Molecular Geometry (Simplified): Study with Video Lessons, Practice Problems & Examples

Molecular geometry describes the shape of a molecule based on electron group repulsion. For 2 electron groups, the geometry is linear, as seen in beryllium chloride and carbon dioxide. With 3 electron groups, geometries can be trigonal planar (3 bonding groups) or bent (2 bonding groups, 1 lone pair). For 4 electron groups, geometries include tetrahedral (4 bonding groups), trigonal pyramidal (3 bonding groups, 1 lone pair), and bent (2 bonding groups, 2 lone pairs). Understanding these shapes is crucial for predicting molecular behavior and reactivity.

Molecular Geometry (Simplified) Concept 1

Now the molecular geometry is seen as the true shape of a molecule that takes into account differences in repulsion between lone pairs and surrounding elements. We're going to say because of this, they treat lone pairs and surrounding elements as different. So let's take a look at 2 electron systems. In 2 electron group systems, we're going to say this is where central elements with 2 electron groups have 0 lone pairs to give only one possible molecular geometry. Right. So the number of electron groups is 2, the number of bonding groups would be 2, and the number of lone pairs on the central element would be 0. Here are some good examples. We have Beryllium Chloride (BeCl₂), Carbon Dioxide (CO₂), and Hydrocyanic Acid (HCN). In all of them, the central element is connected to 2 bonding groups. And again, it doesn't matter if you're double bonded, triple bonded, or single bonded to that surrounding element, it only counts once. The visual representation is we have our central element and we have 2 surrounding groups or 2 surrounding elements. Because of this, we're going to say the molecular geometry would be linear. So when it comes to our molecular geometry, where the electron group number is 2, you can only have a linear shape.

Molecular Geometry (Simplified) Concept 2

Here we're going to say that central elements with 3 electron groups can have either 0 or one lone pair, which gives us 2 possible molecular geometries. Alright. So we have 3 electron groups, and the possibilities are you have 3 bonding groups, surrounding elements, and 0 lone pairs. Or you could have 2 bonding groups and 1 lone pair. Remember, adding up your bonding groups and your lone pairs together gives you back the number of electron groups. So with 3 bonding groups and 0 lone pairs, we have this example. Carbon is connected to 3 surrounding elements. The visual representation will be our central element in black and these 3 surrounding elements. Here, this would be trigonal or trigonal planar, or planar. Let's look at when you have 2 bonding groups and 1 lone pair.

Here we have tin in the center. It has 2 bonding groups in the form of those chlorines, and it has 1 lone pair. Here, the representation would look like this. Here, you would have to add in the fact that the lone pair is there, which is causing this bending of the bond. Now this particular shape goes by a few different names. You might see it written as bent. Others might see it written as V-shaped, or you might see it as angular. Now the first two kind of make sense. It's not a straight line; it's bent. Also, V-shape makes sense because it's an upside-down V. Angular is a little bit tricky, but realize here that any one of these three names are associated with this particular shape where you have 2 bonding groups and one lone pair on the central element. And remember, when you have 3 electron groups total, you can either have this as your molecular geometry or this as your molecular geometry.

Molecular Geometry (Simplified) Example 1

Determine the molecular geometry for the following molecule. We have boron and 3 chlorines. Alright. So boron is in group 3A. Chlorine is in group 7A, and there are 3 of them. So that's a total of 24 valence electrons. We place boron in the center, and it's going to form single bonds with the surrounding elements. Remember to provide enough electrons to the surrounding elements so they each follow the octet rule where they have 8 electrons around them. I've just now used all the electrons. I don't have any electrons remaining. Here, we'd say that our central element has what? It has 3 bonding groups, 0 lone pairs. So remember, we have 3 bonding groups and 0 lone pairs. Your molecular geometry would be trigonal planar. Also, I could see that it was going to be trigonal planar beforehand, so that's why I drew it in this way. Remember, the visualizations that I provided for each of the shapes mean you need to draw it that way so that it could really represent the true shape of the molecule. Now if you don't see it initially, go back and readjust your shape to make it look like this. Okay. So it should look something close as possible to the visualization in the videos that I've previously shown. Right? So just remember in this particular case, we're dealing with trigonal planar for this particular molecule.

Molecular Geometry (Simplified) Concept 3

Here, we're going to say that central elements with 4 electron groups can have anywhere from 0 to 2 lone pairs to give us 3 possible molecular geometries. So here we have 4 electron groups, and that can translate to having 4 bonding groups and 0 lone pairs, or 3 bonding groups and 1 lone pair, or 2 of each. Realize here when you add up your bonding groups with your lone pairs, it should give you back a total of 4 for the number of electron groups. Now when we have 4 bonding groups and 0 lone pairs, here we have a good example of CH₄, the visual representation is the way you should draw it when illustrating these different types of molecules. Here, the molecular geometry would be tetrahedral. For the next one, when it's 3 bonding groups and one lone pair, we have a good example in ammonia, which is NH₃. That one lone pair is up here. If you were to visually represent this, it would look like a pyramid, and that's why the name here is trigonal pyramidal. Trigonal because it's 3 corners of this pyramid, and pyramidal because it's a pyramid. Now when it's 2 and 2, this is interesting because when it's 2 and 2, we know that we have lone pairs here which is causing the spending and there was a lone pair up here. Here we've seen these terms before. When it's 2 and 2, it's bent, v-shaped, or angular. So if you watched my previous video, you would know that when we drew, 2 bonding groups and 1 lone pair, it also had the same shapes of bent, v-shaped, and angular. So just remember, it also applies when we have 2 bonding groups and 2 lone pairs on the central element. Now, again, when the more electron groups we have on the central element, the more possible shapes that arise. When we have 4 electron groups, we have these 3 possible shapes that arise. And for that last one, you could use any of those three names bent, v-shaped, or angular. They all mean the same thing.

Molecular Geometry (Simplified) Example 2

Determine the molecular geometry for the following ion of NH₄⁺, otherwise known as the ammonium ion. So, nitrogen is in group 5A, so we have 5 valence electrons. Hydrogen is in group 1A and there are 4 of them. Then, remember a plus one charge means we've lost an electron, so we'd subtract this from the total. That means we have 8 total valence electrons. Nitrogen will go in the center because hydrogen can never go in the center. Hydrogen only makes single bonds. Remember, each covalent bond that we just drew represents 2 valence electrons. So, we've just used all 8 of our electrons leaving us with none left. Because it's an ion, we put it in brackets and the charge in the top right corner.

At this point, we would say that our central element has connected to it 4 bonding groups, and it has 0 lone pairs. Remember, when you have 4 bonding groups and 0 lone pairs, your molecular geometry would be tetrahedral. So, we'd say here that the molecular geometry for the ammonium ion is tetrahedral. If we wanted to draw it in the correct way with our visual representation, it would look closer to something like this. The shape's a little bit weird in drawing when you look at it, so it's okay if you want to draw it as the first option that I drew. And again, you'd still put it in brackets with the charge in the top right corner.

Determine the molecular geometry for the following molecule: FSSF.

• A. Tetrahedral
• B. Trigonal planar
• C. Bent
• D. Linear

Here’s what students ask on this topic:

What is the molecular geometry of a molecule with 2 electron groups?

The molecular geometry of a molecule with 2 electron groups is linear. This is because the two electron groups will position themselves as far apart as possible, resulting in a straight line. Examples of such molecules include beryllium chloride (BeCl₂) and carbon dioxide (CO₂). In these cases, the central atom is bonded to two surrounding atoms, and there are no lone pairs on the central atom, leading to a linear shape.

How does the presence of lone pairs affect molecular geometry?

The presence of lone pairs affects molecular geometry by causing greater repulsion compared to bonding pairs. This repulsion forces the bonding pairs closer together, altering the shape of the molecule. For example, in a molecule with 3 electron groups, if there are 2 bonding pairs and 1 lone pair, the geometry changes from trigonal planar to bent or V-shaped. Similarly, in a molecule with 4 electron groups, if there are 2 bonding pairs and 2 lone pairs, the geometry is also bent or V-shaped.

What are the possible molecular geometries for a molecule with 3 electron groups?

For a molecule with 3 electron groups, there are two possible molecular geometries: trigonal planar and bent. If all three electron groups are bonding pairs, the geometry is trigonal planar, as seen in molecules like boron trifluoride (BF₃). If there are 2 bonding pairs and 1 lone pair, the geometry is bent or V-shaped, as seen in molecules like sulfur dioxide (SO₂).

What is the molecular geometry of methane (CH₄)?

The molecular geometry of methane (CH₄) is tetrahedral. Methane has 4 electron groups around the central carbon atom, all of which are bonding pairs. These electron groups arrange themselves as far apart as possible, resulting in a tetrahedral shape. This geometry is characterized by bond angles of approximately 109.5°.

How do you determine the molecular geometry of a molecule?

To determine the molecular geometry of a molecule, follow these steps: 1) Draw the Lewis structure of the molecule. 2) Count the total number of electron groups (bonding pairs and lone pairs) around the central atom. 3) Use the VSEPR (Valence Shell Electron Pair Repulsion) theory to predict the arrangement of these electron groups. 4) Identify the molecular geometry based on the positions of the bonding pairs, considering the repulsion caused by lone pairs. For example, a molecule with 4 electron groups and no lone pairs will have a tetrahedral geometry.