CHAPTER 9 COVALENT BONDING: ORBITALS
9.1 Hybridization
Hybridization – mixing of 2 or more atomic orbitals of unequal energy to create new, hybrid orbitals of equal energy
Example: The carbon atom in methane
Number of orbitals being mixed must equal the number of hybrid orbitals formed:
Atomic Orbitals being mixed
Hybrid Orbitals formed
The easy way to predict the hybridization is to look at the number of electron domains that surround the central
atom. The number of electron domains = the number of atomic orbitals used for hybridization
Electron Domains
2
Hybridization
3
4
sigma bond – shared electrons likely to be found between the nuclei of the 2 bonded atoms.
pi bond – shared electrons likely to be found above and below the nuclei. Created by the overlapping
of nonhybridized p orbitals.
What types of bond(s) are in a single bond?
Double bond?
Triple bond?
Example:
Draw the Lewis structure for C2H4. Identify the hybridization for C and give the number of sigma and pi bonds.
Examples: Identify the hybridization of the center atom and give the number of pi and sigma bonds in the molecule or ion. Also
identify the molecular and electronic geometries.
1) PCl3
3) SO3
2) SiF62-
9.2 Molecular Orbital Model
So far, we have described bonding and molecular geometry in terms of valence bond theory. In valence bond theory, we postulate that
bonds result from the sharing of electrons in overlapping orbitals of different atoms. The orbitals may be pure atomic orbitals (i.e. s, p,
d, f, etc) or hybridized atomic orbitals of individual atoms (i.e sp, sp2, or sp3). We describe electrons in overlapping orbitals of
different atoms as being localized in the bonds between the two atoms involved. We can then use the VSEPR theory to determine the
molecular geometries.
However, the valence bond theory does not explain why in resonance structures we see equivalent bond lengths or why some
molecules such as O2 are paramagnetic and not diamagnetic. In molecular orbital theory, we postulate that the combination of
atomic orbitals on different atoms form molecular orbitals (MOs) so that electrons in them belong to the molecule as a whole
(delocalized electrons) and not to any one particular atoms.
Recall the wave behavior of electrons. With waves, we can have constructive and destructive interference. Mathematically,
constructive interference corresponds to addition and destructive interference corresponds to subtraction. When atomic orbitals
interact, we can either add them to obtain a bonding orbital or subtract them to obtain an antibonding orbital. Antibonding orbitals are
denoted with an asterisk (*).
Example: H2 molecule
Combination of Hydrogen 1s Atomic Orbitals to form MOs
One of the greatest successes of MO theory is that it accounts for electron delocalization in resonance. We have seen that some
molecules require resonance structures to be represented accurately. In all such cases, electrons are delocalized over several
bonds/atoms. One main drawback of the VB model is that it assigns electrons to specific atoms/bonds and therefore breaks down
when it comes to explaining delocalized electrons.
HOMEWORK: HYBRIDIZATION (Separate Sheet of Paper)
For each of the following molecules, write the Lewis structure, predict the molecular geometry (including bond angles) and electronic
geometry, give the expected hybridization of the central atom, and predict overall polarity.
(a) CF4
(c) TeF4
(e) IF3
(b) OF2
(d) AsF5
AP FRQ (Separate sheet of Paper)
1) The species represented below all have the same number of chlorine atoms attached to the central atom:
GeCl4
SeCl4
ICl4-
ICl4+
(a) Draw the Lewis Structure of each of the four species. Show all valence electrons in your structure.
(b) On the basis of the Lewis structures drawn in part (a), answer the following questions about the particular species indicated.
(i) What is the hybridization of Ge in GeCl4?
(ii) Is SeCl4 polar? Explain.
(iii) What is the Cl-I-Cl bond angle in ICl4-?
(iv) What is the geometric shape formed by the atoms in ICl4+?
2) Answer the following questions that relate to chemical bonding. (This hw problem should be done on separate paper!)
(a) Draw the complete Lewis structure for each of the three molecules represented below:
CF4
PF5
SF4
(b) On the basis of the Lewis structures drawn above, answer the following questions about the particular molecule
indicated.
(i) What is the F-C-F bond angle in CF4?
(ii) What is the geometry of the electron domains in PF5?
(iii) What is the geometric shape formed by the atoms in SF4?
3. Answer the following questions about the structure of ions that contain only sulfur and fluorine.
(a) The compounds SF4 and BF3 react to form an ionic compound according to the following equation:
SF4 + BF3  SF3BF4
(i) Draw a complete Lewis structure for the SF3+ cation in SF3BF4.
(ii) Identify the type of hybridization exhibited by sulfur in the SF 3+ cation.
(iii) Identify the geometry of the SF3+ cation that is consistent with the Lewis structure drawn in part (a)(i).
(iv) Predict whether the F-S-F bond angle in the SF3+ cation is larger than, equal to, or smaller than 109.5 o.
Justify your answer.
(b) The compounds SF4 and CsF react to form an ionic compound according to the following equation.
SF4 + CsF  CsSF5
(i) Draw the complete Lewis structure for the SF5- anion in CsSF5.
(ii) Identify the geometry of the SF5- anion that is consistent with the Lewis structure drawn in part (b)(i).
(iii) Identify the oxidation number of sulfur in the compound CsSF 5.