• Molecule- A group of atoms.
  • Chemical bond- It is an attractive force which hold various constituents (atoms, ions, etc.) together in different chemical species.
  • The formation of chemical compound takes place by combination of various atoms in different ways.

KOSSEL-LEWIS APPROACH TO CHEMICAL BONDING-

  • In 1916, Kossel and Lewis explain successfully that formation of chemical bond in terms of electron. First they provide logical explanation of valance which based on inertness of noble gas.
  • Lewis postulated the atoms achieve the stable octet when they are linked by chemical bonds. Every elements has maximum 8 electrons in outer shell. In case of Na and Cl, they form ions (Na+ and Cl–) and in case of molecule like H2, Cl2, etc. they form bond by sharing the electron (i.e covalent bond)
  • Lewis Symbols-
    • In formation of molecule, only outer shell take part in combination that is electron of valance shell or valance electron
    • Valance Shell- Outermost shell of atoms.
    • Valance Electron- Electron of valance shell.
    • Lewis Symbol- It represent an atom ‘s valence electrons as dots around its chemical symbol, showing its bonding capacity. (Ex- H·, He:, Li., Be:, B· · · , etc.)
      • The number of dots around the chemical symbol represent the number of valance electrons which help to calculate group valence of element.
    • Some important facts-
      • In periodic table, majority electronegative halogens and majority electropositive alkali metals are separated by noble gases.
      • The formation of anion form halogen and cation form alkali atoms is associated with gain and loss of electrons.
      • The negative and positive ion are stable by electrostatic force of attraction.
      • The bond formed by electrostatic force of attraction between positive and negative ion was termed as electrovalent bond.
      • Electrovalence- It is equal to the number of unit charge(s) on the ion.

RULE OF OCTET-

1. Octet Rule

The Octet Rule was introduced by Kossel and Lewis in 1916. It states that atoms combine by transferring or sharing electrons to achieve a stable configuration, often resembling the electron configuration of noble gases. This rule mainly applies to atoms in the second period of the periodic table.

2. Covalent Bond

  • Refinement by Langmuir (1919): Langmuir introduced the concept of the covalent bond, where atoms share electrons instead of transferring them completely.
  • Covalent Bond Formation: A covalent bond is formed when atoms share one or more pairs of electrons. For example, two chlorine atoms share one electron each to form a Cl₂ molecule, satisfying the octet rule for both.
  • Single, Double, and Triple Bonds:
    • Single Bond: Involves sharing one electron pair, e.g., Cl₂.
    • Double Bond: Two electron pairs are shared, e.g., in CO₂ (carbon dioxide).
    • Triple Bond: Three electron pairs are shared, e.g., in N₂ (nitrogen molecule) or C₂H₂ (acetylene).

3. Lewis Dot Structures

  • These represent the bonding in molecules and ions using dots to show the valence electrons.
  • Steps to Draw Lewis Structures:
    1. Count the total valence electrons from the atoms.
    2. Adjust for charges: add one electron per negative charge, subtract one per positive charge.
    3. Arrange the atoms, typically with the least electronegative atom in the center.
    4. Distribute the electrons to form bonds, ensuring each atom attains the octet configuration (or duplet for hydrogen).
    Example: For CO₂, the carbon atom shares two pairs of electrons with each oxygen atom to form double bonds, completing the octet for both.

4. Formal Charge

Formal charge helps to track the distribution of electrons in a molecule. It is calculated as:Formal Charge=Valence electrons−Non-bonding electrons−12×Bonding electrons\text{Formal Charge} = \text{Valence electrons} – \text{Non-bonding electrons} – \frac{1}{2} \times \text{Bonding electrons}Formal Charge=Valence electrons−Non-bonding electrons−21​×Bonding electrons

  • Purpose: Formal charges help to identify the most stable Lewis structure, where the charges are minimized and as close to zero as possible.

5. Limitations of the Octet Rule

  • Incomplete Octet: Some elements like lithium (Li), beryllium (Be), and boron (B) have fewer than eight electrons in their valence shells, e.g., LiCl, BeH₂, BCl₃.
  • Odd-electron Molecules: Some molecules (e.g., NO, NO₂) have an odd number of electrons and cannot follow the octet rule.
  • Expanded Octet: Elements beyond the second period, like phosphorus (P), sulfur (S), and chlorine (Cl), have d-orbitals and can accommodate more than eight electrons in their valence shell, e.g., PF₅, SF₆, H₂SO₄.

6. Drawbacks of the Octet Rule

  • Noble Gas Exceptions: Some noble gases (e.g., xenon, krypton) can form compounds and don’t follow the octet rule.
  • Doesn’t Explain Molecular Shape: The octet rule doesn’t account for the actual three-dimensional shape of molecules, which is explained by theories like VSEPR (Valence Shell Electron Pair Repulsion).
  • Doesn’t Explain Stability: It doesn’t explain why some molecules are more stable than others, as it doesn’t consider molecular energy.

Formation of Ionic Bonds:

  1. Ionic Bond Formation:
    • Ionic bonds form through two main processes:
      1. Ionization: Removing electrons to form positive ions (cation).
      2. Electron Gain: Adding electrons to form negative ions (anion).
    • The energy changes involved in these processes are:
      • Ionization Enthalpy: Energy required to remove an electron from a neutral atom to form a positive ion.
      • Electron Gain Enthalpy: Energy released or absorbed when an atom gains an electron to form a negative ion.
    • Ionic bonds form easily when:
      • The element can easily form a positive ion (low ionization enthalpy).
      • The element can easily form a negative ion (high electron gain enthalpy).
  2. Ionic Compounds:
    • Typically, ionic compounds consist of metallic cations (positive ions) and non-metallic anions (negative ions).
    • The ammonium ion (NH₄⁺) is an exception where both the cation and anion are from non-metals.
  3. Crystal Structure:
    • In the solid state, ionic compounds form a crystal lattice, where the ions are arranged in a regular, repeating pattern.
    • The crystal structure is stabilized by Coulombic interactions (the attraction between positive and negative charges).
  4. Lattice Enthalpy:
    • Lattice Enthalpy is the energy needed to separate one mole of solid ionic compound into its gaseous ions.
    • Example: For NaCl, 788 kJ/mol energy is needed to separate one mole of NaCl into Na⁺ and Cl⁻ ions.
    • The lattice structure stabilizes the ionic compound by releasing energy (negative lattice enthalpy) when the ions come together.

Bond Characteristics:

  1. Bond Length:
    • Bond Length is the distance between the nuclei of two bonded atoms in a molecule.
    • Measured using spectroscopic or X-ray techniques.
    • The covalent radius is half the distance between two identical atoms in a covalent bond.
  2. Bond Angle:
    • Bond Angle is the angle between the orbitals containing bonding electron pairs around a central atom in a molecule or complex ion.
    • It gives insight into the molecular shape and orbital distribution around the central atom.
  3. Bond Enthalpy:
    • Bond Enthalpy is the energy required to break one mole of bonds between two atoms in a gaseous molecule.
    • The larger the bond dissociation enthalpy, the stronger the bond.
    • For example:
      • The bond enthalpy of the H-H bond is 435.8 kJ/mol.
      • The bond enthalpy of the O=O bond is 498 kJ/mol.
    • For polyatomic molecules like water (H₂O), the bond enthalpy can vary due to different chemical environments around each bond.
  4. Average Bond Enthalpy:
    • For polyatomic molecules, the bond enthalpy of similar bonds can differ due to the surrounding chemical environment. The average bond enthalpy is the mean energy required to break a bond in a molecule.
    • Example for water:
      • Breaking the first O-H bond requires 502 kJ/mol.
      • Breaking the second O-H bond requires 427 kJ/mol.

Additional Concepts:

  • Covalent Radius: The size of an atom in a covalent bond is called the covalent radius. It’s the radius of the atom when it forms a bond.
  • Van der Waals Radius: Represents the size of an atom when it’s not bonded, including its valence shell.

Key Takeaways:

  • The formation of ionic compounds depends on the ease of ionization and electron gain for the involved elements, along with the stabilizing effect of the lattice enthalpy.
  • Bond length, bond angle, and bond enthalpy are fundamental properties that define the structure and stability of molecules.
  • In polyatomic molecules, bond enthalpies can vary depending on the molecular environment.

1. Bond Order

  • Bond Order refers to the number of bonds between two atoms in a molecule. It is related to the stability of the molecule.
  • Examples: In H₂ (1 bond), O₂ (2 bonds), and N₂ (3 bonds). The higher the bond order, the stronger the bond.
  • The bond order can also be used to compare molecules. For instance, the bond order of CO (3) is the same as that of N₂ (3).
  • General Rule: As bond order increases, bond enthalpy (strength) increases and bond length decreases.

2. Resonance Structures

  • Sometimes a single Lewis structure doesn’t fully represent a molecule’s structure. For example, ozone (O₃) can be represented by two structures, but both are not accurate on their own.
  • Resonance involves multiple structures (called canonical forms) that together describe the molecule more accurately. The actual structure is called the resonance hybrid.
  • Important Points about Resonance:
    • Canonical forms have no real existence; they are hypothetical.
    • The molecule does not shift between forms; it exists as a hybrid of them.
    • Resonance stabilizes molecules by lowering their energy.

3. Polarity of Bonds

  • Ionic and Covalent Bonds: No bond is purely ionic or covalent. Even covalent bonds like in H₂ have some ionic character.
  • Polar Covalent Bonds: When atoms with different electronegativities form a bond (e.g., in HF), the shared electrons are drawn more toward the more electronegative atom, creating a polar bond with partial charges (dipole).
    • The Dipole Moment (µ) measures the separation of charges and is given by: Dipole moment(µ)=charge (Q)×distance of separation (r)\text{Dipole moment} (µ) = \text{charge (Q)} \times \text{distance of separation (r)}Dipole moment(µ)=charge (Q)×distance of separation (r)
    • Dipole moment is expressed in Debye (D) units.
    • In polyatomic molecules, the dipole moment is the vector sum of individual bond dipoles. For example, in H₂O, the O-H bond dipoles combine to give a net dipole moment.
    • Zero Dipole Moment: In symmetric molecules like CO₂, the dipole moments of individual bonds cancel each other out.

4. Fajans’ Rules (Covalent Character of Ionic Bonds)

  • Covalent character in ionic bonds increases if:
    • The cation is small and the anion is large.
    • The cation has a high charge.
    • Transition metal cations (which have a specific electronic configuration) are more polarizing than alkali/alkaline earth metal cations.

5. VSEPR Theory (Valence Shell Electron Pair Repulsion Theory)

  • VSEPR Theory helps predict the shapes of molecules by considering electron pair repulsions.
  • Key Principles:
    • The shape of a molecule depends on the number of electron pairs (bonded or lone pairs) around the central atom.
    • Electron pairs arrange themselves to minimize repulsion, with lone pairs taking up more space than bonded pairs.
    • Multiple bonds are treated as a single electron pair for geometry prediction.
    • The repulsion between electron pairs is in the order: Lone pair-lone pair > Lone pair-bond pair > Bond pair-bond pair.
  • Molecular Geometry:
    • For molecules with no lone pairs: Linear (AB₂), Trigonal Planar (AB₃), Tetrahedral (AB₄), etc.
    • For molecules with lone pairs, deviations from ideal geometry occur due to increased repulsion of lone pairs.
  • VSEPR Models work well for molecules with p-block elements and provide accurate geometry predictions.

These are the simplified key points, designed to help understand basic concepts for competitive exams and beyond.

Limitations of Lewis Approach and VSEPR Theory:

  • The Lewis approach helps in drawing molecular structures but cannot explain why chemical bonds form, or why bond dissociation enthalpies and bond lengths vary in molecules like H₂ and F₂.
  • The VSEPR theory can predict the shape of molecules, but it lacks a theoretical basis and has limited applications.

Valence Bond (VB) and Molecular Orbital (MO) Theories:

  • To overcome the limitations of previous theories, VB and MO theories were developed based on quantum mechanics.
  • VB theory was introduced by Heitler and London in 1927 and further developed by Pauling. It uses atomic orbitals, electronic configurations, overlap of orbitals, and hybridization to explain bond formation.

Formation of Hydrogen Molecule (H₂):

  • When two hydrogen atoms approach each other, attractive and repulsive forces emerge. The attractive forces between the atoms lead to bond formation, and the repulsive forces push them apart.
  • At a specific distance, the attractive forces outweigh the repulsive forces, and the system reaches a minimum energy state, resulting in a stable H₂ molecule.
  • The energy released during bond formation is called bond enthalpy, and 435.8 kJ is needed to break one mole of H₂ molecules.

Orbital Overlap Concept:

  • A covalent bond forms when atomic orbitals overlap, and electrons pair up. The greater the overlap, the stronger the bond.
  • Covalent bonding occurs by the overlap of electrons in the valence shell of atoms. The overlap can occur with opposite spins of electrons.

Directional Properties of Bonds:

  • Covalent bonds have directional properties due to orbital overlap. In molecules like CH₄, NH₃, and H₂O, the shape and bond angles are determined by the overlap and hybridization of atomic orbitals.

Types of Orbital Overlaps:

  • Sigma (σ) bond: Formed by end-to-end (head-on) overlap of atomic orbitals along the internuclear axis.
    • Examples: s-s, s-p, and p-p overlap.
  • Pi (π) bond: Formed by sideways (side-to-side) overlap of atomic orbitals. It has two charged clouds above and below the internuclear axis.
  • Sigma bonds are generally stronger than pi bonds due to greater overlap.

Hybridization:

  • Hybridization is the process of combining orbitals to form new orbitals of equivalent energy and shape.
  • The number of hybrid orbitals equals the number of atomic orbitals involved.
  • Hybridized orbitals are more effective in bond formation compared to pure orbitals and are oriented to minimize repulsion.
  • Hybridization is crucial for understanding the geometry of molecules (e.g., tetrahedral shape in CH₄, pyramidal in NH₃, bent in H₂O).

Important Features and Conditions for Hybridization:

  • The number of hybrid orbitals is equal to the number of atomic orbitals that undergo hybridization.
  • Hybrid orbitals are of equal energy and shape.
  • Hybrid orbitals form more stable bonds than pure atomic orbitals.
  • The hybrid orbitals are directed in space to minimize repulsion, leading to a stable molecular structure.
  • For hybridization to occur, the orbitals involved should have similar energy levels. Promotion of electrons is not necessary for hybridization, and even fully filled orbitals can take part.

Types of Hybridisation:

Hybridisation involves the mixing of atomic orbitals to form new hybrid orbitals that influence molecular geometry. The primary types are sp, sp², and sp³ hybridisation, each involving different numbers of s and p orbitals.

1. sp Hybridisation:

  • Involves: One s orbital and one p orbital.
  • Formation: Results in two sp hybrid orbitals with 50% s-character and 50% p-character.
  • Geometry: Linear geometry, with a bond angle of 180°.
  • Example: BeCl₂
    • Process: The beryllium (Be) atom promotes an electron from 2s to 2p, forming two sp hybrid orbitals. These orbitals overlap axially with chlorine atoms to form two sigma (σ) bonds.
    • Shape: Linear (180° angle between Be-Cl bonds).

2. sp² Hybridisation:

  • Involves: One s orbital and two p orbitals.
  • Formation: Three sp² hybrid orbitals, each with 33.3% s-character and 66.7% p-character.
  • Geometry: Trigonal planar, with bond angles of 120°.
  • Example: BCl₃
    • Process: The boron atom promotes an electron to form three unpaired electrons. These three orbitals hybridise to form three sp² orbitals, which overlap with chlorine’s p orbitals to form three sigma bonds.
    • Shape: Trigonal planar (120° bond angle between B-Cl bonds).

3. sp³ Hybridisation:

  • Involves: One s orbital and three p orbitals.
  • Formation: Four sp³ hybrid orbitals, each with 25% s-character and 75% p-character.
  • Geometry: Tetrahedral, with bond angles of 109.5°.
  • Example: CH₄ (Methane)
    • Process: Carbon forms four sp³ hybrid orbitals, each overlapping with hydrogen’s 1s orbital to form four sigma bonds.
    • Shape: Tetrahedral (109.5° bond angle between C-H bonds).
  • Other Examples:
    • NH₃ (Ammonia): Nitrogen forms sp³ hybrid orbitals, three of which form bonds with hydrogen, and one contains a lone pair. The bond angle is reduced to 107° due to lone pair-bond pair repulsion, making the shape pyramidal.
    • H₂O (Water): Oxygen undergoes sp³ hybridisation. Two of the hybrid orbitals bond with hydrogen atoms, while the other two contain lone pairs, creating a bent (V-shape) geometry and bond angle of 104.5°.

4. Other Examples of sp, sp², sp³ Hybridisation:

  • sp³ Hybridisation in C₂H₆ (Ethane):
    • Process: Both carbon atoms are sp³ hybridised. One sp³ orbital from each carbon overlaps axially to form a sigma bond (C-C). The other sp³ orbitals form sigma bonds with hydrogen.
    • Bond Lengths: C-C = 154 pm, C-H = 109 pm.
  • sp² Hybridisation in C₂H₄ (Ethene):
    • Process: Each carbon forms three sp² hybrid orbitals. The sp² orbitals overlap to form a sigma bond (C-C), and two are used for C-H bonds. The remaining unhybridised p orbitals form a pi bond between the two carbon atoms.
    • Bond Angles: H-C-H = 117.6°, H-C-C = 121°.
  • sp Hybridisation in C₂H₂ (Ethyne):
    • Process: Each carbon undergoes sp hybridisation, forming a sigma bond between the carbon atoms and two pi bonds using unhybridised p orbitals.
    • Bond Formation: The triple bond between the carbon atoms consists of one sigma bond and two pi bonds.

5. Hybridisation Involving d Orbitals:

  • Elements with d Orbitals: Elements in the third period (and beyond) have d orbitals that can mix with s and p orbitals to form hybrid orbitals.
  • Common Hybridisations:
    • sp³d (Trigonal Bipyramidal): Example: PF₅, PCl₅.
    • sp³d² (Octahedral): Example: SF₆.
  • Example – PCl₅ (sp³d Hybridisation):
    • Process: Phosphorus undergoes sp³d hybridisation with one s, three p, and one d orbital, resulting in five sp³d hybrid orbitals arranged in a trigonal bipyramidal geometry.
    • Bond Angles: Equatorial bonds (120°), Axial bonds (90°).
  • Example – SF₆ (sp³d² Hybridisation):
    • Process: Sulfur undergoes sp³d² hybridisation using one s, three p, and two d orbitals, forming six sp³d² hybrid orbitals that overlap with fluorine’s orbitals to form six sigma bonds.
    • Shape: Octahedral.

6. Molecular Orbitals (MO) Theory:

  • Electrons in Molecules: Electrons are present in molecular orbitals (MOs), just as in atomic orbitals (AOs).
  • Formation of MOs: When atomic orbitals of comparable energies combine, they form molecular orbitals (one bonding and one antibonding).
  • Bonding MOs: Lower in energy, leading to greater stability.
  • Antibonding MOs: Higher in energy, leading to instability.
  • Aufbau Principle: MOs fill according to the same principles as atomic orbitals, following the Pauli Exclusion Principle and Hund’s Rule.
  • Polycentric Nature of MOs: Unlike atomic orbitals (monocentric), MOs are influenced by multiple nuclei.

Linear Combination of Atomic Orbitals (LCAO)

  • Atomic orbitals (wave functions, ψ) represent the amplitude of electron waves.
  • Schrödinger’s wave equation is used to describe the wave functions, but it’s complex for multi-electron systems.
  • To solve this, LCAO is used to approximate molecular orbitals by combining atomic orbitals (AOs).

2. Formation of Molecular Orbitals (MOs)

  • Example: In a hydrogen molecule (H₂), each hydrogen atom has one electron in a 1s orbital.
  • The atomic orbitals combine to form molecular orbitals through addition (bonding MO) and subtraction (antibonding MO):
    • σ = ψₐ + ψᵦ (bonding MO)
    • σ = ψₐ – ψᵦ* (antibonding MO)
  • The bonding molecular orbital stabilizes the molecule, while the antibonding orbital destabilizes it.
  • Electrons in bonding orbitals lower the energy, while electrons in antibonding orbitals raise the energy.

3. Constructive and Destructive Interference

  • In bonding MOs, electron waves reinforce each other (constructive interference), creating electron density between the nuclei.
  • In antibonding MOs, destructive interference causes electron density to be located away from the nucleus, increasing repulsion.

4. Conditions for Combining Atomic Orbitals

The combination of atomic orbitals to form molecular orbitals depends on:

  1. Similar energy levels: Atomic orbitals should have nearly the same energy.
  2. Same symmetry: Orbitals must have the same symmetry about the molecular axis.
  3. Maximum overlap: Greater overlap between atomic orbitals leads to greater electron density between nuclei.

5. Types of Molecular Orbitals

Molecular orbitals are classified as:

  • σ (sigma) orbitals: Symmetrical around the bond axis.
  • π (pi) orbitals: Non-symmetrical around the bond axis.
  • δ (delta) orbitals: Less common in simpler molecules.

Example:

  • The combination of 1s orbitals results in σ1s (bonding) and σ*1s (antibonding).
  • The combination of 2pz orbitals results in σ2pz and σ*2pz.
  • The combination of 2px and 2py orbitals results in π and π orbitals*.

6. Energy Level Diagram for Molecular Orbitals

Energy levels for various molecular orbitals are:

  • Bonding MOs: σ2s, σ2pz, π2px, π2py.
  • Antibonding MOs: σ2s, σ2pz, π2px, π2py.
  • For molecules like O₂ and F₂, the energy sequence is:
    σ1s < σ1s < σ2s < σ2s < σ2pz < (π2px = π2py) < (π2px = π2py) < σ*2pz.

7. Electronic Configuration and Molecular Behavior

  • The distribution of electrons in molecular orbitals is called electronic configuration.
  • Stability of Molecules: The molecule is stable if the number of electrons in bonding orbitals (Nb) is greater than those in antibonding orbitals (Na).
  • Bond Order:
    Bond Order=12(Nb−Na)\text{Bond Order} = \frac{1}{2} (Nb – Na)Bond Order=21​(Nb−Na) A positive bond order indicates stability, while zero or negative bond order indicates instability.
  • Magnetic Properties:
    • Diamagnetic: All electrons are paired.
    • Paramagnetic: One or more electrons are unpaired, attracted by a magnetic field.

8. Example: Hydrogen Molecule (H₂)

  • The H₂ molecule has 2 electrons in σ1s orbital.
  • Bond order = ½(2 – 0) = 1, meaning a single bond.
  • H₂ is diamagnetic because all electrons are paired.

9. Types of Bonds in Diatomic Molecules

  • Bond Order: Corresponds to the number of bonds (single, double, triple).
  • Bond Length: Shorter bond lengths correspond to higher bond orders.
  • Bond Energy: The energy needed to break the bond.

10. Hydrogen Bonding

  • Hydrogen bonds occur when a hydrogen atom is bonded to an electronegative atom (F, O, or N), creating a partial positive charge on H and a partial negative charge on X.
  • These bonds are weaker than covalent bonds but play a crucial role in the physical properties of substances.

11. Types of Hydrogen Bonds

  • Intermolecular hydrogen bond: Between different molecules (e.g., water, HF).
  • Intramolecular hydrogen bond: Within the same molecule (e.g., o-nitrophenol).

Summary

The study of molecular orbitals and bonding provides insights into molecular stability, bond order, magnetic properties, and the role of hydrogen bonding in molecular interactions. Understanding the conditions for combining atomic orbitals and the symmetry of orbitals is crucial for predicting the behavior of molecules in chemical reactions and their physical properties.

These all are the notes of chapter 4 chemistry. And after some time you get important questions and NCERT solutions HERE. *#THANKS FOR VISITING, VISIT AGAIN#* 😊