Ion electron method (Half reaction method)

Ion–Electron Method (Half Reaction Method) in Redox Reactions

The Ion–Electron Method, also known as the Half Reaction Method, is a systematic way to balance redox reactions. It is especially useful in aqueous solutions and is widely used in electrochemistry, titration calculations, and competitive examinations such as NEET and JEE.

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What is a Redox Reaction?

A redox reaction is a chemical reaction in which oxidation and reduction occur simultaneously.

• Oxidation → Loss of electrons
• Reduction → Gain of electrons

In such reactions, one species loses electrons while another gains electrons. The Ion–Electron Method helps us balance both mass and charge properly.

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Why Do We Need the Ion–Electron Method?

In many redox reactions, especially those occurring in solution, balancing by simple inspection becomes difficult. The Ion–Electron Method provides a step-by-step scientific approach to balance:

• All atoms
• All charges
• Electrons transferred

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Steps to Balance Redox Reaction in Acidic Medium

Let us understand the method with an example:

Example Reaction:

MnO₄^- + Fe²⁺ → Mn²⁺ + Fe³⁺

Step 1: Separate into Two Half Reactions

Oxidation Half:

Fe²⁺ → Fe³⁺

Reduction Half:

MnO₄^- → Mn²⁺

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Step 2: Balance Atoms Other Than Oxygen and Hydrogen

Check if atoms except O and H are balanced. In this case, manganese and iron are already balanced.

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Step 3: Balance Oxygen Using Water (H₂O)

MnO₄^- contains 4 oxygen atoms. Add 4H₂O to the right side:

MnO₄^- → Mn²⁺ + 4H₂O

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Step 4: Balance Hydrogen Using H⁺

There are 8 hydrogen atoms on the right side. Add 8H⁺ to the left:

8H⁺ + MnO₄^- → Mn²⁺ + 4H₂O

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Step 5: Balance Charge Using Electrons (e^-)

Calculate total charge on both sides:

• Left side charge = +8 -1 = +7
• Right side charge = +2

To balance charge, add 5 electrons to the left side:

8H⁺ + MnO₄^- + 5e^- → Mn²⁺ + 4H₂O

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Step 6: Balance Oxidation Half

Fe²⁺ → Fe³⁺ + e^-

To equalize electrons (since reduction uses 5 electrons), multiply the oxidation half by 5:

5Fe²⁺ → 5Fe³⁺ + 5e^-

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Step 7: Add Both Half Reactions

Now add the two half reactions and cancel electrons:

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

✔ All atoms balanced
✔ Charges balanced
✔ Electrons canceled

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Balancing in Basic Medium

For basic medium, follow these steps:

1. First balance the reaction in acidic medium.
2. Add equal number of OH^- ions to both sides to neutralize H⁺.
3. Combine H⁺ and OH^- to form H₂O.
4. Cancel extra water molecules if present on both sides.

This converts the equation into basic medium conditions.

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Advantages of Ion–Electron Method

• Scientifically accurate
• Works for complex reactions
• Useful in electrochemistry
• Important for competitive exams
• Helps in understanding electron transfer clearly

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Applications

• KMnO₄ titration
• K₂Cr₂O₇ reactions
• Electrochemical cells
• Galvanic and electrolytic cells
• Industrial redox processes

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Conclusion

The Ion–Electron Method is one of the most reliable and systematic techniques to balance redox reactions. By separating oxidation and reduction processes, balancing atoms, and finally balancing charges using electrons, we ensure that both mass and charge are conserved.

Mastering this method will strengthen your understanding of electrochemistry and help you perform well in board examinations as well as competitive exams like NEET and JEE.

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Practice Tip: Try solving 4–5 redox reactions daily using this method to gain confidence.

Ionic Product under Ionic Equilibrium

1. Introduction to Ionic Equilibrium

Ionic equilibrium is an important part of physical chemistry that deals with the equilibrium established in electrolyte solutions due to partial or complete ionization of substances. When acids, bases, or salts are dissolved in water, they produce ions. In many cases, this ionization is reversible, and a dynamic equilibrium is established between ions and undissociated molecules.

The concept of ionic equilibrium helps in understanding the behavior of weak electrolytes, strength of acids and bases, solubility of salts, and precipitation reactions. One of the most important ideas derived from ionic equilibrium is the ionic product.

2. Meaning of Ionic Product

The ionic product of a solution is defined as the product of the molar concentrations of the ions present in solution, each raised to the power of its stoichiometric coefficient at any given moment.

For a general salt:

AxBy ⇌ xA^y+ + yB^x−

The ionic product (IP) is given by:

IP = [A^y+]^x [B^x−]^y

Ionic product represents the current state of the solution and does not necessarily indicate equilibrium conditions.

3. Ionic Product and Solubility Product

Ionic product is closely related to solubility product (Ksp), but the two are not the same. Solubility product is a constant value for a sparingly soluble salt at a given temperature, whereas ionic product can have different values depending on the concentrations of ions present in solution.

Ionic product can be less than, equal to, or greater than the solubility product. Comparison of IP with Ksp helps in predicting whether a precipitate will form or not.

4. Relationship between IP and Ksp

Case 1: IP < Ksp

The solution is unsaturated. More solute can dissolve, and no precipitation occurs.

Case 2: IP = Ksp

The solution is saturated and the system is in equilibrium. No precipitation occurs.

Case 3: IP > Ksp

The solution is supersaturated. Excess ions combine to form a solid precipitate.

5. Ionic Product of Water

Water is a weak electrolyte and undergoes slight ionization as shown below:

2H₂O ⇌ H₃O⁺ + OH⁻

The ionic product of water is given by:

K_w = [H⁺][OH⁻]

At 25°C, the value of Kw is:

K_w = 1.0 × 10⁻¹⁴

This constant is useful in determining pH, acidity, and basicity of solutions.

6. Ionic Product in Weak Electrolytes

Weak acids and weak bases do not ionize completely in solution. For a weak acid HA:

HA ⇌ H⁺ + A⁻

The ionic product is:

IP = [H⁺][A⁻]

At equilibrium, this ionic product becomes the acid dissociation constant (Ka). Similarly, for weak bases, ionic product leads to the base dissociation constant (Kb).

7. Ionic Product and Precipitation

When two electrolyte solutions are mixed, precipitation may occur if the ionic product exceeds the solubility product of the resulting salt.

Example:

Ag⁺ + Cl⁻ → AgCl(s)

Ionic product:

IP = [Ag⁺][Cl⁻]

If IP is greater than Ksp of AgCl, precipitation takes place.

8. Selective Precipitation

Selective precipitation is the method of separating ions in a mixture based on differences in their solubility products.

For example, silver chloride precipitates before lead chloride when chloride ions are added slowly because AgCl has a much smaller Ksp value.

This method is widely used in qualitative analysis.

9. Common Ion Effect and Ionic Product

The addition of a common ion increases the ionic product of a solution. This shifts the equilibrium in accordance with Le Chatelier’s principle and reduces the solubility of the salt.

Example:

AgCl(s) ⇌ Ag⁺ + Cl⁻

Addition of NaCl increases chloride ion concentration, thereby increasing IP and decreasing solubility of AgCl.

10. Numerical Example

If the concentration of Ag⁺ is 2 × 10⁻⁴ M and Cl⁻ is 1 × 10⁻⁶ M, then:

IP = (2 × 10⁻⁴) × (1 × 10⁻⁶) = 2 × 10⁻¹⁰

Since this value is greater than Ksp of AgCl, precipitation occurs.

11. Importance of Ionic Product

The concept of ionic product is important in predicting precipitation, understanding solubility, explaining common ion effect, qualitative salt analysis, environmental chemistry, and industrial processes.

12. Common Errors by Students

Students often confuse ionic product with solubility product, ignore stoichiometric coefficients, or use incorrect concentrations. Careful application of the concept avoids these errors.

13. Conclusion

Ionic product is a fundamental concept of ionic equilibrium that helps in understanding the behavior of electrolyte solutions. By comparing ionic product with solubility product, it becomes possible to predict precipitation and solubility behavior accurately. This concept is essential for board examinations and competitive exams like JEE and NEET.

Mendius Reaction and Gattermann Reaction

Mendius reaction and Gattermann reaction are important named reactions of organic chemistry. These reactions are commonly studied in Class 12, NEET, and JEE syllabi. Both reactions are used for the preparation of important organic compounds.

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1. Mendius Reaction

Mendius reaction is the reduction of nitriles to primary amines using sodium metal in alcohol. It is an important method for the preparation of aliphatic primary amines.

General Reaction

R–C≡N + 4[H] → R–CH₂–NH₂

Reagents Used

• Sodium metal (Na)
• Alcohol (ethanol or methanol)

Example

CH₃–CN + Na / C₂H₅OH → CH₃–CH₂–NH₂

Important Points

• Nitrile group is reduced to primary amine
• Carbon chain length increases by one carbon atom
• Sodium acts as a reducing agent
• Reaction proceeds through nascent hydrogen

Uses

• Preparation of primary aliphatic amines
• Used in organic synthesis

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2. Gattermann Reaction

Gattermann reaction is used for the introduction of aldehyde group into an aromatic ring. It is also known as a formylation reaction of aromatic compounds.

General Reaction

Ar–H + HCN + HCl → Ar–CHO

Catalysts Used

• Anhydrous aluminium chloride (AlCl₃)
• Cuprous chloride (CuCl)

Example

C₆H₆ + HCN + HCl → C₆H₅–CHO

In this reaction, benzene is converted into benzaldehyde.

Important Points

• Aldehyde (–CHO) group is introduced into aromatic ring
• Reaction mainly occurs at ortho and para positions
• Used for preparation of aromatic aldehydes

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Difference Between Gattermann and Gattermann–Koch Reaction

Gattermann Reaction	Gattermann–Koch Reaction
Uses HCN and HCl	Uses CO and HCl
Uses AlCl3 and CuCl as catalyst	Uses AlCl3 and CuCl as catalyst
Less commonly used	More commonly used

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One-Line Examination Answers

Mendius Reaction: Reduction of nitriles to primary amines using sodium metal in alcohol.

Gattermann Reaction: Formylation of aromatic compounds using HCN and HCl in presence of AlCl₃ and CuCl.

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Prepared for academic and examination purposes.

Reactions in Which Sodium Mercury Amalgam is Used

Reactions in Which Sodium–Mercury Amalgam (Na–Hg) is Used

Sodium–mercury amalgam, written as Na–Hg, is an important reagent used in both organic and inorganic chemistry. It mainly acts as a mild and controlled reducing agent. Because of its gentle action, it is preferred over metallic sodium in many reactions.

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What is Sodium–Mercury Amalgam?

Sodium amalgam is a solution of sodium metal in mercury. Mercury controls the high reactivity of sodium and allows the gradual release of electrons or nascent hydrogen.

Due to this property, sodium amalgam is widely used in laboratory reductions.

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1. Reduction of Aldehydes

Sodium amalgam reduces aldehydes to primary alcohols in aqueous or alcoholic medium.

General Reaction:

R–CHO + [H] → R–CH₂OH

Here, sodium amalgam acts as a source of nascent hydrogen.

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2. Reduction of Ketones

Ketones are reduced to secondary alcohols using sodium amalgam.

General Reaction:

R–CO–R′ + [H] → R–CHOH–R′

This reaction is useful when mild reduction conditions are required.

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3. Reduction of Nitro Compounds

Nitro compounds are reduced to amines using sodium amalgam.

General Reaction:

R–NO₂ + Na–Hg + H₂O → R–NH₂

This reaction is preferred when strong reducing agents like tin and hydrochloric acid are not suitable.

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4. Reduction of Unsaturated Compounds

Alkenes and alkynes can be reduced by sodium amalgam under controlled conditions.

Example:

R–CH=CH–R → R–CH₂–CH₂–R

The hydrogen produced from sodium amalgam adds across the multiple bond.

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5. Reduction of Oximes and Imines

Oximes and imines are reduced to corresponding amines.

Example:

R–CH=NOH + Na–Hg → R–CH₂–NH₂

This reaction is important in organic synthesis.

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6. Reduction of Carbonyl Compounds in Aqueous Medium

Sodium amalgam reduces carbonyl compounds such as aldehydes and ketones in aqueous medium without affecting sensitive functional groups.

Hence, it is useful for selective reduction reactions.

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7. Reduction of Metal Ions (Inorganic Chemistry)

Sodium amalgam is also used to reduce metal ions in solution.

Example:

Fe³⁺ + Na–Hg → Fe²⁺

This reaction is commonly used in redox chemistry experiments.

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8. Preparation of Alcohols from Acyl Chlorides

Acyl chlorides can be reduced to primary alcohols using sodium amalgam.

General Reaction:

R–COCl + Na–Hg + H₂O → R–CH₂OH

This method avoids harsh reducing conditions.

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Why Sodium–Mercury Amalgam is Preferred

• It is safer than metallic sodium
• It provides controlled reduction
• It is a mild reducing agent
• It is suitable for selective reactions

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One-Line Examination Answer

Sodium–mercury amalgam is used as a mild reducing agent in reactions such as reduction of aldehydes, ketones, nitro compounds, oximes, imines, and metal ions.

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Educational content prepared for chemistry students.

Why Chemical Equation is Used to Find Molecularity and Strength of KMnO4

Why Chemical Equation is Used to Find Molecularity and Strength of KMnO₄ Using M/20 Mohr’s Salt Solution

In volumetric analysis, especially in Class 11 and Class 12 chemistry practicals, potassium permanganate (KMnO₄) is commonly standardized using M/20 Mohr’s salt solution. A very common question asked in exams is:

“Why is a chemical equation required to find the molecularity and strength of KMnO₄?”

The answer lies in the stoichiometry of the reaction and the nature of KMnO₄.

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1. Importance of Chemical Equation in Volumetric Analysis

In titration calculations, we do not depend only on the volume of solutions. All calculations are based on the balanced chemical equation.

KMnO₄ does not react with Mohr’s salt in a simple 1:1 ratio. Therefore, without a chemical equation:

• The reaction ratio cannot be known
• The number of electrons transferred cannot be determined
• The equivalent weight and n-factor cannot be calculated

Hence, a chemical equation is compulsory to calculate the molecularity and strength of KMnO₄.

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2. Why Mohr’s Salt is Used

Mohr’s salt (FeSO₄·(NH₄)₂SO₄·6H₂O) is used because:

• It is a primary standard
• It has high purity
• It has stable composition
• Its molarity (M/20) is accurately known

KMnO₄ is not a primary standard because it decomposes on standing and may contain impurities. Therefore, KMnO₄ solution must be standardized using Mohr’s salt.

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3. Chemical Equation of the Reaction

The titration is carried out in acidic medium using dilute sulphuric acid. The balanced ionic equation is:

MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

From the above equation, we observe that:

• 1 mole of KMnO₄ reacts with 5 moles of Fe²⁺
• KMnO₄ acts as an oxidizing agent
• Mohr’s salt acts as a reducing agent

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4. Determination of Molecularity (Molarity / Normality)

From the chemical equation, it is clear that one mole of KMnO₄ gains five electrons. Therefore:

n-factor of KMnO₄ = 5

Using the titration formula:

N₁V₁ = N₂V₂

The normality of Mohr’s salt is known, and the volumes are measured experimentally. Using this equation, the normality of KMnO₄ is calculated.

Molarity of KMnO₄ is then obtained using:

M = N / n-factor

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5. Determination of Strength of KMnO₄

Once the molarity of KMnO₄ is known, its strength is calculated by:

Strength (g/L) = Molarity × Molar Mass

Molar mass of KMnO₄ = 158 g mol⁻¹

Thus, the strength of KMnO₄ solution can be accurately determined.

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6. Why Chemical Equation is Essential (Exam Point)

• To know the exact reaction ratio
• To calculate the n-factor of KMnO₄
• To determine molarity and strength correctly
• Because the reaction is not 1:1

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7. One-Line Board Answer

Chemical equation is used because KMnO₄ reacts with Mohr’s salt in a fixed stoichiometric ratio, which is necessary to calculate the molecularity and strength of KMnO₄.

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Prepared for educational and examination purposes.

Named Reactions of Class XI and XII Organic Chemistry

Named reactions are very important in organic chemistry. These reactions are frequently asked in CBSE board exams, NEET, JEE, and other competitive examinations. This article provides a complete chapter-wise list of important named reactions from Class XI and Class XII organic chemistry.

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CLASS XI – ORGANIC CHEMISTRY

1. General Organic Chemistry (GOC)

• Wurtz Reaction
• Fittig Reaction
• Wurtz–Fittig Reaction
• Kolbe’s Electrolytic Reaction
• Frankland Reaction

2. Alkanes

• Wurtz Reaction
• Kolbe’s Electrolysis
• Frankland Reaction
• Hunsdiecker Reaction
• Hofmann Degradation

3. Alkenes

• Dehydration of Alcohols
• Saytzeff (Zaitsev) Rule
• Hoffmann Rule
• Markovnikov’s Rule
• Anti-Markovnikov Addition (Peroxide Effect / Kharasch Effect)
• Baeyer’s Test
• Ozonolysis

4. Alkynes

• Ozonolysis of Alkynes
• Acidic Nature of Terminal Alkynes
• Polymerisation Reactions

5. Aromatic Hydrocarbons

• Friedel–Crafts Alkylation
• Friedel–Crafts Acylation
• Wurtz–Fittig Reaction
• Nitration Reaction
• Sulphonation Reaction
• Halogenation Reaction

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CLASS XII – ORGANIC CHEMISTRY

6. Haloalkanes and Haloarenes

• Wurtz Reaction
• Fittig Reaction
• Wurtz–Fittig Reaction
• Finkelstein Reaction
• Swarts Reaction
• Sandmeyer Reaction
• Gattermann Reaction
• Hunsdiecker Reaction
• Dow’s Process
• Ullmann Reaction

7. Alcohols, Phenols and Ethers

• Williamson Ether Synthesis
• Reimer–Tiemann Reaction
• Kolbe–Schmitt Reaction
• Dehydration of Alcohols
• Lucas Test
• Esterification Reaction

8. Aldehydes and Ketones

• Aldol Condensation
• Cross Aldol Condensation
• Cannizzaro Reaction
• Clemmensen Reduction
• Wolff–Kishner Reduction
• Rosenmund Reduction
• Stephen Reaction
• Haloform Reaction
• Perkin Reaction
• Benzoin Condensation

9. Carboxylic Acids

• Hell–Volhard–Zelinsky (HVZ) Reaction
• Esterification Reaction
• Decarboxylation Reaction
• Kolbe’s Electrolytic Reaction
• Reduction to Alcohol using LiAlH₄

10. Amines

• Hofmann Bromamide Reaction
• Gabriel Phthalimide Synthesis
• Carbylamine Reaction
• Hinsberg Test
• Diazotisation Reaction
• Sandmeyer Reaction
• Gattermann Reaction
• Coupling Reaction (Azo Dye Formation)

11. Biomolecules

• Peptide Bond Formation
• Hydrolysis of Proteins
• Mutarotation
• Fermentation Reaction
• Glycosidic Bond Formation

12. Polymers

• Addition Polymerisation
• Condensation Polymerisation
• Free Radical Polymerisation
• Ziegler–Natta Polymerisation

13. Chemistry in Everyday Life

This chapter mainly focuses on applications of chemistry. Very few named reactions are directly asked from this chapter.

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Most Important Exam-Oriented Named Reactions

• Wurtz Reaction
• Kolbe’s Electrolysis
• Friedel–Crafts Reaction
• Markovnikov’s Rule
• Anti-Markovnikov Addition
• Aldol Condensation
• Cannizzaro Reaction
• Clemmensen Reduction
• Wolff–Kishner Reduction
• HVZ Reaction
• Reimer–Tiemann Reaction
• Kolbe–Schmitt Reaction
• Sandmeyer Reaction
• Hofmann Bromamide Reaction
• Gabriel Synthesis
• Williamson Ether Synthesis

Conclusion: Learning named reactions helps students quickly identify reaction pathways, reagents, and products. Proper revision of these reactions can significantly improve performance in board exams and competitive examinations.

How to Find Cation and Anion in a Given Salt

In qualitative inorganic analysis, the identification of a salt involves determining two components:

• Cation (Basic radical)
• Anion (Acid radical)

This systematic analysis is commonly followed in CBSE and other board practical chemistry laboratories. The identification is carried out step by step using preliminary tests, dry tests, and wet confirmatory tests.

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Step 1: Preliminary Examination (Dry Tests)

(a) Physical Observation

Colour of the salt:

• Blue colour – Copper (Cu²⁺)
• Green colour – Iron (Fe²⁺) or Nickel (Ni²⁺)
• White colour – Zinc (Zn²⁺), Calcium (Ca²⁺), Sodium (Na⁺), Potassium (K⁺)

Smell of the salt:

• Ammonia smell – Ammonium ion (NH₄⁺)
• Rotten egg smell on heating – Sulphide ion (S²⁻)

(b) Action of Heat

Heat a small amount of the salt in a dry test tube and observe:

• Crackling sound – Presence of water of crystallization
• Evolution of gas:
  • Carbon dioxide – Carbonate (CO₃²⁻)
  • Sulphur dioxide – Sulphite (SO₃²⁻)
  • Brown fumes – Nitrate (NO₃⁻)

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Step 2: Solubility Test

The salt is tested for solubility in:

• Cold water
• Hot water
• Dilute hydrochloric acid (HCl)
• Dilute nitric acid (HNO₃)

This helps in selecting suitable reagents for further analysis.

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Step 3: Identification of Anion (Acid Radical)

(A) Test with Dilute Hydrochloric Acid

Add dilute HCl to the salt solution and observe:

Observation	Inference (Anion)
Effervescence; CO2 turns lime water milky	Carbonate (CO32−)
Rotten egg smell	Sulphide (S2−)
Pungent smelling gas (SO2)	Sulphite (SO32−)
No reaction	Proceed to next test

(B) Test with Concentrated Sulphuric Acid

Add concentrated H₂SO₄ to solid salt:

Observation	Anion
Brown fumes	Nitrate (NO3−)
White fumes with pungent smell	Chloride (Cl−)
Red vapours	Bromide (Br−)
Violet vapours	Iodide (I−)

(C) Confirmatory Tests for Anions

• Chloride ion: Add AgNO₃ solution → White precipitate soluble in NH₄OH
• Sulphate ion: Add BaCl₂ solution → White precipitate insoluble in HCl
• Nitrate ion: Brown ring test confirms nitrate

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Step 4: Identification of Cation (Basic Radical)

Cations are identified by systematic group analysis using specific reagents.

Group 0: Ammonium Ion

• Warm salt with NaOH → Ammonia gas evolved (Turns red litmus blue)

Group I: Silver, Lead, Mercurous Ions

• Add dilute HCl → White precipitate formed

Group II: Copper, Cadmium, Bismuth etc.

• Pass H₂S gas in acidic medium → Coloured precipitate

Group III: Iron, Aluminium, Chromium

• Add NH₄OH in presence of NH₄Cl → Precipitate formed

Group IV: Zinc, Manganese, Nickel, Cobalt

• Pass H₂S gas in basic medium → Precipitate formed

Group V: Calcium, Barium, Strontium

• Add (NH₄)₂CO₃ → White precipitate

Group VI: Sodium, Potassium, Magnesium

• Identified using flame test and special tests

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Step 5: Flame Test

Flame Colour	Cation
Golden yellow	Sodium (Na+)
Lilac	Potassium (K+)
Brick red	Calcium (Ca2+)
Apple green	Barium (Ba2+)
Blue-green	Copper (Cu2+)

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Final Result

After performing all tests and confirmatory reactions, the salt is reported as:

Cation present: __________
Anion present: __________

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Important Practical Tips

• Always identify the anion before the cation
• Write observation and inference clearly
• Confirm results using confirmatory tests
• Flame test is very useful for Group VI cations

This method ensures accurate identification of cations and anions in an unknown salt during qualitative inorganic analysis.