1. Molecular Shapes – VSEPR Theory
The more you study chemistry, the more you will come to appreciate that it’s all about the electron.
You have probably already learned that when two atoms form a covalent bond, they are in essence sharing a pair of electrons.
Molecules are made up of atoms bonded together this way, and most small molecules – in the 3-7 atom range – are comprised of one central atom with all the other atoms bonded to it, like the ones shown below:

2. Molecular Shapes – VSEPR Theory
But what exactly determines the shape of these molecules? What causes some to be straight, others to be bent? Some to be flat and others to be more 3-dimensional? Again, the answer lies in the electrons. Recall that electrons are negatively charged and therefore repel one another. Since the bonds are made up of electrons, these bonds always repel one another and orient themselves as far away from each other as possible.

3. Molecular Shapes – VSEPR Theory
This is known as “VSEPR” which stands for “Valence Shell Electron Pair Repulsion Theory.” Essentially it states that molecular shapes are determined by the repulsive forces acting between the electron pairs in the central atom’s outermost (valence) level.

4. Molecular Shapes – VSEPR Theory
These electron pairs include not only the bonding pairs that hold the surrounding atoms to the central atom, but also the nonbonding pairs that reside in the central atom’s valence level.
Lewis structures (AKA “electron dot structures”) are diagrams that illustrate exactly how electrons are being shared – and not shared – within a molecule.
Let’s look at some examples.

5. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BeF2.
Notice how the central atom (the Be) has just two electron regions on it. F
180° F Be
These electron regions would naturally repel one another and orient themselves as far away from each other as they can get.
This forces the molecule into what is called a “linear” shape.
The bond angle is 180°.

6. Molecular Shapes – VSEPR Theory
Since a linear molecule is flat with all bonded atoms existing in the same plane as the central atom, drawing one is fairly easy: F Be F Be F
Drawing:
Let’s look at another example:

7. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BF3.
Notice how the central atom (the B) has three electron regions on it. F
120° F B F
These electron regions would naturally repel one another and orient themselves as far away from each other as they can get.
This forces the molecule into a “trigonal planar” shape.
The bond angles are now 120°.

8. Molecular Shapes – VSEPR Theory
Trigonal planar molecules are also completely flat, so drawing one is rather simple. F B F F B F
Drawing:
Let’s look at another example:

9. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SO2.
Notice how the central atom (the S) again has three electron regions on it. O S O
These three regions repel one another,
but with no atom on top, the shape is simply called “bent” or “V-shaped.”
In fact, the nonbonding pair on top repels the bonding pairs more than the bonding pairs repel each other.

10. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SO2.
Notice how the central atom (the S) again has three electron regions on it. S
<120° O O
These three regions repel one another, but with no atom on top, the shape is simply called “bent” or “V-shaped.”
In fact, the nonbonding pair on top repels the bonding pairs more than the bonding pairs repel each other.
This squeezes the bond angle to something less than 120°.

11. Molecular Shapes – VSEPR Theory
A bent molecule like this is also flat and easy to draw: S O O O S
Drawing:
Let’s look at another example:

12. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for CH4.
Notice how this time the central atom has four electron regions on it. H H C H H
You might be thinking that this “+ shape” with all 90° bond angles would be the most spread out arrangement possible,
but a “+ shape” doesn’t take advantage of all the space available.
Keep in mind: no one said that molecules had to be flat.

13. Molecular Shapes – VSEPR Theory
The repelling force between the four bonding electron pairs forces the hydrogen atoms into the following arrangement: H H
109.5° C H H
This shape is known as “tetrahedral,” since it has the form of a symmetrical four-sided geometric solid known as a tetrahedron.
Its bond angles are all ° – or simply 109.5°.

14. Molecular Shapes – VSEPR Theory
Since a tetrahedral molecule is not flat, drawing one is a little trickier:
Note this dotted line bond. Dotted lines are used to show that a bond is receding away from you – behind the plane of the central atom. H
Note how these two other bonds are drawn as normal lines since they are still in the same plane as the central atom. C H H H
Note this wedge shaped bond. Wedges are used to show that a bond is coming out toward you – in front of the plane of the central atom. H C
Drawing:
Also, drawing the closer “H” a little larger and the more distant “H” a little smaller can help add to the three-dimensional appearance of the drawing.
Let’s look at another example:

15. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for NH3.
Again, notice how the central atom has four electron regions on it. H N H H
And though this Lewis structure makes the molecule look T-shaped,
remember: a Lewis structure is not meant to convey the shape of a molecule.
These four regions spread themselves out as they did before…

16. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for NH3.
Again, notice how the central atom has four electron regions on it. H N H H
And though this Lewis structure makes the molecule look T-shaped,
remember: a Lewis structure is not meant to convey the shape of a molecule.
These four regions spread themselves out as they did before…

17. Molecular Shapes – VSEPR Theory
And, like before, the nonbonding electron pair repels the bonding electron pairs more than the bonding electron pairs repel each other: N
<109.5° H H H
This pushes the H’s down into a shape known as “trigonal pyramidal.”
This causes the bond angles to be somewhat smaller than 109.5°.

18. Molecular Shapes – VSEPR Theory
Trigonal pyramidal molecules are drawn just like tetrahedral one, but with a nonbonding electron pair on the top N
Again notice the dotted line for the bond receding away from you and the wedge shape for the bond coming out toward you. H H H
Drawing: H N
Let’s look at another example:

19. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for H2O.
Once again, notice how the central atom has four electron regions on it. H H O
And these electron pairs again repel one another as they did before,
but with only two of the four electron regions involved in bonds, the molecule takes on a “bent” or “V-shaped” arrangement – similar to the SO2 molecule we saw before.

20. Molecular Shapes – VSEPR Theory
And once again, the nonbonding electron pairs repel the bonding pairs more than they repel one another… H O
<109.5° H
So the two H atoms get squeezed a little closer together.
And this gives this bent molecule a bond angle that is significantly less than 109.5°.

21. Molecular Shapes – VSEPR Theory
This bent molecule is like the previous one, except this one has two nonbonding electron pairs instead of just one: H O
It helps to put loops around the nonbonding electron pair and to have the closer pair drawn a little larger and overlapping the pair that is farther away. H H O
Drawing:
Let’s look at another example:

22. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for PF5.
Notice how the central atom now has five electron regions on it. F F F
120° P
90° F F
These electron regions repel one another and orient themselves as far away from each other as possible.
This gives rise to a shape known as “trigonal bipyramidal.”
And the bond angles are a combination of 90° and 120°.

23. Molecular Shapes – VSEPR Theory
With a trigonal bipyramidal molecule, the top, bottom and left hand atoms can all be drawn in the same plane as the central atom: F P F F F
This shape may be thought of as a combination of a linear (going straight up and down) F
and a trigonal planar perpendicular to the linear. F P
Drawing:
Let’s look at another example:

24. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SF4.
Again, the central atom has a total of five electron regions on it. F F S F F
These electron regions repel one another and orient themselves as far away from each other as they can get.
This arrangement is commonly known as “see-saw shaped.”

25. Molecular Shapes – VSEPR Theory
And, once again, the nonbonding electron pair on the left repels the bonding pairs more than they repel each other. F
<90° S F F
<120° F
This pushes the four F atoms to the right a bit.
This makes both the 90° and 120° bond angles a little smaller.

26. Molecular Shapes – VSEPR Theory
See-saw molecules are drawn just like trigonal bi-pyarmidal molecules, but with a nonbonding electron pair on the left: F S F F F F S
Drawing:
Let’s look at another example:

27. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BrF3.
Once again, the central atom has five electron regions on it.
I pitty the fool who can’t remember T-shaped! F F Br F
These five electron regions repel one another as they did before.
And this time the resulting shape is completely flat and is known as “T-shaped.”

28. Molecular Shapes – VSEPR Theory
Again, the nonbonding electron pairs on the right repel the bonding pairs more than the bonding pairs repel each other. F
<90° Br F F
That pushes the top and bottom F to the left.
And it squeezes the bond angles to a little less than 90°.

29. Molecular Shapes – VSEPR Theory
A T-shaped molecule is completely flat and therefore needs no dotted line or wedge-shaped bonds. It is drawn like the trigonal bipyramidal but with two nonbonding pairs on the right: F Br F F F Br
Drawing:
Let’s look at another example:

30. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for KrF2.
Once again, the central atom has five electron regions on it. F Kr F
These five electron regions repel one another as they did before.
And because all three nonbonding electron pairs arrange themselves around the perimeter, the resulting molecule is actually “linear.”

31. Molecular Shapes – VSEPR Theory
As before, the nonbonding electron pairs repel the bonding pairs more than the bonding pairs repel one another. F
180° Kr F
But since they are pushing equally from all sides, these extra repulsive forces have no effect on the bond angle. It remains a precise 180°.

32. Molecular Shapes – VSEPR Theory
This linear molecule is drawn the same way as the last one, only this one has three evenly spaced nonbonding electron pairs wrapped around the central atom: F Kr F F Kr
Drawing:
Let’s look at another example:

33. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SF6.
Notice how the central atom now has six electron regions on it. F
Notice how the bond angles are all now precisely 90°. 3 F F 4 2
90° S 1 F F 7 8 6 F 5
These electron regions would naturally repel one another and orient themselves as far away from each other as they can get.
This pushes the atoms into a shape known as “octahedral.”
An octahedron is a geometric solid with eight identical faces.

34. Molecular Shapes – VSEPR Theory
An octahedral molecule is best drawn with the top and bottom atoms drawn in the same plane as the central atom, then two atoms coming out toward us and two others receding away: F F F S F F F S F
Drawing:
Let’s look at another example:

35. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BrF5.
Once again, the central atom has six electron regions on it. F F F Br F F
These regions repel one another as they did for octahedral,
but with one of the atoms replaced by a non-bonding pair, the resulting shape is called “square pyramidal”
since it resembled a square based pyramid”
– like the ones in Egypt.

36. Molecular Shapes – VSEPR Theory
Once again, the nonbonding electron pair on bottom repel the five bonding pairs more than they repel one another. F F F
<90° Br F F
This pushes the four nearby F atoms up a bit.
This makes all the bond angles a little less than 90°.

37. Molecular Shapes – VSEPR Theory
A square pyramidal molecule is drawn just like the octahedral one, but with a nonbonding pair of electrons in place of the bottom atom. F F F F Br F Br F
Drawing:
Let’s look at one final example:

38. Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for KrF4.
Once again, the central atom has six electron regions on it. F F Kr F F
These regions repel one another as they did for octahedral,
but with the top and bottom atoms replaced by non-bonding pairs, the resulting shape is called “square planar.”

39. Molecular Shapes – VSEPR Theory
Once again, the nonbonding electron pairs on top and bottom repel the four bonding pairs more than they repel one another. F F Kr
90° Kr F F
But since they are pushing equally from both sides, these extra repulsive forces have no effect on the bond angle. So it remains a precise 90°.

40. Molecular Shapes – VSEPR Theory
A square planar molecule is drawn just like the octahedral one, but with nonbonding pairs of electrons in place of the top and bottom atoms. F F Kr Kr F F Kr F
Drawing:
So now let’s try some sample problems.

41. The next 15 slides will give you practice in drawing these molecular shapes.
For each slide, draw and name the shape and give the bond angle on a sheet of scrap paper – or, if you prefer, directly on the slide using the inking tool.
The answer will then be given.
If your answer does not precisely match the one given, go back and try it again.
Remember: the more you practice, the better you will get.

42. Cl P Cl Cl P Cl 1. phosphorus trichloride:
Explanation: Since the central atom has four electron regions on it, think tetrahedral. But one of these regions is a NEP (nonbonding electron pair), so the shape is trigonal pyramidal. Cl P
trigonal pyramidal, <109.5°

43. S C S C S 2. carbon disulfide: linear, 180°
Explanation: Since the central atom has two electron regions on it, neither of which is a NEP, the shape is linear. (It doesn’t matter that these electron regions happen to be double bonds: a double bond counts as just one region.) C S
linear, 180°

44. O O S O S O 3. sulfur trioxide: trigonal planar, 120°
Explanation: Since the central atom has three electron regions on it, none of which are NEP’s, the shape is trigonal planar.
trigonal planar, 120°

45. Se Br Se Br Br Br Br 4. selenium tetrabromide:
Explanation: Since the central atom has five electron regions on it, think trigonal bipyramidal. But one of these regions is a NEP, so the shape is see-saw.
see-saw, <90°, <120°

46. F S F F S 5. sulfur difluoride: Bent, <109.5°
Explanation: Since the central atom has four electron regions, think tetrahedral. But two of these regions are NEP’s so the shape is bent.
Bent, <109.5°

47. Xe Xe F F F 6. xenon difluoride: linear, 180°
Explanation: Since the central atom has five electron regions on it, think trigonal bipyramidal. But three of these regions are NEP’s, so the shape is linear.
linear, 180°

48. Cl B Cl Cl B Cl 7. boron trichloride: trigonal planar, 120°
Explanation: Since the central atom has three electron regions on it, none of which are NEP’s, the shape is trigonal planar.
trigonal planar, 120°

49. H H N H H N H 1+ 1+ 7. ammonium ion: tetrahedral, 109.5°
Explanation: Since the central atom has four electron regions on it, none of which are NEP’s, the shape is tetrahedral. The fact that it is an ion rather than a neutral molecule does not have any impact on its shape.
tetrahedral, 109.5°

50. F F F Rn F F Rn 8. radon tetrafluoride: square planar, 90°
Explanation: Since the central atom has six electron regions on it, think octahedral. But two of these regions are NEP’s, so the shape is square planar.
square planar, 90°

51. As As I I I I I I 9. arsenic pentaiodide:
Explanation: Since the central atom has five electron regions on it, none of which are NEP’s, the shape is trigonal bipyramidal.
trigonal bipyramidal, 90°, 120°

52. Te Cl Te Cl Cl Cl Cl Cl Cl 10. tellurium hexachloride: octahedral, 90°
Explanation: Since the central atom has six electron regions on it, none of which are NEP’s, the shape is octahedral.
octahedral, 90°

53. O Se O O Se 11. selenium dioxide: bent/V-shaped, <120°
Explanation: Since the central atom has three electron regions on it, think trigonal planar. But one of these regions is a NEP, so the shape is bent/V-shaped.
bent/V-shaped, <120°

54. You really ought to remember T-shaped!
12. iodine trichloride: Cl I Cl Cl
You really ought to remember T-shaped! Cl I
Explanation: Since the central atom has five electron regions on it, think trigonal bipyramidal. But two of these regions are NEP’s, so the molecule is T-shaped.
T-shaped, <90°

55. H H Si H H Si H 13. silicon tetrahydride (silane): tetrahedral, 109.5°
Explanation: Since the central atom has four electron regions on it, none of which are NEP’s, the shape is tetrahedral.
tetrahedral, 109.5°

56. F As F F As F 14. arsenic trifluoride: trigonal pyramidal, <109.5°
Explanation: Since the central atom has four electron regions on it, think tetrahedral. But one of these regions is a NEP, so the shape is trigonal pyramidal. F As
trigonal pyramidal, <109.5°

57. Br Cl Cl Cl Cl Cl Cl Br 15. bromine pentachloride:
Explanation: Since the central atom has six electron regions on it, think octahedral. But one of these regions is a NEP, so the shape is square pyramidal.
square pyramidal, <90°