Glycerol + excess HI → 2-iodopropane (isopropyl iodide) + I₂ + H₂O

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Explanation

• First, –OH groups are replaced by –I
• Then elimination and rearrangement occur
• Finally, the most stable secondary alkyl iodide is formed

Ethylene glycol (ethane-1,2-diol)

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• Ethylene glycol has two –OH groups, so it forms strong hydrogen bonding.
• It lowers the freezing point of water (freezing point depression).
• It also raises the boiling point, preventing overheating.

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✔ Used as antifreeze because it prevents water from freezing in cold conditions.

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Other than ethylene glycol, there are a few compounds used as antifreeze:

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Other Antifreeze Compounds

1. Propylene glycol (propane-1,2-diol)

• Less toxic than ethylene glycol
• Used in safer antifreeze formulations (food-grade systems)

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2. Glycerol (glycerine, propane-1,2,3-triol)

• Has three –OH groups → strong H-bonding
• Good freezing point depression
• Used in older or special applications

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3. Methanol / Ethanol

• Also lower freezing point
• But ❌ not preferred because:
  • Highly volatile
  • Flammable
  • Evaporate easily

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Step 1: Calculate DU

Formula:$$DU = \frac{2C + 2 – H}{2}$$

For C₆H₁₄O (ignore O in DU):$$DU = \frac{2(6) + 2 – 14}{2} = \frac{12 + 2 – 14}{2} = 0$$

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Meaning of DU = 0

• No double bond
• No triple bond
• No ring

So, all structures are open-chain saturated alcohols

1. Picric acid has three –NO₂ groups which are strong electron-withdrawing groups.
2. These groups stabilize the conjugate base (picrate ion) by resonance and –I effect.
3. Formic acid has only one –COOH group, so its conjugate base is less stabilized.
4. Greater stabilization of conjugate base ⇒ higher acidity, so picric acid is stronger.

HCOOH (Picric Acid)

Picric acid is soluble in NaHCO₃ solution?

Picric acid (2,4,6-trinitrophenol) is a strong acid due to the presence of three –NO₂ groups, which strongly withdraw electrons and stabilize the phenoxide ion. Because of this high acidity, it can react even with a weak base like NaHCO₃.

Reaction:
Picric acid + NaHCO₃ → Sodium picrate (soluble) + CO₂ + H₂O

So, it dissolves in NaHCO₃ with effervescence of CO₂ gas.

The line corresponding to the transition where the electron comes from an infinitely large orbit (n = ∞) to a fixed lower energy level.

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🔹 What it means simply

• When an electron falls from very high energy levels (n = very large), the spectral lines get closer and closer.
• Finally, at n = ∞, they merge into one line → this is the series limit (limiting line).

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🔹 Formula (Rydberg equation)

$\frac{1}{\lambda} = R \left( \frac{1}{n_1^2} – \frac{1}{n_2^2} \right)$

For limiting line$$\frac{1}{\lambda} = \frac{R}{n_1^2}$$

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Examples of limits

1. Lyman series (n₁ = 1)
   • Limit: electron falls from ∞ → 1
   • Region: Ultraviolet (UV)
2. Balmer series (n₁ = 2)
   • Limit: ∞ → 2
   • Region: Visible
3. Paschen series (n₁ = 3)
   • Limit: ∞ → 3
   • Region: Infrared

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Limiting line = shortest wavelength (maximum energy) line of a series.

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Lyman Alpha (α) Line

The Lyman α (alpha) line is the first spectral line of the Lyman series.

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1. Acidic Oxides

Definition: Oxides that react with water to form acids or react with bases to give salt + water.

Examples:

• CO₂, SO₂, SO₃, P₂O₅, N₂O₅, Cl₂O₇
• Non-metal oxides

Reactions:

• CO₂ + H₂O → H₂CO₃
• SO₃ + 2NaOH → Na₂SO₄ + H₂O

Trend:

• Non-metallic character ↑ ⇒ acidity ↑
• Across a period: acidity increases
• Down a group: acidity decreases

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2. Basic Oxides

Definition: Oxides that react with water to form bases or react with acids to give salt + water.

Examples:

• Na₂O, CaO, MgO, K₂O
• Mostly metal oxides

Reactions:

• Na₂O + H₂O → 2NaOH
• CaO + 2HCl → CaCl₂ + H₂O

Trend:

• Metallic character ↑ ⇒ basicity ↑
• Down a group: basicity increases
• Across a period: basicity decreases

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3. Amphoteric Oxides

Definition: Oxides that react with both acids and bases.

Examples (VERY IMPORTANT for exams):

• Al₂O₃, ZnO, BeO, SnO, PbO, Cr₂O₃

Reactions:

• ZnO + 2HCl → ZnCl₂ + H₂O (acts basic)
• ZnO + 2NaOH → Na₂ZnO₂ + H₂O (acts acidic)

Trend:

• Found near metal–nonmetal boundary
• Small size + high charge density → amphoteric behavior

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4. Neutral Oxides

Definition: Oxides that do not react with acids or bases.

Examples:

• CO, NO, N₂O

Trend:

• Mostly simple molecules
• No acidic or basic nature

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IMPORTANT PERIODIC TREND (MOST ASKED)

Across a period (Left → Right):

Basic → Amphoteric → Acidic

Example (Period 3):
Na₂O → MgO → Al₂O₃ → SiO₂ → P₂O₅ → SO₃ → Cl₂O₇

Key shift:

• Metals → Non-metals
• Basic → Acidic

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QUICK REVISION TABLE

Type	Nature	Examples
Basic	Metal oxides	Na₂O, CaO
Amphoteric	Both	Al₂O₃, ZnO
Acidic	Non-metal oxides	SO₃, CO₂
Neutral	No reaction	CO, NO

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EXAM TIPS

• Al₂O₃, ZnO → always remember (amphoteric)
• CO ≠ acidic (neutral!)
• Higher oxidation state ⇒ more acidic oxide
  • Example: SO₂ < SO₃ (acidity increases)

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MCQs

Q1. Arrange in increasing acidic nature:

A. CO₂ < SiO₂ < SO₃
B. SiO₂ < CO₂ < SO₃
C. SO₃ < CO₂ < SiO₂
D. CO₂ < SO₃ < SiO₂

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Q2. Arrange in decreasing basic strength:

A. Na₂O > MgO > Al₂O₃
B. Al₂O₃ > MgO > Na₂O
C. MgO > Na₂O > Al₂O₃
D. Na₂O > Al₂O₃ > MgO

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Q3. Arrange in increasing amphoteric character:

A. Na₂O < MgO < Al₂O₃
B. Al₂O₃ < MgO < Na₂O
C. MgO < Na₂O < Al₂O₃
D. Na₂O < Al₂O₃ < MgO

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Q4. Arrange in increasing acidic strength:

A. SO₂ < SO₃ < Cl₂O₇
B. Cl₂O₇ < SO₃ < SO₂
C. SO₃ < SO₂ < Cl₂O₇
D. SO₂ < Cl₂O₇ < SO₃

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Q5. Arrange in increasing basic character:

A. CaO < MgO < Na₂O
B. Na₂O < MgO < CaO
C. MgO < CaO < Na₂O
D. CaO < Na₂O < MgO

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Q6. Arrange in increasing acidic nature:

A. P₂O₅ < SO₃ < Cl₂O₇
B. Cl₂O₇ < SO₃ < P₂O₅
C. SO₃ < P₂O₅ < Cl₂O₇
D. P₂O₅ < Cl₂O₇ < SO₃

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Q7. Arrange in increasing basic strength:

A. BeO < MgO < CaO
B. CaO < MgO < BeO
C. MgO < BeO < CaO
D. BeO < CaO < MgO

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Q8. Arrange in increasing acidic character:

A. N₂O₃ < NO₂ < N₂O₅
B. N₂O₅ < NO₂ < N₂O₃
C. NO₂ < N₂O₃ < N₂O₅
D. N₂O₃ < N₂O₅ < NO₂

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ANSWERS

1. B → SiO₂ < CO₂ < SO₃
2. A → Na₂O > MgO > Al₂O₃
3. A → Na₂O < MgO < Al₂O₃
4. A → SO₂ < SO₃ < Cl₂O₇
5. C → MgO < CaO < Na₂O
6. A → P₂O₅ < SO₃ < Cl₂O₇
7. A → BeO < MgO < CaO
8. A → N₂O₃ < NO₂ < N₂O₅

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SHORT TRICKS (VERY IMPORTANT)

• Across period: Basic ↓, Acidic ↑
• Down group: Basic ↑
• Higher oxidation state ⇒ more acidic
• Metal oxides → basic
• Non-metal oxides → acidic
• Border elements (Be, Al, Zn) → amphoteric

🔹 Definition

Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a chemical bond.

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Important Points

• Applies only to bonded atoms
• It is a relative (dimensionless) value
• Most commonly used scale → Pauling scale

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Trends in Periodic Table

Across a Period (→)

Increases

• Reason: Increase in nuclear charge → stronger pull on electrons

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Down a Group (↓)

Decreases

• Reason: Increase in size → weaker attraction

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🔹 Important Order

Fluorine (F) is the most electronegative element$$\textbf{F > O > N > Cl > Br > I}$$

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Factors Affecting Electronegativity

1. Atomic Size
   • Smaller atom → higher electronegativity
2. Nuclear Charge
   • More protons → stronger attraction
3. Shielding Effect
   • More shielding → lower electronegativity
4. Hybridisation
   • sp > sp² > sp³

More s-character → electrons closer to nucleus

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Important Exceptions

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❗ Noble Gases

• Usually no electronegativity
• (Because they rarely form bonds)

Pauling Scale (Electronegativity) — JEE/NEET Concept

The Pauling scale is the most commonly used scale to measure electronegativity, proposed by Linus Pauling.

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Basic Idea

Electronegativity is calculated based on:

✔ Bond energies (bond dissociation enthalpies)

• If a bond A–B is stronger than expected, it means:
  A and B have different electronegativities

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Formula (Conceptual)

Pauling related electronegativity difference to bond energy:$$\chi_A – \chi_B \propto \sqrt{D_{AB} – \frac{D_{AA} + D_{BB}}{2}}$$

Where:

• $D_{AB}$ = bond energy of A–B
• $D_{AA}, D_{BB}$​ = bond energies of A–A and B–B

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Key Points

1. Fluorine is highest
   • Value = 4.0 (maximum on scale)
2. Values are relative
   • Not absolute, just comparison
3. Dimensionless
   • No unit

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🔹 Example Values

Element	Electronegativity
F	4.0
O	3.5
N	3.0
Cl	3.0
H	2.1

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Important Applications

1. Predict Bond Type

• ΔEN ≈ 0 → Non-polar covalent
• ΔEN small → Polar covalent
• ΔEN large → Ionic

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2. Predict Polarity

• Larger difference → more polar bond

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Mulliken Scale — Electronegativity (JEE/NEET)

The Mulliken scale was proposed by Robert S. Mulliken.

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Basic Idea

Electronegativity depends on:

✔ Ionisation Enthalpy (IE)
✔ Electron Gain Enthalpy (EGE / Electron Affinity)

So, it considers both:

• Tendency to lose electron (IE)
• Tendency to gain electron (EGE)

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🔹 Formula

$$\chi = \frac{\text{IE} + \text{EGE}}{2}$$(In some books, electron affinity is used instead of EGE)

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🔹 Meaning

• Higher IE → atom doesn’t lose electrons easily
• More negative EGE → atom gains electrons easily

✔ So, higher value = higher electronegativity

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🔹 Key Points

1. Absolute scale
   • Based on measurable energies (unlike Pauling)
2. Applies to isolated atoms
   • Not bond-based
3. Units initially in energy
   • Often converted to dimensionless values

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🔹 Comparison with Pauling Scale

Feature	Mulliken	Pauling
Based on	IE + EGE	Bond energy
Type	Absolute	Relative
Concept	Atomic property	Bond property

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Basic Relation

The two scales are related approximately by:$$\chi_{\text{Pauling}} \approx \frac{\chi_{\text{Mulliken}}}{2.8}$$

🔹 Definition

Electron Gain Enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom.

• Usually negative (energy released)
• More negative = greater tendency to gain electron

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🔹 Representation

$$X(g) + e^- \rightarrow X^-(g)$$

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🔹 General Trends

Across a Period (Left → Right)

EGE becomes more negative

Reason:

• Increase in nuclear charge → stronger attraction for incoming electron

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Down a Group (Top → Bottom)

EGE becomes less negative

Reason:

• Increase in size → electron added far from nucleus → less attraction

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IMPORTANT EXCEPTIONS (Very Important for JEE/NEET)

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❗ Exception 1: Be, Mg (Group 2)

EGE ≈ zero or slightly positive

Reason:

• Stable ns² configuration
• Incoming electron enters higher energy p-orbital

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❗ Exception 2: N (Group 15)

EGE is less negative than expected

Reason:

• Stable half-filled configuration (np³)
• Adding electron causes electron-electron repulsion

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❗ Exception 3: Noble Gases (Group 18)

EGE is positive

Reason:

• Completely filled orbitals → very stable
• Electron must enter new shell → requires energy

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❗ Exception 4: Fluorine vs Chlorine

👉 Cl has more negative EGE than F ❗

❌ Expected: F > Cl
✅ Actual: Cl > F

Reason:

• F is very small → strong electron-electron repulsion in 2p orbital
• Cl has larger size → less repulsion → easier electron addition

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🔹 Important Order Example

Among halogens:
Cl > F > Br > I

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Successive Electron Gain Enthalpy

👉 Adding 2nd electron:$$X^- + e^- \rightarrow X^{2-}$$

❗ Always positive
Reason:

• Electron added to already negative ion → strong repulsion

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🔹 Elements in Group 16

• Oxygen (O)
• Sulfur (S)
• Selenium (Se)
• Tellurium (Te)
• Polonium (Po)

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Increases down the group
O < S < Se < Te

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Why is the 2nd Electron Gain Enthalpy of Oxygen Positive?

Consider the process:$$\text{O}^- (g) + e^- \rightarrow \text{O}^{2-} (g)$$

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🔹 Key Reason: Electron–Electron Repulsion

After gaining one electron:

• Oxygen becomes O⁻ (negatively charged)

Now, adding another electron:

• Incoming electron is repelled by the already negative ion

✔ So, energy must be supplied → EGE becomes positive