Author: Sarvan Kumar

Sarvan Kumar http://sarvankumar.wordpress.com My name is Sarvan Kumar, I am post graduate in chemistry (MSc). besides teaching chemistry at various schools and coaching centres I have been giving home tuitions to students for 10 years. I have helped students score good marks in chemistry not only in board examination but also JEE, NEET, SAT, IGCSE and IB examinations. Nearly 90% of my students have scored more than 95% in their CBSE board examination. Moreover they have also secured a seat in prestigious engineering and medical college. Since every student is different hence my unique style of teach

Leaving Group Ability Order

Leaving Group Ability (Best → Worst)

$\boxed{ \text{Sulfonium} > \text{Quaternary ammonium} > \text{Sulfonate (OTs/OMs)} > I^- > Br^- > Cl^- > F^- > H_2O > ROH > OH^- > OR^- > NH_2^- }$

Sulfonate ester

Sulfonium ion structure

Why HO⁻ and RO⁻ are Poor Leaving Groups

HO⁻ (hydroxide ion) and RO⁻ (alkoxide ion) are:
- Strong bases
- High-energy species
A good leaving group should be:
- Stable after leaving
- Weak base

Since HO⁻ and RO⁻ are strong bases, they are unstable when free, so they do not want to leave.

2. Comparison with Alkyl Halides

In alkyl halides (R–X), the leaving group is X⁻ (Cl⁻, Br⁻, I⁻):
- These are weak bases
- Highly stable ions

Consequence in Reactions

Alcohols (R–OH) and ethers (R–O–R’):
- Do not undergo SN1/SN2/E1/E2 easily
- Because –OH or –OR cannot leave directly

So they are less reactive than alkyl halides.

4. Activation of Alcohols & Ethers

To make them reactive, we convert the poor leaving group into a good one.

Common activation methods:

(A) Protonation (Acidic Medium)

In presence of acid (like H₂SO₄ or HCl):

$R-OH + H^+ \rightarrow R-OH_2^+$

Now leaving group = H₂O (water)
Water is neutral and stable → good leaving group

Conversion to Better Leaving Groups

Convert –OH into:
- Tosylate (–OTs)
- Mesylate (–OMs)

These are excellent leaving groups.

$R-OH + HCl \rightarrow R-Cl + H_2O$

Summary (Exam Ready)

HO⁻ and RO⁻ = strong bases → poor leaving groups
Alkyl halides = weak base leaving groups → high reactivity
Alcohols/ethers = less reactive
Need activation (protonation or conversion) before reaction

How Niels Bohr found radius of hydrogen-like atom

According to Bohr, electrons move in stationary orbits where they do not radiate energy, hence energy of an orbit remains constant with time.

Bohr assumed stationary orbits (not proved mathematically) This explains atomic stability Energy is quantized and constant in each orbit.

Electrostatic force = Centripetal force

Step 1: Electron is attracted to nucleus, so: $\frac{mv^2}{r} = \frac{1}{4\pi \varepsilon_0} \cdot \frac{Z e^2}{r^2}$

Step 2 : Quantization of angular momentum

$mvr = \frac{nh}{2\pi}$

Step 3: Substitute v into Step 1

$m \left( \frac{nh}{2\pi mr} \right)^2 = \frac{1}{4\pi \varepsilon_0} \cdot \frac{Z e^2}{r}$

Step 4: Simplify

$\frac{n^2 h^2}{4\pi^2 m r^2} = \frac{1}{4\pi \varepsilon_0} \cdot \frac{Z e^2}{r}$

Step 5: Solve for radius r

$r = \frac{4\pi \varepsilon_0 n^2 h^2}{4\pi^2 m Z e^2}$

Simplify: $r = \frac{\varepsilon_0 h^2}{\pi m e^2} \cdot \frac{n^2}{Z}$

Final Result

$r_n = \frac{n^2 a_0}{Z}$

where, $a_0 = \frac{\varepsilon_0 h^2}{\pi m e^2} = 0.529 \, \text{Å}$

Key Results

$r_n \propto n^2$
$r_n \propto \frac{1}{Z}$
Valid for hydrogen-like species (H, He⁺, Li²⁺)

Bohr model, hydrogen-like atom, radius derivation, angular momentum quantization, electrostatic force

The limiting line (or limit) of a spectral series (like Lyman, Balmer..)

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

🔹 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).

🔹 Formula (Rydberg equation)

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

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

Examples of limits

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

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

Lyman Alpha (α) Line

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

Period 2 elements — Anomalous Behaviour

Elements: Li, Be, B, C, N, O, F

These show properties different from the rest of their groups.

Main Reasons

Very small atomic size
High ionization enthalpy
High electronegativity
No d-orbitals available (cannot expand octet)
High charge density / polarising power

Important Examples (Exam Focus)

Li (vs other Group 1)

Forms Li₃N (others don’t easily)
LiCl shows covalent character
Li₂CO₃ decomposes on heating

Be (vs other Group 2)

Compounds are covalent
BeO, Be(OH)₂ are amphoteric
Strong complex formation

B (vs other Group 13)

Always covalent
Does not form B³⁺ ion easily
Forms electron-deficient compounds (e.g., BF₃)

🔹 C, N, O, F

Show strong pπ–pπ bonding
Form multiple bonds (C=C, C≡C, N≡N, etc.)
Higher electronegativity than heavier elements

Diagonal relationship and its Reason

Diagonal relationship is said to exist between certain pairs of diagonally adjacent elements in the second and the third period of the periodic table

Reason: Due to similar charge density (charge/size ratio) of diagonal elements.

Li and Mg Similarities

Be and Al Similarities

Acidic Behaviour Across a Period (Left → Right)

Acidity increases across the period and down the grouop

✔️ Example

CH₄ < NH₃ < H₂O < HF

Electronegativity Increases

2. SiH₄ < PH₃ < H₂S < HCl

3. HF < HCl < HBr < HI

Hydration Enthalpy (ΔHₕyd) JEE/NEET Concepts

Definition:

Hydration enthalpy is the energy released when 1 mole of gaseous ions gets surrounded by water molecules.

Always negative (exothermic)

Example:

Na⁺(g) → Na⁺(aq) + energy released

Factors Affecting Hydration Enthalpy

1. Size of Ion

Smaller ion → higher hydration enthalpy

✔️ Reason:

Higher charge density
Strong attraction with water

Order:
Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

2. Charge on Ion

Higher charge → higher hydration enthalpy

✔️ Example:
Mg²⁺ > Na⁺

Group Trends

Group 1 elements

Decreases down the group
Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

Group 2 elements

Decreases down the group
Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺

Comparison (Very Important)

Group 2 > Group 1 (hydration enthalpy)

✔️ Example:
Mg²⁺ > Na⁺

👉 Reason:

Higher charge (+2)
Stronger attraction with water

Applications

✔️ 1. Solubility

Higher hydration enthalpy → more soluble

✔️ 2. Stability of Ions in Solution

Small, highly charged ions are more stable in water

Melting point (MP) & boiling point (BP) trends

1.Melting / Boiling Point Of Group 1 elements

Decreases down the group

Order (MP/BP):
Li > Na > K > Rb > Cs

Bigger atoms = weaker metallic bond = lower MP/BP”

2. Be >Ca > Sr > Ba> Mg (irregular pattern M.P)

Be >Ba >Ca> Sr>Mg ( B.P)

Melting point (MP) and boiling point (BP) of Group 2 elements are higher than the corresponding Group 1 element

Fagan’s Rule (Polarisation Concept)

Definition:

Fajans’ Rule predicts whether a bond will be more ionic or covalent based on polarisation of ions.

More polarisation ⇒ more covalent character

What is Polarisation?

Cation distorts the electron cloud of the anion
This distortion = polarisation

Factors Affecting Polarisation

1. Size of Cation

Smaller cation → more polarising power → more covalent

✔️ Example:

Li⁺ > Na⁺ > K⁺ (in polarising power)

2. Charge on Cation

Higher charge → more polarisation

✔️ Example:

Al³⁺ > Mg²⁺ > Na⁺

3. Size of Anion

Larger anion → more easily polarised → more covalent

✔️ Example:

I⁻ > Br⁻ > Cl⁻ > F⁻

4. Charge on Anion

Higher negative charge → more polarisation

✔️ Example:

N³⁻ > O²⁻ > F⁻

5. Electronic Configuration (Important Exception Concept)

Cations with pseudo noble gas configuration (d¹⁰) show more polarisation than noble gas config

✔️ Example:

Cu⁺ (d¹⁰) > Na⁺ (noble gas)

✔️ 1. Compare Covalent Character:

LiCl > NaCl > KCl (LiCl most covalent)

✔️ 2. Melting Point Trend:

More covalent → lower melting point

✔️ Example:

AlCl₃ (covalent, low MP)
NaCl (ionic, high MP)

✔️ 3. Solubility:

Ionic → soluble in water
Covalent → soluble in organic solvents

Metallic & non-metallic trends in the periodic table

Metallic Character (Electropositive Nature)

Trend:

Down a group: ↑ increases
Across a period (L → R): ↓ decreases

Reason:

Down the group → atomic size ↑ → valence electrons easily lost
Across period → effective nuclear charge ↑ → electrons tightly held

✔️ Order Example:

Group 1: Li < Na < K < Rb < Cs
Period 3: Na > Mg > Al > Si > P > S > Cl

Exceptions in Metallic Character

1. Be vs Al

Normally metallic character ↓ across period
But: Al > Be (unexpected)

Reason:

Be has very high ionization energy
Al can lose electron easily (3p¹ electron is less tightly held)

2. Ga vs Al

Expected: Ga > Al (down group)
Actual: Al > Ga

Reason:

Poor shielding by d-electrons in Ga → higher effective nuclear charge → harder to lose electron

Non-Metallic Character (Electronegativity / Electron Gain)

Trend:

Down a group: ↓ decreases
Across a period (L → R): ↑ increases

Reason:

Across period → size ↓, attraction for electrons ↑
Down group → size ↑, attraction ↓

✔️ Order Example:

Group 17: F > Cl > Br > I
Period 2: Li < Be < B < C < N < O < F