CHEMICAL BONDING: CHAPTER 4: CLASS 11

 

The Nature of Chemical Bonds and Theories of Chemical Bonding

The Nature of Chemical Bonds and Theories of Chemical Bonding

1. Introduction to Matter and Elements:
• Matter consists of distinct elements.
• Under normal conditions, elements exist as independent atoms except for noble gases.
2. Formation of Molecules:
• Atoms group together to form molecules.
• Molecules are collections of atoms with characteristic properties.
• A molecule is held together by a force known as a chemical bond.
3. The Concept of Chemical Bonds:
• Chemical bonds are attractive forces between atoms, ions, etc.
• Chemical bonds are responsible for holding together the constituents of different chemical species.
4. Questions about Chemical Bonding:
• The process of forming chemical compounds from combinations of atoms raises questions.
• Why do atoms combine? Why are specific combinations possible?
• What determines why certain atoms combine while others do not?
• Why do molecules have definite shapes?
5. Theories and Concepts of Chemical Bonding:
• Kössel-Lewis Approach:
• A theory explaining the transfer of electrons between atoms to achieve stable electron configurations.
• Focuses on achieving noble gas electron configurations through electron transfer.
• Valence Shell Electron Pair Repulsion (VSEPR) Theory:
• Describes molecular shapes based on the repulsion between valence electron pairs.
• Explains the three-dimensional arrangement of atoms in molecules.
• Valence Bond (VB) Theory:
• Explains chemical bonding in terms of overlapping atomic orbitals.
• Emphasizes the role of unpaired electrons in the formation of bonds.
• Molecular Orbital (MO) Theory:
• Describes chemical bonding using molecular orbitals formed by the combination of atomic orbitals.
• Electrons are treated as wave-like entities, leading to a more comprehensive understanding of bonding.
6. Relation to Atomic Structure and Periodic Table:
• Development of bonding theories closely linked to advancements in atomic structure understanding.
• Electronic configurations of elements and their placement in the periodic table influence bonding behavior.
7. Stability and Bonding:
• Systems tend to achieve stability.
• Bonding is a natural process that reduces the energy of a system, leading to increased stability.

The nature of chemical bonds and the theories of chemical bonding have evolved over time to address fundamental questions about atomic combinations, molecule formation, shapes, and stability. The theories, including Kössel-Lewis, VSEPR, Valence Bond, and Molecular Orbital theories, provide different perspectives on the forces that hold atoms together in molecules, considering electron configurations, atomic orbitals, and molecular shapes. These theories have been influenced by our understanding of atomic structure and the periodic table, ultimately explaining why atoms combine, the possible combinations, and the shapes of molecules. Bonding serves as nature's mechanism to achieve stability within systems.

 

KOSSEL-LEWIS APPROACH TO CHEMICAL BONDING

Kössel-Lewis Approach to Chemical Bonding

1. Introduction:
• In 1916, independent efforts by Kössel and Lewis brought about a satisfactory explanation for chemical bonding based on electron interactions.
• Their approach was grounded in understanding valence and drew inspiration from the inert properties of noble gases.
2. Lewis's Model of the Atom:
• Lewis conceptualized atoms as having a "Kernel" comprising the nucleus and inner electrons, surrounded by an outer shell.
• The outer shell could accommodate up to eight electrons, distributed at the corners of a cube that enveloped the "Kernel."
• This arrangement formed a stable octet of electrons, promoting stability in the atom.
3. Stable Octet and Chemical Bonds:
• Lewis proposed that atoms attain stability by achieving the octet electron configuration.
• For atoms to achieve this stable state, they form chemical bonds.
• In the example of sodium (Na) and chlorine (Cl), sodium donates an electron to chlorine, resulting in the formation of Na⁺ and Cl^– ions.
• The transfer of an electron from one atom to another leads to the formation of ionic bonds.
4. Covalent Bonds and Octet Rule:
• In molecules like Cl₂, H₂, F₂, etc., atoms form bonds by sharing pairs of electrons.
• This sharing of electrons enables each atom to complete its outer shell and attain the stable octet configuration.
• The concept of the octet rule underlines the tendency of atoms to seek eight electrons in their outer shell, mirroring noble gas electron configurations.
5. Lewis Symbols:
• In molecular formation, only the outer shell electrons (valence electrons) partake in chemical bonding.
• Inner shell electrons are shielded and usually remain uninvolved in bonding.
• G.N. Lewis, an American chemist, introduced a simplified notation to represent valence electrons in atoms, known as Lewis symbols.

Advantages and Insights of Kössel-Lewis Approach:

• Explanation of Valence: Kössel and Lewis provided a logical explanation for valence based on the stable octet configuration.
• Inertness of Noble Gases: The concept of achieving noble gas-like electron configurations explained why noble gases are chemically inert.
• Ionic and Covalent Bonds: The approach distinguished between ionic and covalent bonds, elucidating how electron transfer and sharing contribute to bond formation.
• Octet Rule: The octet rule became a guiding principle for understanding chemical behavior and predicting molecular stability.

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The Kössel-Lewis approach to chemical bonding, developed by Kössel and Lewis, introduced a model of atoms based on the outer shell's electron arrangement. This approach elucidated the significance of achieving a stable octet configuration for atoms through chemical bonding. The concept of Lewis symbols simplified the representation of valence electrons. The approach's insights into ionic and covalent bonding, the octet rule, and the inertness of noble gases contributed significantly to the understanding of chemical interactions and molecular stability.

 

SIGNIFICANCE OF LEWIS SYMBOLS

Significance of Lewis Symbols and Kössel's Contributions to Chemical Bonding

1. Lewis Symbols and Valence Electrons:
• Lewis symbols represent an element's valence electrons as dots around its chemical symbol.
• The number of dots in a Lewis symbol corresponds to the number of valence electrons.
• This representation aids in calculating the common or group valence of the element.
2. Group Valence Calculation:
• Group valence of an element is often either equal to the number of dots in its Lewis symbol or 8 minus the number of dots (valence electrons).
3. Kössel's Observations in Chemical Bonding:
• Kössel's contributions to chemical bonding shed light on critical observations:
• The periodic table separates highly electronegative halogens and highly electropositive alkali metals with noble gases in between.
• Formation of negative ions from halogen atoms and positive ions from alkali metal atoms involves electron loss and gain, respectively.
• Negative and positive ions formed acquire stable noble gas electronic configurations, especially the octet (eight electrons) in the outer shell (ns²np⁶).
• The stability of noble gas configurations implies that ions aim to achieve similar electron arrangements.
4. Stabilization through Electrostatic Attraction:
• The interaction between positive and negative ions formed due to electron transfer results in electrostatic attraction.
• This type of bond was termed the "electrovalent bond."
• Electrovalence corresponds to the number of unit charges on the ion. For instance, calcium carries a positive electrovalence of two, while chlorine bears a negative electrovalence of one.
5. Implications and Applications of Kössel's Postulations:
• Kössel's ideas laid the groundwork for modern concepts related to ion formation through electron transfer and the creation of ionic crystalline compounds.
• These insights contributed significantly to understanding and systematizing ionic compounds.
• While Kössel's concepts were valuable, they also recognized that certain compounds deviated from these ideas.

Key Contributions and Insights:

• Electronegativity and Electropositivity: Kössel highlighted the contrasting properties of halogens and alkali metals, separated by noble gases in the periodic table.
• Ionic Formation: Kössel's observations explained the electron gain and loss during ion formation, leading to stable noble gas-like configurations.
• Stable Outer Shell: Kössel emphasized that negative and positive ions achieve stability by acquiring noble gas outer shell electron configurations.
• Electrovalent Bonds: The electrostatic attraction between oppositely charged ions was termed as electrovalent bonding, with electrovalence representing ion charge.
• Ionic Compounds Understanding: Kössel's ideas significantly enhanced the comprehension and organization of ionic compounds.

Kössel's contributions revolutionized the understanding of chemical bonding, emphasizing the connection between electron transfer and ion formation. His insights into electronegativity, electron gain and loss, stable electron configurations, and electrovalent bonding laid the foundation for modern concepts. While his ideas provided a profound understanding of ionic compounds, they also acknowledged that certain compounds didn't conform to these concepts, reflecting the evolving nature of chemical understanding.

 

THE OCTET RULE IN CHEMICAL BONDING

The Octet Rule in Chemical Bonding

1. Introduction:
• In 1916, Kössel and Lewis introduced the electronic theory of chemical bonding.
• This theory explains how atoms combine through the transfer or sharing of valence electrons.
2. Fundamental Idea - Octet Rule:
• Atoms can attain stability by having a complete outer electron shell.
• The octet rule states that atoms tend to combine in a way that allows them to achieve a stable configuration of eight electrons in their valence shell, resembling the noble gases' electronic configuration.
3. Two Modes of Combination:
• Electron Transfer (Ionic Bonding):
• Involves the transfer of valence electrons from one atom to another.
• One atom gains electrons to fill its valence shell, becoming negatively charged (anion), while the other loses electrons and becomes positively charged (cation).
• The electrostatic attraction between oppositely charged ions leads to the formation of ionic compounds.
• Electron Sharing (Covalent Bonding):
• Atoms share pairs of electrons to complete their valence shells.
• By sharing electrons, each atom achieves a stable configuration similar to noble gases.
• Covalent bonding is common in molecular compounds and forms when atoms have similar electronegativities.
4. Significance of the Octet Rule:
• The octet rule guides the formation of various chemical compounds by dictating how atoms will interact to achieve stability.
• It explains why atoms either gain, lose, or share electrons during bonding.
5. Stability and Noble Gas Configuration:
• Noble gases possess a stable electron configuration with eight electrons in their valence shell (except helium, which has two).
• Other elements aim to emulate this stable state by following the octet rule during chemical bonding.
6. Predictive Power:
• The octet rule aids in predicting the types of bonds that will form between different elements.
• It also provides insights into the properties of resulting compounds.
7. Limitations and Exceptions:
• While the octet rule is a useful guideline, it doesn't explain every type of chemical bonding.
• Some molecules and compounds don't strictly adhere to the octet rule, especially for elements with d or f orbitals that can accommodate more than eight electrons.

Implications of the Octet Rule:

• Ionic and Covalent Bond Types: The octet rule underlies the fundamental distinction between ionic and covalent bonds based on electron transfer and sharing, respectively.
• Stability and Noble Gas Mimicry: The pursuit of the octet configuration drives atoms towards enhanced stability by resembling the noble gas electron arrangements.
• Predictive Tool: The octet rule is an invaluable predictive tool, aiding in explaining the behavior of various elements during bonding.
• Beyond Octet Rule: While essential, the octet rule has limitations, especially for elements with more complex electron configurations.

The octet rule, developed by Kössel and Lewis, forms the cornerstone of chemical bonding theories. It elucidates how atoms combine by either transferring or sharing valence electrons to achieve a stable configuration with eight electrons in their valence shell. While it provides a strong basis for understanding most bonding scenarios, exceptions exist due to the intricate electron arrangements of certain elements.

 

COVALENT BONDING AND LEWIS-LANGMUIR THEORY

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Covalent Bonding and Lewis-Langmuir Theory

1. Introduction:
• Langmuir (1919) built upon Lewis's ideas, refining the concept of chemical bonding.
• Langmuir introduced the term "covalent bond" and expanded on Lewis's octet rule.
2. Evolution from Lewis to Lewis-Langmuir Theory:
• Lewis's idea of atoms bonding through electron sharing was enhanced by Langmuir.
• Langmuir abandoned the concept of a fixed cubic arrangement of the octet and introduced the term "covalent bond."
3. Chlorine Molecule Example - Cl₂:
• The Cl atom has the electronic configuration [Ne]3s²3p⁵, lacking one electron for the argon configuration.
• Cl₂ formation involves the sharing of a pair of electrons between two chlorine atoms.
• Each chlorine atom contributes one electron to the shared pair.
• The result is that both chlorine atoms achieve the outer shell octet of argon, leading to a stable configuration.
4. Representation of Covalent Bonds - Lewis Dot Structures:
• Lewis dot structures use dots to represent electrons in atoms and molecules.
• These structures provide visual insights into the arrangement of shared electrons.
• Lewis dot structures are applicable to different molecules with identical or different combining atoms.
5. Key Conditions of Covalent Bonding:
• Electron Pair Sharing: Covalent bonds form by sharing an electron pair between atoms.
• Contribution of Electrons: Each atom involved contributes at least one electron to the shared pair.
• Achievement of Stable Configuration: Combining atoms achieve noble gas configurations due to shared electrons.
6. Examples of Covalent Bonds:
• Water (H₂O) and Carbon Tetrachloride (CCl₄): In these molecules, atoms share electron pairs to form covalent bonds, fulfilling the conditions of electron sharing and noble gas configurations.
7. Multiple Bonds:
• Single Covalent Bond: Formed when two atoms share one electron pair.
• Double Bond: Two pairs of electrons are shared between atoms. Example: Carbon dioxide (CO₂).
• Triple Bond: Three pairs of electrons are shared between atoms. Examples: Nitrogen gas (N₂), ethyne (C₂H₂).

Implications of the Lewis-Langmuir Theory:

• Advancement of Covalent Bond Concept: Langmuir's covalent bond concept refined Lewis's ideas, focusing on electron sharing and arrangement.
• Visualization through Lewis Dot Structures: Lewis dot structures provide a visual representation of electron distribution in molecules.
• Universal Application: The Lewis-Langmuir theory's principles apply to various compounds, guiding our understanding of covalent bonding.
• Explanation of Multiple Bonds: The theory explains the formation of single, double, and triple bonds through shared electron pairs.

Langmuir's extension of Lewis's ideas gave rise to the Lewis-Langmuir theory of covalent bonding. This theory emphasizes electron sharing between atoms to achieve stable electron configurations, visualized through Lewis dot structures. Examples like the chlorine molecule showcase how covalent bonds lead to noble gas-like outer shell configurations. Additionally, the theory accounts for multiple bonds, elucidating the formation of double and triple bonds through shared electron pairs.

 

The Nature of Chemical Bonds and Theories of Chemical Bonding

Lewis Representation of Simple Molecules (Lewis Structures)

1. Introduction to Lewis Dot Structures:
• Lewis dot structures visually depict molecular and ionic bonding through shared electron pairs and adherence to the octet rule.
• While not exhaustive, they greatly aid in understanding molecule formation and properties.
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2. Importance and Utility:
• Lewis dot structures contribute to understanding bonding patterns and properties of molecules, albeit not providing a complete explanation.
• They are valuable tools in explaining and predicting molecular behavior.
3. Steps to Write Lewis Dot Structures:
• Determine Total Electrons: Calculate the total electrons required for the structure by summing the valence electrons of combining atoms.
• Account for Charges: For anions, add electrons equal to the negative charge, while for cations, subtract electrons equal to the positive charge.
• Distribute Electrons: Using chemical symbols of the atoms and knowledge of the compound's skeletal structure, distribute electrons as shared pairs proportionally among atoms.
• Central Atom Placement: Typically, the least electronegative atom occupies the central position within the molecule or ion.
• Terminal Atom Placement: More electronegative atoms usually occupy the terminal positions.
• Multiple Bonds and Octet Rule: Distribute remaining electron pairs after accounting for single bonds. These pairs may contribute to multiple bonds or remain as lone pairs, ensuring each bonded atom achieves an octet of electrons.
4. Illustrative Examples:
• CH4 Molecule: In methane (CH₄), eight valence electrons are available (4 from carbon, 4 from hydrogen) for bonding.
• Anions and Cations: Anions gain electrons according to their negative charge, while cations lose electrons based on their positive charge.
• Central Atom Position: Generally, the least electronegative atom becomes the central atom. In compounds like NF₃ and CO₃²⁻, nitrogen and carbon serve as central atoms, while fluorine and oxygen occupy terminal positions.
• Octet Rule and Electron Distribution: The distribution of electrons in Lewis structures ensures that each atom achieves an octet of electrons, promoting stability.

Significance and Implications:

• Visual Bonding Representation: Lewis dot structures visually convey the bonding arrangement in molecules and ions.
• Octet Rule Reinforcement: Lewis structures emphasize adherence to the octet rule, reflecting the tendency of atoms to attain noble gas configurations.
• Predictive Tool: These structures help predict molecular behavior and properties based on electron distribution.
• Central Atom Positioning: The choice of the central atom and terminal atoms is influenced by electronegativity differences.
• Balancing Charges: Lewis structures effectively accommodate anions and cations, adjusting the number of electrons as per charge.

Lewis dot structures serve as valuable tools to visualize bonding in molecules and ions, showcasing shared electron pairs and the octet rule's importance. While not exhaustive, these structures provide insights into molecular properties and behavior, offering a foundational understanding of chemical bonding.

 FORMAL CHARGE IN LEWIS DOT STRUCTURES:

Formal Charge in Lewis Dot Structures

1. Limitations of Lewis Dot Structures:
• Lewis dot structures are valuable tools for understanding molecular bonding but do not directly depict actual molecular shapes.
• Especially in polyatomic ions, the net charge is distributed across the ion as a whole, rather than being concentrated on a specific atom.
2. Assignment of Formal Charge:
• To address the distribution of charge in polyatomic ions, the concept of formal charge is introduced.
• Formal charge is assigned to individual atoms within a molecule or ion to evaluate their electron distribution.
3. Defining Formal Charge:
• Calculation Basis: Formal charge quantifies the difference between an atom's valence electrons in its isolated (free) state and the electrons allocated to it in the Lewis structure.
• Formula: Formal Charge = Valence Electrons (Free Atom) - Assigned Electrons (Lewis Structure)
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4. Purpose of Formal Charge:
• Formal charge helps in assessing the electron distribution among atoms in a molecule or ion, considering their Lewis structure.
• It aids in understanding the extent to which an atom has gained or lost electrons, relative to its neutral state.
5. Application to Polyatomic Molecules and Ions:
• Polyatomic Ions: In polyatomic ions, the net charge is distributed among multiple atoms, making it necessary to consider individual atoms' charges.
• Formal Charge Allocation: By assigning formal charges to each atom in the ion's Lewis structure, a clearer picture of electron distribution and charge is obtained.

Significance and Implications of Formal Charge:

• Addressing Charge Distribution: Formal charge rectifies the limitation of Lewis structures by considering how charge is distributed across atoms in polyatomic ions.
• Determining Electron Distribution: Formal charges assist in understanding which atoms carry a surplus or deficit of electrons within a molecule or ion.
• Comparing Alternative Lewis Structures: Formal charge evaluation guides the selection of the most appropriate Lewis structure when there are multiple possibilities.

Formal charge, a concept used alongside Lewis dot structures, offers a solution to the challenges of representing charge distribution in polyatomic ions. It calculates the difference between valence electrons of an atom in its free state and the electrons assigned to it within the Lewis structure. While Lewis structures are limited in illustrating molecular shapes, the addition of formal charges enables a more comprehensive understanding of charge distribution and electron allocation within complex molecules and ions.

 LIMITATIONS OF THE OCTET RULE:

Limitations of the Octet Rule:

1. Limited Applicability:
• The octet rule is not universally applicable.
• Primarily applies to second-period elements in the periodic table.
• Most useful for understanding the structures of organic compounds.
2. Incomplete Octet:
• Some compounds do not follow the octet rule, especially when the central atom has fewer than four valence electrons.
• Examples include LiCl, BeH₂, and BCl₃, as Li, Be, and B have 1, 2, and 3 valence electrons, respectively.
• Additional examples include AlCl₃ and BF₃.
3. Odd-Electron Molecules:
• Molecules with an odd number of electrons, such as nitric oxide (NO) and nitrogen dioxide (NO₂), do not satisfy the octet rule for all atoms.
4. Expanded Octet:
• Elements in and beyond the third period of the periodic table have 3d orbitals in addition to 3s and 3p orbitals available for bonding.
• Some compounds of these elements exceed the octet rule, termed as an "expanded octet."
• Examples include PF₅, SF₆, H₂SO₄, and various coordination compounds.
• Notably, sulfur forms compounds where it both follows and deviates from the octet rule, as seen in sulfur dichloride (SCl₂).

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Other Drawbacks of the Octet Theory:

• Based on the chemical inertness of noble gases, but some noble gases like xenon and krypton form compounds like XeF₂, KrF₂, XeOF₂, etc.
• Does not provide insights into molecular shapes.
• Fails to explain the relative stability of molecules and does not address molecular energy considerations.

 FORMATION OF IONIC OR ELECTROVALENT BOND:

Formation of Ionic or Electrovalent Bonds:

1. Ion Formation:
• Ionic bond formation depends on the ease of forming positive and negative ions from neutral atoms.
• Positive ions (cations) are formed by ionization, which involves removing electrons from neutral atoms.
• Negative ions (anions) are formed by adding electrons to neutral atoms.
2. Enthalpy Changes:
• Ionization enthalpy refers to the energy required to remove an electron from a neutral atom, which is always endothermic.
• Electron gain enthalpy represents the energy change when a gas-phase atom gains an electron, and it can be either exothermic or endothermic.
• Electron affinity is the negative of the energy change during electron gain.
3. Factors in Ionic Bond Formation:
• Ionic bonds are more likely to form between elements with:
• Low ionization enthalpies (easy electron removal).
• High negative values of electron gain enthalpy (favorable electron addition).
4. Composition of Ionic Compounds:
• Most ionic compounds consist of cations from metallic elements and anions from non-metallic elements.
• Exception: The ammonium ion (NH⁴⁺) forms the cation in some ionic compounds.
5. Crystal Structure:
• Ionic compounds in the crystalline state have three-dimensional arrangements of cations and anions held together by coulombic interaction energies.
• The crystal structure varies based on ion sizes, packing arrangements, and other factors.
6. Energy Considerations:
• The sum of electron gain enthalpy and ionization enthalpy may be positive, but ionic compounds stabilize due to energy released during crystal lattice formation.
• Example: NaCl formation, where ionization enthalpy for Na⁺(g) and electron gain enthalpy for Cl^– (g) are compensated by the enthalpy of lattice formation for NaCl(s).
7. Qualitative Measure of Stability:
• The stability of an ionic compound is assessed by its enthalpy of lattice formation, not solely by achieving an octet of electrons around the ionic species in the gaseous state.
8. Role of Lattice Enthalpy:
• Lattice enthalpy is crucial in the formation of ionic compounds and determines their stability.
• Understanding lattice enthalpy is important for studying the energetics of ionic bond formation.

LATTICE ENTHALPY:

Lattice Enthalpy:

1. Definition:
• Lattice enthalpy refers to the energy required to completely separate one mole of a solid ionic compound into its constituent ions in the gaseous state.
• For example, the lattice enthalpy of NaCl is 788 kJ mol–1, indicating that 788 kJ of energy is needed to separate one mole of solid NaCl into one mole of Na⁺ (g) and one mole of Cl^– (g), moving them to an infinite distance apart.
2. Energy Components:
• Lattice enthalpy considers both attractive forces between ions of opposite charges and repulsive forces between ions of like charges within the solid crystal structure.
3. Three-Dimensional Structure:
• Ionic solids have a three-dimensional crystal structure, making it impossible to calculate lattice enthalpy solely from ion-ion interactions.
• Factors related to the crystal's geometry and arrangement must be taken into account.
4. Calculation Complexity:
• Calculating lattice enthalpy involves a combination of ion-ion interactions, crystal geometry, and intermolecular forces.
• These complex interactions require advanced computational methods and experimental data.

Lattice enthalpy represents the energy needed to separate ionic compounds into their constituent ions in the gaseous state, accounting for both attractive and repulsive forces, along with the geometric factors within the crystal lattice. Calculating lattice enthalpy requires a comprehensive understanding of these interactions and often involves sophisticated techniques.

BOND PARAMETERS:  BOND LENGTH

Bond Parameters: Bond Length

1. Definition:
• Bond length is the equilibrium distance between the nuclei of two atoms that are bonded together in a molecule.
• Techniques such as spectroscopy, X-ray diffraction, and electron diffraction are used to measure bond lengths, typically in more advanced chemistry studies.
2. Contribution of Atoms:
• In a molecule, each atom contributes to the overall bond length.
• For covalent bonds, the contribution of each atom is referred to as the covalent radius of that atom.
3. Covalent Radius:
• The covalent radius is approximately determined as the radius of an atom's core that comes into contact with the core of an adjacent atom in a bonded state.
• It is calculated as half of the distance between two similar atoms joined by a covalent bond within the same molecule.
4. Van der Waals Radius:
• The van der Waals radius represents the overall size of an atom, including its valence shell, in a nonbonded situation.
• It is calculated as half of the distance between two similar atoms in separate molecules in a solid.
5. Illustration of Covalent and Van der Waals Radii:
• Covalent and van der Waals radii are illustrated for chlorine in Figure
6. Typical Bond Lengths:
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  Some average bond lengths for single, double, and triple bonds are provided in Table
• Table presents bond lengths for common molecules.
7. Covalent Radii of Common Elements:
• Table lists the covalent radii of some common elements.

Bond length is the distance between the nuclei of bonded atoms in a molecule. It is a fundamental parameter in understanding molecular structure and is determined through various measurement techniques. The covalent radius represents an atom's contribution to the bond length in a molecule, while the van der Waals radius accounts for the overall size of an atom in nonbonded situations. Specific bond lengths for different types of bonds and common molecules are available in reference tables.

 

BOND PARAMETERS:  BOND ANGLE

Bond Angle:

1. Definition:
• Bond angle is the angle measured in degrees between the orbitals containing bonding electron pairs around the central atom in a molecule or complex ion.
2. Measurement:
• Bond angles can be experimentally determined using spectroscopic methods.
3. Importance:
• Bond angles provide insights into the distribution of orbitals around the central atom within a molecule or complex ion.
• Bond angles are crucial in determining the overall shape of a molecule.
4. Example:
• For instance, the H–O–H bond angle in water (H₂O) can be represented as an example:
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Bond angle is the angle between the orbitals containing bonding electron pairs around a central atom within a molecule or complex ion. It helps determine the molecular or ion's shape and is measured in degrees using spectroscopic techniques.

BOND PARAMETERS:  BOND ENTHALPY

Bond Enthalpy:

1. Definition:
• Bond enthalpy is defined as the amount of energy required to break one mole of bonds of a specific type between two atoms in a gaseous state.
• The unit of bond enthalpy is kJ mol^–1.
2. Example:
• For instance, the H–H bond enthalpy in a hydrogen molecule is 435.8 kJ mol^–1, represented as: H₂(g) → H(g) + H(g); ∆_aH = 435.8 kJ mol–1
3. Strength of Bonds:
• Bond enthalpy is a measure of bond strength; the larger the bond dissociation enthalpy, the stronger the bond in the molecule.
• For molecules with multiple bonds, such as O₂ and N₂, the bond enthalpies are as follows:
• O₂ (O = O) (g) → O(g) + O(g); ∆_aH = 498 kJ mol^–1
• N₂ (N ≡ N) (g) → N(g) + N(g); ∆_aH = 946.0 kJ mol^–1
4. Heteronuclear Diatomic Molecules:
• In the case of heteronuclear diatomic molecules like HCl, the bond enthalpy is determined as follows:
•  HCl (g) → H(g) + Cl(g); ∆_aH = 431.0 kJ mol–1
5. Polyatomic Molecules:
• Measuring bond strength in polyatomic molecules is more complex.
• For example, in H₂O, breaking the two O–H bonds doesn't require the same amount of energy:
• H₂O(g) → H(g) + OH(g); ∆_aH₁ = 502 kJ mol^–1
• OH(g) → H(g) + O(g); ∆_aH₂ = 427 kJ mol^–1
6. Average Bond Enthalpy:
• To account for variations in bond strengths within polyatomic molecules, the concept of average bond enthalpy is used.
• It is calculated by dividing the total bond dissociation enthalpy by the number of bonds broken.
• For example, in water (H₂O): Average bond enthalpy = (502 + 427) / 2 = 464.5 kJ mol–1

Bond enthalpy quantifies the energy required to break a specific type of bond in a gaseous state. It is an indicator of bond strength, with larger bond dissociation enthalpies corresponding to stronger bonds. In polyatomic molecules, variations in bond strength are addressed by calculating the average bond enthalpy.

BOND PARAMETERS:  BOND ORDER

Bond Order:

1. Definition:
• In the Lewis description of a covalent bond, bond order refers to the number of bonds between two atoms within a molecule.
• It quantifies the sharing of electron pairs between the bonded atoms.
2. Examples of Bond Orders:
• Bond order is often represented as an integer:
• In H₂, with one shared electron pair, the bond order is 1.
• In O₂, with two shared electron pairs, the bond order is 2.
• In N₂, with three shared electron pairs, the bond order is 3.
• In CO, with three shared electron pairs between C and O, the bond order is 3.
3. Isoelectronic Molecules and Ions:
• Isoelectronic molecules and ions have identical bond orders.
• For example, F₂ and O₂^2– both have a bond order of 1.
• N₂, CO, and NO⁺ all have a bond order of 3.
4. Relationship with Bond Enthalpy and Bond Length:
• There is a general correlation between bond order, bond enthalpy, and bond length in molecules.
• As bond order increases, bond enthalpy (the energy required to break the bond) typically increases, while bond length (the distance between bonded atoms) decreases.
5. Example of High Bond Order:
• N₂ has a bond order of 3 and a high bond enthalpy (946 kJ mol^–1), making it one of the highest for a diatomic molecule.

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Bond order quantifies the number of bonds between two atoms in a molecule and is a fundamental concept in understanding the nature of covalent bonds. It can provide insights into bond strength and molecular stability, with higher bond orders generally corresponding to stronger bonds and shorter bond lengths.

BOND PARAMETERS:  RESONANCE STRUCTURES

Resonance Structures:

1. Need for Multiple Structures:
• In many cases, a single Lewis structure is insufficient to accurately represent a molecule based on its experimentally determined properties.
• For instance, consider the ozone (O₃) molecule, which can be represented by two Lewis structures, I and II.
2. Ozone Example:
• Both structures I and II depict an O–O single bond and an O=O double bond.
• Normal O–O and O=O bond lengths are 148 pm and 121 pm, respectively.
• Experimentally determined oxygen-oxygen bond lengths in O₃ are identical (128 pm), indicating the bonds are intermediate between double and single bonds.
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• Neither of the two Lewis structures accurately represents this situation.
3. Introduction of Resonance:
• The concept of resonance is introduced to address such cases where a single Lewis structure falls short.
• Resonance suggests that when a single Lewis structure is inadequate, multiple structures with similar energy, nuclear positions, and electron pairs are considered as canonical structures to describe the molecule accurately.
4. Resonance Hybrid:
• In the case of O₃, structures I and II constitute the canonical structures or resonance structures.
• The hybrid of these structures, represented as III, provides a more accurate depiction of O₃ and is termed the resonance hybrid.
• Resonance is denoted by a double-headed arrow (⇌).
5. Examples:
• Other examples of molecules with resonance structures include the carbonate ion and carbon dioxide.
6. Stabilization and Averaging:
• Resonance serves two key purposes:
• It stabilizes the molecule as the energy of the resonance hybrid is lower than that of any single canonical structure.
• It averages the characteristics of bonds throughout the molecule.
7. Energy Consideration:
• The energy of the O₃ resonance hybrid is lower than either of the two individual canonical forms I and II.

Resonance structures are used when a single Lewis structure cannot accurately describe a molecule's properties. Resonance involves multiple canonical structures with similar characteristics, and their hybrid provides a more precise representation. Resonance stabilizes molecules and helps average bond characteristics, leading to a more accurate depiction of the molecule's behavior.

BOND PARAMETERS:  MISCONCEPTIONS ABOUT RESONANCE

Dispelling Misconceptions about Resonance:

1. Non-Existence of Canonical Forms:
• Canonical forms (resonance structures) do not have a real, independent existence as separate entities.
2. No Fractional Existence:
• The molecule does not exist for a fraction of time in one canonical form and then switch to another canonical form. It doesn't physically shift between these forms.
3. Absence of Equilibrium:
• There is no equilibrium between canonical forms, as is seen in tautomeric forms (such as keto and enol) in tautomerism. Resonance does not involve dynamic interconversion between forms.
4. Single Molecular Structure:
• The molecule has a single, well-defined structure, which is the resonance hybrid of the canonical forms.
• This resonance hybrid cannot be accurately represented by a single Lewis structure.

It's crucial to understand that resonance involves a single molecule with a resonance hybrid structure derived from multiple canonical forms. There is no physical transition or equilibrium between these forms, and the canonical forms themselves are not separate, independently existing entities. Resonance is a tool for more accurately describing the distribution of electrons within a molecule.

BOND PARAMETERS:  POLARITY OF BONDS

Polarity of Bonds:

1. Idealized Bond Types:
• In reality, no bond or compound is purely ionic or purely covalent; they represent idealized situations.
2. Nonpolar Covalent Bonds:
• When covalent bonds form between identical atoms, like H₂, O₂, Cl₂, N₂, or F₂, the shared electron pair is equally attracted by both atoms.
• Electron pair is situated exactly between the two identical nuclei.
• Such bonds are termed nonpolar covalent bonds.
3. Polar Covalent Bonds:
• In heteronuclear molecules, like HF, the shared electron pair is displaced more towards the more electronegative atom, like fluorine.
• This results in a polar covalent bond, where one end is more electron-rich (negative) and the other end is more electron-deficient (positive).
4. Dipole Moment:
• Polarization leads to the molecule having a dipole moment (µ), expressed as µ = charge (Q) × distance of separation (r).
• Dipole moments are usually measured in Debye units (D), with 1 D = 3.33564 × 10^–30 C m.
• Dipole moment is a vector quantity, depicted as a crossed arrow ( ⇀ ) with the cross at the positive end and the arrowhead at the negative end.
5. Contribution in Polyatomic Molecules:
• In polyatomic molecules, the dipole moment depends not only on individual bond dipoles but also on their spatial arrangement.
• For example, in H₂O, the net dipole moment results from the orientation of the two O–H bond dipoles.
6. Canceling Dipole Moments:
• In some cases, like BeF₂, bond dipoles point in opposite directions and cancel each other, resulting in a zero net dipole moment.
7. NH₃ vs. NF₃: 
   

• Let us study an interesting case of NH₃ and NF₃ molecule. Both the molecules have pyramidal shape with a lone pair of electrons on nitrogen atom. Although fluorine is more electronegative than nitrogen, the resultant, dipole moment of NH₃ (4.90 × 10^–30 C m) is greater than that of NF₃ (0.8 × 10^–30 C m). This is because, in case of NH₃ the orbital dipole due to lone pair is in the same direction as the resultant dipole moment of the N – H bonds, whereas in NF₃ the orbital dipole is in the direction opposite to the resultant dipole moment of the three N–F bonds. The orbital dipole because of lone pair decreases the effect of the resultant N – F bond moments, which results in the low dipole moment of NF₃ as represented below:

• Interesting cases like NH₃ and NF₃ demonstrate the influence of lone pairs on dipole moments.
• Despite fluorine being more electronegative, NH₃ has a higher dipole moment due to the direction of the lone pair's orbital dipole.
8. Partial Covalent Character in Ionic Bonds:
• Both covalent and ionic bonds have some partial characteristics of the other type.
• Fajans' rules explain the partial covalent character of ionic bonds:
• Smaller cations and larger anions increase covalent character.
• Greater charge on cations increases covalent character.
• Transition metal cations with (n-1)dⁿns^o electronic configuration are more polarizing than noble gas configuration (ns²np⁶).
9. Factors Influencing Covalent Character:
• Polarizing power of the cation, polarisability of the anion, and the extent of distortion (polarisation) of the anion determine the percentage of covalent character in an ionic bond.

No bond is entirely ionic or covalent, and the polarity of a bond depends on the electronegativity of the atoms involved. Dipole moments reflect bond polarity, and they can be influenced by lone pairs and molecular geometry in polyatomic molecules. Ionic bonds also have some degree of covalent character, which can be explained by Fajans' rules, considering factors like cation size, charge, and electronic configuration.

BOND PARAMETERS:  VSEPR THEORY

Origin and Development: Proposed by Sidgwick and Powell in 1940 and later refined by Nyholm and Gillespie in 1957. Developed to address the limitations of the Lewis concept, which couldn't explain molecular shapes. Basic Premise: VSEPR theory is based on the repulsive interactions of electron pairs in the valence shell of atoms. It assumes that electron pairs (both bonding and non-bonding) around a central atom arrange themselves in a way that minimizes electrostatic repulsion to achieve the most stable geometry. Electron Pair Types: VSEPR considers two main types of electron pairs: bonding pairs and non-bonding pairs (lone pairs). Geometric Arrangement: The theory predicts the three-dimensional arrangement of atoms around a central atom. It provides a simple procedure to predict the shape of covalent molecules. Steric Number: The steric number of an atom is the sum of the number of bonded atoms and the number of lone pairs around it. The steric number is used to determine the molecular geometry. Molecular Geometries: VSEPR theory predicts various molecular geometries based on the steric number. Common geometries include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral, among others. Electron Pair Repulsion: The theory postulates that electron pairs (both bonding and lone pairs) repel each other. Lone pairs exert stronger repulsion than bonding pairs, leading to specific bond angles and shapes. Bond Angles: VSEPR theory provides specific bond angles for different molecular geometries. For example, in a tetrahedral molecule like methane (CH4), the bond angle is approximately 109.5 degrees. Application: VSEPR theory is widely used in chemistry to predict the shapes of molecules and polyatomic ions. It helps in understanding and explaining the physical and chemical properties of compounds. Limitations: VSEPR theory provides a qualitative understanding of molecular shapes but doesn't consider the influence of other factors like bond polarity. It may not accurately predict the shapes of molecules with complex or unusual bonding patterns. Advancements: While VSEPR theory laid the foundation for understanding molecular shapes, more advanced theories like molecular orbital theory and hybridization have been developed to provide a deeper insight into molecular structure and bonding.

Valence Shell Electron Pair Repulsion (VSEPR) Theory:

1. Introduction:
• The Lewis concept is insufficient for explaining the shapes of molecules.
• Sidgwick and Powell (1940) introduced the VSEPR theory based on repulsive interactions of electron pairs in the valence shell.
• Nyholm and Gillespie (1957) further refined the theory.
2. Main Postulates:
• The shape of a molecule is determined by the number of valence shell electron pairs (bonded or nonbonded) around the central atom.
• Electron pairs in the valence shell repel each other due to their negative charge.
• Electron pairs position themselves to minimize repulsion and maximize distance.
• The valence shell is treated as a sphere, with electron pairs localizing on its surface.
• Multiple bonds are treated as single electron pairs, and multiple bond electron pairs are treated as a single super pair.
• VSEPR is applicable to any resonance structure when multiple structures exist.

• Dependence on Valence Shell Electron Pairs: The shape of a molecule is determined by the number of valence shell electron pairs (both bonded and nonbonded) surrounding the central atom. Electron Pair Repulsion: Electron pairs in the valence shell exert mutual repulsion because their electron clouds carry a negative charge. Minimization of Repulsion: Electron pairs arrange themselves in three-dimensional space to minimize repulsive interactions, thus maximizing the distance between them. Spherical Valence Shell: The valence shell of the central atom is treated as a spherical surface, with electron pairs localizing on this spherical surface at the maximum possible distance from each other. Treatment of Multiple Bonds: Multiple bonds (e.g., double and triple bonds) are treated as if they were single electron pairs, with the two or three electron pairs of a multiple bond considered as a single "super pair" for the purpose of geometry prediction. Applicability to Resonance Structures: When multiple resonance structures can represent a molecule, the VSEPR model is applicable to any such structure.

3. Repulsion Strength:
• Electron pair repulsion decreases in the order: Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair (bp) > Bond pair (bp) – Bond pair (bp).
• Lone pairs occupy more space than bonding pairs, leading to increased repulsion.
4. Influence on Molecular Shapes:
• Repulsion effects lead to deviations from idealized shapes and alterations in bond angles.
5. Categories for Prediction:
• To predict molecular shapes using VSEPR, molecules are categorized into two types: (i) Central atom with no lone pairs (ii) Central atom with one or more lone pairs.
6. Arrangement of Electron Pairs:
• The table below depicts the arrangement of electron pairs around a central atom A (without lone pairs) and the resulting geometries of molecules/ions of the type AB.

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1. Shapes with Lone Pairs:
• Table 4.7 shows shapes of molecules/ions where the central atom has one or more lone pairs.
  
  
  
  
  
  
  
  
  
2. Reasons for Distortions:
• Table 4.8 explains the reasons for deviations from ideal geometries in molecules.
3. Examples:
• The VSEPR theory successfully predicts the geometries of molecules, especially those of p-block elements.
• It works well even when energy differences between possible structures are small.
4. Theoretical Basis:
• The VSEPR theory's theoretical foundation regarding the impact of electron pair repulsions on molecular shapes is not entirely clear and remains a topic of discussion.

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The VSEPR theory is a powerful tool for predicting the shapes of molecules based on the repulsion between electron pairs in the valence shell. It successfully describes a wide range of molecular geometries, especially for p-block elements, although its theoretical basis is still a subject of debate.