Molecular Models
EXPERIMENT#3
Name of Members: Ryv Gabor, Ian Mauro, Orlino Salvacion
Date Performed: September 26, 2018
I.
Date Submitted: October 2,2018
Objective
1. See the different shapes that molecules have.
2. Understand that polar bonds could be nonpolar.
3. See how different molecules’ bond angles vary.
4. Understand why the bond angles of different molecules vary.
II.
Material
1. small balls made of modeling clay
2. one box of multicolored toothpick
3. ruler
4. scissors
5. protractor
6. tape
7. sheet of notebook paper
8. sheet of construction paper
III.
IV.
Data & Results
Toothpicks
Red and
Blue
Red and
Green
Red and
Yellow
Blue and
Green
Blue and
Yellow
Green and
Yellow
Distance
Bold Angle
11 cm
109°
11 cm
109°
11 cm
109°
11 cm
109°
11 cm
109°
11 cm
109°
Answers to Questions
1. For methane, compare the bond angles in your model to those in the actual
molecule.
The actual methane molecule has a tetrahedral shape, which has 4 electron
bonds, which has a central atom is located at the center with four substituents
that are located at the corners of a tetrahedron. The bond angle is equal
109.4712206...° ≈ 109.5° when all four substituents are the same, nearly the
same as the molecular model we made.
2. Compare your models of methane, ammonia and water. Should the bond angles
in these models be equal? Why or why not?
In water (H2O), ammonia (NH3) and methane (CH4), we have four electron
pairs on the central atom. Four electron pairs should get as far apart as possible
giving a pyramidal shape, with a bold angle of 109.5°. This should correspond
with sp3 hybridization. However, in ammonia there is one lone pair and in water
two lone pairs. This creates variations in the geometry of the molecule. This
gives water a bent shape with the bond angle of 105°, and ammonia a pyramidal
shape, with the bond angle of 107°.
3. In part B, which are the polar molecules and which are nonpolar molecules?
Can a molecule with polar bonds be nonpolar? Explain.
Ammonia (NH3), Water (H2O) are polar, while Carbon Dioxide (CO2),
Oxygen (O2) and Boron Trichloride (BCl3) are. A molecule can possess polar
bonds and still be nonpolar. If the polar bonds are evenly (or symmetrically)
distributed, the bond dipoles cancel and do not create a molecular dipole. For
example, the three bonds in a molecule of BF3 are significantly polar, but they
are symmetrically arranged around the central boron atom. No side of the
molecule has more negative or positive charge than another side, and so the
molecule is nonpolar.
V.
Conclusion
Although Valence Shell Electron Pair Repulsion (VSEPR) allows us to predict the
shape of a molecule, it should still be verified by conducting experiments. This can be
proven in the 3rd question because a molecule with polar bonds can still be a non-polar
molecule.