Inorganic Chemistry: Chemical Bonding

Comprehensive vocabulary flashcards covering chemical bonding concepts, including bond classifications, VBT, hybridization types, VSEPR theory, bond parameters, dipole moments, MOT, hydrogen bonding, van der Waals interactions, and Fajan's rules of polarization.

Chemical Bond

A force that acts between two or more atoms to hold them together as a stable molecule, involving the redistribution of electrons accompanied by a decrease in energy.

Strong Bond

An interatomic bond with an energy of approximately $200\,kJ/mole$, including ionic bonds, covalent bonds, co-ordinate bonds, and metallic bonds.

Weak Bond

An intermolecular bond with an energy of approximately $2-42\,kJ/mole$, including hydrogen bonds and van der Waals bonds.

Potential Energy Curve of Bond Formation

A diagram demonstrating that as two atoms approach, attractive forces lower the system's potential energy to a minimum at the bond internuclear distance, making bond formation an exothermic process ($\Delta H = -ve$).

Valency

The combining capacity of an element, defined modernly by its electronic configuration: equal to the number of valence shell electrons for groups IA to IVA, and equal to $|8 - (\text{number of valence } e^-)|$ for groups VA to zero group.

Lewis Octet Rule

The tendency of atoms to achieve eight electrons in their outermost shell (or a duplet in the case of $H$ and $Li$) to acquire a stable inert gas configuration by sharing, losing, or gaining electrons.

Hypovalent Molecules

Incomplete octet or electron-deficient molecules in which the central atom has fewer than $8\,e^-$ in its outermost shell, such as $BeCl_2$ ($4\,e^-$), $BF_3$, $BCl_3$, and $AlCl_3$.

Hypervalent Molecules

Electron-efficient molecules in which the central atom possesses an expanded octet containing more than $8\,e^-$ in its outermost shell, such as $PCl_5$ ($10\,e^-$), $SF_6$ ($12\,e^-$), and $IF_7$ ($14\,e^-$).

Odd Electron Molecules

Molecules wherein the central atom contains an odd number of electrons ($7\,e^-$, $11\,e^-$, etc.) in its outermost shell, exemplified by $NO$, $NO_2$, and $ClO_2$.

Covalent Bond

A bond formed by the mutual sharing of electrons with opposite spins between two atoms of electronegative elements to complete their octet or duplet.

Covalency

The number of covalent bonds that an atom forms in a molecule in either its ground state or an excited state.

Promotion Energy

The excitation energy required to promote paired lone pair electrons into vacant orbitals of another subshell within the same main energy level to maximize the number of unpaired electrons.

Promotion Rule

The principle stating that the excitation of electrons during variable valency expansion takes place only between different subshells belonging to the very same principal shell.

Valence Bond Theory (VBT)

A theory introduced by Heitler and London (and extended by Pauling and Slater) stating that covalent bonds form via the directional overlap of half-filled valence shell atomic orbitals.

Sigma ($\sigma$) Bond

A strong, directional bond formed by coaxial or end-to-end overlapping of half-filled orbitals along the internuclear axis, permitting only one such bond between any two atoms.

Pi ($\pi$) Bond

A weaker, less directional bond formed by the sidewise or lateral (collateral) partial overlapping of atomic orbitals oriented perpendicular to the internuclear axis.

Co-axial Overlapping Strength Order

The order of covalent bond strength arising from directional properties of overlapping orbitals: $p-p > s-p > s-s$.

Hybridisation

The phenomenon proposed by Pauling involving the intermixing of atomic orbitals of different shapes and nearly equal energy from the central atom to form an equal number of new equivalent orbitals of identical shape and energy.

Structure of Hybrid Orbital

An asymmetric orbital possessing two lobes—a large major lobe where overlapping and bond formation take place, and a small minor lobe.

sp Hybridisation

The mixing of one $s$ and one $p$ orbital to produce two equivalent hybrid orbitals arranged collinearly at a bond angle of $180^\circ$ with $50\%$ $s$-character, yielding linear geometry.

sp2 Hybridisation

The mixing of one $s$ and two $p$ orbitals yielding three equivalent hybrid orbitals oriented at an angle of $120^\circ$ in a trigonal planar arrangement with $33.3\%$ $s$-character.

sp3 Hybridisation

The mixing of one $s$ and three $p$ orbitals giving four equivalent hybrid orbitals directed toward the corners of a regular tetrahedron at an angle of $109^\circ 28'$ ($109.5^\circ$).

sp3d Hybridisation

The mixing of one $s$, three $p$, and one $d_{z^2}$ orbital producing five equivalent hybrid orbitals configurated in a trigonal bipyramidal geometry with three equatorial angles ($120^\circ$) and two axial angles ($90^\circ$).

Axial Bond Elongation in $PCl_5$

The physical condition in trigonal bipyramidal $PCl_5$ where the two axial $P-Cl$ bonds are longer and weaker than the three equatorial $P-Cl$ bonds due to greater electron repulsion from equatorial bond pairs.

sp3d2 Hybridisation

The mixing of one $s$, three $p$, and two $d$ orbitals ($d_{x^2-y^2}$ and $d_{z^2}$) producing six equivalent hybrid orbitals oriented toward octahedral (square bipyramidal) positions at adjacent angles of $90^\circ$.

sp3d3 Hybridisation

The mixing of one $s$, three $p$, and three $d$ orbitals ($d_{z^2}$, $d_{x^2-y^2}$, $d_{xy}$) producing seven hybrid orbitals in a pentagonal bipyramidal geometry with five $72^\circ$ angles and ten $90^\circ$ angles.

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory; states that the geometric shape of a molecule depends on the total number of electron pairs (bonded and non-bonded) surrounding the central atom, which orient to minimize repulsions ($\text{lp-lp} > \text{lp-bp} > \text{bp-bp}$).

Bond Angle

The angle formed between orbitals containing bonding electron pairs around the central atom in a molecule or complex ion, measured in degrees ($^\circ$), minutes ('$), and seconds (''$$).

Bond Length

The equilibrium distance between the nuclei of two bonded atoms, typically expressed in Angstroms ($\text{\AA}$) and inversely proportional to bond order and percent $s$-character.

Bond Dissociation Energy

The amount of energy required to dissociate or break one mole of a specific chemical bond between two atoms in a gaseous state.

Co-ordinate Bond

A dative covalent bond in which both shared electrons originate from a single donor atom (Lewis base) possessing a complete octet and lone pair, donated into the vacant orbital of an acceptor atom (Lewis acid).

Dipole Moment ($\mu$)

A vector quantity measuring the extent of polarity in a chemical bond or molecule, calculated as the product of charge magnitude ($q$) and separation distance ($d$): $\mu = q \times d$.

Debye (D)

The standard unit of dipole moment, defined as $1\,\text{Debye} = 1 \times 10^{-18}\,\text{esu}\,\text{cm} = 3.33 \times 10^{-30}\,\text{C}\,\text{m}$.

Molecular Orbital Theory (MOT)

An approach introduced by Hund and Mulliken where atomic orbitals combine linearly (LCAO) across the whole molecule to produce polycentric bonding and antibonding molecular orbitals.

Bonding Molecular Orbital (BMO)

A molecular orbital formed by the constructive interference (addition) of atomic orbital wavefunctions ($\psi_b = \psi_A + \psi_B$), resulting in increased electron density between nuclei, lower potential energy, and molecule stabilization.

Antibonding Molecular Orbital (ABMO)

A higher-energy molecular orbital produced by the destructive interference (subtraction) of wavefunctions ($\psi_a = \psi_A - \psi_B$), featuring a nodal plane with zero electron density between the nuclei and destabilizing the system.

Bond Order

A quantitative value reflecting bond strength and stability, calculated by the equation $\text{Bond order} = \frac{1}{2}(N_b - N_a)$, where $N_b$ is the number of bonding electrons and $N_a$ is the number of antibonding electrons.

Resonance

The delocalization of $\pi$-electrons across polyatomic species represented by multiple canonical structures to explain observed equivalence in bond lengths and increased thermodynamic stability.

Formal Charge

A book-keeping charge assigned to an individual atom within a Lewis formula, calculated on average for surrounding atoms as $\text{Average formal charge} = \frac{\text{Total charge on surrounding atoms}}{\text{Number of surrounding atoms}}$.

Hydrogen Bonding

An electrostatic dipole-dipole attractive force established between a hydrogen atom covalently bonded to a small, highly electronegative atom ($F$, $O$, $N$, or $sp$-carbon) and an electronegative atom of an adjacent or the same molecule.

Intermolecular Hydrogen Bonding

Hydrogen bonding occurring between two or more separate molecules of the same compound (homo) or different compounds (hetero), causing molecular association and raising boiling points and viscosity.

Intramolecular Hydrogen Bonding

Hydrogen bonding occurring internally within the same molecule between a hydrogen atom and a neighboring electronegative atom, resulting in ring chelation, decreased volatility, and lower boiling points.

Van der Waals Forces

Weak, non-directional, non-valence electrical attractive forces operating among neutral atoms or molecules due to temporary or induced dipole polarization.

Keesom Force

An orientation-dependent van der Waals attractive force acting between two permanent dipole molecules, such as in $HCl$, $H_2O$, and $NH_3$.

Debye Force

An induction-based van der Waals attraction that occurs when a molecule with a permanent dipole induces a temporary dipole in an adjacent non-polar species.

London Dispersion Force

The weakest van der Waals attraction, caused by instantaneous fluctuations in electron clouds that create momentary dipoles and induce complementary dipoles in adjacent atoms, such as noble gases.

Ionic Bond

A non-directional chemical bond formed through the complete transfer of one or more valence electrons from an electropositive metal to an electronegative non-metal, held together by electrostatic attractions.

Electrovalency

The total number of electrons lost or gained by an atom during the formation of ions in an ionic compound.

Lattice Energy ($U$)

The quantity of energy released when one mole of an ionic crystal lattice forms from its separate constituent gaseous ions, or the energy needed to dismantle one mole of solid lattice into gaseous ions.

Born-Haber Cycle

A thermochemical cycle applying Hess's Law to relate sublimation energy, ionization potential, bond dissociation energy, electron affinity, and lattice energy to the overall enthalpy of formation ($\Delta H_f$) of an ionic solid.

Polarisation (Fajan's Rule)

The deformation and distortion of the electron cloud of an anion by the electrostatic field of an adjacent cation, which promotes electron sharing and introduces covalent character into an ionic bond.

Ionic Potential ($\phi$)

The ratio of a cation's positive charge to its ionic radius ($\phi = \frac{\text{Charge on cation}}{\text{Size of cation}}$), serving as a quantitative measure of its polarizing power.

Polarizability

The ease with which the electron cloud of an anion can be distorted and drawn toward a cation, increasing directly with the anion's size and negative charge.