Formal Charge- class 11 chemistry

 

Formal Charge-

Formal Charge is a hypothetical charge assigned to an atom in a molecule, assuming that the electrons in a chemical bond are equally shared between atoms, regardless of their electronegativity.

* It helps to identify the most stable Lewis structure for a molecule or ion.

Why is Formal Charge Important?

* Determines the most stable resonance structure.

* Useful in predicting molecular structure and bonding.

 Formal Charge Formula

Where:

V = Number of valence electrons in the free atom

L = Number of lone pair electrons on the atom

B = Number of bonding (shared) electrons

 

 Step-by-Step Method to Calculate Formal Charge

* Draw the correct Lewis structure of the molecule or ion.

* Count valence electrons (V) for each

* Count lone pair electrons (L) on each atom.

* Count bonding electrons (B) shared with other atoms.

* Use the formal charge formula for each atom.

*Sum of all formal charges should equal the overall charge of the molecule/ion.

Example : Formal Charge on Each Atom in NO₃⁻ (Nitrate Ion)

Structure has many  resonance:

Take one resonance form:

 

O = N - O⁻

    |

    O

N:         V = 5, L = 0, B = 8 (4 bonds)

O (double bond):       V = 6, L = 4, B = 4

O (single bond with -1 charge):      V = 6, L = 6, B = 2

Total charge = -1,    matches given.

So, this is stable structure.

 

What is a Covalent Bond?

 

What is a Covalent Bond?

A covalent bond is a type of chemical bond formed by the mutual sharing of electrons between two atoms to achieve stable electronic configuration . (usually octet).

* Covalent bonding generally occurs between non-metallic elements.

* The shared electrons are counted in the valence shell of both atoms.

Properties of Covalent Compounds

Physical State         -   Usually gases or liquids, some solids

Melting/Boiling Point     - Low (except giant covalent structures like diamond)

Electrical Conductivity    - Poor conductors (no free ions/electrons)

Solubility     -  Soluble in non-polar solvents (e.g., benzene)

Important Terms

Bond Pair:  Electron pair involved in bonding.

Lone Pair: Electron pair not involved in bonding.

Electronegativity: Tendency of an atom to attract shared electrons.

Formation of Covalent Bond

Example: Hydrogen Molecule (H₂)

Each hydrogen atom has 1 electron and needs 1 more to complete its duplet.

Two hydrogen atoms share 1 pair of electrons:

H • + • H → H:H or H–H

Example: Oxygen Molecule (O₂)

Each oxygen atom has 6 valence electrons, needs 2 more.

Share 2 pairs of electrons (double bond):

O = O

 

Lewis Structures

* Lewis symbols represent valence electrons as dots.

* A Lewis structure shows:

                   Bond pairs (shared electrons)

                   Lone pairs (unshared electrons)

Example: Water (H₂O)

 

Coordinate (Dative) Covalent Bond

A coordinate covalent bond is a type of covalent bond in which both electrons in the shared pair are donated by only one atom.

* Represented by an arrow (→)

* Still behaves like a covalent bond

Example: Ammonium Ion (NH₄⁺)

NH₃ + H⁺ → NH₄⁺

 

Non-Polar Covalent Bonds

Definition:

A non-polar covalent bond is formed when the electrons are equally shared between two atoms.

* Equal sharing of electrons

* Usually between identical atoms or atoms with very small difference in electronegativity (∆EN ≈ 0)

* No partial charges on atoms

* Molecule is electrically neutral and symmetrical

Examples:

Molecule     Structure        Electronegativity 

                                             Difference (∆EN)

H₂                      H–H                          0

Cl₂                    Cl–Cl                          0

O₂                    O=O                          0

CH₄                 C–H                       0.4 

                       (considered nearly non-polar)

  Polar Covalent Bonds

 Definition:

A polar covalent bond is formed when the electrons are unequally shared between two atoms with a significant difference in electronegativity.

Key Features:

* usually ∆EN between 0.4 and 1.7.

* Unequal sharing of electrons .The atom with higher electronegativity attracts the shared electrons more strongly

* Leads to partial charges:

δ⁻ (partial negative charge) on more electronegative atom

δ⁺ (partial positive charge) on less electronegative atom

* Molecule may have a net dipole moment

Examples:

Definition of Electronegativity

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

* Depends on          -  Atomic size, nuclear charge, 

                                   and shielding

* Units  -                   No units (relative scale)

* Highest value           Fluorine (F) – 4.0

* Lowest                   Cesium (Cs) – 0.7

Factors Affecting Electronegativity

1.  Atomic Size-   Smaller atom → higher EN

Electrons are closer to the nucleus → stronger attraction.

2.  Nuclear Charge (Z)

More protons = stronger pull on electrons = higher EN

3. Electron Shielding / Screening Effect

More inner electrons = more shielding → lower EN

4. Hybridization

Greater s-character → higher EN

EN of sp >  EN of sp² >  EN of sp³

Electronegativity difference (∆EN) to predict bond type:

∆EN Range                  Bond Type

0 – 0.4                           Non-Polar Covalent

0.4 – 1.7                         Polar Covalent

> 1.7                                 Ionic Bond

 

Electronegativity of Common Elements (Pauling Scale)

Element                   Electronegativity

Hydrogen                   H                     2.1

Carbon                        C                      2.5

Nitrogen                     N                     3.0

Oxygen                        O                     3.5

Fluorine                       F                4.0 (highest)

Chlorine                      Cl                    3.0

Bromine                      Br                    2.8

Iodine                          I                       2.5

Sulfur                          S                      2.5

Phosphorus                P                     2.1

Sodium                      Na                    0.9

Magnesium              Mg                   1.2

Aluminum                 Al                    1.5

Silicon                         Si                     1.8

Calcium                     Ca                    1.0

Potassium                  K                      0.8

Lithium                       Li                     1.0

Boron                          B                       2.0

Quick Notes:

* Fluorine (F) has the highest electronegativity: 4.0

* Metals (like Na, K, Ca) generally have low electronegativity

* Non-metals (like O, N, Cl) have high electronegativity

* Noble gases usually don’t have electronegativity values (as they don’t form bonds easily)

Electronegativity Order (for non-metals)

F > O > N ≈ Cl > Br > C ≈ S > H > P

 

What is a Chemical Bond?

 

What is a Chemical Bond?

A chemical bond is a force that holds two or more atoms together in a molecule of compound.

Atoms bond together to achieve a stable electronic configuration, usually the octet (8 electrons in their outermost shell), like the noble gases.

Why Do Atoms Form Bonds?

Atoms form chemical bonds to achieve stability by reducing their potential energy.

Main Assumptions of Kossel-Lewis Theory

(i) Atoms tend to attain octet configuration (8 electrons in valence shell).

Noble gases (except He) have stable octet configurations, which is why they are chemically inert.

(ii) Atoms can achieve octet by:

* Losing electrons → forming positive ions (cations)

* Gaining electrons → forming negative ions (anions)

* Sharing electrons → forming covalent bonds

 

Kossel’s Contribution – Ionic Bonding

Ionic bond formation involves:

* Metal atoms (with 1–3 valence electrons) lose electrons to form cations.

* Non-metal atoms (with 5–7 valence electrons) gain electrons to form anions.

* Resulting oppositely charged ions attract each other to form ionic bonds.

Example: Formation of NaCl

Na (2,8,1) → Na⁺ + e⁻

Cl (2,8,7) + e⁻ → Cl⁻

Na⁺ and Cl⁻ combine to form NaCl via electrostatic attraction.

Lewis’s Contribution – Covalent Bonding

Covalent bond formation involves:

* Atoms sharing valence electrons to attain octet/duplet configuration.

* Shared electron pair belongs to both atoms.

* Lewis represented atoms using Lewis dot symbols:

* Valence electrons shown as dots around the element symbol.

Example: Formation of H₂ molecule

Each H atom has 1 electron (needs 2 for duplet)

They share one pair of electrons: H:H

Each H attains a duplet → stable H₂ molecule

Octet Rule

Atoms tend to gain, lose, or share electrons to attain 8 electrons in their outermost shell.

·        Helium (He) is an exception (only 2 electrons – duplet rule).

Limitations of Octet Rule:

Doesn’t apply to:

* Odd-electron molecules (e.g., NO, NO₂)

* Incomplete octet (e.g., BCl₃)

* Expanded octet (e.g., PCl₅, SF₆)

 

Modes Of Chemical Combination

·        Electrovalent bond or ionic bond

·        Covalent bond

·        Co-ordinate bond

 

Electrovalent Bond / Ionic Bond

Definition

An electrovalent bond is a chemical bond formed by the complete transfer of electrons from one atom to another. It also known as an ionic bond.

* The atom that loses electrons becomes a positively charged ion (cation).

* The atom that gains electrons becomes a negatively charged ion (anion).

* These oppositely charged ions are held together by strong electrostatic force of attraction.

Example: Formation of NaCl (Sodium Chloride)

Step 1: Electron transfer

Na (2,8,1) → Na⁺ + e⁻

Cl (2,8,7) + e⁻ → Cl⁻

Step 2: Formation of ions

Na⁺ (cation), Cl⁻ (anion)

Step 3: Attraction

Na⁺ and Cl⁻ attract each other and form NaCl

Na  + Cl  → Na⁺ + Cl⁻    → NaCl

 

Energy Changes During Ionic Bond Formation

Forming an ionic bond involves multiple energy steps:

(i) Ionization Energy (IE)

Energy required to remove electron(s) from a metal atom.

(ii) Electron Affinity (EA)

Energy released when a non-metal gains electron(s).

(iii) Lattice Energy

Energy released when oppositely charged ions come together to form a solid ionic lattice.

Higher lattice energy = more stable ionic compound

Example: MgO has high lattice energy → very stable compound.

 

Conditions for Formation of Ionic Bond

To form an ionic bond, the following conditions are generally required:

(i) Low Ionization Energy of the metal

So that it can easily lose electrons and form a cation.(e.g., Na, Mg, Ca)

(ii) High Electron Affinity of the non-metal

So that it can easily gain electrons and form an anion. (e.g., Cl, O, F)

(iii) Large Electronegativity Difference

Difference in electronegativity between metal and non-metal should be greater than 1.7.

Factors Affecting Strength of Ionic Bond

(i) Charge on Ions

Greater charge = stronger bond

e.g., MgO (Mg²⁺ and O²⁻) has stronger bonding than NaCl (Na⁺ and Cl⁻)

(ii) Size of Ions

Smaller ions = stronger attraction

e.g., LiF has stronger bonding than CsI

Characteristics of Ionic Compounds

(i) Physical State

Generally exist as crystalline solids.

(ii) Melting and Boiling Points

High due to strong electrostatic forces.

(iii) Solubility

* Soluble in polar solvents like water.

* Insoluble in non-polar solvents like benzene.

(iv) Electrical Conductivity

* Good conductors in molten or aqueous state (ions are free to move).

* Poor conductors in solid state (ions are fixed).

(v) Bond Strength

Strong bonds due to strong attraction between oppositely charged ions.

 Advantages of Ionic Bonding

* Explains the formation and properties of salts.

* Helps understand reactivity of metals and non-metals.

Limitations of Ionic Bonding

* Ionic bonding does not explain the properties of compounds with covalent character.

* Cannot be formed if electronegativity difference is too small.

* Compounds may show partial covalent character (Fajans’ Rule).

 

Fajans’ Rule (Covalent Character in Ionic Compounds)

* It explains why some ionic compounds show covalent character.

e.g  ,AlCl₃ shows covalent nature despite being made of metal and non-metal.

* It is based on the concept of polarization.Example:

Polarization

Polarization is the distortion of the electron cloud of an anion by a cation.

* The greater the polarization, the greater the covalent character in the compound.

Factors affecting Covalent Character (Fajans’ Rule)

(i) Size of the Cation

Smaller cation → greater polarizing power → more covalent character

Example: Al³⁺ (small) > Na⁺ → AlCl₃ more covalent than NaCl

(ii) Charge on the Cation

Higher charge → stronger attraction to electrons → more distortion

Example: Fe³⁺ > Fe²⁺ → FeCl₃ more covalent than FeCl₂

(iii) Size of the Anion

Larger anion → loosely held electrons → easily distorted → more covalent

Example: I⁻ > Br⁻ > Cl⁻ > F⁻

So, NaI > NaBr > NaCl > NaF (in covalent character)

(iv) Charge on the Anion

Higher negative charge → more electrons to polarize → more covalent character

Example: O²⁻ > F⁻

→ MgO is more covalent than MgF₂

 

Factors Affecting Covalent Character

Factor                     Trend for More Covalent Character

Cation size             Smaller

Cation charge        Higher

Anion size               Larger

Anion charge           Higher

 

Examples Based on Fajans’ Rule

Compound   Why It Is More Covalent?          Nature

AlCl₃   Al³⁺ is small and highly charged → strong polarization       More covalent

NaCl   Na⁺ is large, Cl⁻ is small → less polarization    Mostly ionic

BeCl₂  Be²⁺ is very small → high polarizing power    Covalent

MgI₂   I⁻ is large → easily polarized       Partially covalent

FeCl₃   Fe³⁺ more polarizing than Fe²⁺   More covalent

 

Application of Fajans’ Rule

1.Predicts the degree of covalent character in ionic compounds.

2. Explains why:

* Some ionic compounds are insoluble in water.

* Some ionic compounds have low melting points.

* Transition metal halides often show covalent nature.

* AlCl₃ is covalent and exists as a dimer in vapor phase.

Crystalline solids

Crystalline solids are a type of solid material where the atoms, ions, or molecules are arranged in a highly ordered and repeating pattern extending in all three spatial dimensions.

Characteristics of Crystalline Solids:

1. Definite Geometric Shape

Due to the regular arrangement of particles, crystalline solids often form flat surfaces and sharp edges—these are called crystal faces.

2. Long-Range Order

The internal structure is consistent and periodic across large distances.

3. Sharp Melting Point

Crystalline solids melt at a specific temperature, because all bonds of the same type break at once.

4. Cleavage Planes

They tend to break along specific planes where atomic bonding is weaker.

Examples of Crystalline Solids:

* Table salt (NaCl) – Ionic crystal with cubic symmetry

* Diamond – Covalent network crystal with tetrahedral bonding

* Quartz (SiO₂) – Covalent network with repeating tetrahedra

* Copper (Cu) – Metallic crystal with good electrical conductivity

 

Class 11 chemistry Atomic structure

 

Principal Quantum Number (n)

It indicates the main energy level or shell in which an electron is located.

* The principal quantum number is denoted by the symbol n. n can have positive integer values: 1, 2, 3, 4, ...

Shells are often labeled with letters:

n = 1 → K

n = 2 → L

n = 3 → M

n = 4 → N, etc.

Importance of Principal Quantum Number (n)-

( a ) Determines Energy and Distance

Higher n means:

Electron is at a higher energy level

Electron is located farther from the nucleus

As n increases, the size of the orbital increases.

( b ) Determines Maximum Number of Electrons in a Shell-

The maximum number of electrons a shell can hold is given by: 2n²

For n = 1 → 2(1)² = 2 electrons

For n = 2 → 2(2)² = 8 electrons

For n = 3 → 2(3)² = 18 electrons

 

 

Azimuthal Quantum Number ( l ):

It defines the type of subshell in which the orbitals resides. it  is denoted by l.

The Azimuthal Quantum Number, also known as the angular momentum quantum number.

* It Depends on Principal Quantum Number (n)-

For a given value of n, l  can have integer values from 0 to (n − 1).

If n = 3, then ℓ = 0, 1, 2.

Importance of Azimuthal Quantum Number ( l ):

( a ) It determines Subshell/Orbital Type-

Each value of ℓ corresponds to a type of subshell:

ℓ = 0 → s- subshell

ℓ = 1 → p- subshell

ℓ = 2 → d- subshell

ℓ = 3 → f- subshell

( b ) It determines Number of Orbitals per Subshell

number of orbitals in that subshell= 2ℓ + 1

s (ℓ = 0) → 1 orbital

p (ℓ = 1) → 3 orbitals

d (ℓ = 2) → 5 orbitals

f (ℓ = 3) → 7 orbitals

 

( c ) It determines  Orbital Shape-

ℓ = 0 (s) → spherical shape

ℓ = 1 (p) → dumbbell shape

ℓ = 2 (d) → cloverleaf shape

ℓ = 3 (f) → complex shape

( d ) It determines  Orbital Angular Momentum-

Angular momentum of an electron is given by:

√[ℓ(ℓ + 1)] × ħ,

where ħ is the reduced Planck’s constant.

 

It give Subshell Notation-

The combination of n and ℓ gives subshell names:

n = 2, ℓ = 1 → 2p

n = 3, ℓ = 2 → 3d, etc.

(e) It tells Maximum Electrons in a Subshell-

Each type of orbital can hold a maximum number of electrons:

s → 2 electrons

p → 6 electrons

d → 10 electrons

f → 14 electrons

 

Magnetic Quantum Number (m )

The Magnetic Quantum Number describes the orientation of an orbital in space within a given subshell.

It depends on Azimuthal Quantum Number (ℓ)-

For a given value of ℓ, m can have integer values from −ℓ to +ℓ, including 0.

Example: If ℓ = 1 → m = −1, 0, +1

Importance of Magnetic Quantum Number ( m ):

(a) It determines Orbital Orientation:

It Specifies the direction or orientation of the orbital around the nucleus.

For p-orbitals (ℓ = 1):

m = −1 → px

m = 0 → py

m = +1 → pz

Each orbital defined by a value of m can hold up to 2 electrons (with opposite spins).

(b) Affects Electron Behavior in Magnetic Fields

The value of m determines how orbitals are affected when placed in an external magnetic field (basis of Zeeman effect).

 

Spin Quantum Number (s) Definition:

The spin quantum number describes the intrinsic angular momentum (spin) of an electron within an atom.

*It is represented by s

*An electron can have only two possible spin values:

+½ (spin-up),  −½ (spin-down)

*Electron spin resonance and the splitting of spectral lines in magnetic fields (Zeeman effect).

 

What is Hund’s Rule?

Hund’s Rule states:

Electrons fill empty orbitals of the same energy with parallel spins first before pairlng in the same orbital.

In simpler terms:Hund’s Rule tells us:

*Place one electron in each orbital before pairing up.

*Make sure all have parallel spins (same direction).

 

Why is Hund’s Rule Important?

Hund's Rule helps us to:

1. Predict How Electrons Fill Orbitals

It gives specific arrangement of electrons in p, d, and f subshells.

2. Determine Magnetic Properties

Atoms with unpaired electrons exhibit magnetic behavior. Hund's Rule helps us identify which atoms are: Paramagnetic or Diamagnetic

• Paramagnetic: Attracted to a magnetic field (due to unpaired electrons).
• Diamagnetic: Not attracted to a magnetic field (all electrons are paired).

 Example: Nitrogen (N)

Atomic Details:

• Atomic Number: 7
• Number of Electrons: 7
• Electron Configuration: 1s² 2s² 2p³

We focus on the 2p subshell (the highest-energy occupied subshell here), which contains 3 electrons.

Applying Hund’s Rule:

The 2p subshell has three orbitals:

• px, py, and pz

So, the electron filling looks like this:

px: ↑    py: ↑    pz: ↑

Each orbital has one unpaired electron.

---

Magnetic Properties of Nitrogen

Since nitrogen’s 2p electrons are unpaired, the atom is:

• Paramagnetic
• It will be attracted to a magnetic field

This is a direct consequence of Hund’s Rule.

---

  

Aufbau Principle-

The term "Aufbau" is derived from the German word meaning "building up."

The Aufbau Principle states that:

Electrons occupy the lowest energy orbitals first before filling higher energy orbitals.

*This principle helps determine the electron configuration of atoms in their ground state.

Order of Orbital Filling

Electrons are added to atomic orbitals in a specific sequence based on increasing energy levels. The typical order in which orbitals are filled is:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

Exceptions to the Aufbau Principle

There are notable exceptions to the Aufbau Principle, especially among transition metals. These exceptions arise because:

Half-filled (e.g., d⁵) and Fully-filled (e.g., d¹⁰) subshells offer extra stability to atoms.

Examples:

Chromium (Cr):

Expected: [Ar] 3d⁴ 4s²

Actual: [Ar]  3d⁵ 4s¹

 

Copper (Cu):

Expected: [Ar] 3d⁹ 4s²

Actual: [Ar] 3d¹⁰ 4s¹

These exceptions occur because such configurations lower the overall energy of the atom.

 

Pauli Exclusion Principle-

The Pauli Exclusion Principle states:

No two electrons in the same atom can have the same set of four quantum numbers.

* Each electron in an atom is described by four quantum numbers:

n (principal), l (azimuthal), m (magnetic), and m(spin).

* If two electrons are in the same orbital (i.e., have the same values for n, l, and m), then they must have opposite spins (mₛ = +½ and -½).

This principle explains the maximum of 2 electrons per orbital.

 

Example:

In a 1s orbital:

n = 1, l = 0, m = 0

It can hold 2 electrons, but only if:

One electron has spin +½

The other has spin -½