Air Oxidation of the Reducing Agent
Reducing agents such as Fe2+ may be oxidized by oxygen in the air during storage or handling. If Fe2+ is oxidized to Fe3+, the amount of Fe2+ titrated by KMnO4 decreases.
| Sample Condition | Fe2+ Remaining | KMnO4 Titration Volume | Calculated Fe2+ Concentration | Direction of Discussion |
|---|---|---|---|---|
| Immediately after preparation | 100% | 15.36 mL | 0.06144 mol/L | Reference |
| Left for a while | 96% | 14.75 mL | 0.05900 mol/L | Partially oxidized |
| Left for a long time | 90% | 13.82 mL | 0.05528 mol/L | Underestimated |
If the reducing agent in the sample is oxidized before titration, the concentration determined by titration becomes lower than the concentration at the time of preparation.
Changes in the Concentration of KMnO4 Standard Solution During Storage
The concentration of KMnO4 solution may change during storage due to decomposition or reactions with impurities.
| Storage Condition | Apparent Concentration | Concentration Determined by Standardization | Direction of Discussion |
|---|---|---|---|
| Immediately after preparation / stored protected from light | 0.0200 mol/L | 0.01998 mol/L | Almost equal to the nominal value |
| Stored for several days | 0.0200 mol/L | 0.01984 mol/L | Slight decrease |
| Long-term storage | 0.0200 mol/L | 0.01940 mol/L | Standardization required |
| Precipitate observed | 0.0200 mol/L | Unstable | Attention to MnO2 formation and similar changes |
Because an inaccurate KMnO4 standard-solution concentration affects all concentration calculations, it is important to use the standardized value.
Main Sources of Error
| Source of Error | Effect on the Measurement | Trend in the Result | Improvement / Check |
|---|---|---|---|
| Deviation in KMnO4 standard-solution concentration | Affects all titration calculations | Concentration is overestimated or underestimated | Use the standardized concentration |
| Over-titration at the endpoint | Titration volume becomes larger | Amount of reducing agent is overestimated | Use a pale reddish-purple color as the endpoint |
| Insufficient acidic conditions | The reaction equation changes | Inaccurate due to MnO2 formation and similar effects | Maintain sufficiently acidic conditions |
| Air oxidation of the reducing agent | The sample changes before titration | Fe2+ amount is underestimated | Measure immediately after preparation |
| Slow reaction rate | It takes time for the color to disappear | The endpoint may be misidentified | Use appropriate temperature and stirring |
| Error in reading titration volume | The measured volume shifts | Directly affects concentration calculations | Read the burette correctly |
| No blank correction | Includes consumption caused by reagents | Sample concentration is overestimated | Perform a blank test |
Example of How to Write the Results
A 25.00 mL sample containing Fe2+ was titrated with 0.0200 mol/L KMnO4 standard solution. The average of three satisfactory titration volumes was 15.36 mL. The reaction equation under acidic conditions is MnO4− + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O, and the reaction ratio between MnO4− and Fe2+ is 1:5.
The amount of KMnO4 was 0.0200 × 0.01536 = 3.072×10−4 mol. Therefore, the amount of Fe2+ was five times this value, 1.536×10−3 mol. Dividing by the sample volume of 25.00 mL gave an Fe2+ concentration of 0.06144 mol/L.
The endpoint was taken as the point at which a pale reddish-purple color from KMnO4 remained for a while. Because KMnO4 itself has a purple color, the endpoint can be determined without using a separate indicator. However, if too much is added in an attempt to make the color more clearly visible, the titration volume becomes larger and the Fe2+ concentration is overestimated.
Points to Connect to the Discussion
In a discussion of permanganate titration, it is important not only to calculate concentration from the titration volume, but also to explain the redox reaction equation, number of electrons, acidic conditions, endpoint determination, standardization, and oxidation of the sample in relation to one another.
- Can you explain the half-reaction of MnO4− under acidic conditions?
- Do you understand that MnO4− is an oxidizing agent that accepts five electrons?
- Can you use the 1:5 reaction ratio with Fe2+ in the calculation?
- Can you explain that the reaction ratio with oxalic acid is 2:5?
- Can you explain why standardization of the KMnO4 standard solution is necessary?
- Can you discuss the relationship between the pale reddish-purple endpoint and over-titration?
- Can you explain that insufficient acidic conditions may change the reaction, such as by producing MnO2?
- Can you discuss that air oxidation of reducing agents such as Fe2+ lowers the measured value?
- Can you explain blank correction, burette reading, reaction rate, and temperature conditions as sources of error?
Example Discussion Text
In this experiment, the Fe2+ concentration was determined by redox titration using a potassium permanganate standard solution. Under acidic conditions, MnO4− accepts five electrons and is reduced to Mn2+. On the other hand, Fe2+ loses one electron and is oxidized to Fe3+. Therefore, 5 mol of Fe2+ react with 1 mol of MnO4−.
The average titration volume was 15.36 mL, and when 0.0200 mol/L KMnO4 was used, the amount of MnO4− was 3.072×10−4 mol. From the 1:5 reaction ratio, the amount of Fe2+ was 1.536×10−3 mol, and the Fe2+ concentration in the sample was determined to be 0.06144 mol/L. In this calculation, it is important to use the coefficients of the redox reaction equation correctly.
In KMnO4 titration, the purple color of the permanganate ion itself can be used to determine the endpoint. During titration, the added MnO4− is immediately consumed by the reducing agent, so the purple color disappears, but after the endpoint is exceeded, a pale reddish-purple color from slightly excessive MnO4− remains. Attempting to make the endpoint color dark results in over-titration and overestimation of the amount of reducing agent, so the endpoint must be taken as the point at which a pale color remains for a certain period of time.
Acidic conditions are also important. If the amount of acid is insufficient, MnO4− may not be reduced all the way to Mn2+, and a brown MnO2 precipitate may form. In this case, a reaction different from the ideal reaction equation proceeds, reducing the reliability of the concentration determined from the titration volume. Therefore, maintaining sufficiently acidic conditions is necessary for accurate titration.
Possible sources of error include changes in the concentration of the KMnO4 standard solution, over-titration at the endpoint, insufficient acidic conditions, air oxidation of the reducing agent, slow reaction rate, and errors in reading the titration volume. In particular, Fe2+ may be oxidized to Fe3+ by oxygen in the air, and if oxidation proceeds before titration, the Fe2+ concentration is underestimated. In addition, because the concentration of KMnO4 standard solution may change during storage, it is desirable to use the actual concentration determined by standardization.
Summary
In potassium permanganate titration, MnO4− acts as a strong oxidizing agent that accepts five electrons under acidic conditions. A reaction ratio of 1:5 is used for the determination of Fe2+, while a reaction ratio of 2:5 is used for the determination of oxalic acid.
This reference example covered Fe2+ titration, standardization of KMnO4 using oxalic acid, determination of an oxalic acid sample, the pale red endpoint color, insufficient acidic conditions, heating, blank correction, over-titration, air oxidation of the reducing agent, and changes in standard-solution concentration during storage. In a report, it is useful to discuss the reaction equation, number of electrons, reaction ratio, endpoint determination, and sources of error in relation to one another.
How to Think About Reaction Equations and Quantitative Relationships
In redox titration, it is important to correctly use the coefficients of the reaction equation. Unlike acid-base titration, the substances do not always react in a simple 1:1 ratio, and the coefficients are determined by the number of electrons exchanged between the oxidizing agent and reducing agent.
For example, under acidic conditions, permanganate ions accept electrons and are reduced to manganese(II) ions. In the reaction with the reducing agent in the sample, the amount of substance is calculated according to the coefficients of the reaction equation.
The reaction equation and coefficients differ depending on the type of reducing agent used in the experiment. In the report, check the reaction equation given in the laboratory manual and calculate the amount of substance using the coefficients.
Example Discussion:
In redox titration, the reaction ratio between the oxidizing agent and reducing agent must be determined from the reaction equation. Permanganate ions act as a strong oxidizing agent under acidic conditions and oxidize reducing substances in the sample. Therefore, when determining the sample concentration from the titration volume, it is necessary not only to use the amount of titrant but also to account for the coefficients of the reaction equation.
Error in Endpoint Determination
In potassium permanganate titration, the endpoint is taken as the point at which a pale reddish-purple color remains. However, because the color change is judged visually, individual differences occur in endpoint determination. In particular, because a faint persistent color is used as the endpoint, the judgment may be too early or too late.
If too much potassium permanganate solution is added beyond the endpoint, the titration volume becomes larger than the actual value. As a result, the amount of reducing substance in the sample may be overestimated.
Example Discussion:
Endpoint determination is performed by visually confirming the persistence of a pale reddish-purple color, so individual differences may occur in judgment. If potassium permanganate solution is added beyond the endpoint, the titration volume becomes larger than the amount actually required. As a result, the amount of reducing substance in the sample is overestimated and the calculated concentration may become higher than the actual value.
Discussion of Variation in Titration Volumes
If titration volumes vary among multiple measurements, consider endpoint determination, addition rate, burette-reading errors, sample volume, and the mixing condition of the solution. In redox titration, reaction rate and temperature may also affect the results.
| Cause | What Happens | Effect on the Result |
|---|---|---|
| Differences in endpoint determination | Judgment of the remaining reddish-purple color differs | Titration volumes vary |
| Addition rate is too fast | The endpoint is easily exceeded | Titration volume becomes larger |
| Burette-reading error | Initial or endpoint values are misread | Error occurs in the titration volume |
| Insufficient stirring | The reaction proceeds locally | Judgment of color disappearance or persistence becomes unstable |
| Differences in temperature conditions | Reaction rate changes | May affect endpoint determination |
Example Discussion:
One possible cause of variation in the titration volumes is differences in endpoint determination. In potassium permanganate titration, the endpoint is taken as the point at which a pale reddish-purple color remains, but judgment of the color intensity and persistence time may differ between individuals. In addition, if the addition rate near the endpoint was too fast, excess titrant could easily have been added and the titration volume may have become larger.
Error Due to the Concentration of the Potassium Permanganate Solution
If the concentration of the potassium permanganate solution used as the titrant is inaccurate, an error also occurs in the determined sample concentration. The concentration of potassium permanganate solution may change depending on storage conditions and preparation conditions. Therefore, standardization may be performed as necessary to determine the accurate concentration before use.
Example Discussion:
One possible reason the determined sample concentration differed from the theoretical value is that the concentration of the potassium permanganate solution was inaccurate. If the titrant concentration differed from the actual value, the sample concentration calculated from the titration volume would be incorrect even if the titration volume itself was accurate. Therefore, it is important to confirm the concentration of potassium permanganate solution by standardization before use.
Error When Acidic Conditions Are Insufficient
In potassium permanganate titration, the reaction may not proceed as expected if the acidic conditions are insufficient. For example, if a brown precipitate such as manganese dioxide forms, the reaction ratio differs from that when permanganate ions are reduced to manganese(II) ions. As a result, an error occurs in the titration calculation.
Example Discussion:
If a brown precipitate or turbidity was observed during titration, the acidic conditions may have been insufficient and the permanganate ions may not have been reduced all the way to manganese(II) ions as expected. In this case, the quantitative relationship of the reaction equation changes, causing an error in the sample concentration determined from the titration volume. Therefore, it is important to maintain appropriate acidic conditions in potassium permanganate titration.
Error Due to the Oxidation or Reduction State of the Sample
If the reducing substance in the sample has already been partially oxidized by oxygen in the air or other factors, the amount of reducing substance that reacts during titration becomes smaller than originally present. As a result, less potassium permanganate is consumed and the amount of reducing substance in the sample may be underestimated.
Example Discussion:
If the reducing substance in the sample had been partially oxidized by oxygen in the air before measurement, the amount capable of reacting with potassium permanganate would decrease. Therefore, the volume of potassium permanganate solution required for titration would become smaller and the concentration of the reducing substance in the sample could be underestimated.
Error Due to Insufficient Stirring
In potassium permanganate titration, the added permanganate ions immediately react with reducing substances in the sample solution. If stirring is insufficient, permanganate ions may become locally excessive and color may temporarily remain. As a result, the endpoint may be judged earlier than the actual endpoint.
Example Discussion:
If stirring during titration was insufficient, the added potassium permanganate solution may not have mixed uniformly throughout the sample solution and reddish-purple color may have remained locally. Therefore, the endpoint may have been incorrectly judged even though reducing substances still remained. In this case, the titration volume becomes smaller than the actual value and the calculated sample concentration may be underestimated.
Effect of Temperature
In some redox reactions, the reaction rate changes greatly with temperature. Under conditions where the reaction is slow, color may remain immediately after adding potassium permanganate but disappear after some time. In this case, if the endpoint is judged too quickly, the titration volume may become smaller.
Depending on the experiment, a specific temperature condition may be designated to facilitate the reaction. In the report, the effects of reaction rate and temperature conditions on endpoint determination can be discussed.
Example Discussion:
Under conditions where the reaction rate is slow, the reddish-purple color may remain immediately after potassium permanganate is added but disappear after the reaction proceeds for a while. If this state is judged as the endpoint, the titration may be stopped with a smaller titration volume than the actual value. Therefore, temperature conditions and reaction rate are considered to affect endpoint determination.
Burette-Reading Error
In redox titration, burette-reading errors occur in the same way as in acid-base titration. If the meniscus is not read correctly when the initial and endpoint values are recorded, the titration volume becomes larger or smaller than the actual value.
Example Discussion:
If the line of sight did not align with the burette scale when reading it, an error may have occurred in the initial or endpoint value. Because the titration volume is determined from the difference between the initial and endpoint values, this reading error directly affects the sample-concentration calculation. Therefore, the burette scale must be read at the same height as the meniscus.
How to Write About Excluding a Titration Value
When titration is performed multiple times, a value that clearly exceeded the endpoint or differed greatly from the other values may be excluded. However, when a value is excluded, the reason must always be based on observed facts.
Example Discussion:
In the first titration, the solution showed a dark reddish-purple color at the endpoint and the titration volume was larger than the other measured values. This is considered to have occurred because excess potassium permanganate solution was added beyond the endpoint. Therefore, the first value was excluded from the calculation of the average titration volume, and the second and third values, for which endpoint determination was appropriate, were used.
When the Result Can Be Considered Good
A good result in potassium permanganate titration is indicated when multiple titration volumes are close to one another, the pale reddish-purple color at the endpoint can be clearly confirmed, and the determined concentration is close to the theoretical or labeled value. It is also important that no precipitate or turbidity occurs during titration.
Example Discussion:
The titration volumes from multiple trials were close to one another, and no large variation was observed. This suggests that the reproducibility of endpoint determination was relatively high. In addition, because a persistent pale reddish-purple color was observed at the endpoint and no abnormalities such as brown precipitate were seen, the redox reaction with potassium permanganate is considered to have proceeded approximately appropriately.
Example Discussion When the Experiment Did Not Go Well
If the titration values varied greatly, the endpoint color was difficult to judge, a brown precipitate formed, or the determined concentration differed greatly from the theoretical value, check endpoint determination, acidic conditions, the standard solution, reaction rate, and the condition of the sample.
Example Discussion:
One possible cause of variation in the titration volumes is the difficulty of endpoint determination. In potassium permanganate titration, the endpoint is taken as the point at which a pale reddish-purple color remains, but judgment of the color intensity and persistence time may differ between individuals. In addition, if stirring was insufficient, reddish-purple color may have remained locally and the endpoint may have been judged earlier than the actual point. These factors are considered to have caused errors in the determined concentration.
How to Write Points for Improvement
In a discussion of redox titration, including points for improvement as well as sources of error makes the report easier to organize. It is important to write improvements that correspond to the actual sources of error considered.
Methods for Improving Endpoint Determination
- Add the potassium permanganate solution one drop at a time near the endpoint
- Use the point at which a pale reddish-purple color persists for a certain period of time as the endpoint
- Do not continue adding titrant until a dark reddish-purple color appears
- Titrate while thoroughly stirring the solution
- Perform multiple titrations to confirm reproducibility
Methods for Improving Reaction Conditions
- Maintain the acidic conditions specified in the laboratory manual
- Observe whether any precipitate or turbidity appears
- Maintain the required temperature conditions
- Confirm the concentration of the potassium permanganate solution by standardization
- Prepare and store the sample solution appropriately before measurement
Example of How to Write Points for Improvement:
To reduce variation in titration volumes, the potassium permanganate solution should be added one drop at a time near the endpoint and the point at which a pale reddish-purple color remains for a certain period of time should be carefully confirmed. In addition, to avoid local persistence of color caused by insufficient stirring, it is important to thoroughly mix the solution during titration. Furthermore, appropriate acidic conditions must be maintained so that the reaction proceeds as expected.
Difference Between a Superficial Discussion and a Good Discussion
In a discussion of redox titration, simply writing that “the endpoint was difficult to determine” or “there was an error in the titration volume” results in a superficial discussion. A persuasive discussion can be produced by explaining which condition affected which value and in which direction.
| Superficial Discussion | Good Discussion |
|---|---|
| The endpoint was difficult to determine. | In potassium permanganate titration, a pale reddish-purple color is used as the endpoint, so individual differences occur in judging the color intensity and persistence time. If titrant is added beyond the endpoint, the titration volume becomes larger than the actual value and the amount of reducing substance in the sample may be overestimated. |
| The acidic conditions were poor. | If the acidic conditions were insufficient, permanganate ions may not be reduced all the way to manganese(II) ions as expected, and precipitates such as manganese dioxide may form. In this case, the quantitative relationship of the reaction changes and an error occurs in the titration calculation. |
| The values varied. | Variation in the titration volumes may have been caused by differences in endpoint determination, addition rate, insufficient stirring, and burette-reading errors. In particular, if the addition rate near the endpoint is too fast, excess addition may cause the titration volume to become larger. |
Examples of Expressions That Can Be Used in Reports
The following expressions can be used when writing the results and discussion of redox titration and potassium permanganate titration. Adjust the necessary parts according to your own experimental results.
- Because permanganate ions show a reddish-purple color, they can be used as a self-indicator for endpoint determination.
- While reducing substances remain in the sample, the added permanganate ions are reduced and the reddish-purple color disappears.
- The point at which a pale reddish-purple color persisted for a certain period of time was judged to be the endpoint.
- If titrant is added beyond the endpoint, the titration volume becomes larger than the actual value and the sample concentration may be overestimated.
- If the acidic conditions were insufficient, a reaction different from the assumed reaction may occur and cause an error in the titration calculation.
- If a brown precipitate was observed, manganese dioxide may have formed.
- Possible causes of variation in the titration volume include differences in endpoint determination and the effect of the addition rate.
- If stirring was insufficient, permanganate ions may have become locally excessive and the endpoint may have been judged too early.
- If the concentration of the potassium permanganate solution was inaccurate, a systematic error would also occur in the determined sample concentration.
- The quantitative relationship between the titrant and sample must be determined by accounting for the coefficients of the reaction equation.
Points to Check When Discussing Redox Titration
Checking the following points before writing the report makes the discussion easier to write.
- Can you explain why potassium permanganate acts as a self-indicator?
- How did you judge the pale reddish-purple endpoint?
- Did you avoid adding titrant beyond the endpoint?
- Was there variation in the titration volumes?
- If a measured value was excluded, did you state the reason?
- Were the acidic conditions appropriate?
- Was any brown precipitate or turbidity observed?
- Was the concentration of the potassium permanganate solution accurate?
- Did you correctly calculate using the coefficients of the reaction equation?
- Did you consider the effects of insufficient stirring and addition rate?
- Did temperature and reaction rate affect endpoint determination?
- Did you explain whether each error would make the sample concentration larger or smaller?
Summary
Redox titration is an analytical method used to determine the amount or concentration of a substance in a sample by utilizing electron transfer between an oxidizing agent and a reducing agent. A major characteristic of potassium permanganate titration is that permanganate ions themselves have a reddish-purple color and can therefore act as a self-indicator for endpoint determination.
In a discussion of potassium permanganate titration, it is important to consider determination of the pale reddish-purple endpoint, the need for acidic conditions, coefficients in the reaction equation, variation in titration volume, concentration of the standard solution, and the effects of stirring and addition rate. In particular, if the acidic conditions are insufficient, a reaction different from the assumed reaction may occur and cause errors in the titration calculation.
In a report, do not simply write that “the titration volume was different.” Specifically explain how endpoint determination and reaction conditions affected the measured values. More reliable results can be obtained by carefully adding titrant near the endpoint, thoroughly stirring the solution, and performing multiple measurements.
