Worksheet: Chemical Bonding and Molecular Structure

The Kössel-Lewis approach explains chemical bonding through the sharing and transfer of valence electrons aimed at achieving a stable octet. It discusses how atoms energetically prefer configurations similar to those of noble gases. The approach forms the basis for understanding ionic and covalent bonds, using Lewis symbols to illustrate molecular structure. In sodium chloride (NaCl), the electron transfer between Na and Cl illustrates ionic bonding.

The octet rule states that atoms tend to form compounds in ways that give them eight valence electrons, achieving a noble gas electron configuration. Limitations include exceptions like incomplete octets in BeCl2 and BF3, odd-electron molecules like NO, and expanded octets in SF6. These exceptions highlight the complexity of electronic frameworks in larger or more electronegative atoms and ions.

A covalent bond is formed when two atoms share pairs of electrons. In Cl2, each chlorine atom contributes one electron to form a single bond, resulting in a shared electron pair. The bond in Cl2 is nonpolar since both atoms have the same electronegativity, leading to an equal sharing of electrons. The Lewis structure shows Cl—Cl with a pair of dots representing the shared electrons.

The VSEPR theory states that the geometry of molecules is determined by repulsions between electron pairs around a central atom. For NH3, the molecule adopts a trigonal pyramidal shape due to three bonding pairs and one lone pair, resulting in bond angles slightly less than 109.5 degrees. H2O has two bonding pairs and two lone pairs, adopting a bent shape with a reduced bond angle of approximately 104.5 degrees.

Valence bond theory describes how covalent bonds form through the overlap of atomic orbitals. Hybridization is the process of combining atomic orbitals to create new equivalent orbitals. Types include sp (linear geometry), sp2 (trigonal planar geometry), and sp3 (tetrahedral geometry). For instance, in CH4, one s and three p orbitals hybridize to form four equivalent sp3 orbitals.

Molecular orbital theory posits that electrons in O2 occupy bonding and antibonding molecular orbitals, leading to its stability. With six electrons in bonding orbitals (σ2s, π2p), and two in antibonding orbitals (π*2p), the bond order for O2 is calculated as [(10 bonding - 6 antibonding) / 2], which equals 2. This reflects a double bond, confirming O2's paramagnetic nature due to unpaired electrons in the antibonding orbitals.

Hydrogen bonding occurs when a hydrogen atom covalently bonded to a highly electronegative atom (like O, N, or F) is attracted to another electronegative atom. This interaction significantly raises the boiling point of water compared to similar-sized molecules without hydrogen bonds. The presence of hydrogen bonds also enables water's unique properties like high surface tension and solvent capabilities.

Hybridization involves the mixing of atomic orbitals to form new hybrid orbitals that dictate molecular geometry. Examples include sp3 hybridization in CH4, resulting in tetrahedral geometry, and sp2 in BF3 leading to trigonal planar geometry. The geometry can be predicted based on the type of hybridization and the number of bonding and lone pairs present.

Resonance occurs when a single Lewis structure cannot adequately represent a molecule; multiple structures collectively illustrate the true distribution of electrons. Examples include the carbonate ion (CO3^2–) and benzene (C6H6), where resonance explains bond delocalization and stability. Resonance structures help predict reactivity and stability in various compounds.

Kössel and Lewis proposed that atoms bond to achieve a noble gas configuration, promoting stability. Ionic bonds form through electron transfer, such as NaCl formation, while covalent bonds arise from electron sharing (e.g., H₂). Limitations include the inability to account for certain molecular behaviors and the octet rule's exceptions.

Water has a bent shape due to two lone pairs on oxygen, reducing the H-O-H bond angle to approximately 104.5°. In CO₂, no lone pairs are present, resulting in a linear structure with a 180° bond angle. The repulsion between lone pairs in water leads to a greater distortion from the ideal tetrahedral angle than in CO₂.

Ethylene's sp² hybridization leads to a planar structure with 120° bond angles, consisting of one sigma bond and one pi bond between carbon atoms. Ethyne’s sp hybridization results in a linear arrangement with 180° bond angles, comprising a sigma bond and two pi bonds. The sigma bond forms from orbital overlap, while pi bonds arise from side-to-side overlap of unhybridized p orbitals.

The octet rule does not apply to molecules with incomplete octets (e.g., BeCl₂, BF₃) and those with expanded octets (e.g., SF₆, PCl₅). It also fails for odd-electron species such as NO and ClO₃−. The rule is mainly applicable to second-period elements, limiting its usefulness for heavier elements.

Hydrogen bonding occurs when hydrogen atoms bonded to highly electronegative atoms (F, O, N) interact with other electronegative atoms. In water, hydrogen bonds lead to high boiling/melting points, unique density properties, and high solvent capabilities compared to nonpolar liquids due to weaker van der Waals forces. This pervasive networking of hydrogen bonds is critical for biological systems.

The formal charge helps identify the most stable Lewis structure by calculating charge distribution among atoms. In O₃, structures are drawn to minimize formal charges, resulting in resonance structures with equal bond lengths. The best structure minimizes formal charge and adheres to octet rules, illustrating resonance hybridization.

Bond order = (number of bonding electrons - number of antibonding electrons) / 2. For N₂, it is 3 (10 bonding - 5 antibonding); for O₂, it is 2 (10 bonding - 6 antibonding). An increasing bond order indicates increasing bond strength and shorter bond lengths.

Ionic bonds form through electron transfer and result in coulombic attraction between ions, while covalent bonds form through shared electrons. Lattice energy stabilizes ionic compounds by compensating ionization and electron affinity energies, making the compound energetically favorable over isolated ions.

Resonance indicates that a molecule can be represented by multiple structures, showing electron delocalization. For CO₃²−, the equivalent structures suggest that bond orders across C–O bonds are equal, elucidating real structure as an average of contributing forms, influencing reactivity and stability.

Both SF₆ (octahedral) and PCl₅ (trigonal bipyramidal) involve sp³d and sp³d² hybridizations. The spatial arrangement is dictated by lone pair-bond pair and bond pair-bond pair repulsions, which affect bond angles and geometry.