5.2 Bonding and Structures in p-Block Compounds

Bonding Theories

Covalent and Ionic Bonding

p-Block compounds display a wide range of bonding, from purely covalent to substantially ionic. The bonding type determines structure, physical properties, and reactivity, so getting this right matters for everything that follows.

Covalent bonding involves the sharing of electron pairs between atoms. It typically forms between nonmetals and produces either discrete molecules (like $\text{PCl}_3$) or extended network covalent structures (like diamond or $\text{SiO}_2$). Bond strength increases with the number of shared pairs: triple bonds are shorter and stronger than double bonds, which are stronger than single bonds. A key p-block trend is that elements in the second period (C, N, O) readily form multiple bonds through effective $p\pi$–$p\pi$ overlap, while heavier p-block elements (Si, P, S) generally prefer single bonds due to poorer orbital overlap at longer bond distances.

Ionic bonding results from electrostatic attraction between oppositely charged ions. It typically occurs between metals and highly electronegative nonmetals (e.g., $\text{NaCl}$, $\text{CaF}_2$) and produces crystalline lattice structures rather than discrete molecules. Lattice energy, which governs the strength of the ionic interaction, increases with higher ion charges and smaller ionic radii.

In practice, most p-block compounds fall on a spectrum between purely covalent and purely ionic. Fajan's rules help predict where a compound sits: small, highly charged cations polarize anions and introduce covalent character into nominally ionic bonds (e.g., $\text{AlCl}_3$ is far more covalent than you'd expect from a metal-nonmetal pair).

Molecular Geometry Theories

VSEPR theory predicts molecular shapes by minimizing repulsion between electron domains (both bonding pairs and lone pairs) around a central atom. Lone pairs occupy more angular space than bonding pairs, so they compress bond angles. For example:

• $\text{BF}_3$: 3 bonding pairs, 0 lone pairs → trigonal planar, 120°
• $\text{NH}_3$: 3 bonding pairs, 1 lone pair → trigonal pyramidal, ~107°
• $\text{H}_2\text{O}$: 2 bonding pairs, 2 lone pairs → bent, ~104.5°

To apply VSEPR:

1. Draw the Lewis structure of the molecule.
2. Count the total number of electron domains (bonding + lone pairs) around the central atom. A double or triple bond counts as one domain.
3. Determine the electron-domain geometry from the total count (2 = linear, 3 = trigonal planar, 4 = tetrahedral, 5 = trigonal bipyramidal, 6 = octahedral).
4. Identify the molecular geometry by considering only the positions of the atoms (ignoring lone pairs in the shape name).

Hybridization complements VSEPR by explaining geometry through orbital mixing. An atom combines its atomic orbitals to form a set of equivalent hybrid orbitals:

The number of hybrid orbitals always equals the number of atomic orbitals combined, and each hybrid orbital can hold a bonding pair or a lone pair.

Advanced Bonding Models

Molecular orbital (MO) theory goes beyond Lewis structures and hybridization by treating electrons as delocalized across the entire molecule. Atomic orbitals combine to form bonding MOs (lower energy, stabilizing) and antibonding MOs (higher energy, destabilizing). Bond order is calculated as:

$\text{Bond Order} = \frac{(\text{electrons in bonding MOs}) - (\text{electrons in antibonding MOs})}{2}$

MO theory is especially useful for p-block diatomics. For instance, it correctly predicts that $\text{O}_2$ is paramagnetic (two unpaired electrons in $\pi^*$ orbitals), something Lewis structures cannot explain. For homonuclear diatomics of the second period, note that $\text{B}_2$, $\text{C}_2$, and $\text{N}_2$ have a different MO energy ordering than $\text{O}_2$ and $\text{F}_2$ due to $s$–$p$ mixing, which raises the $\sigma_{2p}$ orbital above the $\pi_{2p}$ orbitals.

The octet rule states that atoms tend to gain, lose, or share electrons to achieve 8 valence electrons, mimicking a noble gas configuration. This works well for second-period elements but has important limitations in p-block chemistry:

• Electron-deficient compounds: Some elements form stable compounds with fewer than 8 electrons. $\text{BF}_3$ has only 6 electrons around boron, making it a strong Lewis acid.
• Expanded octets: Elements in period 3 and beyond can accommodate more than 8 electrons around the central atom. Whether this involves actual d-orbital participation or is better described through multi-center bonding and hypervalency is an active discussion in inorganic chemistry (see the hypervalent compounds section below).

Molecular Structure and Properties

Electronic Distribution and Polarity

Electronegativity measures an atom's ability to attract shared electrons in a bond. On the Pauling scale, fluorine is the most electronegative element (4.0), and electronegativity generally increases left to right across a period and decreases down a group. The electronegativity difference between two bonded atoms determines bond polarity: a difference greater than ~1.7 typically indicates significant ionic character.

Dipole moment ($\mu$) is a vector quantity that reflects the magnitude and direction of charge separation in a bond or molecule. Molecular dipole moments are the vector sum of all individual bond dipoles. This means symmetrical molecules can have polar bonds but zero net dipole moment. For example, $\text{CO}_2$ is linear, so its two $\text{C=O}$ bond dipoles cancel. In contrast, $\text{SO}_2$ is bent, so its bond dipoles don't cancel, giving it a net dipole moment.

Polarity directly affects physical properties: polar molecules tend to have higher boiling points (due to dipole-dipole interactions) and dissolve better in polar solvents.

Structural Representations

Lewis structures depict valence electrons as dots and lines (bonding pairs). To draw one:

1. Count the total valence electrons for all atoms in the molecule or ion.
2. Place the least electronegative atom as the central atom (hydrogen and fluorine are always terminal).
3. Connect atoms with single bonds first, then distribute remaining electrons as lone pairs to satisfy octets on terminal atoms.
4. If the central atom lacks an octet, convert lone pairs on adjacent atoms into double or triple bonds.
5. Calculate formal charges ($\text{FC} = \text{valence e}^- - \text{lone pair e}^- - \frac{1}{2}\text{bonding e}^-$) and choose the structure that minimizes them.

Resonance structures arise when a single Lewis structure can't adequately represent the electron distribution. The real molecule is a weighted average (resonance hybrid) of all valid structures. Resonance stabilizes molecules by delocalizing electron density. In $\text{NO}_3^-$, for example, three equivalent resonance structures give each N–O bond a bond order of 4/3, and all three bond lengths are identical (rather than one short double bond and two long single bonds).

Molecular Arrangement

Isomerism describes compounds that share the same molecular formula but differ in structure or spatial arrangement. While isomerism is more commonly associated with organic chemistry, it appears in p-block inorganic compounds as well:

• Structural (constitutional) isomers have different atom connectivity. For example, $\text{N}_2\text{O}$ could theoretically be N–N–O or N–O–N (though N–N–O is the actual stable form).
• Geometric (cis-trans) isomers differ in the arrangement of groups around a rigid structural feature like a double bond or a ring. In $\text{N}_2\text{F}_2$, the cis isomer has both fluorines on the same side of the N=N bond, while the trans isomer has them on opposite sides.
• Optical isomers are non-superimposable mirror images. This is less common in simple p-block compounds but can occur in certain coordination environments.

Advanced Bonding Concepts

Hypervalent Compounds

Hypervalent compounds have a central atom bonded to more atoms than the traditional octet would allow. They are common for p-block elements in period 3 and below because these larger atoms can accommodate higher coordination numbers.

Classic examples:

• $\text{PCl}_5$: Phosphorus forms 5 bonds, giving a trigonal bipyramidal geometry. The axial bonds (longer, ~214 pm) are slightly weaker than the equatorial bonds (shorter, ~202 pm) due to different orbital interactions.
• $\text{SF}_6$: Sulfur forms 6 bonds in an octahedral arrangement. This compound is remarkably inert despite having 12 electrons around sulfur, largely because the six fluorines sterically shield the sulfur center.
• $\text{XeF}_4$: Xenon forms 4 bonds and retains 2 lone pairs, giving a square planar molecular geometry (the electron-domain geometry is octahedral).

The traditional explanation invokes d-orbital participation ($sp^3d$ or $sp^3d^2$ hybridization). However, modern computational studies suggest that d-orbital involvement is minimal. A more accurate description uses 3-center 4-electron (3c-4e) bonding: in the axial bonds of $\text{PCl}_5$, for instance, three atoms share four electrons across three molecular orbitals (one bonding, one nonbonding, one antibonding), with only the bonding and nonbonding MOs occupied. This model explains why axial bonds are typically longer and weaker than equatorial bonds without requiring significant d-orbital hybridization.

Factors that favor hypervalency:

• A central atom from period 3 or below (larger atomic radius accommodates more ligands)
• Highly electronegative terminal atoms like F or Cl (they stabilize the extra electron density)
• Low steric strain around the central atom

Hypervalent compounds are not just curiosities. $\text{SF}_6$ is widely used as an electrical insulator in high-voltage equipment, and $\text{PCl}_5$ is a common chlorinating agent in synthesis.