1.10 Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur

Hybridization and Bonding in Nitrogen, Oxygen, Phosphorus, and Sulfur

Sp3 hybridization in nitrogen and oxygen

Carbon isn't the only atom that hybridizes. Nitrogen and oxygen both commonly adopt sp3 hybridization in organic molecules, but their lone pairs change the picture compared to carbon.

Nitrogen has five valence electrons. One s orbital and three p orbitals mix to form four sp3 hybrid orbitals. Three of those orbitals each hold one electron and can form covalent bonds, while the fourth orbital holds a nonbonding (lone) pair. That's why nitrogen typically forms three bonds in neutral molecules like amines ($\text{R-NH}_2$). However, nitrogen can form four bonds if it donates its lone pair to an empty orbital, as in ammonium ions ($\text{NH}_4^+$). In that case, all four sp3 orbitals are involved in bonding.

Oxygen has six valence electrons. The same four sp3 hybrid orbitals form, but now two of them are filled with lone pairs, leaving only two orbitals with single electrons available for bonding. This is why oxygen typically forms two covalent bonds, as in alcohols ($\text{R-OH}$) and water ($\text{H}_2\text{O}$).

Keep in mind that nitrogen and oxygen can also be sp2 or sp hybridized depending on context (think of a carbonyl oxygen or a nitrile nitrogen). The hybridization depends on how many electron domains surround the atom, not just the atom's identity.

Bonding capabilities of phosphorus and sulfur

Phosphorus and sulfur sit in the third row of the periodic table, which means they have access to 3d orbitals. This allows them to exceed the octet rule in certain compounds, something second-row elements like nitrogen and oxygen cannot do.

• Phosphorus has five valence electrons and can form up to five covalent bonds. When it does, one s, three p, and one d orbital hybridize to form five sp3d hybrid orbitals arranged in a trigonal bipyramidal geometry. An example is phosphorus pentachloride ($\text{PCl}_5$). In many organic molecules, though, phosphorus forms three bonds with one lone pair (sp3), similar to nitrogen. Phosphines ($\text{R}_3\text{P}$) are a common example.
• Sulfur has six valence electrons and can form up to six covalent bonds. When it does, one s, three p, and two d orbitals hybridize to form six sp3d2 hybrid orbitals arranged in an octahedral geometry. Sulfur hexafluoride ($\text{SF}_6$) is the classic example. In organic chemistry, you'll also see sulfur with an expanded octet in compounds like sulfoxides and sulfones.

It's worth noting that the role of d orbitals in hypervalent bonding is debated in modern chemistry. Some computational studies suggest that charge-separated resonance structures can explain these molecules without invoking d-orbital participation. For an introductory organic course, the sp3d/sp3d2 model is the standard framework you'll be expected to use.

Nonbonding pairs in organic molecules

Nonbonding electron pairs (lone pairs) are valence electrons that aren't shared in a covalent bond. They sit in hybrid orbitals just like bonding electrons do, and they have a major influence on both shape and reactivity.

Lone pairs matter for two main reasons:

1. They affect molecular geometry and polarity. Lone pairs occupy space around an atom and repel bonding pairs, compressing bond angles. This is why the shape around nitrogen in an amine is trigonal pyramidal (not tetrahedral) and the shape around oxygen in water is bent. Lone pairs also contribute to a molecule's overall dipole moment.

2. They serve as electron donors in reactions. A lone pair is a ready source of electron density. Atoms with lone pairs can act as nucleophiles (attacking electron-poor centers) or as bases (donating electrons to a proton). This is central to organic reaction mechanisms.

Common examples in organic molecules:

• Amines ($\text{R-NH}_2$): nitrogen has one lone pair, making amines both basic and nucleophilic
• Alcohols ($\text{R-OH}$): oxygen has two lone pairs, which can donate to protons or participate in reactions
• Phosphines ($\text{R}_3\text{P}$): phosphorus has one lone pair; phosphines are strong nucleophiles because the lone pair is in a larger, more diffuse orbital
• Thiols ($\text{R-SH}$): sulfur has two lone pairs; sulfur's larger size makes thiols more nucleophilic than alcohols in many contexts

Molecular structure and bonding

VSEPR theory (Valence Shell Electron Pair Repulsion) predicts the 3D shape around any atom by counting all electron domains, both bonding pairs and lone pairs. Electron domains repel each other and arrange themselves to maximize distance.

• Two domains → linear (180°)
• Three domains → trigonal planar (120°)
• Four domains → tetrahedral (109.5°)
• Five domains → trigonal bipyramidal (90°/120°)
• Six domains → octahedral (90°)

Lewis structures are your starting point for determining hybridization. Draw the Lewis structure, count the electron domains (bonds + lone pairs) around the atom in question, and that count tells you the hybridization:

• 4 electron domains → sp3
• 3 electron domains → sp2
• 2 electron domains → sp

Finally, electronegativity differences between bonded atoms determine bond polarity. Oxygen and nitrogen are both quite electronegative, which is why C–O and C–N bonds are polar. This polarity, combined with lone pairs, drives much of the reactivity you'll see in later units.