Preparation, Properties, and Uses of Potassium Dichromate (K₂Cr₂O₇) and Potassium Permanganate (KMnO₄)

Potassium Dichromate (K₂Cr₂O₇)

Preparation of K₂Cr₂O₇

Potassium dichromate is primarily prepared from chromite ore (FeCr₂O₄). The process involves several steps:

1. Fusion with Sodium Carbonate: Chromite ore is heated in a furnace with sodium carbonate (Na₂CO₃) in the presence of air (oxygen). This converts chromium(III) oxide (Cr₂O₃) to sodium chromate (Na₂CrO₄). Magnesium oxide (MgO) and aluminum oxide (Al₂O₃) present as impurities do not react.

   Chemical equation:

   4FeCr₂O₄ + 8Na₂CO₃ + 7O₂ → 8Na₂CrO₄ + 2Fe₂O₃ + 8CO₂

2. Conversion of Sodium Chromate to Sodium Dichromate: The solution of sodium chromate is filtered and acidified by passing carbon dioxide (CO₂) through it. Sodium chromate is converted to sodium dichromate (Na₂Cr₂O₇).

   Chemical equation:

   2Na₂CrO₄ + CO₂ → Na₂Cr₂O₇ + Na₂CO₃

   Alternatively, acidification with sulfuric acid can be used:

   2Na₂CrO₄ + H₂SO₄ → Na₂Cr₂O₇ + Na₂SO₄ + H₂O

3. Conversion of Sodium Dichromate to Potassium Dichromate: Sodium dichromate is more soluble in water than potassium dichromate. Therefore, when a solution of sodium dichromate is treated with potassium chloride (KCl), potassium dichromate (K₂Cr₂O₇) precipitates out due to its lower solubility.

   Chemical equation:

   Na₂Cr₂O₇ + 2KCl → K₂Cr₂O₇↓ + 2NaCl

Properties of K₂Cr₂O₇

• Physical State: It is a crystalline solid, typically orange-red in color.

• Solubility: It is soluble in water, forming an orange-colored solution.

• Acidic Nature of Solution: The aqueous solution is acidic due to hydrolysis.

• Oxidizing Agent: Potassium dichromate is a strong oxidizing agent, especially in acidic medium. The dichromate ion (Cr₂O₇^2-) gets reduced to chromium(III) ions (Cr³⁺).

  In acidic medium (e.g., with H₂SO₄), the reduction half-reaction is:

  Cr₂O₇^2- + 14H⁺ + 6e^- → 2Cr³⁺ + 7H₂O (E° = +1.33 V)

  Examples of Oxidation Reactions:

  • Oxidation of ferrous ions (Fe²⁺) to ferric ions (Fe³⁺):

  Cr₂O₇^2- + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O

  • Oxidation of iodide ions (I^-) to iodine (I₂):

  Cr₂O₇^2- + 6I^- + 14H⁺ → 2Cr³⁺ + 3I₂ + 7H₂O

  • Oxidation of sulfur dioxide (SO₂) to sulfuric acid (H₂SO₄):

  Cr₂O₇^2- + 3SO₂ + 2H⁺ → 2Cr³⁺ + 3SO₄^2- + H₂O

  • Oxidation of alcohols to aldehydes or ketones, and further to carboxylic acids.

• Effect of pH: The color of the dichromate solution depends on pH. In alkaline solution, dichromate ions (Cr₂O₇^2-, orange) are converted to chromate ions (CrO₄^2-, yellow).

  Cr₂O₇^2- + 2OH^- → 2CrO₄^2- + H₂O (Orange to Yellow)

  In strongly acidic solution, the chromate ions are converted back to dichromate ions.

  2CrO₄^2- + 2H⁺ → Cr₂O₇^2- + H₂O (Yellow to Orange)

• Thermal Decomposition: When heated strongly, potassium dichromate decomposes to potassium chromate, chromium(III) oxide, and oxygen.

  4K₂Cr₂O₇ → 4K₂CrO₄ + 2Cr₂O₃ + 3O₂

Uses of K₂Cr₂O₇

• It is used as a strong oxidizing agent in many industrial processes and in the laboratory.

• It is used in the preparation of other chromium compounds.

• It is used in dyeing and printing textiles.

• It is used in the manufacture of matches and firecrackers.

• It is used in tanning of leather.

• It is used in electroplating for chromium plating.

• It is used as a cleaning agent for glassware, especially for removing organic matter.

• It is used in the petroleum industry to purify kerosene.

Potassium Permanganate (KMnO₄)

Preparation of KMnO₄

Potassium permanganate is prepared from manganese dioxide (MnO₂) in two steps:

1. Oxidation of MnO₂ to Potassium Manganate: Manganese dioxide is fused with potassium hydroxide (KOH) in the presence of air or oxygen. This reaction is carried out at high temperatures (around 513 K or 240 °C) in an iron or nickel tray. This produces potassium manganate (K₂MnO₄), which is green in color.

   Chemical equation:

   2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O

2. Oxidation of Potassium Manganate to Potassium Permanganate: Potassium manganate is then oxidized to potassium permanganate. This can be done by:

   • Electrolytic Oxidation: This is the most common and efficient method. An aqueous solution of potassium manganate is electrolyzed. At the anode, manganate ions (MnO₄^2-) are oxidized to permanganate ions (MnO₄^-).

     Anode reaction: MnO₄^2- → MnO₄^- + e^-

     Cathode reaction: 2H₂O + 2e^- → H₂ + 2OH^-

     Overall reaction: 2K₂MnO₄ + 2H₂O → 2KMnO₄ + 2KOH + H₂

     The KOH formed is recycled back to the first step.

   • Chemical Oxidation: Potassium manganate can also be oxidized using chlorine gas or ozone in neutral or slightly alkaline solution.

     Using Chlorine:

     2K₂MnO₄ + Cl₂ → 2KMnO₄ + 2KCl

     Using Ozone:

     2K₂MnO₄ + O₃ + H₂O → 2KMnO₄ + 2KOH + O₂

Properties of KMnO₄

• Physical State: It is a dark purple, crystalline solid. It is sparingly soluble in water.

• Oxidizing Agent: Potassium permanganate is a very strong oxidizing agent. Its oxidizing power depends on the pH of the medium.

  In Acidic Medium: It is a very strong oxidizing agent. The permanganate ion (MnO₄^-) is reduced to manganese(II) ion (Mn²⁺), which is almost colorless.

  Reduction half-reaction:

  MnO₄^- + 8H⁺ + 5e^- → Mn²⁺ + 4H₂O (E° = +1.51 V)

  Examples:

  • Oxidation of ferrous ions (Fe²⁺) to ferric ions (Fe³⁺):

  MnO₄^- + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

  • Oxidation of iodide ions (I^-) to iodine (I₂):

  2MnO₄^- + 10I^- + 16H⁺ → 2Mn²⁺ + 5I₂ + 8H₂O

  • Oxidation of oxalate ions (C₂O₄^2-) to carbon dioxide (CO₂):

  2MnO₄^- + 5C₂O₄^2- + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O

  This reaction is used in the quantitative estimation of KMnO₄ (titration with oxalic acid).

  In Neutral or Weakly Alkaline Medium: It acts as an oxidizing agent and is reduced to manganese dioxide (MnO₂), a brown precipitate.

  Reduction half-reaction:

  MnO₄^- + 2H₂O + 3e^- → MnO₂(s) + 4OH^-

  Examples:

  • Oxidation of ferrous ions (Fe²⁺) to ferric ions (Fe³⁺):

  MnO₄^- + 3Fe²⁺ + 7H₂O → MnO₂(s) + 3Fe³⁺ + 7OH^-

  • Oxidation of iodide ions (I^-) to iodine (I₂):

  2MnO₄^- + 3I^- + H₂O → 2MnO₂(s) + 3I₂ + 4OH^-

  In Strongly Alkaline Medium: It is reduced to manganate ions (MnO₄^2-), which are green.

  Reduction half-reaction:

  MnO₄^- + e^- → MnO₄^2-

  Example:

  • Oxidation of ferrous ions (Fe²⁺) to ferric ions (Fe³⁺):

  MnO₄^- + Fe²⁺ + 2OH^- → MnO₄^2- + Fe³⁺ + H₂O

• Thermal Decomposition: When heated strongly, potassium permanganate decomposes to potassium manganate, manganese dioxide, and oxygen.

  2KMnO₄(s) $\xrightarrow{\Delta}$ K₂MnO₄(s) + MnO₂(s) + O₂(g)

Uses of KMnO₄

• It is a powerful oxidizing agent used in qualitative and quantitative analysis (titrations).

• It is used as a disinfectant and germicide, often in dilute solutions (e.g., for sterilizing water, treating skin infections).

• It is used in the preparation of oxygen in the laboratory by heating with hydrogen peroxide or by reaction with concentrated sulfuric acid.

• It is used in the purification of water, oxidizing impurities.

• It is used in the de-colorization of certain products in the cosmetic and pharmaceutical industries.

• It is used in the synthesis of organic compounds.

Lanthanoids and Actinoids

Lanthanoids

The lanthanoid series comprises 14 elements from cerium (Ce, Z=58) to lutetium (Lu, Z=71). These elements are characterized by the successive filling of the 4f orbitals. They are also called "rare earth elements," although they are not particularly rare. They are placed in Group 3 of the periodic table, below the actinides, and are usually shown as a separate block.

Electronic Configurations of Lanthanoids

The general electronic configuration of lanthanoids is [Xe] 4f^1-14 5d^0-1 6s².

The filling of the 4f orbitals is the defining characteristic. However, there are some exceptions due to the relative stability of half-filled (f⁷) and completely filled (f¹⁴) subshells, as well as the close energy levels of 4f, 5d, and 6s orbitals.

General Trend:

• All lanthanoids end with 6s².
• Most have 5d⁰, but some have 5d¹.
• The 4f subshell is filled from 4f¹ to 4f¹⁴.

Exceptions:

• Gadolinium (Gd, Z=64): [Xe] 4f⁷ 5d¹ 6s² (Half-filled 4f subshell is very stable).
• Lutetium (Lu, Z=71): [Xe] 4f¹⁴ 5d¹ 6s² (Completely filled 4f subshell is very stable). Although the 5d orbital has one electron, lutetium is considered the last element of the series as the 4f subshell is completely filled. Some texts might show Lu as [Xe] 4f¹⁴ 6s², considering the 5d electron as an anomaly.

Electronic Configurations of Lanthanoids:

Element	Symbol	Atomic Number (Z)	Electronic Configuration
Cerium	Ce	58	[Xe] 4f1 5d1 6s2
Praseodymium	Pr	59	[Xe] 4f3 6s2
Neodymium	Nd	60	[Xe] 4f4 6s2
Promethium	Pm	61	[Xe] 4f5 6s2
Samarium	Sm	62	[Xe] 4f6 6s2
Europium	Eu	63	[Xe] 4f7 6s2
Gadolinium	Gd	64	[Xe] 4f7 5d1 6s2
Terbium	Tb	65	[Xe] 4f9 6s2
Dysprosium	Dy	66	[Xe] 4f10 6s2
Holmium	Ho	67	[Xe] 4f11 6s2
Erbium	Er	68	[Xe] 4f12 6s2
Thulium	Tm	69	[Xe] 4f13 6s2
Ytterbium	Yb	70	[Xe] 4f14 6s2
Lutetium	Lu	71	[Xe] 4f14 5d1 6s2

Oxidation States of Lanthanoids

Lanthanoids exhibit multiple oxidation states, but +3 is the most common and stable oxidation state for all of them. This is because it corresponds to the loss of the two 6s electrons and one 4f electron, resulting in a stable configuration with either a half-filled (f⁷) or a completely filled (f¹⁴) 4f subshell.

Common Oxidation State: +3

• All lanthanoids show +3 oxidation state. For example:
  • Ce³⁺ is [Xe] 4f¹
  • Eu³⁺ is [Xe] 4f⁶
  • Yb³⁺ is [Xe] 4f¹³
  • Lu³⁺ is [Xe] 4f¹⁴

Other Oxidation States:

Some lanthanoids exhibit +2 and +4 oxidation states. These are usually observed when the resulting electronic configuration is particularly stable (e.g., f⁰, f⁷, or f¹⁴).

• +2 Oxidation States:
  • Samarium (Sm): [Xe] 4f⁶ 6s² → Sm²⁺ [Xe] 4f⁶ (not particularly stable)
  • Europium (Eu): [Xe] 4f⁷ 6s² → Eu²⁺ [Xe] 4f⁷ (very stable, half-filled f subshell)
  • Ytterbium (Yb): [Xe] 4f¹⁴ 6s² → Yb²⁺ [Xe] 4f¹⁴ (very stable, completely filled f subshell)
• +4 Oxidation States:
  • Cerium (Ce): [Xe] 4f¹ 5d¹ 6s² → Ce⁴⁺ [Xe] (very stable, empty f subshell)
  • Praseodymium (Pr): [Xe] 4f³ 6s² → Pr⁴⁺ [Xe] 4f¹ (relatively stable)
  • Neodymium (Nd): [Xe] 4f⁴ 6s² → Nd⁴⁺ [Xe] 4f² (relatively stable)
  • Terbium (Tb): [Xe] 4f⁹ 6s² → Tb⁴⁺ [Xe] 4f⁷ (very stable, half-filled f subshell)

Summary of Oxidation States:

Element	Common State	Other States
La	+3	-
Ce	+3	+4
Pr	+3	+4
Nd	+3	+4
Pm	+3	-
Sm	+3	+2
Eu	+3	+2
Gd	+3	-
Tb	+3	+4
Dy	+3	-
Ho	+3	-
Er	+3	-
Tm	+3	+2
Yb	+3	+2
Lu	+3	-

Actinoids

The actinoid series comprises 14 elements from thorium (Th, Z=90) to lawrencium (Lr, Z=103). These elements are characterized by the successive filling of the 5f orbitals. All actinoids are radioactive, and most of them are synthetic.

Electronic Configurations of Actinoids

The general electronic configuration of actinoids is [Rn] 5f^1-14 6d^0-1 7s².

Similar to lanthanoids, there are exceptions due to the close energy of 5f, 6d, and 7s orbitals. The filling of 5f orbitals is the defining characteristic.

General Trend:

• All actinoids end with 7s².
• The 5f and 6d orbitals are filled.

Electronic Configurations of Actinoids:

Element	Symbol	Atomic Number (Z)	Electronic Configuration
Actinium	Ac	89	[Rn] 6d1 7s2
Thorium	Th	90	[Rn] 6d2 7s2 (or [Rn] 5f0 6d2 7s2)
Protactinium	Pa	91	[Rn] 5f2 6d1 7s2 (or [Rn] 5f1 6d2 7s2)
Uranium	U	92	[Rn] 5f3 6d1 7s2
Neptunium	Np	93	[Rn] 5f4 6d1 7s2
Plutonium	Pu	94	[Rn] 5f6 7s2 (or [Rn] 5f5 6d1 7s2)
Americium	Am	95	[Rn] 5f7 7s2
Curium	Cm	96	[Rn] 5f7 6d1 7s2
Berkelium	Bk	97	[Rn] 5f9 7s2
Californium	Cf	98	[Rn] 5f10 7s2
Einsteinium	Es	99	[Rn] 5f11 7s2
Fermium	Fm	100	[Rn] 5f12 7s2
Mendelevium	Md	101	[Rn] 5f13 7s2
Nobelium	No	102	[Rn] 5f14 7s2
Lawrencium	Lr	103	[Rn] 5f14 6d1 7s2 (or [Rn] 5f13 6d2 7s2)

Oxidation States of Actinoids

Actinoids exhibit a wider range of oxidation states compared to lanthanoids. This is because the energies of the 5f, 6d, and 7s electrons are very close, making it easier to lose varying numbers of these electrons.

Common Oxidation States:

• +3: This is a common oxidation state for most actinoids, analogous to lanthanoids.
• +4: This is also a common and stable oxidation state for many actinoids, particularly Thorium (Th).

Other Oxidation States:

Actinoids can exhibit oxidation states ranging from +3 to +7.

• +5: Seen in Np (e.g., NpO₂⁺) and Pu.
• +6: Seen in U (e.g., UO₂²⁺) and Pu.
• +7: Seen in Np (e.g., NpO₅^3-) and Am in strongly oxidizing conditions.

Variability:

• The earliest actinoids (Th, Pa, U) show a greater variety of oxidation states, including higher ones like +4, +5, and +6.
• As the atomic number increases, the +3 oxidation state becomes more dominant, similar to the trend in lanthanoids.
• Americium (Am) is the last element to show a +7 oxidation state.
• Nobelium (No) and Lawrencium (Lr) show +2 and +3 states, respectively, with +2 being stable for No due to the f¹⁴ configuration.

Comparison with Lanthanoids:

• Actinoids show a greater number and wider range of oxidation states than lanthanoids.
• The 5f orbitals are less effectively shielded than the 4f orbitals, leading to greater involvement of 5f electrons in bonding.
• Actinoids are generally more reactive than lanthanoids.

Electronic Configuration and Stability:

Similar to lanthanoids, oxidation states are favored when they lead to stable electronic configurations (f⁰, f⁷, f¹⁴).

• Th⁴⁺ is [Rn] 5f⁰.
• Am³⁺ is [Rn] 5f⁷.
• No²⁺ is [Rn] 5f¹⁴.