11.5: Multimedia Resources for Chapter 11
Copy and distribute the lesson worksheets. Ask students to complete the worksheets alone or in pairs as a review of lesson content.
Molecular Geometry Worksheet
CK-12 Foundation Chemistry
Name______________________ Date_________
Lewis structures only show how many bonding pairs of electrons, and unshared pairs of electrons, surround a given atom on a flat page. The molecules are actually three dimensional which is not shown by Lewis structures. To convey a sense of three dimensionality, we use “ball and stick” models.
There is a correlation between the number of electron pairs, (sigma bonds plus non-shared pairs) around the central atom of a molecule, and the electronic geometry of that molecule.
The idea that allows us to predict the electronic geometry is that each pair of electrons (shared or unshared) repels all the other electron pairs. The electron pairs move as far apart as possible, but since they are all tied to the central atom, they can only orient themselves in such a way that they make the angles between them as large as possible. This is the essence of the Valence Shell Electron Pair Repulsion (VSEPR) Theory for predicting molecular shapes.
To use VSEPR theory, we must first be able to determine the number of valence shell electron pairs around the central atom. These pairs consist of all sigma bond pairs and all unshared pairs of electrons. Pi bond electrons are excluded because the electrons are not placed between bonding atoms and therefore, do not contribute to electronic geometry.
Electron Pairs | Image |
---|---|
To visualize the electron pairs that contribute to electronic geometry, imagine them situated on the surface of a sphere with the central atom at the center. | |
If there are only two pairs of electrons in the valence shell of the central atom, the two pairs can avoid each other best if they are \begin{align*}180^ \circ \end{align*} apart. This means that the two pairs and the central atom are in a straight line; the arrangement is linear. | |
If a third pair of electrons is added, the three pairs push around to the shape shown at right. The angles between electron pairs would be \begin{align*}120^ \circ\end{align*} and we call the shape trigonal planar. The three pairs of electrons and the central atom are all in a single plane. | |
A fourth pair of electrons causes the electrons to push around into the shape shown at right, the tetrahedron. The angles in this shape are \begin{align*}109.5^ \circ \end{align*}. | |
A fifth pair of electrons produces the shape known as trigonal bipyramidal. The angles between the three pairs of electrons around the center is \begin{align*}120^ \circ\end{align*} and the angles between the pairs around the center and the pairs on the ends is \begin{align*}90^ \circ\end{align*}. | |
Finally, the sixth pair of electrons produces the octahedral shape shown at right. All angles in this shape are \begin{align*}90^ \circ\end{align*}. |
Once the number of electron pairs surrounding the central atom is determined, the electronic geometry is known.
Electron Pairs Around the Central Atom | Electronic Geometry |
---|---|
1 pair | Linear |
2 pairs | Linear |
3 pairs | Trigonal Planar |
4 pairs | Tetrahedral |
5 pairs | Trigonal Bipyramidal |
6 pairs | Octahedral |
The molecular geometry may be different from the electronic geometry because many times, not all the electron pairs are shared. An unshared electron pair will not have an atom in that position of the electronic geometry. In order to determine molecular geometry, we must recognize which pairs of electrons have an atom attached and which are lone pairs. The overall shape of the molecule is determined by how many pairs of electrons are around the central atom, and how many of these have atoms attached.
It is sometimes difficult for students to recognize the difference between the orientation of electron pairs (called electronic geometry) and the overall shape of the molecule (called molecular geometry). We will look at an example that shows the difference between electronic and molecular geometry. Consider the following four molecules: hydrogen chloride, \begin{align*}HCl\end{align*}; water, \begin{align*}H_2O\end{align*}; ammonia, \begin{align*}NH_3\end{align*}; and methane, \begin{align*}CH_4\end{align*}.
Shared Pairs | Molecular Geometry |
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The central atom of each of these molecules is surrounded by four pairs of electrons. According to VSEPR theory, these four pairs will be oriented in three-dimensional space to be as far away from each other as possible. The four pairs will point to the corners of the geometrical shape known as a tetrahedron. The angles between the electron pairs will be approximately \begin{align*}109.5^ \circ\end{align*}. In all four cases, the electronic geometry is tetrahedral but only one of the molecules will have tetrahedral molecular geometry. | |
In the case of \begin{align*}HCl\end{align*}, even though there are four pairs of electrons around the chlorine atom, three of them are not shared. There is no atom attached to them. These spaces are empty. Since there are only two atoms joined by a bond, the molecular geometry will be linear. | |
In the water molecule, two electron pairs are shared and two are unshared. So while the electronic geometry is tetrahedral, the molecular geometry is bent (aka angular, aka V-shaped). | |
In the ammonia molecule, one pair of electrons is unshared and the other three are shared. This results in a molecular shape called pyramidal. | |
In the methane molecule, all four pairs of electrons are shared, and so not only is the electronic geometry tetrahedral but the molecular geometry is also tetrahedral. |
Central Atom Electron Pairs | Electronic Geometry | Bonding Pairs | Molecular Geometry | Sketch |
---|---|---|---|---|
\begin{align*}2\end{align*} | Linear | \begin{align*}2\end{align*} | Linear | |
\begin{align*}3\end{align*} | Trigonal Planar | \begin{align*}1\end{align*} | Linear | |
\begin{align*}3\end{align*} | Trigonal Planar | \begin{align*}2\end{align*} | Bent | |
\begin{align*}3\end{align*} | Trigonal Planar | \begin{align*}3\end{align*} | Trigonal Planar | |
\begin{align*}4\end{align*} | Tetrahedral | \begin{align*}1\end{align*} | Linear | |
\begin{align*}4\end{align*} | Tetrahedral | \begin{align*}2\end{align*} | Bent | |
\begin{align*}4\end{align*} | Tetrahedral | \begin{align*}3\end{align*} | Pyramidal | |
\begin{align*}4\end{align*} | Tetrahedral | \begin{align*}4\end{align*} | Tetrahedral | |
\begin{align*}5\end{align*} | Trigonal Bipyramidal | \begin{align*}1\end{align*} | Linear | |
\begin{align*}5\end{align*} | Trigonal Bipyramidal | \begin{align*}2\end{align*} | Linear | |
\begin{align*}5\end{align*} | Trigonal Bipyramidal | \begin{align*}3\end{align*} | T-shape | |
\begin{align*}5\end{align*} | Trigonal Bipyramidal | \begin{align*}4\end{align*} | Distorted Tetrahedron | |
\begin{align*}5\end{align*} | Trigonal Bipyramidal | \begin{align*}5\end{align*} | Trigonal Bipyramidal | |
\begin{align*}6\end{align*} | Octahedral | \begin{align*}1\end{align*} | Linear | |
\begin{align*}6\end{align*} | Octahedral | \begin{align*}2\end{align*} | Linear | |
\begin{align*}6\end{align*} | Octahedral | \begin{align*}3\end{align*} | T-shape | |
\begin{align*}6\end{align*} | Octahedral | \begin{align*}4\end{align*} | Square Planar | |
\begin{align*}6\end{align*} | Octahedral | \begin{align*}5\end{align*} | Square Pyramidal | |
\begin{align*}6\end{align*} | Octahedral | \begin{align*}6\end{align*} | Octahedral |
In order to choose the correct molecular geometry, you must keep in mind that only electron pairs involved in sigma bonds and unshared pairs contribute to electronic geometry. Pi bonds are not directed bonds, and those electron pairs do not contribute to electronic geometry. In the Lewis structure for the carbon dioxide molecule (shown at right), it is clear that the central atom is carbon, and the carbon atom is surrounded by 4 pairs of electrons. But these fours pairs of electrons are involved in two sigma bonds and two pi bonds. Therefore, the electronic geometry of carbon dioxide is based on two pairs of electrons around the central atom, and will be linear. Since both pairs of electrons are shared, the molecular geometry will also be linear.
The Lewis structure for the carbonate ion, shown at right, shows the central atom is carbon and it is surrounded by \begin{align*}4\end{align*} electron pairs. One of those pairs, however, is a pi bond, and therefore the electronic geometry of the carbonate ion is based on 3 pairs of electrons around the central atom. Thus, the electronic geometry is trigonal planar and since all three pairs are shared, the molecular geometry is also trigon planar.
Polarity
Bonds between atoms that are of the same element are non-polar bonds. Molecules composed of all the same atom such as \begin{align*}Cl_2\end{align*}, \begin{align*}O_2\end{align*}, \begin{align*}H_2\end{align*}, \begin{align*}S_8\end{align*}, \begin{align*}P_4\end{align*}, have no polar bonds and therefore do not have dipoles. That is, the molecules will be non-polar. A molecule that does have polar bonds can still be non-polar. If the polar bonds are symmetrically distributed, the bond dipoles cancel and do not produce a molecular dipole.
Molecular Shape | Symmetry |
---|---|
Linear | Symmetrical |
Bent | Non-Symmetrical |
Trigonal Planar | Symmetrical |
Pyramidal | Non-Symmetrical |
Tetrahedral | Symmetrical |
T-shaped | Non-Symmetrical |
Distorted Tetrahedron | Non-Symmetrical |
Trigonal Bipyramidal | Symmetrical |
Square Planar | Symmetrical |
Square Pyramidal | Non-Symmetrical |
Octahdral | Symmetrical |
Exercises
Fill in the table with electronic geometry, molecular geometry, and indicate whether the molecular will be polar or non-polar.
Formula | Electronic Geometry | Molecular Geometry | Polarity |
---|---|---|---|
\begin{align*}AsH_3\end{align*} | |||
\begin{align*}BCl_3\end{align*} | |||
\begin{align*}IF_3\end{align*} | |||
\begin{align*}SiBr_4\end{align*} | |||
\begin{align*}SeH_4\end{align*} | |||
\begin{align*}XeI_4\end{align*} | |||
\begin{align*}OF_2\end{align*} | |||
\begin{align*}KrF_2\end{align*} | |||
\begin{align*}ICl_5\end{align*} | |||
\begin{align*}CCl_2F_2\end{align*} |
Answers to Worksheets
- The worksheet answer keys are available upon request. Please send an email to teachers-requests@ck12.org to request the worksheet answer keys.
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