Preparation and Properties of K2Cr2O7 and KMnO4 — Explained

Detailed Explanation

The preparation and properties of potassium dichromate ($K_2Cr_2O_7$) and potassium permanganate ($KMnO_4$) are fundamental topics in inorganic chemistry, particularly within the context of transition elements. These compounds exemplify the high oxidation states achievable by d-block elements and their profound utility as powerful oxidizing agents.

I. Potassium Dichromate ($K_2Cr_2O_7$)

A. Preparation of Potassium Dichromate

Potassium dichromate is primarily prepared from chromite ore, which has the chemical formula $FeCr_2O_4$. The process involves several key steps:

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Concentration and Roasting of Chromite Ore:

The finely powdered chromite ore ($FeCr_2O_4$) is mixed with sodium carbonate ($Na_2CO_3$) or potassium carbonate ($K_2CO_3$) and lime ($CaO$) and then roasted in a reverberatory furnace in the presence of excess air (oxygen) at high temperatures (around $1000-1200^circ C$).

The iron in the ore is oxidized to iron(III) oxide, and chromium is oxidized from +3 to +6, forming sodium chromate ($Na_2CrO_4$).

$$4FeCr_2O_4 (s) + 8Na_2CO_3 (s) + 7O_2 (g) \xrightarrow{Heat} 8Na_2CrO_4 (s) + 2Fe_2O_3 (s) + 8CO_2 (g)$$

The lime ($CaO$) is added to prevent the fusion of the mass and to make it porous, facilitating the reaction.

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Extraction and Acidification of Sodium Chromate:

The roasted mass is then extracted with water, which dissolves the soluble sodium chromate ($Na_2CrO_4$), leaving behind insoluble iron(III) oxide ($Fe_2O_3$). The yellow solution of sodium chromate is then filtered.

To convert sodium chromate to sodium dichromate ($Na_2Cr_2O_7$), the solution is acidified with concentrated sulfuric acid ($H_2SO_4$). Chromate ions ($CrO_4^{2-}$), which are yellow, exist in equilibrium with dichromate ions ($Cr_2O_7^{2-}$), which are orange.

The equilibrium shifts towards dichromate in acidic medium.

$$2Na_2CrO_4 (aq) + H_2SO_4 (aq) \rightarrow Na_2Cr_2O_7 (aq) + Na_2SO_4 (aq) + H_2O (l)$$

Alternatively, carbon dioxide can be passed through the solution to achieve acidification, though sulfuric acid is more common for industrial scale.

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Conversion to Potassium Dichromate:

Sodium dichromate is highly soluble in water. To obtain potassium dichromate, which is less soluble, the solution of sodium dichromate is treated with potassium chloride ($KCl$). Potassium dichromate crystallizes out upon cooling due to its lower solubility compared to sodium dichromate.

$$Na_2Cr_2O_7 (aq) + 2KCl (aq) \rightarrow K_2Cr_2O_7 (s) + 2NaCl (aq)$$

The orange crystals of $K_2Cr_2O_7$ are then separated by filtration and purified by recrystallization.

B. Properties of Potassium Dichromate

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Physical Properties:

* It is an orange-red crystalline solid. * It is soluble in water, forming an orange solution. * It has a melting point of $398^circ C$.

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Chemical Properties:

* Effect of Heat: On strong heating, $K_2Cr_2O_7$ decomposes to potassium chromate, chromium(III) oxide, and oxygen.

$$4K_2Cr_2O_7 (s) \xrightarrow{Heat} 4K_2CrO_4 (s) + 2Cr_2O_3 (s) + 3O_2 (g)$$

* Acid-Base Equilibrium (Chromate-Dichromate Interconversion): The chromate ($CrO_4^{2-}$, yellow) and dichromate ($Cr_2O_7^{2-}$, orange) ions are in equilibrium, which is pH-dependent.

$$2CrO_4^{2-} (aq) + 2H^+ (aq) \rightleftharpoons Cr_2O_7^{2-} (aq) + H_2O (l)$$

In acidic medium, the equilibrium shifts to the right, forming orange dichromate ions. In alkaline medium, the equilibrium shifts to the left, forming yellow chromate ions.

$$Cr_2O_7^{2-} (aq) + 2OH^- (aq) \rightleftharpoons 2CrO_4^{2-} (aq) + H_2O (l)$$

* Oxidizing Nature: Potassium dichromate is a powerful oxidizing agent, especially in acidic solutions. The $Cr_2O_7^{2-}$ ion is reduced to $Cr^{3+}$ (green) ions.

The standard electrode potential for this reduction is $+1.33,V$.

$$Cr_2O_7^{2-} (aq) + 14H^+ (aq) + 6e^- \rightarrow 2Cr^{3+} (aq) + 7H_2O (l)$$

This means one mole of dichromate ion accepts 6 electrons.

The equivalent weight of $K_2Cr_2O_7$ in acidic medium is $rac{Molar,Mass}{6}$. * Oxidation of Iodides: Oxidizes $I^-$ to $I_2$.

$$Cr_2O_7^{2-} + 14H^+ + 6I^- \rightarrow 2Cr^{3+} + 3I_2 + 7H_2O$$

* Oxidation of Iron(II) salts: Oxidizes $Fe^{2+}$ to $Fe^{3+}$.

$$Cr_2O_7^{2-} + 14H^+ + 6Fe^{2+} \rightarrow 2Cr^{3+} + 6Fe^{3+} + 7H_2O$$

* Oxidation of Sulfites: Oxidizes $SO_3^{2-}$ to $SO_4^{2-}$.

$$Cr_2O_7^{2-} + 8H^+ + 3SO_3^{2-} \rightarrow 2Cr^{3+} + 3SO_4^{2-} + 4H_2O$$

* Oxidation of Hydrogen Sulfide: Oxidizes $H_2S$ to $S$.

$$Cr_2O_7^{2-} + 8H^+ + 3H_2S \rightarrow 2Cr^{3+} + 3S + 7H_2O$$

* Chromyl Chloride Test: This is a characteristic test for chloride ions. When a mixture of a chloride salt, $K_2Cr_2O_7$, and concentrated $H_2SO_4$ is heated, deep red vapors of chromyl chloride ($CrO_2Cl_2$) are evolved.

These vapors, when passed through water, give a yellow solution which turns yellow-orange with $NaOH$ and gives a yellow precipitate with lead acetate ($Pb(CH_3COO)_2$).

C. Structure of Dichromate Ion ($Cr_2O_7^{2-}$)

In the dichromate ion, two $CrO_4$ tetrahedra share one corner, with a $Cr-O-Cr$ bond. The $Cr-O$ bond lengths are not all equal. The $Cr-O-Cr$ angle is approximately $126^circ$.

D. Uses of Potassium Dichromate

As an oxidizing agent in volumetric analysis (e.g., estimation of $Fe^{2+}$).
In leather tanning.
Preparation of chrome alum ($K_2SO_4 cdot Cr_2(SO_4)_3 cdot 24H_2O$).
As a primary standard in analytical chemistry.
In photography (hardening of gelatin).
In pigments (chrome yellow, chrome red).

II. Potassium Permanganate ($KMnO_4$)

A. Preparation of Potassium Permanganate

Potassium permanganate is prepared from pyrolusite ore ($MnO_2$). The process involves two main steps:

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Conversion of Pyrolusite to Potassium Manganate:

Finely powdered pyrolusite ore ($MnO_2$) is fused with potassium hydroxide ($KOH$) in the presence of an oxidizing agent like air (oxygen), potassium nitrate ($KNO_3$), or potassium chlorate ($KClO_3$). This reaction oxidizes manganese from +4 to +6, forming green potassium manganate ($K_2MnO_4$).

$$2MnO_2 (s) + 4KOH (s) + O_2 (g) \xrightarrow{Heat} 2K_2MnO_4 (s) + 2H_2O (l)$$

(Using $KNO_3$ as oxidizer: $MnO_2 + 2KOH + KNO_3 \rightarrow K_2MnO_4 + KNO_2 + H_2O$)

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Oxidation of Potassium Manganate to Potassium Permanganate:

The green potassium manganate ($K_2MnO_4$) is then oxidized to purple potassium permanganate ($KMnO_4$). This can be done by chemical oxidation or electrolytic oxidation. * Chemical Oxidation: This involves disproportionation of manganate ions in acidic or neutral medium.

Manganate ($MnO_4^{2-}$) is stable only in strongly alkaline solutions. In neutral or acidic solutions, it disproportionates into permanganate ($MnO_4^-$) and manganese dioxide ($MnO_2$).

$$3MnO_4^{2-} (aq) + 4H^+ (aq) \rightarrow 2MnO_4^- (aq) + MnO_2 (s) + 2H_2O (l)$$

This can be achieved by passing $CO_2$ or chlorine gas through the manganate solution.

$$2K_2MnO_4 (aq) + Cl_2 (g) \rightarrow 2KMnO_4 (aq) + 2KCl (aq)$$

$$3K_2MnO_4 (aq) + 2CO_2 (g) \rightarrow 2KMnO_4 (aq) + MnO_2 (s) + 2K_2CO_3 (aq)$$

* Electrolytic Oxidation: This is the preferred industrial method.

The green manganate solution is electrolyzed between iron electrodes. At the anode, manganate ions are oxidized to permanganate ions.

$$MnO_4^{2-} (aq) \xrightarrow{Anode} MnO_4^- (aq) + e^-$$

The purple $KMnO_4$ crystallizes out upon concentration.

B. Properties of Potassium Permanganate

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Physical Properties:

* It is a dark purple (almost black) crystalline solid. * It is moderately soluble in water, forming a deep purple solution. * It has a metallic luster.

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Chemical Properties:

* Effect of Heat: On heating, $KMnO_4$ decomposes to potassium manganate, manganese dioxide, and oxygen.

$$2KMnO_4 (s) \xrightarrow{Heat} K_2MnO_4 (s) + MnO_2 (s) + O_2 (g)$$

* Oxidizing Nature: Potassium permanganate is an extremely strong oxidizing agent.

Its oxidizing power and the reduction products vary significantly with the pH of the medium. * **In Acidic Medium (Strongly Acidic, e.g., with $H_2SO_4$):** The permanganate ion ($MnO_4^-$) is reduced to colorless manganese(II) ions ($Mn^{2+}$).

The change in oxidation state is from +7 to +2, meaning it accepts 5 electrons. The standard electrode potential is $+1.51,V$.

$$MnO_4^- (aq) + 8H^+ (aq) + 5e^- \rightarrow Mn^{2+} (aq) + 4H_2O (l)$$

The equivalent weight of $KMnO_4$ in acidic medium is $rac{Molar,Mass}{5}$.

* Oxidation of Iron(II) salts:

$$MnO_4^- + 8H^+ + 5Fe^{2+} \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O$$

* Oxidation of Oxalates:

$$2MnO_4^- + 16H^+ + 5C_2O_4^{2-} \rightarrow 2Mn^{2+} + 10CO_2 + 8H_2O$$

* Oxidation of Iodides:

$$2MnO_4^- + 16H^+ + 10I^- \rightarrow 2Mn^{2+} + 5I_2 + 8H_2O$$

* Oxidation of Hydrogen Sulfide:

$$2MnO_4^- + 6H^+ + 5H_2S \rightarrow 2Mn^{2+} + 5S + 8H_2O$$

* Oxidation of Nitrites:

$$2MnO_4^- + 6H^+ + 5NO_2^- \rightarrow 2Mn^{2+} + 5NO_3^- + 3H_2O$$

* In Neutral or Weakly Alkaline Medium: The permanganate ion ($MnO_4^-$) is reduced to brown manganese dioxide ($MnO_2$).

The change in oxidation state is from +7 to +4, meaning it accepts 3 electrons.

$$MnO_4^- (aq) + 2H_2O (l) + 3e^- \rightarrow MnO_2 (s) + 4OH^- (aq)$$

The equivalent weight of $KMnO_4$ in neutral/weakly alkaline medium is $rac{Molar,Mass}{3}$.

* Oxidation of Thiosulfate:

$$2MnO_4^- + H_2O + 3S_2O_3^{2-} \rightarrow 2MnO_2 + 3SO_4^{2-} + 2OH^-$$

* In Strongly Alkaline Medium: The permanganate ion ($MnO_4^-$) is reduced to green manganate ions ($MnO_4^{2-}$).

The change in oxidation state is from +7 to +6, meaning it accepts 1 electron.

$$MnO_4^- (aq) + e^- \rightarrow MnO_4^{2-} (aq)$$

The equivalent weight of $KMnO_4$ in strongly alkaline medium is $rac{Molar,Mass}{1}$.

This reaction is often used to prepare manganate from permanganate.

C. Structure of Permanganate Ion ($MnO_4^-$)

The permanganate ion has a tetrahedral geometry, with the manganese atom at the center and four oxygen atoms at the corners. All $Mn-O$ bond lengths are identical due to resonance.

D. Uses of Potassium Permanganate

As a powerful oxidizing agent in volumetric analysis (e.g., estimation of $Fe^{2+}$, oxalates, nitrites).
As a disinfectant and antiseptic (e.g., in well water to kill bacteria).
In organic synthesis as an oxidizing agent (e.g., oxidation of alkenes to diols, alcohols to carboxylic acids).
As a bleaching agent for wool, cotton, and other textiles.
In the purification of water.

III. Common Misconceptions and NEET-Specific Angle

pH Dependence: — A common mistake is to confuse the reduction products and electron changes for $KMnO_4$ in different media. Always remember: acidic ($Mn^{2+}$, 5e-), neutral/weakly alkaline ($MnO_2$, 3e-), strongly alkaline ($MnO_4^{2-}$, 1e-).
Color Changes: — Students often mix up the colors. $K_2Cr_2O_7$ is orange, $CrO_4^{2-}$ is yellow, $Cr^{3+}$ is green. $KMnO_4$ is purple, $K_2MnO_4$ is green, $Mn^{2+}$ is colorless, $MnO_2$ is brown.
Equivalent Weight: — The equivalent weight of $KMnO_4$ is not constant; it depends on the reaction medium due to varying n-factors. For $K_2Cr_2O_7$, it's generally 6 in acidic medium.
Balancing Redox Reactions: — NEET frequently tests the ability to balance redox reactions involving these compounds, often in different media. Mastering half-reaction method is crucial.
Preparation Steps: — Questions can be asked about the reagents used at each step of preparation (e.g., role of $Na_2CO_3$ and $O_2$ in chromite roasting, or $KOH$ and $O_2$ in pyrolusite fusion).
Chromyl Chloride Test: — This specific test for chlorides using $K_2Cr_2O_7$ is a recurring conceptual question.

Understanding these compounds requires a solid grasp of oxidation states, redox principles, and the influence of pH on chemical equilibria and reaction pathways. Their distinct colors and transformations make them excellent visual aids for learning these concepts.