2.7: VSEPR and Bond Hybridization

Learning Targets/ Content Objectives

SAP-4.C Based on the relationship between Lewis diagrams, VSEPR theory, bond orders, and bond polarities:

a. Explain structural properties of molecules.

b. Explain electron properties of molecules.

Essential Knowledge

SAP-4.C.1 VSEPR theory uses the Coulombic repulsion between electrons as a basis for predicting the arrangement of electron pairs around a central atom.
SAP-4.C.2 Both Lewis diagrams and VSEPR theory must be used for predicting electronic and structural properties of many covalently bonded molecules and polyatomic ions, including the following:

Molecular geometry
Bond angles
Relative bond energies based on bond order
Relative bond lengths (multiple bonds, effects of atomic radius)
Presence of a dipole moment
Hybridization of valence orbitals of the molecule

SAP-4.C.3 The terms “hybridization” and “hybrid atomic orbital” are used to describe the arrangement of electrons around a central atom. When the central atom is sp hybridized, its ideal bond angles are 180^o; for sp2 hybridized atoms the bond angles are 120^o; and for sp3 hybridized atoms the bond angles are 109.5^o.
SAP-4.C.4 Bond formation is associated with overlap between atomic orbitals. In multiple bonds, such overlap leads to the formation of both sigma and pi bonds. The overlap is stronger in sigma than pi bonds, which is reflected in sigma bonds having greater bond energy than pi bonds. The presence of a pi bond also prevents the rotation of the bond and leads to structural isomers.

Molecular Structure: the vsepr model

The valence shell electron-pair repulsion (VSPER) theory predicts the geometries of molecules and polyatomic ions. The shape, or geometry, of a molecule is determined by lone pairs or bonds on the central atom of a molecule as these areas of electron density or “charge clouds” will minimize electron–electron repulsions by positioning themselves as far apart as possible. Lone pairs of electrons repel more than bonds and tend to compress the angle between bonding atoms.

Steps to determine molecular geometry:

Draw the Lewis structure of the molecule.

Tip: Hydrogens are always terminal and the least electronegative atom is usually central.

Count the bonds and lone pairs around the central atom. Lone pairs and bonds are “change clouds.”For structures with double or triple bonds, count the multiple bond(s) as one charge cloud. If there is a single electron (as in an odd electron molecule such as NO3), count the odd electron as one charge cloud.

Use the arrangement of the lone pairs and bonds to determine the molecular geometry from the chart provided. *You must memorize these geometries and bond angles for the AP Chem exam as this chart will not be provided.

If there is more than one central atom we determine the geometric shape of each central atom individually by following the same steps above for each central atom

Bond Hybridization

To explain molecular geometries, we assume that the atomic orbitals on an atom mix to form hybrid orbitals. The shape of a hybrid orbital is a mix of the shapes of the original atomic orbitals such as s (spherical) and p(dumbbell). The total number of atomic orbitals on an atom remains constant, so the number of hybrid orbitals on an atom equals the number of atomic orbitals that are mixed. In methane, (CH4), the 2s and three 2p orbitals of carbon mix to form four sp3 hybrid orbitals. In general, the sum of the superscripts on the hybrid orbitals equals the number of electron clouds around the central atom. The carbon atom in CH4 has 4 electron clouds around it. The sum of the superscripts in sp3 equals 4. You will only need to identify sp, sp2 and sp3 hybridizations in AP Chemistry.

Below are summary tables for All possible geometric shapes and hybridization of molecules

Handout_Lewis_VSEPR.pdf

VSEPR Summary .pdf

Sigma (σ ) and pi (π) bonds

Overlapping orbitals from a single bond are known as sigma bonds represented by the greek letter (σ ) which are very strong bonds.
Double and triple bonds are formed from unhybridized p orbitals and are called pi bonds represented by the greek letter ( π)
A double bond contains one sigma (1σ) and one pi (1π ) bond.
A triple bond contains one sigma (1σ) and two pi (2π ) bonds.

Bond Length, Strength and Energy

As the number of bonds between two atoms increases, the bond increases in strength and energy but decreases in length. Pi bonds, present in double and triple bonds, pull the atoms closer together. Triple bonds are the strongest and shortest bonds with the highest energy, while single bonds are the longest and weakest bonds with the lowest energy. Double bonds are stronger and shorter and higher in energy than single bonds.

Bond Order

Bond order is the number of bonds between two atoms. When the bond order increases, the bond lengthdecreases and the bond energy increases.

For a molecule that exhibits resonance, we see that the bonds that have resonance are experimentally determined to be the same length. In ozone, O3, pictured below, we would expect the single bond to be shorter than the double bond, but that is not the case. They both have a bond length that is halfway between a single and double bond, so the bond order is halfway between a single and double bond order, and it’s bond order is 1.5. Another way to think about this is that there are three bonds shared between two atoms, and 3/2 = 1.5.

In the nitrate ion, NO3-, shown below, there are four bonds shared by the nitrogen with each of the three oxygenatoms.The bond order is 4/3 = 1.33

Bond Polarity

When electrons are shared in a covalent bond, the shared electrons spend more time around the more electronegative element in the bond. This gives the more electronegative element a partial negative charge (δ-) and the less electronegative element a slightly positive charge (δ+) resulting in a polar molecule. The arrow in the diagram below points to the more electronegative atom, and where the electrons tend to spend more time. HCl is a polar molecule because there is a dipole moment that isn’t canceled out

If a molecule is polar, it has a dipole moment.To know if a molecule is polar, you need to know if the bonds are polar and the overall shape of the molecule.

Symmetrical molecules tend to be nonpolar, and asymmetrical molecules tend to be polar.

Example (a), it must first be determined that carbon dioxide,CO2, is linear. Drawing in the arrows pointing to the more electronegative element in equal and opposite directions demonstrates that the molecule has polar bonds but is overall a nonpolar molecule.

Example (b) shows a water molecule. Because the molecular geometry is bent, drawing the arrows indicates that the dipoles do not cancel, therefore water is a polar molecule.

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