01 · Explore
The Kössel-Lewis Approach to Chemical Bonding
The first logical explanation of chemical bonding was provided independently by Kössel and Lewis in 1916, based on the chemical inertness of noble gases.
Lewis pictured the atom as a positively charged 'Kernel' (nucleus plus inner electrons) surrounded by an outer shell that can hold up to eight electrons. He assumed these electrons occupy the corners of a cube surrounding the Kernel. Atoms achieve a stable 'octet' of electrons when they are linked by chemical bonds, mimicking the stable electronic configuration of noble gases (ns²np⁶).
Kössel noted that highly electronegative halogens and electropositive alkali metals are separated by noble gases. The formation of ions involves the gain or loss of electrons to attain stable noble gas configurations. For example, sodium (Na) loses an electron to become Na⁺, while chlorine (Cl) gains one to become Cl⁻. The resulting electrostatic attraction between these ions forms an electrovalent or ionic bond.
Lewis symbols use dots around an element's symbol to represent valence electrons. These symbols help determine the 'group valence' of an element, which is either the number of dots or eight minus the number of dots. Only outer shell electrons participate in bonding; inner electrons are well-protected and generally uninvolved.
| Element | Lewis Symbol | Valence Electrons |
|---|---|---|
| Lithium (Li) | Li· | 1 |
| Beryllium (Be) | ·Be· | 2 |
| Boron (B) | ·B·· | 3 |
| Carbon (C) | ·C··· | 4 |
| Nitrogen (N) | ··N·· | 5 |
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What is the 'Kernel' in Lewis's atomic model?
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How is an electrovalent bond formed?
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02 · Explore
The Octet Rule and Covalent Bonds
The octet rule states that atoms combine by transferring or sharing valence electrons to achieve a stable octet in their valence shells.
In 1919, Langmuir refined Lewis's ideas by introducing the term 'covalent bond.' A covalent bond is formed when two atoms share at least one pair of electrons, with each atom contributing one electron to the pair. This sharing allows both atoms to attain a stable noble gas configuration. For instance, in a chlorine molecule (Cl₂), each Cl atom shares one electron to complete its octet.
Multiple bonds occur when atoms share more than one pair of electrons. A double bond involves sharing two pairs (as in CO₂ or ethene), while a triple bond involves sharing three pairs (as in N₂ or ethyne). Lewis dot structures provide a pictorial representation of these shared pairs and any remaining lone pairs.
While the octet rule is highly useful for organic compounds and second-period elements, it has significant limitations. These include the formation of compounds with incomplete octets (like LiCl, BeH₂ or BCl₃), odd-electron molecules (like NO and NO₂), and expanded octets in elements beyond the third period (like PF₅, SF₆ and H₂SO₄) where d-orbitals are available for bonding.
Formation of Water (H₂O)
2H· + ··O·· -> H:O:H
Each hydrogen atom shares its single electron with oxygen. Oxygen shares two of its six valence electrons (one with each H). Oxygen attains an octet (8 electrons), and each Hydrogen attains a duplet (2 electrons, like Helium).
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Why does SF₆ violate the octet rule?
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03 · Explore
Formal Charge and Resonance Structures
Formal charge helps track valence electrons in polyatomic species, while resonance explains molecules that cannot be represented by a single Lewis structure.
The formal charge (F.C.) of an atom is the difference between its valence electrons in a free state and the electrons assigned to it in a Lewis structure. It is calculated as: F.C. = [Total valence e⁻] - [Total non-bonding e⁻] - 1/2 [Total bonding e⁻]. Assigning formal charges helps identify the most stable (lowest energy) Lewis structure, which is typically the one with the smallest formal charges.
Resonance occurs when a single Lewis structure cannot accurately describe a molecule's experimental properties, such as bond length. Instead, the molecule is represented as a 'resonance hybrid' of several 'canonical structures.' For example, in ozone (O₃), the two O-O bonds are identical in length (128 pm), intermediate between a single (148 pm) and double (121 pm) bond.
It is a common misconception that a molecule flips between canonical forms. In reality, the molecule has a single, constant structure—the resonance hybrid—which is more stable than any individual canonical form. Resonance stabilizes the molecule by lowering its total energy.
Formal Charge on Central Oxygen in O₃
F.C. = 6 - 2 - (1/2 * 6) = +1
Oxygen has 6 valence electrons. In the central position of the ozone Lewis structure, it has 1 lone pair (2 electrons) and 3 bonds (6 shared electrons). The calculation results in a formal charge of +1.
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What does a double-headed arrow indicate in resonance?
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04 · Explore
Ionic Bonding and Lattice Enthalpy
Ionic bonds form through the transfer of electrons, resulting in a crystalline lattice stabilized by energy release.
The formation of an ionic compound depends on the ease of creating ions and their arrangement in a solid lattice. Ionization enthalpy (energy to remove an electron) is always endothermic, while electron gain enthalpy can be exothermic. An ionic bond forms most easily between elements with low ionization enthalpies (metals) and those with high negative electron gain enthalpies (non-metals).
Ionic solids exist as orderly three-dimensional arrangements of cations and anions called a crystal lattice. The stability of these compounds is provided by the 'Lattice Enthalpy,' defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous constituent ions. For NaCl, this value is 788 kJ/mol.
According to Fajans' rules, no bond is 100% ionic. Covalent character increases if the cation is small, the anion is large, or the charges on the ions are high. Cations with transition metal electronic configurations [(n-1)dⁿ ns⁰] are more polarizing than those with noble gas configurations, further increasing covalent character.
Energy Balance in NaCl Formation
- 1
Ionization
Na(g) requires 495.8 kJ/mol to become Na⁺(g).
- 2
Electron Gain
Cl(g) releases 348.7 kJ/mol to become Cl⁻(g).
- 3
Lattice Formation
Na⁺ and Cl⁻ ions arrange into a solid, releasing 788 kJ/mol (Lattice Enthalpy).
- 4
Stability
The energy released during lattice formation compensates for the net energy absorbed during ion formation.
The formation of ionic crystals is driven by the massive release of energy when ions aggregate into a lattice.
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Define Lattice Enthalpy.
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05 · Explore
Bond Parameters and Polarity
Physical properties of bonds, such as length, angle, and enthalpy, determine the behavior and geometry of molecules.
Bond length is the equilibrium distance between nuclei of two bonded atoms. In covalent bonds, it is the sum of the covalent radii of the atoms. Bond angle is the angle between orbitals containing bonding electron pairs around a central atom, helping determine molecular shape. Bond enthalpy is the energy required to break one mole of a specific bond type in the gaseous state; higher bond enthalpy indicates a stronger bond.
Bond order is the number of bonds between two atoms. As bond order increases, bond enthalpy increases and bond length decreases. For example, N₂ has a bond order of 3, a very high bond enthalpy (946 kJ/mol), and a short bond length (109 pm).
Bonds between different atoms are 'polar' because the shared electrons are pulled toward the more electronegative atom. This creates a dipole moment (μ), calculated as the product of charge (Q) and distance (r). Dipole moment is a vector quantity; in symmetrical molecules like BF₃ or CCl₄, individual bond dipoles cancel out, resulting in a net zero dipole moment.
| Molecule | Bond Type | Bond Length (pm) | Bond Enthalpy (kJ/mol) |
|---|---|---|---|
| H₂ | H-H | 74 | 435.8 |
| O₂ | O=O | 121 | 498 |
| N₂ | N≡N | 109 | 946 |
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Why is the dipole moment of CO₂ zero?
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06 · Explore
VSEPR Theory and Molecular Geometry
The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on the idea that electron pairs repel each other.
The central postulate of VSEPR is that electron pairs (both bonding and lone pairs) around a central atom arrange themselves as far apart as possible to minimize repulsion. The magnitude of repulsion follows the order: Lone Pair-Lone Pair (lp-lp) > Lone Pair-Bond Pair (lp-bp) > Bond Pair-Bond Pair (bp-bp).
Molecules are categorized by whether the central atom has lone pairs. If there are no lone pairs (type ABₙ), the geometry is 'ideal': AB₂ is linear, AB₃ is trigonal planar, AB₄ is tetrahedral, AB₅ is trigonal bipyramidal, and AB₆ is octahedral.
When lone pairs are present, they occupy more space and squeeze the bonding pairs, distorting the ideal angles. For example, CH₄, NH₃, and H₂O all have four electron pairs, but while CH₄ is a perfect tetrahedron (109.5°), NH₃ (1 lone pair) is pyramidal (107°) and H₂O (2 lone pairs) is bent (104.5°).
| Molecule Type | Bonding Pairs | Lone Pairs | Shape | Example |
|---|---|---|---|---|
| AB₂E | 2 | 1 | Bent | SO₂ |
| AB₃E | 3 | 1 | Trigonal Pyramidal | NH₃ |
| AB₂E₂ | 2 | 2 | Bent | H₂O |
| AB₄E | 4 | 1 | See-saw | SF₄ |
| AB₃E₂ | 3 | 2 | T-shape | ClF₃ |
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Why is the bond angle in H₂O smaller than in NH₃?
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07 · Explore
Valence Bond Theory and Orbital Overlap
Valence Bond (VB) theory explains bond formation through the overlapping of atomic orbitals and the pairing of electrons with opposite spins.
When two atoms approach, attractive forces (between one nucleus and another's electrons) and repulsive forces (between nuclei or between electrons) operate. A bond forms when the atoms reach a distance where attractive forces balance repulsive forces, resulting in a minimum energy state. For H₂, this equilibrium distance is 74 pm.
The strength of a covalent bond depends on the extent of orbital overlap; greater overlap results in a stronger bond. Orbitals must have the same sign (phase) and orientation for 'positive overlap' to occur. Overlap can be axial (forming a sigma bond) or sidewise (forming a pi bond).
A sigma (σ) bond is formed by end-to-end overlap along the internuclear axis (s-s, s-p, or p-p). A pi (π) bond is formed by the sidewise overlap of p-orbitals, where the electron clouds are above and below the plane of the nuclei. Sigma bonds are stronger than pi bonds because the extent of overlap is greater.
Sigma vs Pi Bond in Ethene (C₂H₄)
C=C bond = 1 σ + 1 π
The carbon atoms are joined by one strong sigma bond (axial overlap) and one weaker pi bond (sidewise overlap of unhybridized p-orbitals).
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Which is stronger: a sigma bond or a pi bond?
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08 · Explore
Hybridisation of Atomic Orbitals
Hybridisation is the process of intermixing atomic orbitals of slightly different energies to form a new set of equivalent 'hybrid orbitals.'
Pauling introduced hybridisation to explain the observed geometries and bond angles of polyatomic molecules. The number of hybrid orbitals produced always equals the number of atomic orbitals that were mixed. These hybrid orbitals are equivalent in energy and shape and are directed in space to minimize electron repulsion.
In sp hybridisation (e.g., BeCl₂), one s and one p orbital mix to form two linear orbitals at 180°. In sp² hybridisation (e.g., BCl₃), one s and two p orbitals form three trigonal planar orbitals at 120°. In sp³ hybridisation (e.g., CH₄), one s and three p orbitals form four tetrahedral orbitals at 109.5°.
Elements in the third period and beyond can involve d-orbitals in hybridisation. Examples include sp³d (trigonal bipyramidal, e.g., PCl₅) and sp³d² (octahedral, e.g., SF₆). In PCl₅, the axial bonds are slightly longer and weaker than equatorial bonds because they experience more repulsion from the equatorial bond pairs.
The Process of sp³ Hybridisation in Carbon
- 1
Ground State
Carbon has configuration 2s² 2p².
- 2
Excited State
One 2s electron is promoted to the vacant 2p orbital, giving 2s¹ 2p³.
- 3
Intermixing
The one 2s and three 2p orbitals mix together.
- 4
Hybrid Orbitals
Four equivalent sp³ hybrid orbitals are formed, directed toward tetrahedral corners.
Hybridisation explains how carbon forms four identical C-H bonds in methane despite having different types of valence orbitals.
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What is the bond angle in sp² hybridisation?
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09 · Explore
Molecular Orbital Theory (MOT)
Molecular Orbital Theory describes electrons in a molecule as occupying orbitals that belong to the entire molecule rather than individual atoms.
Developed by Hund and Mulliken, MOT uses the Linear Combination of Atomic Orbitals (LCAO) method. When two atomic orbitals (AOs) combine, they form two molecular orbitals (MOs): a bonding MO (lower energy, stable) and an antibonding MO (higher energy, unstable, denoted with an asterisk *).
Electrons fill MOs following the Aufbau principle, Pauli exclusion principle, and Hund's rule. The stability of a molecule is determined by its Bond Order: B.O. = 1/2 (N_b - N_a), where N_b is the number of electrons in bonding orbitals and N_a is the number in antibonding orbitals. A positive bond order indicates a stable molecule.
MOT successfully explains properties that other theories cannot, such as the paramagnetism of O₂. Oxygen has two unpaired electrons in its π*2p antibonding orbitals, which makes it attracted to magnetic fields. It also explains why molecules like He₂ (Bond Order = 0) do not exist.
Bond Order of Hydrogen (H₂)
B.O. = 1/2 (2 - 0) = 1
H₂ has 2 electrons, both in the σ1s bonding orbital. The bond order of 1 indicates a single covalent bond.
Bond Order of Helium (He₂)
B.O. = 1/2 (2 - 2) = 0
He₂ would have 4 electrons: 2 in σ1s and 2 in σ*1s. The bond order of 0 means the molecule is unstable and does not exist.
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Why is O₂ paramagnetic?
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10 · Explore
Hydrogen Bonding
Hydrogen bonding is a unique attractive force that occurs when hydrogen is bonded to highly electronegative elements like Nitrogen, Oxygen, or Fluorine.
When H is covalently bonded to a highly electronegative atom (X), the shared electron pair shifts toward X, giving H a partial positive charge (δ+) and X a partial negative charge (δ-). This polar molecule can then form an electrostatic link with another electronegative atom. This link is the hydrogen bond, represented by a dotted line (---).
There are two types of hydrogen bonds. Intermolecular hydrogen bonds occur between different molecules (e.g., in H₂O, HF, or alcohols). Intramolecular hydrogen bonds occur within the same molecule when two electronegative atoms are positioned appropriately (e.g., in o-nitrophenol).
Hydrogen bonds are weaker than covalent or ionic bonds but significantly stronger than van der Waals forces. They have a powerful effect on the physical properties of compounds, such as increasing boiling points and influencing the crystal structures of solids like ice.
Types of Hydrogen Bonding
- 1
Polarization
H-F bond is polar; F pulls electrons, becoming δ-.
- 2
Attraction
The δ+ H of one HF molecule is attracted to the δ- F of another.
- 3
Intermolecular
A chain forms: H-F---H-F---H-F.
- 4
Intramolecular
In o-nitrophenol, H from the -OH group bonds to O from the -NO₂ group in the same molecule.
Hydrogen bonds act as bridges between electronegative atoms, either linking separate molecules or stabilizing single ones.
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Which elements can form hydrogen bonds with Hydrogen?
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