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Lab 7: Bonding in Molecules and Lewis Structures.

Lab 7: Bonding in Molecules and Lewis Structures.

 

Objectives:

  1. Analyze Lewis structures of compounds and how that compares to ionic bonds (MO 6.1)
  2. Visualize the three dimensional geometry of different bonds and what causes them to deviate from the ideal bond ngle (MO 7.2)

Lab Objectives:

At the conclusion of this lab, students will be able to write the Lewis structure representations of the bonding and valence electrons in molecules. Students will also be able to utilize the VSEPR model and utilizing common household items, predict the molecular geometries (shapes) of molecules and determine whether a molecule is polar.

 

Materials:

  • Marshmallows (mini work best)
  • Toothpicks
  • Marker

 

Description

Molecular compounds are formed by sharing electrons between non-metal atoms. A useful theory for understanding the formation of molecular compounds, shapes of molecules and several other properties is called a Lewis dot structure. We will explore the use of Lewis dot theory to generate structures of molecular compounds and to build 3D models in order to predict shapes, and polarity of molecules.

 

Lewis structures are based on the concept that all atoms want filled orbitals. Usually atoms want to duplicate the 8 valence (or outer) electrons of the noble gases. This is commonly called the Octet Rule. This can be done in two ways. One is to exchange (or transfer) electrons, which is done in the formation of ionic bonds. Another is to share electrons in covalent bonds. Do molecules ever violate the Octet Rule? In this experiment, we will look at shared electrons in covalent bonds.

PROCEDURE:

For each covalent molecule listed in column 1 of Table 1, you will:

 

Step 1: Count the number of valence (or outer) electrons available for bonding and record in Table 1 (column 2).

 

Step 2: For every atom-to-atom connection, make a single bond (a shared electron pair).

 

Step 3: Distribute the remaining electrons available first to nonbonded pair positions (a maximum of six nonbonded pair electrons per atom, except H, which cannot have nonbonded pairs). Why can H not have nonbonded electrons?

 

Step 4: Identify atoms that need to have an octet but do not yet have one.

 

Step 5: Rearrange electrons from nonbonded pairs of adjacent atoms to multiple bonds to satisfy the octet rule on atoms identified in step 4. Count the electrons to make sure that you still have the same number as you started with! Draw your Lewis structure in Table 1 (column 3).

 

Step 7: Determine the molecular geometry or shape of the molecule and record in Table 1 (column 4).

 

Step 8: Calculate the formal charge of each atom and record in Table 1 (column 5).

 

Step 9: Analysis. Are the bonds nonpolar or polar covalent? Does the molecule have an overall dipole moment? How many resonance structures can you draw? Answer each in Table 1 (columns 6 and 7).

 

Step 10: Select ten molecules and build using marshmallows, toothpicks, and a marker.

  1. Label each atom (marshmallow) using a marker with the element’s symbol (see example below):
  1. Use toothpicks for each bond (one toothpick for a single bond, two toothpicks for a double bond, etc.).
  1. Build the molecule and clearly label

(example with methane):

 

METHANE

 

 

 

 

 

 

 

 

 

 

  1. Take a picture of the ten molecules that you built and submit a photo of each with your lab report.

 

 

 

Data Sheet

Submit Only Table 1 and Pictures

Submit pictures of the ten molecules that you built (10 points)

Table 1. Bonding in Molecules and Lewis Structures (10 points)

 

1. Molecule 2. Number of valence electrons 3. Lewis structure 4. Molecular Shape 5. Formal Charge (each element) 6. Polar molecule? 7. Resonance?
CH4

 

CH2Cl2

 

 

CH4O

 

H2O

 

H3O+

 

HF

 

N2

 

NH3

 

C2H4

 

C2H2

 

C2H2Br2

 

C2H6

 

SO42-

 

NO3

 

CO2

 

SCN

 

CO32-

 

HNO3

 

 

 

 

Lab 7: Bonding in Molecules and Lewis Structures.

 

 

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