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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

Isoelectronic species are:

Atoms, ions, or molecules that have the same number of electrons.

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🔹 Examples

• N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺

All of these have 10 electrons → so they are isoelectronic

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How to identify?

Just count total electrons

Example:

• O²⁻ → 8 + 2 = 10 electrons
• F⁻ → 9 + 1 = 10 electrons
• Na⁺ → 11 − 1 = 10 electrons

✔ Same electrons → Isoelectronic

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🔹 Important Trend (VERY IMPORTANT for JEE/NEET)

In an isoelectronic series:

Size decreases as nuclear charge (Z) increases

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Example Order of Size

For species with 10 electrons:

N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺

Reason:

• All have same electrons
• But increasing protons (Z) pull electrons closer → size decreases

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🔹 Key Exam Line

“Isoelectronic species have same number of electrons but different nuclear charge.”

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The Van der Waals radius is defined as:

Half of the distance between the nuclei of two non-bonded identical atoms when they are just touching each other.

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Simple Understanding

• When atoms are not chemically bonded (no covalent/ionic bond), they still can come close due to weak intermolecular forces.
• The distance at this closest approach = Van der Waals distance
• So,
  Van der Waals radius = (Van der Waals distance) / 2

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Example

If two neon atoms (not bonded) are 3.2 Å apart:

• Van der Waals radius of neon = 3.2 / 2 = 1.6 Å

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Correct Order of Atomic Radii:

Van der Waals radius > Metallic radius > Covalent radius

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🔹 Why this order?

1. Covalent radius (smallest)
   • Atoms are strongly bonded
   • Nuclei pull shared electrons → atoms come closer
2. Metallic radius (middle)
   • Metal atoms are packed in a lattice
   • Bonding is weaker than covalent → atoms are slightly farther apart
3. Van der Waals radius (largest)
   • No bonding, only weak attraction
   • Atoms stay far apart

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🔹 Comparison Table

Radius Type	Condition	Smaller/Larger
Covalent radius	Bonded atoms	Smaller
Metallic radius	Metal lattice	Medium
Van der Waals	Non-bonded atoms	Largest

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