VSEPR and Bond HybridizationVSEPR 与杂化

From a flat diagram to a 3D molecule

A Lewis diagram shows how atoms are connected but not the shape of the molecule in space. Lesson 2.4.6 used the VSEPR model to predict shapes for 2 to 4 electron domains and treated 5 and 6 as an extension. AP topic 2.7 makes 5 and 6 required too, and asks you to judge polarity from shape, explain shapes with hybrid orbitals, and count σ and π bonds.

Counting electron domains

Each electron domain价层电子对 around the central atom counts once: every bond (single, double, or triple) is one domain, and every lone pair is one domain. The domains repel each other and spread as far apart as possible, which sets the electron geometry; looking only at where the atoms are gives the molecular shape.

2 to 6 electron domains: one table

Some patterns:

• Lone pairs take up more room. A lone pair is attracted by only one nucleus, repels more strongly, and squeezes the bonding pairs closer together: $\ce{CH4}$ (109.5°) > $\ce{NH3}$ (107°) > $\ce{H2O}$ (104.5°). $\ce{BF3}$ and $\ce{BH3}$, with only 3 electron domains and no lone pair, have the largest angle, 120°.
• With 5 domains, lone pairs go equatorial first. An equatorial position has only two neighbors at 90°, so repulsion is smallest there. That’s why $\ce{ClF3}$ is T-shaped and $\ce{XeF2}$ is linear.
• With 6 domains, two lone pairs sit opposite each other. That’s why $\ce{XeF4}$ is square planar.
• Which molecules are planar? In linear, bent, trigonal planar, T-shaped, and square planar molecules, all the atoms lie in one plane; trigonal pyramidal (like $\ce{PCl3}$), tetrahedral, seesaw, trigonal bipyramidal, and octahedral molecules are not planar.

Choose a molecule and drag the model to view it from different angles; turn on “Electron domains” to compare the electron geometry with the molecular shape:

Molecular polarity

Whether a molecule is polar depends on whether its bond dipoles add up to zero (2.4.7):

• Polar bonds but a nonpolar molecule: when identical outer atoms are arranged symmetrically, the bond dipoles cancel, as in $\ce{CO2}$, $\ce{BF3}$, $\ce{CCl4}$, $\ce{PF5}$, $\ce{SF6}$, and $\ce{XeF4}$.
• A lone pair on the central atom breaks the symmetry: $\ce{BF3}$ is trigonal planar and nonpolar, but N in $\ce{NF3}$ has a lone pair, so the molecule is trigonal pyramidal and the three N–F bond dipoles don’t cancel: $\ce{NF3}$ is polar. Likewise, $\ce{CO2}$ is linear and nonpolar, while S in $\ce{SO2}$ has a lone pair, giving 3 electron domains and a bent, polar molecule.
• Different outer atoms: in molecules such as $\ce{CHCl3}$ and HCN, the bond dipoles aren’t equal, so the molecule is polar even though its shape looks regular.

Note: the reason $\ce{NF3}$ is polar is that “the lone pair makes the molecule trigonal pyramidal, so the bond dipoles can’t cancel,” not that “the lone pair is the negative end of the molecule.”

Hybrid orbitals

The ground-state valence configuration of C is 2s²2p², yet the four C–H bonds in $\ce{CH4}$ are identical and point to the corners of a tetrahedron. To explain molecular shapes in terms of atomic orbitals, chemists picture the central atom’s s and p orbitals mixing during bonding to form a set of equal-energy hybrid orbitals杂化轨道 that point in the VSEPR directions. The number of hybrid orbitals equals the number of electron domains:

Hybridization depends only on the number of electron domains, and lone pairs count: the central atoms of $\ce{NH3}$ and $\ce{H2O}$ are both sp³ hybridized.

σ bonds and π bonds

• σ bondσ 键: two orbitals overlap end to end along the line between the two nuclei. Every single bond is a σ bond.
• π bondπ 键: two unhybridized p orbitals overlap side by side, above and below that line. A double bond is 1 σ bond + 1 π bond; a triple bond is 1 σ bond + 2 π bonds.

Because a π bond uses unhybridized p orbitals, an atom with a π bond must be sp or sp² hybridized. π bonds also stop the atoms at either end of a double bond from rotating freely, which is why all six atoms of $\ce{C2H4}$ stay in one plane.

Click an atom to see its electron domains and hybridization; check the box below to highlight the π bonds:

Hybridization can change in a reaction

When ethyne (acetylene) adds chlorine, one π bond of the C≡C triple bond breaks, and each C atom forms a new σ bond to a Cl atom:

Before the reaction, each C has 2 electron domains (sp, H–C–C angle 180°); afterward, each C has 3 (sp², angle about 120°).

In larger molecules and ions, judge each central atom separately. In the zwitterion of glycine, ⁺H₃N–CH₂–COO⁻, the N of $\ce{NH3+}$ has 4 σ bonds, so it is sp³ with angles of about 109.5°; the middle C is also sp³; and the carboxylate C has 3 electron domains, so it is sp² with an O–C–O angle of about 120°, and this C lies in one plane with its two O atoms.

Example 1: The shape and polarity of ClF₃

Predict the molecular shape, bond angles, and polarity of $\ce{ClF3}$.

Example 2: Hybridization and σ and π bonds in HCN

Identify the hybridization of C in HCN and the H–C–N bond angle, and count the σ and π bonds.

Common mistakes

Quick check

AP-style free response

Summary

• Number of electron domains = bonds (each counts once) + lone pairs on the central atom; it sets the electron geometry, and the positions of the atoms give the molecular shape (see the table for 2 to 6 domains).
• Lone pairs repel more strongly and shrink bond angles; with 5 domains lone pairs go equatorial, and with 6 domains two lone pairs sit opposite each other.
• Symmetric molecules whose bond dipoles cancel are nonpolar; a lone pair on the central atom or different outer atoms often make a molecule polar.
• 2, 3, and 4 electron domains correspond to sp, sp², and sp³ hybridization; a single bond is 1 σ, a double bond 1 σ + 1 π, and a triple bond 1 σ + 2 π.