Chemical Bonding and Molecular Structure

Vocabulary practice flashcards covering chemical bonding concepts, including bond energetic criteria, Lewis structures, octet exceptions, covalent and ionic bonding characteristics, thermodynamics of ionic lattices, and Fajan's rules.

Chemical Bond

The attractive force that acts between two or more atoms to hold them together as a stable molecule, accompanied by a decrease in potential energy ($PE$).

Bond Length

The internuclear distance between two bonded atoms at minimum potential energy ($PE$), where the attractive forces between the atoms equal the repulsive forces.

Bond Energy

The energy released during the formation of a chemical bond between atoms, or the energy required to break that bond, where stability is inversely related to potential energy ($\text{Stability} \propto \frac{1}{PE}$).

Lewis Octet Rule

The principle that atoms tend to gain, lose, or share electrons in order to achieve a stable valence shell of $8\,e^-$ (or a duplet of $2\,e^-$ in the case of hydrogen), identical to a noble gas configuration.

Strong Bond

An interatomic bond—such as an ionic, covalent, coordinate covalent, or metallic bond—with an energy of approximately $200\,\text{kJ\,mol}^{-1}$.

Weak Bond

An intermolecular attractive interaction—such as a hydrogen bond or van der Waals forces (including dipole-dipole, ion-dipole, and London dispersion forces)—with an energy of approximately $2-42\,\text{kJ\,mol}^{-1}$.

Odd-Electron Molecule

A compound containing an odd total number of valence electrons (such as $NO$ and $NO_2$) that cannot satisfy the octet rule.

Hypovalent Compound

An electron-deficient molecule containing fewer than $8\,e^-$ in the valence shell of its central atom, such as $BeCl_2$, $BH_3$, and $AlCl_3$.

Hypervalent Compound

A molecule characterized by an expanded octet containing more than $8\,e^-$ around the central atom, such as $PCl_5$, $SF_6$, and $H_2SO_4$.

Formal Charge

The difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure, given by $\text{FC} = \text{Valence } e^- - \text{Non-bonding } e^- - \frac{1}{2}(\text{Bonding } e^-)$.

Resonance

The representation of a single molecule by two or more valid Lewis structures when one structure alone cannot adequately describe its true bonding properties.

Covalent Bond

A chemical bond formed by the mutual sharing of localized electron pairs with opposite spins between two electronegative non-metal atoms via the overlap of half-filled atomic orbitals.

Covalency

The total number of covalent bonds formed by an atom, which can increase upon the excitation of paired electrons into vacant orbitals (such as vacant $d$-orbitals in periods 3 and higher).

Covalent Solid

A substance composed of atoms linked together by a continuous, giant three-dimensional network of covalent bonds, such as diamond, graphite, carborundum ($SiC$), and aluminum nitride ($AlN$).

Molecular Solid

A solid composed of discrete, distinct molecules held together in the solid state by weak intermolecular forces of attraction, such as solid $I_2$ and solid $CO_2$ (dry ice).

Ionic Bond (Electrovalent Bond)

A non-directional bond formed by the complete transfer of electrons between a strongly electropositive metal and a strongly electronegative non-metal, held together by electrostatic forces of attraction.

Lattice Energy

The energy released during the formation of one mole of an ionic crystalline solid from its constituent separated gaseous ions, directly proportional to ionic charge and inversely proportional to interionic distance.

Hess's Law

The thermodynamic principle stating that the total enthalpy change for a chemical reaction is identical whether the reaction takes place in one step or in a series of steps ($\Delta H = \Delta H_1 + \Delta H_2 + \Delta H_3$).

Born-Haber Cycle

An enthalpy cycle based on Hess's Law used to calculate the lattice energy of an ionic solid by relating it to sublimation energy, bond dissociation energy, ionization energy, electron gain enthalpy, and standard enthalpy of formation.

Hydration Energy

The exothermic enthalpy released when gaseous ions become surrounded by water molecules through ion-dipole interactions, proportional to $\frac{\text{charge}}{\text{size}}$.

Dielectric Constant

A property of a solvent indicating its capacity to separate cations and anions and reduce the electrostatic attraction between them in solution (for example, $120$ for $HF$, $81$ for $H_2O$, and $79$ for $D_2O$).

Ionic Mobility

The speed at which an ion moves through a solution under an applied potential difference, which is inversely proportional to its hydrated radius and directly determines ionic conductivity.

Isomorphism

The phenomenon in which different ionic compounds display identical crystalline structures and contain ions with similar outer electronic configurations, such as $ZnSO_4 \cdot 7H_2O$ and $FeSO_4 \cdot 7H_2O$.

Polarisation

The distortion of the electron cloud of an anion caused by the electrostatic attraction of an adjacent cation's nucleus, introducing partial covalent character into an ionic bond.

Fajan's Rules

A set of guidelines that predicts the degree of covalent character in an ionic compound based on the polarising power of the cation and the polarisability of the anion.

Polarising Power (Ionic Potential)

The capability of a cation to distort the electron cloud of an anion, quantified by the ionic potential ratio $\phi = \frac{\text{charge}}{\text{size}}$.

Polarisability

The susceptibility of an anion's electron cloud to being deformed by a cation, which increases with larger ionic radius and higher negative charge.

Pseudo-Inert Gas Configuration

An outer electronic arrangement consisting of $18\,e^-$ ($ns^2\,np^6\,nd^{10}$, found in cations such as $Cu^+$, $Zn^{2+}$, and $Cd^{2+}$) that exhibits greater polarising power and produces higher covalent character than an inert gas ($8\,e^-$) configuration.