AP CHEMISTRYCHAPTER 9BONDING

Hybridization

When drawing Lewis structures to explain bonding, we have been using the Localized Electron Model of bonding. This assumes that the electrons stay with (or close to) the atoms from which they originated and that bonds are formed by the overlap of atomic orbitals. This model needs to be developed a little further to explain experimental data.

In methane, CH4, carbon has one 2s and three 2p orbitals available for bonding.

Hydrogen has one 1s orbital available. We could imagine that one of the bonds in methane would be formed from the overlap of a hydrogen 1s orbital and a carbon 2s orbital and the other three bonds would be formed from the overlap of a hydrogen 1s orbital and a carbon 2p orbital. This would make 2 different kinds of bonds. Experimental data shows that all four bonds

are identical.

Let’s look at the bonding in methane, CH4. Carbon has an electron configuration of 1s22s22p2. This means that there is one 2s orbital and three 2p orbitals available for bonding. Hydrogen has an electron configuration of 1s1. It uses a 1s orbital for bonding. Remember that s orbitals are spherical and p orbitals are “peanut” or “dumbbell” shaped. If these orbitals are used in their normal form for bonding, the result would be as shown below. One bond is different from the other three.

Bonding orbitals for C

However, experimental evidence shows that all 4 bonds in CH4 are identical! This model does not work!

Hybridization is an addition to the localized electron model that explains this. Carbon needs 4 identical orbitals to form 4 identical bonds. Imagine throwing the one 2s orbital and the three 2p orbitals from carbon into a blender. Whrrrrr!! Out comes four new orbitals that are a homogenous blend of the four

that went in!!

The new orbitals have properties of the s and of the p orbitals. Since they were made from one s and three p orbitals, we can call them sp3 orbitals. We put four atomic orbitals in and got four hybrid orbitals out. The hybrid sp3 orbitals each overlap with the 1s orbitals from each hydrogen atom and make four identical

bonds having 109.5o bond angles.

Atoms with four effective electron pairs have sp3 hybridization, even if they have unshared pairs as in H2O and NH3. The electron pairs have a tetrahedral arrangement.

We can look at the hybridization in the bonding of ethene (C2H4). H H C = C H H

Each carbon atom has three effective

pairs and needs three equal orbitals. To get three hybrid orbitals we throw one s and two p orbitals into the blender. Whrrrr!!! Out come three identical orbitals made from one s and two p orbitals. The three new orbitals are called sp2. This type of hybridization results in trigonal planar electron pair arrangements and 120o bond angles.

The overlap of the 1s orbitals from each hydrogen with the sp2 hybrid orbitals of carbon results in sigma bonds. In a sigma bond, the orbitals overlap head-on and the electrons are in the region between the two nuclei. All single bonds are sigma bonds. Each double bond and each triple bond contains one sigma

bond.

What happened to the p orbital that wasn’t hybridized? The remaining p

orbitals from each carbon atom overlap above and below the plane of the nuclei resulting in a sideways overlap that is called a pi bond.

A pi bond can form only if there is also a sigma bond between the same two atoms. Because the electrons in a pi bond are not directly between the nuclei, a pi bond is weaker than a sigma bond. A double bond contains one sigma plus one pi bond and is thus stronger and shorter than a single bond.

Let’s look at the hybridization and bonding in ethyne (acetylene), C2H2.

H-CC-H

Each carbon in ethyne has two effective electron pairs. We can throw one s and one p orbital from carbon into the blender. Whrrrr!!! Out come two sp orbitals. The s orbital from each hydrogen atom overlaps with one of the sp orbitals to form a sigma bond. The other sp orbital overlaps with one

of the sp orbital on the other carbon to form a sigma bond between the two C atoms.

The remaining two p orbitals from each carbon atom overlap outside the plane of the nuclei to form two pi bonds. A triple bond is one sigma and two pi bonds and is thus stronger and shorter than either a single or a double bond. Linear molecules and 180o bond angles result

from sp hybridization.

We learned that certain elements can exceed the octet rule by utilizing unfilled d

orbitals. These d orbitals may also be involved in hybridization. For example, PF5 has five effective electron pairs. These five pairs require five equal bonds. We must throw one s, three p and one d orbital into the blender to get five dsp3 orbitals. This form of hybridization results in trigonal bipyramidal shapes and 90o and 120o bond angles.

With SF6 we have six effective electron pairs. We must hybridize one s, three p, and two d orbitals to get six d2sp3 orbitals. This results in the octahedral shape with 90o bond angles.

News Flash! Recent experimental data (2011) show that d orbitals are not involved in hybridization. Until textbooks catch up, you may still see dsp3 and d2sp3 in test questions. The predicted shapes have not changed. You will not see dsp3 and d2sp3 on the AP test this year! You will simply not be asked to explain the hybridization of something with 5 or 6 electron pairs. You may still see dsp3 and d2sp3 on other types of tests.

Sigma bonds (σ)- formed by the end-to-end overlap of atomic orbitals between two nuclei

Pi bonds(π)- formed by the side-to-side overlap of unhybridized p orbitals outside the plane of two nuclei

σ bonds are stronger than π bonds because the electrons in a σ bond are between two nuclei.

Single bonds= 1 sigma bond

Double bonds = 1 sigma bond + 1 pi bondDouble bonds are stronger and shorter than single bonds.

Triple bonds = 1 sigma bond + 2 pi bondsTriple bonds are stronger and shorter than double bonds.

Practice:

For each of the following, determine:

a)hybridization of central atom(s)b)bond anglesc)molecular shaped)# of sigma bonds and # of pi bondse) molecular polarity

SbCl5

Cl Cl ClSb

Cl Cl

NO3−

O ON

O

CHCl3

HCl - C - Cl

Cl

:N C CN: C = C :N C CN:

MOLECULAR ORBITAL BONDING THEORY (MO)

The MO model is more complex than the localized electron model but works better in many instances.

A molecular orbital (MO) is formed from valence atomic orbitals. Two atomic orbitals combine to form two MOs. One MO has a higher energy than the atomic orbitals and is called an antibonding MO. The other MO has a lower energy (more stable) and is called a bonding MO.

In a bonding orbital, electrons have the greatest probability of being between the nuclei. In an antibonding orbital, electrons are mainly outside the space between the nuclei.

We can have and MOs. Orbitals are conserved. The number of MOs will be the same as the number of AOs used to construct them.

Bond order may be determined as a measure of the stability of a molecule. Using the localized electron model (Lewis structures), the bond order is equal to the number of bonds. For example, a triple bond has a bond order of 3.

The larger the bond order, the stronger the bond.

PARAMAGNETISM AND DIAMAGNETISM

Substances that contain unpaired electrons are weakly attracted into magnetic fields and are said to be paramagnetic. Substances in which all of the electrons are paired are very weakly repelled by a magnetic field and are said to be diamagnetic. Paramagnetism is much weaker than ferromagnetism (Fe, Co, Ni)

Figure 9.37Diagram of the Kind of Apparatus Used to Measure the Paramagnetism of a Sample

The Lewis structure of a molecule does not always correctly predict paramagnetism. The oxygen molecule is an example of this. O=OExperimental data shows that oxygen is paramagnetic (has unpaired electrons). The more complex molecular orbital model will correctly predict oxygen’s paramagnetism.

Paramagnetic or Diamagnetic?NO2

17 electrons

O = N − O

Resonance-To eliminate the need for resonance

we can use the localized electron model to describe the bonding and the MO model to describe the bonding.

In benzene, each C is sp2 hybridized. The 6 additional p orbitals perpendicular to the ring are used to form MOs. The electrons in these MOs are delocalized above and below the plane of the ring.

Figure 9.48The Pi System for Benzene