Because hydrogen’s valence electron is in the first shell, it cannot hybridize and so it only uses the s orbital. In the second shell, atoms can use their s or p orbitals. Atoms with valance electrons can use s, p and d orbitals. The hybridization choice depends on the number of areas of electron density around the atom. If there are three areas of electron density, for example, The atom will be sp2 hybridized.
 

Please examine the hybridization scheme for carbon shown below. To make these new hybrids, the atom takes the same number of orbitals to make the hybrids. Note that the number of orbitals does not change, just type types. In a case where carbon is attached to three things there will be three areas of electron density and so it will use two p orbitals (px+py) and one s orbital to make three sp2 orbitals leaving the pz alone.
 

	Areas of	orbitals before	orbitals after	Bond
	electron density	hybridization	hybridization	angles
	2	s+px+py+pz	2 sp+py+pz	180°
	3	s+px+py+pz	3 sp2+pz	120°
	4	s+px+py+pz	4 sp3	109.5°
	5	s+px+py+pz+d	4 sp3 d	120° & 90°
	6	s+px+py+pz+d+d	4 sp3 d2	90°

In methane, the Lewis structure shows 4 areas of electron density around the carbon. The VSEPR model shows the tetrahedral electron pair and tetrahedral molecular geometry. Examining the hybridization, each hydrogen is attached to the C through overlap of a sp3 orbital of the carbon and the s orbital of the hydrogen.
 

A single bond involves only a sigma overlap of orbitals between nuclei. The double bond involved sharing 4 electrons. Like a single bond, 2 electrons are shared between the nuclei. The next two electrons must be shared above and below this area. We define the shared electrons between the nuclei as a s bond. We define the shared electrons above and below this area as a p bond.
 
 

Prelab assignment:

1. Bring to lab scissors, scotch tape, a stapler and crayons or markers (or highlighters) with 4 colors, preferably red, green, blue and yellow.

2. Read the lab and chapter 10.1 of the book.

3. Draw the Lewis structures for the following molecules C₂H₆, C₂H₄, C₂H₂, NO₂, SO₂, CO₂. Make sure that you include all formal charges and resonance structures if there are any. Identify the electron pair geometry and the molecular geometry of each.

In lab:

You will be building models using styrofoam balls, construction paper and toothpicks. I plan that you will color coodinate your orbitals as follows:

orbital color

sp pink

sp2 blue

sp3 yellow

s white

p green

Your models will show overlapping bonds etc.

During the Lab:

Your instructor will assign groups of three. Your instructor will assign each person in the group a number one, two or three. During the presentation part of the lab, each person will be assigned a role based on his or her number. Before the presentation part of the lab, there will be a preparation part of the lab.

Preparation Time.

In order to prepare for your presentation, you will get together with another person who is presenting the same material in a different group. You will have to do the following things:

1. Show how you determined the correct Lewis structure including formal charges and resonance structures.
2. How you determined the electron pair geometry, the molecular geometry and the bond angles.
3. You will be required to make a model of your molecule showing ALL the electrons and the orbitals. The orbitals will be made from construction paper and the nuclei will be the Styrofoam balls. You will stick the construction paper to the balls using toothpicks.

Presentation Time.

There will be three roles that will rotate during the presentation time. You are receiving a group grade for this work so if at any time you do not understand what a member of your group is saying or doing, you should ask for a clarification. Similarly, if you see a mistake you should bring it up immediately. Your group report will be graded harshly so it is very important that the information it contains is completely correct.

The presenter is responsible for the stuff above. As presenter you may wish to use a new sheet of paper for your presentation or you may wish to use your laboratory notebook.

The checker after the completion of presenter's talk should verbally check the presenter's work.

1. Show that the Lewis structure and resonance structures are correct. (Obeys octet rule and has correct number of electrons.)
2. Show that all the electrons are in some type of orbital (s, p, sp, sp², sp³)
3. The VB model is consistent with the electron pair geometry, the molecular geometry and the bond angles.

1. Copying down the agreed upon Lewis structure including formal charges and resonance structures.
2. Copying down the agreed upon electron pair geometry, the molecular geometry and the bond angles.
3. Making a drawing of the 3-D valence bond model of the molecule.

The molecules and the role assignments are as follows:
 

	Presenter	Checker	Scribe
C2H6	1	2	3
C2H4	2	3	1
C2H2	3	1	2
NO2	1	2	3
SO2	2	3	1
CO2	3	1	2

Post Lab Questions:

For the following molecules, write down:

1. The Lewis structure including formal charges and resonance structures.
2. The electron pair geometry, the molecular geometry and the bond angles.
3. The 3-D valence bond model of the molecule (in color).

(Hint: Only O₃ has resonance structures)

Of course you will have to tell me which colors you chose for the sp3 d and sp3 d² orbitals.
 

s	p	sp	sp2	sp3	sp3 d	sp3 d2
white	green	pink	blue	yellow