CHEMISTRY ACTIVITY—MOLECULAR GEOMETRY
Go to this website: http://phet.colorado.edu/en/simulation/molecule-shapes-basics
Click on the green "Run Now" button.
1. Start with a central atom (purple), with a single grey atom attached. If your model has
more than one atom attached, remove the others by clicking on the red X in the upper
right hand corner.
2. When you add a second atom, by clicking on the single bonded atom in the upper
right hand corner, what angle do the two grey atoms make?
180o (a line)
3. Add a third bond; what is the angle now? 120o (a triangle)
Is this molecule flat or three-dimensional? 2D (a flat triangle)
4. Add one more atom; how does this change the molecule? (Two answers!)
The bond angles change and it is now a 3D molecule
5. Check the "Options" box to show bond angles. Are all bond angles the same, or are
they different? Click on the "name" box in the lower left to see the name of this shape:
The bond angles are all 109.5o and it is tetrahedral in shape
6. Remove one of the bonds, and determine the shape of the molecule with three bonded
atoms:
The bond angles are all 120o and it is trigonal planar in shape
7. Choose the "Real Molecules" tab, and predict the shape and bond angles of:
BeCl2: linear, 180o (whatever you predict)
BF3: trigonal planar. 120 o (whatever you predict)
CH4: tetrahedral, 109.5o (whatever you predict)
Check the "Name" and "Bond Angle" boxes to check your answers.
8. Go back to "Model" and "Remove All" bonded atoms. Then add two double-bonded
atoms. How is this shape different from your molecule in #1? (HINT: check the bond
angles and molecular geometry).
It isn’t – it is linear with 180o bond angles.
Now, go to this website: http://phet.colorado.edu/en/simulation/molecule-shapes
9. Select the "Real Molecules" tab; uncheck "show lone pairs."
10. Predict the shape and bond angle in H2O:
bent, 120o? bond angles (whatever you guess)
Check "show bond angles," and check "name molecular geometry."
11. Now, report its actual shape and bond angle:
bent, 104.5o bond angles
12. Why are they (possibly) different than your predictions? Check the "show lone pairs"
box, and describe what you see:
The lone pairs change the shape, because they “take up space” as negatively charged
particles.
13. Rotate the molecule by holding down the mouse and swiveling it. Re-examine the
CH4 molecule (pull-down menu). How do bond angles in H2O differ from those in CH4?
CH4 has 109.5o bond angles, a little bigger than water.
14. How many electron pairs are there around the C atom in CH4? (Remember that each
bond is a pair.) How many around the O in H2O?
Four electron pairs are around the Carbon, and 4 are around the Oxygen.
Realize, then, that the 4 electron pairs around the central atom in H2O and in CH4 are
not much different, except that some are unbonded (lone) in H2O. This accounts for the
differences in shape and bond angle, since LONE PAIRS REPEL just like bonding pairs;
in fact, they repel even more than bonding pairs!
15. Look at the NH3 molecule; check and uncheck the "lone pair" box as necessary.
What similarities do you see to H2O and/or CH4? What differences?
There is one lone pair of electrons in NH3; in between what H2O and CH4 have.
16. How does the bond angle in NH3 compare with that of H2O & CH4? Is this molecule
2- or 3-dimensional?
The bond angles are in between H2O and CH4, having one lone pair does change the bond
angles as much as having 2 lone pairs. The molecule is 3D
17. Draw the Lewis Structure for each of CH4, NH3, and H2O. Be sure to include the lone
pairs!
See Table 12.4 on p. 389!
18. How many electron pairs around each central atom in CH4/NH3/H2O? How many of
these pairs are used in bonding? Refer to table 12.4 (p. 389). How do the lone
pairs/bonding pairs affect the shapes?
CH4 has 4 pairs of electrons on the central atom, all of which are involved in bonds. NH3
has 4 pairs of electrons on the central atom, three of which are involved in bonds. H2O
has 4 pairs of electrons on the central atom only two of which are involved in bonds. The
more. As the number of nonbonding pairs increases, the bond angle gets smaller and
the overall shape of the molecule changes from tetrahedral to trigonal pyramidal to bent.
GO TO: http://phet.colorado.edu/en/simulation/molecule-polarity
Select "Two Atoms," view "Bond dipole," and "Partial charges." Select either "Electrostatic
Potential" OR "Electron Density." Electric Field off. Set both atoms to medium
electronegativity values.
19. Decrease the electronegativity value of ONE atom, and record what happens to the
electrostatic potential/electron density (HINT: arrow & partial charges)
The dipole moment arrow and the partial charge symbols get bigger as the
electronegativity of one of the atom decreases. The one with reduced electronegativity is
the “positive” end, and the other atom is the negative end.
REALIZE THAT THESE STEPS ARE NOT POSSIBLE IN REAL LIFE—WE CANNOT
CHANGE THE ELECTRONEGATIVITY OF AN ELEMENT!
20. Increase the other atom's electronegativity, and record your results.
A large/strong dipole moment is created; the one with lowest electronegativity is the
“positive” end, and the one with the highest electronegativity is the negative end.
21. Turn on the electric field and describe how the molecule orients itself paying
attention to the partial charges.
The “positive” end is aligned with the negative side of the field and the negative end is
aligned with the positive side of the field.
22. Reverse the electronegativity values of the atoms and record your results.
When the values were reversed, the molecule flipped itself so that the “positive” end is
aligned with the negative side of the field and the negative end is aligned with the positive
side of the field.
Select "Three Atoms," and check "Bond Dipoles" and "Molecular Dipoles." Set all atoms
to medium electronegativities.
23. Decrease the electronegativity of C to the minimum value and record your
observations of both bond and molecular dipoles. Remember that the head of the arrow
points in the negative direction.
The bond dipole arrow points from C to B, and the molecular dipole arrow points the
same direction but is shown “leaving” B.
24. PREDICT what would happen if you were to minimize the electronegativity of atom A
and C at the same time, leaving B alone.
The bond dipole and the molecular dipole arrows should point straight up, making B the
negative pole, and A and C the positive.
25. Try minimizing A and C while leaving B alone, and record your results.
The bond dipoles point from A to B and from C to B, and the molecular dipole arrow
points straight up, making B the negative pole, and A and C the positive.
26. Now turn on the Electric Field, and record your observations in detail: how does the
molecule respond to the electric field specifically?
The molecule orients itself so that the yellow molecular dipole arrow is pointing straight
at the positive side of the electric field.
Finally, select "Real Molecules," and check both (bond and molecule) dipoles, "atom
electronegativites," and "Atom Labels."
27. Describe the bond polarities and molecular polarities of the following molecules:
MOLECULE
BOND POLARITY (grey arrows)
MOLECULAR POLARITY (yellow arrows)
H2
none
none
O2
none
none
H2O
From each H to O
straight up from O
HF
From H to F
straight out from F
BF3
From B out to each F
none
CO2
From C out to each O
none
NH3
From each H to N
straight up from N
CH4
From each H to C
none
CH3F
From each H to C, from C to F
straight up from F
28. What can you conclude about molecular polarity from your analysis of the molecules
in #25? (HINT: compare the dipoles of CH4 and CH3F.)
The more electronegative element gets the dipole moment but the molecular dipole only
occurs when the geometry of the atom arrangements allows for there to be an “end of the
molecule” with a negative charge.