Chemical Bonding & Molecular Structure Flashcards

Comprehensive vocabulary flashcards covering the key concepts, theories, and laws of Chemical Bonding and Molecular Structure as presented in the lecture notes.

Chemical bond

The attractive force which holds various constituents (atoms, ions, etc.) together in different chemical species.

Kossel-Lewis Approach

Based on the electronic concept of the atom, this theory is called the electronic theory of valence.

Valence electrons

Electrons of the outermost shell of an atom.

Kernel

The part of the atom left after the removal of valence electrons; it is positively charged.

Octet rule

Atoms of elements combine to complete their octet (eight electrons) or duplet (two electrons) in their valence shell.

Lewis symbols

Simple notations, also called electron dot symbols, used to represent the valence electrons of an atom without considering inner shell electrons.

Covalent bond

A bond formed by the mutual sharing of electron pairs among the combining atoms.

Electrovalent Bond

Also known as an ionic bond, it is formed by the transfer of electrons between a metal and a non-metal.

Coordinate bond

Also known as a dative bond, it is a special type of covalent bond in which the shared electron pair is contributed by only one of the combining atoms.

Formal Charge

The difference between the number of valence electrons of an atom in a free state and the number of electrons assigned to that atom in the Lewis structure, calculated as $F.C. = V - L - \frac{B}{2}$.

Lattice Enthalpy

The energy required to completely separate one mole of a solid ionic compound into its gaseous constituent ions.

Born Haeber Cycle

An indirect method used to calculate lattice energy for ionic solids like $NaCl(s)$.

Bond Length

The equilibrium distance between the nuclei of two bonded atoms in a molecule.

Bond Angle

The angle between the orbitals containing bonding electron pairs around the central atom in a molecule.

Bond Enthalpy

The amount of energy required to break one mole of a bond of a particular type between two atoms in a gaseous state.

Bond Order

The number of bonds between two atoms in a molecule as described in a Lewis structure.

Resonance Structures

Canonical structures with similar energy and positions of nuclei used to describe a molecule when a single Lewis structure is inaccurate; the resulting describing structure is the resonance hybrid.

Dipole Moment ($\text{D.M.}$)

A vector quantity measuring the separation of positive and negative charges, expressed in Debye ($1 D = 3.33564 \times 10^{-30} Cm$), calculated as $\text{charge } (Q) \times \text{ distance of separation } (r)$.

Polarisation

The distortion of the electron cloud of an anion when it is attracted by a cation.

Polarising power

The power of a cation to polarise an anion.

Polarisability

The tendency of an anion to get polarised by a cation.

Fazan's Rule

Principles stating that covalent bonds have partial ionic character and vice versa, influenced by factors such as small cation size, large anion size, and large ionic charges.

Valence Bond Theory ($\text{VBT}$)

A theory initially proposed by Heitler and London that explains covalent bond formation through the overlapping of two half-filled valence atomic orbitals.

Sigma bond (\text{\sigma-bond})

A covalent bond formed by the linear, head-on overlap of two atomic orbitals along the inter-nuclear axis.

Pi bond (\text{\pi-bond})

A covalent bond formed by the parallel or sidewise overlap of atomic orbitals where electron density is present above and below the inter-nuclear axis.

Hybridization

The intermixing of two or more pure atomic orbitals of an atom with almost the same energy to produce the same number of identical and degenerate hybrid orbitals.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory, which states that the shape of a molecule depends on the number of valence shell electron pairs (bonded or nonbonded) around the central atom.

Molecular Orbital Theory ($\text{MOT}$)

A theory developed by Robert S. Mullikan where atomic orbitals combine to form molecular orbitals that surround the entire molecule, allowing for predictions regarding magnetism and bond order.

Hydrogen Bonding

The electrostatic attraction between a hydrogen atom bonded to a highly electronegative atom ($N$, $O$, or $F$) and a lone pair of electrons on another electronegative atom.

Intermolecular hydrogen bond

A hydrogen bond formed between two different molecules.

Intramolecular hydrogen bond

A hydrogen bond formed between two different atoms within the same molecule.