The F- And D- Block Elements - Preparation Of Dichromate

Slide 1: The f- and d- block elements - Preparation of Dichromate

The f-block elements, also known as inner transition metals, are located in the 6th and 7th periods of the periodic table.
The d-block elements, also known as transition metals, are located in the 4th and 5th periods of the periodic table.
Dichromate is an important compound that is widely used in various industrial processes.
The preparation of dichromate involves the use of transition metals, specifically chromium.
Let’s explore the process of preparing dichromate in more detail.

Slide 2: Preparation of Dichromate - Step 1

The first step in the preparation of dichromate involves the conversion of chromium(III) ion to chromium(VI) ion.
This conversion is achieved by oxidizing chromium(III) ion using a strong oxidizing agent.
Potassium permanganate (KMnO4) is commonly used as the oxidizing agent.
The reaction can be represented as follows: Cr3+ + MnO4- → CrO42- + Mn2+
This reaction takes place in an acidic medium.
The purple color of the potassium permanganate solution fades during the reaction.

Slide 3: Preparation of Dichromate - Step 2

The second step in the preparation of dichromate involves the oxidation of chromate ion to dichromate ion.
This oxidation is also achieved using a strong oxidizing agent.
Sodium or potassium dichromate (Na2Cr2O7 or K2Cr2O7) is commonly used as the oxidizing agent.
The reaction can be represented as follows:

2 CrO42- + 2 H+ → Cr2O72- + H2O

This reaction also takes place in an acidic medium.
The change in color during the reaction indicates the formation of dichromate ion.

Slide 4: Preparation of Dichromate - Overall Reaction

By combining the two steps, the overall reaction for the preparation of dichromate can be represented as follows: Cr3+ + MnO4- + 4 H+ → Cr2O72- + Mn2+ + 2 H2O
Potassium permanganate is used to oxidize chromium(III) to chromium(VI).
Sodium or potassium dichromate is used to further oxidize chromate ion to dichromate ion.
Both reactions occur in an acidic medium.

Slide 5: Example: Preparation of Dichromate

Let’s consider an example to make the process clear.
We start with an aqueous solution containing chromium(III) ion (Cr3+).
To this solution, we add an excess of potassium permanganate (KMnO4).
The reaction takes place in an acidic medium.
As a result, the solution turns from purple to green, indicating the formation of chromium(VI) ion (CrO42-).

Slide 6: Example: Preparation of Dichromate (Continued)

In the second step, we need to further oxidize the chromate ion (CrO42-) to dichromate ion (Cr2O72-).
For this, we add sodium dichromate (Na2Cr2O7) to the solution.
The reaction takes place in an acidic medium.
The color of the solution changes from green to orange, indicating the formation of dichromate ion.

Slide 7: Equations Involved

Oxidation of chromium(III) using potassium permanganate: Cr3+ + MnO4- → CrO42- + Mn2+ (in acidic medium)

Oxidation of chromate ion to dichromate ion using sodium or potassium dichromate: 2 CrO42- + 2 H+ → Cr2O72- + H2O (in acidic medium)

Overall reaction for the preparation of dichromate: Cr3+ + MnO4- + 4 H+ → Cr2O72- + Mn2+ + 2 H2O

Slide 8: Significance of Dichromate

Dichromate has various important applications in different industries.
It is used in the production of various dyes and pigments.
It is a key component in the manufacturing of chrome-plated metal products.
Dichromate is also used in the tanning of leather.
In addition, it plays a crucial role in laboratory experiments and chemical analysis.

Slide 9: Precautions and Safety Measures

It is important to handle dichromate compounds with caution.
They are toxic and can be harmful if ingested, inhaled, or comes in contact with the skin or eyes.
Protective gear, such as gloves, goggles, and lab coats, should be worn when working with dichromate.
Proper ventilation should be ensured to avoid exposure to harmful fumes.
Waste containing dichromate should be disposed of according to applicable regulations.

Slide 10: Summary

The preparation of dichromate involves the oxidation of chromium(III) to chromium(VI) using potassium permanganate.
The chromate ion is then further oxidized to dichromate ion using sodium or potassium dichromate.
Both reactions occur in an acidic medium.
Dichromate has significant applications in industries like dye production, chrome-plating, and leather tanning.
Safety precautions should be taken when handling dichromate compounds.

Properties of Dichromate

Dichromate has an intense orange color.
It is highly soluble in water.
It is a powerful oxidizing agent.
Dichromate compounds are generally stable and have high melting points.
Solutions of dichromate are acidic in nature.

Oxidizing Power of Dichromate

The ability of dichromate to act as an oxidizing agent is due to the presence of positive charges on the chromium atoms.
The highest oxidation state of chromium (+6) allows it to gain electrons from other species.
Examples of its oxidation reactions include the oxidation of alcohols to carbonyl compounds and the conversion of sulfite ions to sulfate ions.

Application in the Laboratory - Chromic Acid Test

The chromic acid test is commonly used in the laboratory to test for the presence of primary or secondary alcohols.
In this test, the alcohol is oxidized by chromic acid (H2CrO4) to form an aldehyde or ketone.
A color change from orange to green is observed, indicating the formation of the respective carbonyl compound.

Application in the Industry - Leather Tanning

Dichromate compounds are widely used in the leather industry for the process of tanning.
Chromium salts, such as chromium sulfate (Cr2(SO4)3), are applied to the animal hide, forming a complex with collagen proteins.
This complex enhances the strength, durability, and water resistance of the leather.

Environmental Concerns

Dichromate compounds are toxic to both humans and the environment.
They are classified as potential carcinogens.
Proper disposal of waste containing dichromate must be done to prevent environmental contamination.
Regulations are in place to control and limit the use of dichromate in various industries.

Redox Reactions Involving Dichromate

Dichromate participates in a wide range of redox reactions.
It can accept electrons and be reduced to chromium(III) ions.
Examples include the reduction of dichromate by sulfur dioxide (SO2) to form chromium(III) sulfate and sulfite ions.

Half-Reactions and Balancing

In redox reactions, it is useful to write half-reactions for both the oxidation and reduction processes.
By balancing the number of electrons in both half-reactions, we can determine the overall balanced equation.
For example, in the reduction of dichromate by sulfite ions, the balanced half-reactions are: Cr2O7^2- + 14H+ + 6e^- → 2Cr^3+ + 7H2O SO3^2- + H2O → SO4^2- + 2H+ + 2e

Spectator Ions and Net Ionic Equations

In many redox reactions, certain ions do not participate directly and remain unchanged.
These ions are known as spectator ions.
In determining the net ionic equation, the spectator ions are omitted, and only the species involved in the redox process are included.
This simplifies the equation and focuses on the actual redox reaction.

Equivalence Point in Dichromate Titrations

Dichromate is often used as a titrant in redox titrations to determine the concentration of an analyte.
The equivalence point is reached when the moles of dichromate added are chemically equivalent to the moles of the analyte.
In the presence of an indicator, a color change occurs to signify the completion of the titration.
Examples of dichromate titrations include the determination of iron(II) ions and hydrogen peroxide.

Example: Determination of Iron(II) by Dichromate Titrations

Iron(II) ions (Fe^2+) can be determined by titration with dichromate.
The reaction can be represented as follows: 6 Fe^2+ + Cr2O7^2- + 14 H+ → 6 Fe^3+ + 2 Cr^3+ + 7 H2O
Potassium dichromate is used as the titrant, and a suitable indicator, such as diphenylamine, is employed.
The end point is reached when a blue-green color is observed, indicating the excess of dichromate.

Slide 21: Properties of Dichromate

Dichromate has an intense orange color.
It is highly soluble in water.
It is a powerful oxidizing agent.
Dichromate compounds are generally stable and have high melting points.
Solutions of dichromate are acidic in nature.

Slide 22: Oxidizing Power of Dichromate

The ability of dichromate to act as an oxidizing agent is due to the presence of positive charges on the chromium atoms.
The highest oxidation state of chromium (+6) allows it to gain electrons from other species.
Examples of its oxidation reactions include the oxidation of alcohols to carbonyl compounds and the conversion of sulfite ions to sulfate ions.

Slide 23: Application in the Laboratory - Chromic Acid Test

The chromic acid test is commonly used in the laboratory to test for the presence of primary or secondary alcohols.
In this test, the alcohol is oxidized by chromic acid (H2CrO4) to form an aldehyde or ketone.
A color change from orange to green is observed, indicating the formation of the respective carbonyl compound.

Slide 24: Application in the Industry - Leather Tanning

Dichromate compounds are widely used in the leather industry for the process of tanning.
Chromium salts, such as chromium sulfate (Cr2(SO4)3), are applied to the animal hide, forming a complex with collagen proteins.
This complex enhances the strength, durability, and water resistance of the leather.

Slide 25: Environmental Concerns

Dichromate compounds are toxic to both humans and the environment.
They are classified as potential carcinogens.
Proper disposal of waste containing dichromate must be done to prevent environmental contamination.
Regulations are in place to control and limit the use of dichromate in various industries.

Slide 26: Redox Reactions Involving Dichromate

Dichromate participates in a wide range of redox reactions.
It can accept electrons and be reduced to chromium(III) ions.
Examples include the reduction of dichromate by sulfur dioxide (SO2) to form chromium(III) sulfate and sulfite ions.

Slide 27: Half-Reactions and Balancing

In redox reactions, it is useful to write half-reactions for both the oxidation and reduction processes.
By balancing the number of electrons in both half-reactions, we can determine the overall balanced equation.
For example, in the reduction of dichromate by sulfite ions, the balanced half-reactions are:
1. Cr2O7^2- + 14H+ + 6e^- → 2Cr^3+ + 7H2O
2. SO3^2- + H2O → SO4^2- + 2H+ + 2e

Slide 28: Spectator Ions and Net Ionic Equations

In many redox reactions, certain ions do not participate directly and remain unchanged.
These ions are known as spectator ions.
In determining the net ionic equation, the spectator ions are omitted, and only the species involved in the redox process are included.
This simplifies the equation and focuses on the actual redox reaction.

Slide 29: Equivalence Point in Dichromate Titrations

Dichromate is often used as a titrant in redox titrations to determine the concentration of an analyte.
The equivalence point is reached when the moles of dichromate added are chemically equivalent to the moles of the analyte.
In the presence of an indicator, a color change occurs to signify the completion of the titration.
Examples of dichromate titrations include the determination of iron(II) ions and hydrogen peroxide.

Slide 30: Example: Determination of Iron(II) by Dichromate Titrations

Iron(II) ions (Fe^2+) can be determined by titration with dichromate.
The reaction can be represented as follows: 6 Fe^2+ + Cr2O7^2- + 14 H+ → 6 Fe^3+ + 2 Cr^3+ + 7 H2O
Potassium dichromate is used as the titrant, and a suitable indicator, such as diphenylamine, is employed.
The end point is reached when a blue-green color is observed, indicating the excess of dichromate.