Molecular Shapes

As you know from the Wrap-Up section of the last lesson, chemical substances cover a wide variety of bonding types (ionic, covalent, metallic) and structures (molecular, network, and atomic). For the moment, we will focus solely on covalent molecular substances. We will be learning a system for looking at the formula of a molecular compound and determining the shape of its molecules.

It will be useful here to have some examples. To start, use some scratch paper to sketch Lewis structures for the following molecules: CH₄, NH₃, H₂O, SO₃, SO₂, and CO₂. Review the Lewis Structures section from Lesson 8 if you need to.

Electron Groups

Let's focus on the central atom in each Lewis structure. In each case, to specify the shape of the molecule we want to describe the geometric arrangement of the atoms that are bonded to the central atom. This arrangement is what we call the molecule’s shape.

*Note: Larger molecules typically have more than one “central” atom, that is, an atom that is bonded to two or more other atoms. In these cases, we describe the “shape” of the molecule by describing the arrangement of atoms around each of these central atoms individually.

The arrangement of atoms around the central atom is, in turn, determined by the arrangement of the electron groups around the central atom. An electron group can be an unshared pair of electrons, a single bond, a double bond, or a triple bond.

All bonding electrons between two atoms count as one group:

the two electrons in a single bond are one group;
the four electrons in a double bond are one group;
the six electrons in a triple bond are one group

Each lone pair of electrons is also one group. Count the number of groups of electrons around the central atom for the six molecules we are working with, and check your answers below.

There are two places that students commonly get tripped up here.

First, they do not focus solely on the central atom. Note that, in the SO₃ molecule above, the lone pairs on the oxygen atoms are not shown in blue. They are irrelevant here.
Second, they forget that a triple or double bond counts the same as a single bond here. Any type of bond, no matter how many electrons are in it, counts as one group.

Electron Geometries

Because electrons repel each other, the groups of electrons around the central atom get as far away from each other as possible. This is what determines the arrangement of the groups of electrons. The central atom is a sphere, so we want to know what arrangement of two, three, and four groups of electrons gets them as far apart as possible on the surface of a sphere. Each of these arrangements is an electron geometry.

We are not directly interested in the geometry of electrons in a molecule; what we car about are the atoms. But determining the electron geometry is a necessary step in finding the arrangement of the atoms in space.

Two Electron Groups

When there are two groups of electrons, they move to exactly opposite sides of the central atom. This puts the two groups of electrons and the central atom in a straight line, so it is called a linear electron geometry. In such a geometry, the two electron groups are separated by a 180° angle.

There are two possible structures that can lead to a linear electron geometry. One is shown above in CO₂, with its two double bonds. The other possible arrangement is one single bond and one triple bond (like HCN, at left).

Three Electron Groups

When there are three groups of electrons around a central atom, the farthest they can get from each other is to move towards the corners of a flat triangle. This arrangement is called trigonal planar, and the atoms are separated by 120° angles. Important: when identifying this geometry, you must use both words: "trigonal planar."

In atoms that obey the octet rule, there are two possible structures that can produce a trigonal planar electron geometry. One is to have three bonded atoms (ignoring whether the bonds are single, double or triple), like in SO₃ above. The other is to have two bonded atoms and one lone pair, like in SO₂.

Four Electron Groups

When there are four groups of electrons around a central atom, they maximize the distance between them by moving to the corners of a tetrahedron. This arrangement is called tetrahedral. It is a little harder to see in three dimensions, but the atoms in a tetrahedron are separated by 109.5° angles.

In a tetrahedral arrangement, there are three possible configurations: four bonds and no lone pairs (like in CH₄), three bonds and one lone pair (like in NH₃), and two bonds plus two lone pairs (like in H₂O).

Molecular Geometries

The “shape” of a molecule is how the outside atoms are arranged around the central atom. We call this arrangement the molecular geometry. The lone pairs of electrons affect this arrangement, but are not part of it. It’s important to distinguish between the arrangement of groups of electrons and the arrangement of atoms around the central atom. If there are no lone pairs of electrons, these two arrangements will be the same, but if there are lone pairs the arrangement of groups will be different than the arrangement of atoms.

Four Electron Groups

As we have seen above, with four electron groups on a central atom, they arrange themselves in a tetrahedral geometry (reminder: this is the electron geometry). In the case of CH₄, above, where all four groups are single bonds, we also describe the molecular geometry as tetrahedral.

In the case of NH₃, however, we ignore the lone pair of electrons when describing the molecular shape. The pair is still there, repelling the three other electron groups, so it influences the shape of the molecule. But because the lone pair is invisible, we describe the molecular geometry as trigonal pyramidal. This means a pyramid with a triangular base. The three hydrogens form the base of the pyramid, with the nitrogen atom raised up above them. If you see three bonds and a lone pair on a central atom, this will be the molecular geometry.

Moving on to H₂O, we have a similar situation, but with two lone pairs. Those lone pairs, plus two bonds, make four groups, to give us a tetrahedral electron geometry. However, when we ignore the lone pairs and just look at the atoms, they form a bent arrangement. This arrangement is sometimes also known as "angular."

Water is a great example of why you can't just look at a Lewis structure to draw conclusions about molecular shape. If you just looked at the Lewis structure above for water, you might think that it is a linear molecule, with the H atoms at 180° apart from each other. Only by going through, identifying the electron geometry, and counting the lone pairs, can you conclude that water is a bent molecule. And this will be extremely important when analyzing the polarity of water and other molecules in the next section.

Three Electron Groups

With three electron groups on a central atom, they arrange themselves in a trigonal planar geometry.

*You should see now what it is so important to say "trigonal planar" and not just "trigonal." Writing just "trigonal" leaves it unclear whether you mean "trigonal planar" or "trigonal pyramid."

In the case of SO₃, above, without any lone pairs, the molecular geometry is also trigonal planar.

However, with SO₂, we have a lone pair. As with water, ignoring this lone pair leaves is with three atoms in a bent or angular arrangement.

Just as with water, it would be easy to mistakenly conclude SO₂ was linear if you weren't careful.

Two Electron Groups

With two electron groups on a central atom, the molecular geometry is always linear. These are the most simple cases to analyze. You just have to make sure that you truly do have only two groups, and aren't ignoring any lone pairs.

Practice

Your lab workbook has many examples of molecules for you to practice determining electron and molecular geometries. You should do so now and make sure you feel solid about this skill before moving on to molecular polarity.

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