4 - VSEPR Theory and Molecular Geometry
Abstract (TL;DR):
According to the electron configuration of an atom, the shapes that molecules take can be completely different.
——
Hydrogen bonds can slightly change a molecules bond angles.
Image via video by the Canadian Museum of Nature
With the naming of the molecules in the books, I promise you that you’ll never have to go through a similar gauntlet again (really, chemistry is the only discipline with this level of variation - kudos to all chemists for your mastery of the subject).
Molecules can get fairly hairy the larger they get. And we’ll see a few examples when we end up talking about biological topics. But, for now, I want to settle a few of the more bureaucratic odds and ends of molecules. Let’s start with bond angles and lengths.
Lewis Structures and Ball-And-Stick Structures
In the second section of this arc, I made sure to include two representations of most molecules. Those that showed the atoms and the bonds between them were Lewis structures - simplified two-dimensional models that show what’s going on between the atoms and the electrons they share. The other were “ball-and-stick” structures, three-dimensional models that are closer to a general representation of what the molecule actually looks like.
The difference between the two is that you can draw a Lewis structure however you wish, ignoring how those atoms react to each other within the molecule. Ball-and-stick models, on the other hand, have to be more accurate. These models are often computer-generated and used to show the shape of a molecule in a 3D space. An accurate 3D model can inform how two molecules interact with each other in any given space.
For the sake of our understanding, I made sure that both models looked similar. However, in reality, the bond angles and the length of those bonds between atoms are very much varied.
Bond Angles
As early as the Elemental Flow arc, we discussed how the energy of the atom increases the closer electrons get to each other. In response, electrons push each other away, sometimes even entering a different orbital to decrease energy as much as possible. So how, then, would an atom react to another nearby atom?
Well, what happens in the microscopic world sometimes happens in the larger, macroscopic world. This is one of those cases.
Just as electrons enter orbitals - particular places within atoms - to help keep their overall energy low, atoms enter specific positions in a molecule to keep their energy low. We denote those positions according to the angle between those atoms, or the bond angle.
Based on what we’ve already discussed, we know that atoms will try to not interfere with each other outside of forming bonds. Therefore, we can predict many of the bond angles between atoms as we know that atoms will tend to be as far away from each other as possible.
To use these predictive skills, we, first, have to go over a few molecules and acquaint ourselves with their bond angles and molecular geometry.
Molecular Geometry
Methane (Top - Lewis Structure; Bottom - Ball-And-Stick Structure)
First, take a look at these two models. Lewis structures, just by nature of their two-dimensional appearance, can’t give an accurate representation of the angles between each hydrogen and carbon atom in this methane molecule. In the ball-and-stick structure, however, we can see the three-dimensional nature of methane, including its bond angles. Chemists measured the angle between each atom in methane to be 109.5° - a tetrahedral.
Ammonia
Ammonia’s 3D model looks similar to methane’s, so much so that you might think that the loss of hydrogen wouldn’t cause that much of a change. That thought is correct (we’ll explain why in a bit), but the angle is still slightly different.
Figuring Out The Number of Valence Electrons*
Try to master the core-valence form. For all atoms, except transition metals, you can add the spin quantum numbers (the exponents) of each subshell to figure out the amount of valence electrons.
The slight difference is due to the presence of valence electrons in nitrogen. It’s orbital (in core-valence form) is [He]2s22p3 - meaning it has 5 valence electrons*.
That means, it has two valence electrons that are unbonded that aren’t shown in the above Lewis structure. Often, the lone pairs, the pairs of electrons that surround an atom which represent its unbonded valence electrons, are left out of Lewis structures for the sake of neatness - yet another possible pitfall of Lewis structures.
This one is quite insidious. For example, think about the amount of times you’ve seen the Lewis structure of water. Have you ever seen the molecule look like this - its more accurate form?
An accurate Lewis structure of water.
A more accurate representation of ammonia’s Lewis structure would include the pair of electrons on nitrogen. This pair of electrons influences surrounding atoms almost as much as an atom does, causing a bond angle of 107.3° - a trigonal pyramid.
These two molecules help to support the idea that, whether molecules are covalently bonded to atoms or contain lone pairs, the valence electrons determine the geometry of molecule and its bond angles. Given that the two geometries above differ by a little over 2 degrees, we can unite both under one electron-pair geometry, naming the general shape of the molecule according to how electrons are arranged around a central atom.
This organization of electron-pair geometries and their corresponding molecular geometries is called the valence-shell electron-pair repulsion (VSEPR) theory, a very useful way to predict what most molecules look like just by understanding their valence electrons.
| |
|
|
||||||
|---|---|---|---|---|---|---|---|---|
| Linear | 
Linear (Carbon Dioxide) |
~ | ~ | 180° | ||||
| Trigonal Planar | 
Trigonal Planar (Formaldehyde) |
Bent (Sulfur Dioxide) |
~ | 120° | ||||
| Tetrahedral | 
Tetrahedral (Methane) |
Trigonal Pyramid (Ammonia) |
Bent(Hydrogen Sulfide) |
~109° | ||||
Note that, if there are no lone pairs, the molecular geometry and the electron-pair geometry have the same name.
The “Bent” Geometry
Hydrogen Bonds in Water
Image via BioNinja
You might have noticed that both trigonal planar and tetrahedral electron-pair geometries can have a “bent” molecular geometry. Putting aside the ambiguous description of geometry, bent molecules are marked by a flexible bond angle, which can be anywhere from approximately 104° to 120° degrees.
The lone pairs present in the molecule are at fault. In bent molecules, the lone pair not only produce a force that affects a molecule’s bond angle, but can also interact with other molecules in the environment, which shifts the angle over time.
Above, I showed another example of this geometry, one of the most necessary molecule for life on Earth: water.
Water contains an oxygen atom, an atom with high electronegativity that has 6 valence electrons. Since this oxygen uses two of its electrons to form covalent bonds with hydrogen atoms, it has four free valence electrons, which remain unbonded in two lone pairs.
If this water molecule is in its natural, liquid state and it’s alone, then these free valence electrons push against the hydrogen atoms, bending them to create a 104.5° angle. But, in the case where it’s frozen into a solid and surrounded by other water molecules (think an ice cube), those electrons, attracted to hydrogen’s positive dipole, form a unique kind of bond called hydrogen bonds.
Ordinarily, in liquid water, these hydrogen bonds break and form rapidly in the presence of the hydrogen atoms of surrounding water molecules. But, when frozen, the electrons lock in their bond, changes the molecule’s bond angles to 109.5° each - a tetrahedral - as though the hydrogen bonds were stretching and pulling the molecule into place.
Image via video by icelcn
Freezing Liquid Water (striped line represents hydrogen bonds)
Improving the Lewis Structure
The Lewis structures might be worse than ball-and-stick structures when it comes to understanding bond angles, however we can make them better by adding new symbols. Take methane, for example.
Image via Boundless
3D Lewis Structure of Methane
The solid line, which exists in the usual Lewis structure, represents an atom that is in the same plane as the surface that you are drawing. The wedge that juts out to the front, however, represents an atom that is coming out towards the viewer. The dashed lines that points back, on the other hand, represents an atom that is pointing away from the viewer.
The combination of the three creates a “3D” Lewis structure, useful when talking about bond angles or explaining a molecule’s symmetry and shape, without computer programs.
A New Theory?
There’s one more part to these bond angles that will take another section to complete. Remember the connection that I made between molecular geometry for atoms and orbitals for electrons earlier? While the reaction of atoms to nearby atoms is similar to that of electrons near other electrons, orbitals and molecular geometry are still different phenomena. The kicker is that these phenomena interact with each other, creating an entirely new theory.
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