Exceptions to the Octet Rule, Bond Properties, and Reaction Enthalpy

Vocabulary flashcards covering exceptions to the octet rule (incomplete octets, expanded octets, and free radicals), formal charges, resonance structures, bond length and strength relationships, and estimating reaction enthalpies using bond energies.

Lewis Theory

A chemical model used to predict the spatial arrangement of atoms within a molecule and the distribution and sharing of valence electrons across a compound, which helps determine chemical properties.

Resonance Hybrid

The single, actual electronic state of a compound that represents the weighted composite or average of two or more contributing Lewis resonance structures.

Operational Rules for Major Resonance Structures

Criteria dictating that the most stable contributing structure has the maximum number of full valence octets, the minimum magnitude and count of formal charges, and negative formal charges positioned on the more electronegative atoms.

Curved Arrow Push Notation (in Resonance)

A notation indicating the relocation of electron pairs only—specifically transitioning lone pairs to bonds, bonds to bonds, or bonds to lone pairs—without altering atomic positions.

Formal Charge Rule of Stability

The principle stating that molecular configurations are most stable when the maximum number of atoms have a formal charge equal to $0$, and any unavoidable negative formal charges reside on the most electronegative atom(s).

Incomplete Octet

An exception to the octet rule where stable compounds (such as $\text{BeCl}_2$ and $\text{BF}_3$) form with fewer than $8$ valence electrons surrounding the central atom because satisfying the octet creates unfavorable formal charges.

Coordinate Covalent Bond

A covalent linkage formed when an electron-deficient species reacts with a molecule containing a lone pair, with both electrons in the shared bond provided by that single donor species.

Free Radical

A highly reactive chemical species possessing an odd number of valence electrons (such as $\text{NO}$, $\text{NO}_2$, and $\text{ClO}_2$), which typically generates another free radical upon reacting with paired-electron molecules.

Expanded Octet

A bonding situation where a central atom accommodates more than $8$ valence electrons via empty, available $d$ orbitals; this is restricted to elements in Period 3 ($n = 3$) and beyond and is impossible for principal quantum number $n \le 2$.

Tetrachloroiodate Ion ($\text{ICl}_4^-$) Electron Count

A polyatomic anion containing a total of $36$ valence electrons, where the central iodine atom accommodates an expanded octet of $12$ valence electrons utilizing unoccupied $5d$ orbitals.

Bond Enthalpy

The quantity of energy required to break a chemical bond in the gas phase; it is inherently endothermic ($\Delta H > 0$), with higher values signifying stronger, shorter chemical bonds.

Bond Length

The equilibrium internuclear distance between two bonded nuclei at their minimum potential energy, experimentally determined through X-ray, electron, or neutron diffraction.

Bond Order Effects on Length and Strength

As the bond order between two atoms increases from single to double to triple, the covalent bond length decreases while the bond strength and bond enthalpy increase.

Reaction Enthalpy from Bond Energies

An application of Hess's Law estimating overall enthalpy changes via the formula: $\Delta H_{\text{rxn}} = \sum (\text{enthalpy of bonds broken}) - \sum (\text{enthalpy of bonds formed})$.

Lewis Theory

A theoretical model used to predict the arrangement of atoms within a molecule and the sharing and redistribution of valence electrons, which helps predict chemical properties.

Formal Charge

A bookkeeping charge assigned to an individual atom in a Lewis structure; structures are considered more stable when formal charges are minimized to zero and negative charges reside on the most electronegative atoms.

Resonance Hybrid

The actual, single physical state of a molecule that exists as an average or combination of all contributing resonance structures.

Major Resonance Structure

The dominant contributing Lewis structure of a molecule, favored by having the greatest number of full valence octets, the fewest formal charges, and negative formal charges located on the more electronegative atoms.

Curved Arrows ("Push" Arrows)

Symbolic arrows used to generate alternate resonance structures by demonstrating the relocation of electron pairs (lone pair to bond, bond to bond, or bond to lone pair) while keeping atom positions fixed.

Incomplete Octet

An exception to the octet rule where central atoms with few valence electrons (such as beryllium in $BeCl_2$ or boron in $BF_3$) form stable compounds with fewer than eight valence electrons to satisfy formal charge rules.

Coordinate Covalent Bond

A covalent bond formed when an electron-deficient species reacts with a species carrying a lone pair, acquiring an octet by receiving both shared electrons from that single atom.

Free Radical

A highly reactive molecule or atom containing an odd number of valence electrons, leaving at least one electron unpaired (e.g., $NO$, $NO_2$, and $ClO_2$).

Expanded Octet

An octet rule exception occurring exclusively in period 3 ($n \times 1 \rightarrow n \times 3$ or $n \times 1 \rightarrow n \times 4$) elements ($n \times 1 \rightarrow n \times 5$) or higher ($n \times 1 \rightarrow n \times 6$), where a central nonmetal utilizes unoccupied $d$ orbitals to accommodate 10, 12, or more valence electrons.

Bond Enthalpy

The quantity of energy required to break one mole of a specific chemical bond in the gas phase; it is always positive (endothermic), with higher values signifying stronger bonds.

Bond Length

The equilibrium internuclear distance between two bonded atoms corresponding to the minimum potential energy well, experimentally measurable using X-ray, electron, or neutron diffraction.

Bond Order

The number of electron pairs shared between two atoms (single, double, or triple bonds); increasing bond order results in shorter bond lengths and greater bond strengths.

Orbital Overlap

The spatial overlap of valence electron orbitals between two bonded atoms, directly governing the resulting strength and stability of the covalent bond.

Reaction Enthalpy from Bond Energies

The estimated enthalpy change of a reaction calculated via Hess' Law: $\Delta H_{\text{rxn}} = \sum (\text{enthalpy of bonds broken}) - \sum (\text{enthalpy of bonds formed})$.