chapter10

 Chapter 10 Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals

Discussion Terms for Chapter 10: Valence Shell: The outermost shell occupied by electrons in an atom VSEPR Model: (Valence Shell Electron-Pair Repulsion) Model that uses the valence shell of an atom to determine its molecular geometry Dipole Moment: The measure of the polarity of a bond. Can be determined using the molecular geometry and bond polarities Polar Molecule: Any molecule that has a dipole moment is considered to be polar. Diatomic molecules of two different elements (CO, NO, OH, etc.) are always polar. Nonpolar Molecule: Any molecule that does not have a dipole moment is considered to be nonpolar. Diatomic molecules of the same element (F 2, H 2 , O 2 , etc.) are nonpolar. Valence Bond Theory: A theory on molecular bonding stating that a bond between two atoms is caused by the overlapping of two molecular orbitals. Lone Pair: any electron pair on a central atom that is not shared in a bond.

VSEPR Model: The VSEPR models and molecular geometry of a molecule depends 100% on the number of electrons in the valence shell of its element. Consider CO 2 ; the carbon has four electron in its valence shell and each oxygen has six, so the total availible electrons is 16. divide this number by two to obtain the number of electron pairs (8). Now you would set up the picture as such O C O and begin to add the pairs. The goal is to have each atom in the molecule be sharing and holding eight electrons or 4 pairs. CO2 becomes O==C==O and each O atom has two pairs of electrons as well. Because the central atom (C) has no lone pairs, this molecule is considered linear and has bond angles of 180 degrees.

The following are all of the possible VSEPR models for atoms with no lone pairs: Two Bonds or Linear (angles of 180 degrees) Three bonds or Trigonal Planar (angles of 120 degrees) Four Bonds or Tetrahedral (angles of 109.5 degrees) Five Bonds or Trigonal Bipyrimidal (angles of 90 and 120 degrees) Six Bonds or Octahedral (angles of 90 degrees)

When molecules have lone pairs, their geometry and VSEPR model changes. These are all the possible VSEPR models with one lone pair:

Two Bonds and One Lone Pair or Bent (angles of 120 degrees) Three Bonds and One Lone Pair or Trigonal Pyramidal (angles of 109.5 degrees) Four Bonds and One Lone Pair or Seesaw (angles of 120 and 180 degrees) Five Bonds and One Lone Pair or Square Pyramidal (angles of 90 degrees)

The next amount of lone pairs on a central atom is two. When two lone pairs are part of a molecule, the geometry and VSEPR model changes once again. These are all the possible VSEPR models for molecules with two lone pairs: Two Bonds and Two Lone Pairs or Bent (angles of 109.5 degrees) Three Bonds and Two Lone Pairs or T-Shaped (angles of 90 and 180 degrees) Four Bonds and Two Lone Pairs or Square Planar (angles of 90 degrees)

The last type of molecular geometry possible is found in atoms with three lone pairs. There is only one possible VSEPR Model amongst all atoms with this make-up. Two Bonds and Three Lone Pairs (angles of 90 degrees)

Hybridization Hybridization- The mixing of atomic orbitals in an atom (usually the central atom) to generate a set of hybrid orbitals.

When finding the hyrbidization of an molecules, there are a few rules that must be followed. There are four types of orbitals (S, P, D, and F), each of which can hold a different number of electrons. Any S orbital can hold a maximum of two electrons, any P orbital can hold a maximum of six electrons, any D orbital can hold a maximum of ten electrons, and any F orbital can hold a maximum of fourteen electrons. There are mulitple of each orbital level, and these are differentiated by their quantum number. The appearance of the orbitals on the periodic table of elements and the sequence in which these orbitals fill (1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p) is located in the pictures below

When drawing the hyrbridization of a molecule, you use a simple system of drawing to map out the electrons. For each orbital (s,p,d,&f), there is a different drawing. These drawings are as follows: __** sp Hybridization **__ Using the example molecule of BeCl 2, we know that a Be atom has two valence electrons and has a quantum number of two, so it will have this hybridization drawing: Because the Be atom has its 2s orbital full, it would not bond with the two Cl atoms, but hybridization explains this. The electrons in the 2s orbital of the Be atom enter an excited state and split apart leaving one behind in the 2s orbital and adding one electron to the 2p orbital. This leaves what are called two equivalent sp hybrid orbital and two empty 2p orbitals. This configuration gives the two chlorine atoms the opportunity to bond and form a linear molecule because once the orbitals become hybridized the bonds become equal.

**__Other Hybridizations__** When the central atom of molecule has more valence electrons, the hybridization becomes more advanced. There are many types of hybridizations which are: sp, sp2, sp3, sp3d, sp3d2. The hybridizations of sp2 and sp3 are explained on [] and the hybridization on the sp3d and sp3d2 are found on [] and []

When two atoms share electrons the bond that is formed is called a sigma bond. This type of bond is formed between two S orbitals such as in H--H.
 * __Hyrbidization of Double and Triple Bonds__ **

When a bond is formed between two P bonds, it creates a sigma bond and another bond called a Pi bond. This type of bond only occurs in dobule and triple bonding scenarios. In double bonds there are one pi bond and one sigma bond. In triple bonds there are two pi bonds and one sigma bond. The formation of sigma and pi bonds are explained well in this video: [] To see whether a molecule will have sigma and/or pi bonds, a molecular orbit energy level diagram can be completed. Using the vlance shell of the two atoms that form the bond, the electrons first combine to form a sigma bond (σ) then if there are more electrons to be shared they combine in the sigma antibond ( σ*). If the molecules that are bonding have P atomic orbitals then pi bonds will form and the diagram will have an additional level.

__**VSEPR Labratory**__

**Name: _**  __**Lab: Lewis Structures and VSEPR Diagrams**  __**Procedure:**__
 * Period: _**
 * Purpose:** to practice drawing Lewis Structures and assigning VSEPR shapes to molecules.__

__**(1) Draw the Lewis Structure for each molecule.**  __**Observations:**__ __
 * (2) Determine the VSEPR shape for each molecule.**
 * (3) Using the materials provided, create a VSEPR model for each of the molecules**__

__(1) Draw a three dimensional representation of the model for each molecule. Colour the diagram and create a legend. (2) Classify the shape of each molecule Conclusion:__ __What does VSEPR stand for__? What does it show about a molecule?
 * Questions:** (complete on the paper provided) __

Questions and Answers for Chapter 10:

1. What bond angles are present in a Hydronium (H3O+) ion?

An H30+ ion has three bonds and one lone pair. This means it is trigonal pyrimidal and has bond angles of 109.5 degrees.

2. Are the following molecules planar: XeF4? PCl3? BCl3?

XeF4 has four bonds and no lone pairs. This makes it tetrahedral and planar. PCl3 has three bonds and one lone pair. This makes it trigonal pyrimidal and non-planar BCl3 has three bonds and no lone pairs. This makes it trigonal planar and polar.

3. Does the SO2 ion have a dipole moment?

Yes, the SO2 molecule has two bonds and one lone pair. It has a bent shape and is non-polar. All these factors combined cause a dipole moment in the SO2 ion.

4. Which ion has smaller bond angles, SO3 or SO4(-2)?

SO4(-2) has smaller bond angles. When the VSEPR models of the two ions are compared, SO4(-2) is square planar with angles of 109.5 degrees and SO3 has a bent shape with angles of 120 degrees. SO4(-2)SO3

5. Explain why an atom cannot have a permanent dipole moment?

An atom cannot have a permanent dipole moment because the dipole moment is dependent upon the bond moment which is constantly changing.

6. Draw the Molecular Orbit Energy Level Diagram of O2.




 * all information that is uncited in this page is extracted from Chemistry 8th Edition by Raymond Chang**