H2 Chemistry: Chemical Bonding (I)

Flashcards covering the fundamental concepts of ionic, covalent, and dative bonding, orbital overlap, VSEPR theory, and bond characteristics as presented in the H2 Chemistry lecture notes.

Chemical bonds

Electrostatic attractive forces between particles (atoms, ions or molecules) resulting in a lower energy arrangement.

Ionic bonds

The electrostatic attraction between oppositely charged ions.

Covalent bond

The electrostatic attraction between a shared pair of electrons and positively charged nuclei.

Metallic bond

The electrostatic attraction between a lattice of positive ions and delocalised electrons.

Kernel

The nucleus and inner shell electrons of an atom, represented by the element symbol in a dot-and-cross diagram.

Bond pair of electrons

Valence electrons involved in the formation of a covalent bond.

Lone pair of electrons

Valence electrons not involved in bonding that occupy a larger volume around the central atom and belong only to that nucleus.

Octet Rule

The tendency of atoms in covalent molecules to acquire a stable electronic configuration with eight electrons in the valence shell.

Sigma ($\sigma$) bond

A bond formed by the head-on (or end-on) overlap of atomic orbitals ($s$ and $p$) with a region of electron density between the two nuclei.

Pi ($\pi$) bond

A bond formed by the sideways overlap of $p$ orbitals, consisting of two regions of electron density above and below the sigma bond.

Bond length

The inter-nuclear distance between two atoms involved in bonding, representing the optimum bonding distance.

Molecular orbital

A new bonding orbital formed at the optimum distance where maximum overlap of atomic orbitals occurs, containing the shared pair of electrons.

Co-ordinate (dative covalent) bond

A type of covalent bond in which only one of the atoms involved in bonding contributes both electrons to the shared pair.

Dimer

An association of two identical species, such as molecules, linked together (e.g., $Al_2Cl_6$).

Bond energy

The energy, in $kJ\,mol^{-1}$, required to break one mole of covalent bonds between two atoms in the gaseous state.

Bond order

The number of covalent bonds formed between two atoms involved in bonding; higher order corresponds to larger bond energy and shorter bond length.

Electronegativity

A measure of the ability of an atom in a covalent bond to attract the bonding electrons towards itself.

Polarised ionic bond

An ionic bond that exhibits some degree of covalent character due to the positive ion distorting the electron cloud of the negative ion.

Polar covalent bond

A covalent bond where electrons are shared unequally between two atoms of different electronegativities, resulting in a net dipole moment.

Dipole moment

A vector quantity with magnitude and direction, represented by an arrow pointing toward the more electronegative atom in a polar bond.

Valence Shell Electron Pair Repulsion (VSEPR) theory

A theory used to predict molecular shapes based on the principle that valence electron pairs arrange themselves to maximize distance and minimize electrostatic repulsion.

Trigonal planar

A molecular shape with three regions of electron density and zero lone pairs around the central atom, featuring a bond angle of $120^{\circ}$.

Tetrahedral

A molecular shape with four bond pairs and zero lone pairs around the central atom, featuring a bond angle of $109^{\circ}$.

Trigonal pyramidal

A molecular shape derived from a tetrahedral arrangement with three bond pairs and one lone pair, such as in ammonia ($NH_3$), with a bond angle of approximately $107^{\circ}$.

Octahedral

A molecular shape with six bond pairs and zero lone pairs around the central atom, featuring bond angles of $90^{\circ}$.

Hydrogen bonding

The strongest type of intermolecular force of attraction, occurring in molecules containing $-NH$ or $-OH$ groups.