6.4: Molecular Structure and Polarity
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- Scott Van Bramer
- Widener University
Learning Objectives
- Predict the structures of small molecules using valence shell electron pair repulsion (VSEPR) theory
- Explain the concepts of polar covalent bonds and molecular polarity
- Assess the polarity of a molecule based on its bonding and structure
Thus far, we have used two-dimensional Lewis structures to represent molecules. However, molecular structure is actually three-dimensional, and it is important to be able to describe molecular bonds in terms of their distances, angles, and relative arrangements in space (Figure \(\PageIndex{1}\)). A bond angle is the angle between two adjacent bonds. A bond length is the distance between the nuclei of two bonded atoms. Bond distances are measured in Ångstroms (1 Å = 10 –10 m) or picometers (1 pm = 10 –12 m, 100 pm = 1 Å).
VSEPR Theory
Valence shell electron-pair repulsion theory (VSEPR theory) is used to predict the structure around a central atom of a molecule based on the number of bonds and lone electron pairs in the Lewis structure. The VSEPR model assumes that electron pairs in the valence shell of a central atom arrange to minimize repulsions by maximizing the distance between these electron pairs. The electrons in the valence shell of a central atom form either bonding pairs of electrons, located primarily between bonded atoms, or lone pairs. The electrostatic repulsion of these electrons is reduced when the various regions of high electron density assume positions as far from each other as possible.
VSEPR theory predicts the arrangement of electron pairs around each central atom and the arrangement of atoms in a molecule. We should understand, however, that the theory is a first approximation and that fully understanding the structure of a molecule is much more complex.
A gaseous BeF 2 molecule is a simple example to introduce VSEPR. The Lewis structure of BeF 2 (Figure \(\PageIndex{2}\)) shows two electron pairs around the central beryllium atom. There are two bonds and no lone pairs of electrons on this beryllium atom so according to VSEPR the bonds should be as far apart as possible. The electron density from the bonds repel and this repulsion is minimized when the bonds are on opposite sides of the central atom so the bond angle is 180° (Figure \(\PageIndex{2}\)).
Figure \(\PageIndex{3}\) illustrates electron-pair geometries for different structures that minimize the repulsions among regions of high electron density (bonds and/or lone pairs). Two regions of electron density around a central atom in a molecule form a linear geometry; three regions form a trigonal planar geometry; four regions form a tetrahedral geometry; five regions form a trigonal bipyramidal geometry; and six regions form an octahedral geometry. The spatial relationship, diagrams for the structure, names for the structures, and bond angles are all shown in this figure.
Electron-pair Geometry versus Molecular Structure
It is important to note that electron-pair geometry around a central atom is not the same thing as its molecular structure. The electron-pair geometries shown in Figure \(\PageIndex{3}\) describe all regions where electrons are located, bonds as well as lone pairs. Molecular structure describes the location of the atoms , not the electrons.
We differentiate between these two situations by naming the geometry that includes all electron pairs the electron-pair geometry . The structure that includes only the placement of the atoms in the molecule is called the molecular structure . The electron-pair geometries are the same as the molecular structures when there are no lone electron pairs around the central atom. When a central atom has lone pair electrons the molecular structure is different from the electron-pair geometry.
For example, the methane molecule, CH 4 , which is the major component of natural gas, has four bonding pairs of electrons around the central carbon atom; the electron-pair geometry is tetrahedral, as is the molecular structure (Figure \(\PageIndex{4}\)). On the other hand, the ammonia molecule, NH 3 , also has four electron pairs associated with the nitrogen atom, and thus has a tetrahedral electron-pair geometry. One of these regions, however, is a lone pair, which is not included in the molecular structure, and this lone pair influences the shape of the molecule (Figure \(\PageIndex{5}\)). Since the electron-pair geometry and the molecular structure are different for ammonia we need two different names for these geometries.
In the ammonia molecule, the three hydrogen atoms attached to the central nitrogen are not arranged in a flat, trigonal planar molecular structure, but rather in a three-dimensional trigonal pyramid (Figure \(\PageIndex{5}\)) with the nitrogen atom at the apex and the three hydrogen atoms forming the base. The ideal bond angles in a trigonal pyramid are based on the tetrahedral electron pair geometry. There are slight deviations from the ideal because lone pairs occupy larger regions of space than bonding electrons so the H–N–H bond angle in NH 3 is slightly smaller than the 109.5° angle tetrahedron. (Figure \(\PageIndex{6}\)) summarizes the possible elecvtron-pair geometries and the corresponding molecular structures for up to six electron pairs.
Please note that because the lone pair-bonding pair repulsion is greater than the bonding pair-bonding pair repulsion the bond angles for real molecules will differ slightly from those shown in Figure \(\PageIndex{6}\). The ideal molecular structures are predicted based on the electron-pair geometries.
Small distortions from these ideal angles are shown for ammonia in Figure \(\PageIndex{5}\). The distortions result from differences in repulsion between various regions of electron density. VSEPR theory predicts the magnitude of these distortions by establishing an order of repulsions and an order of the space occupied by different types of electron pairs. The order of electron-pair repulsions from greatest to least repulsion is:
lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair
Different types of electron pairs also need different amounts of space. A lone pair of electrons occupies a larger region of space than the electrons in a triple bond; in turn, electrons in a triple bond occupy more space than those in a double bond, and so on. The order of sizes from largest to smallest is:
lone pair > triple bond > double bond > single bond
Taken together the amount of repulsion and the amount of space can help determine the location of the lone pair electrons in the molecular structure and the amount that expected bond angles are distorted in actual molecules. Consider formaldehyde, H 2 CO shown in Figure \(\PageIndex{1}\). This molecule has three regions of high electron density, two single bonds and one double bond. The basic geometry is trigonal planar with predicted 120° bond angles. However, the double bond takes up more space and pushes the single bonds further away so the angle between the single bonds is slightly smaller (118°) than expected.
According to VSEPR theory, the terminal atom locations (Xs in Figure \(\PageIndex{7}\)) are equivalent within the linear, trigonal planar, and tetrahedral electron-pair geometries (the first three rows of the table). It does not matter which X is replaced with a lone pair because the molecules can be rotated and all the positions are equivilent. For trigonal bipyramidal electron-pair geometries, however, there are two distinct X positions (Figure \(\PageIndex{7}\)a): an axial position (if we hold a model of a trigonal bipyramid by the two axial positions, we have an axis around which we can rotate the model) and an equatorial position (three positions form an equator around the middle of the molecule). The axial position is surrounded by bond angles of 90°, whereas the equatorial position has more space available because of the 120° bond angles. In a trigonal bipyramidal electron-pair geometry, lone pairs always occupy equatorial positions because these more spacious positions can more easily accommodate the larger lone pairs.
The molecule ClF 3 is an example of a compound with trigonal bipyramid electron-pair geometry and two sets of lone pair electrons. Figure \(\PageIndex{7}\) shows the three possible arrangements for the three C-F bonds in the ClF 3 molecule. The most stable structure is with the lone pairs in equatorial locations so they have the most space. This gives ClF 3 a T-shaped molecular structure as shown in Figure \(\PageIndex{7}\)b.
Similarly for compounds with octahedral electron-pair geometry, when a central atom has lone electron pairs there are several different possible molecular structures. With two lone pairs they orient on opposite sides of the octahedron (180° apart), giving a square planar molecular structure that minimizes lone pair-lone pair repulsions. With three lone pairs they orient to give a T shaped molecule and with four lone pairs they orient to give a linear molecule. All these options are summarized in Figure \(\PageIndex{6}\).
Predicting Electron Pair Geometry and Molecular Structure
The following procedure uses VSEPR theory to determine the electron pair geometries and the molecular structures:
- Write the Lewis structure of the molecule or polyatomic ion.
- Count the number of regions of electron density (lone pairs and bonds) around the central atom. A single, double, or triple bond counts as one region of electron density.
- Identify the electron-pair geometry based on the number of regions of electron density: linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral (Figure \(\PageIndex{7}\), first column).
- Use the number of lone pairs to determine the molecular structure (Figure \(\PageIndex{7}\) ). If more than one arrangement of lone pairs and chemical bonds is possible, choose the one that will minimize repulsions, remembering that lone pairs occupy more space than multiple bonds, which occupy more space than single bonds. In trigonal bipyramidal arrangements, repulsion is minimized when every lone pair is in an equatorial position. In an octahedral arrangement with two lone pairs, repulsion is minimized when the lone pairs are on opposite sides of the central atom.
The following examples illustrate the use of VSEPR theory to predict the molecular structure of molecules or ions that have no lone pairs of electrons. In this case, the molecular structure is identical to the electron pair geometry.
Example \(\PageIndex{1}\): Predicting Electron-pair Geometry and Molecular Structure
Predict the electron-pair geometry and molecular structure for each of the following:
- carbon dioxide, CO 2 , a molecule produced by the combustion of fossil fuels
- boron trichloride, BCl 3 , an important industrial chemical
Solution
(a) We write the Lewis structure of CO 2 as:
This shows us two regions of high electron density around the carbon atom—each double bond counts as one region, and there are no lone pairs on the carbon atom. Using VSEPR theory, we predict that the two regions of electron density arrange themselves on opposite sides of the central atom with a bond angle of 180°. The electron-pair geometry and molecular structure are identical, and CO 2 molecules are linear.
(b) We write the Lewis structure of BCl 3 as:
Thus we see that BCl 3 contains three bonds, and there are no lone pairs of electrons on boron. The arrangement of three regions of high electron density gives a trigonal planar electron-pair geometry. The B–Cl bonds lie in a plane with 120° angles between them. BCl 3 also has a trigonal planar molecular structure.
The electron-pair geometry and molecular structure of BCl 3 are both trigonal planar. Note that the VSEPR geometry indicates the correct bond angles (120°), unlike the Lewis structure shown above.
Exercise \(\PageIndex{1}\)
Carbonate, \(\ce{CO3^2-}\), is a common polyatomic ion found in various materials from eggshells to antacids. What are the electron-pair geometry and molecular structure of this polyatomic ion?
- Answer
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The electron-pair geometry is trigonal planar and the molecular structure is trigonal planar. Due to resonance, all three C–O bonds are identical. Whether they are single, double, or an average of the two, each bond counts as one region of electron density.
Example \(\PageIndex{2}\): Predicting Electron-pair Geometry and Molecular Structure
Two of the top 50 chemicals produced in the United States, ammonium nitrate and ammonium sulfate, both used as fertilizers, contain the ammonium ion. Predict the electron-pair geometry and molecular structure of the \(\ce{NH4+}\) cation.
Solution
We write the Lewis structure of \(\ce{NH4+}\) as:
Exercise \(\PageIndex{2}\)
Identify a molecule with trigonal bipyramidal molecular structure.
- Answer
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Any molecule with five electron pairs around the central atoms including no lone pairs will be trigonal bipyramidal. \(\ce{PF5}\) is a common example
The next several examples illustrate the effect of lone pairs of electrons on molecular structure.
Example \(\PageIndex{3}\): Lone Pairs on the Central Atom
Predict the electron-pair geometry and molecular structure of a water molecule.
Solution
The Lewis structure of H 2 O indicates that there are four regions of high electron density around the oxygen atom: two lone pairs and two chemical bonds:
Exercise \(\PageIndex{3}\)
The hydronium ion, H 3 O + , forms when acids are dissolved in water. Predict the electron-pair geometry and molecular structure of this cation.
- Answer
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electron pair geometry: tetrahedral; molecular structure: trigonal pyramidal
Example \(\PageIndex{4}\): SF 4 Sulfur tetrafluoride,
Predicting Electron-pair Geometry and Molecular Structure: SF 4 , is extremely valuable for the preparation of fluorine-containing compounds used as herbicides (i.e., SF 4 is used as a fluorinating agent). Predict the electron-pair geometry and molecular structure of a SF 4 molecule.
Solution
The Lewis structure of SF 4 indicates five regions of electron density around the sulfur atom: one lone pair and four bonding pairs:
Exercise \(\PageIndex{4}\)
Predict the electron pair geometry and molecular structure for molecules of XeF 2 .
- Answer
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The electron-pair geometry is trigonal bipyramidal. The molecular structure is linear.
Example \(\PageIndex{4}\): XeF 4
Of all the noble gases, xenon is the most reactive, frequently reacting with elements such as oxygen and fluorine. Predict the electron-pair geometry and molecular structure of the XeF 4 molecule.
Solution
The Lewis structure of XeF 4 indicates six regions of high electron density around the xenon atom: two lone pairs and four bonds:
Exercise \(\PageIndex{4}\)
In a certain molecule, the central atom has three lone pairs and two bonds. What will the electron pair geometry and molecular structure be?
- Answer
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electron pair geometry: trigonal bipyramidal; molecular structure: linear
Molecular Structure for Multicenter Molecules
When a molecule or polyatomic ion has only one central atom, the molecular structure completely describes the shape of the molecule. Larger molecules do not have a single central atom, but are connected by a chain of interior atoms that each possess a “local” geometry. The way these local structures are oriented with respect to each other also influences the molecular shape, but such considerations are largely beyond the scope of this introductory discussion. For our purposes, we will only focus on determining the local structures.
Example \(\PageIndex{5}\): Predicting Structure in Multicenter Molecules
The Lewis structure for the simplest amino acid, glycine, H 2 NCH 2 CO 2 H, is shown here. Predict the local geometry for the nitrogen atom, the two carbon atoms, and the oxygen atom with a hydrogen atom attached:
Solution
Consider each central atom independently. The electron-pair geometries:
- nitrogen––four regions of electron density; tetrahedral
- carbon ( C H 2 )––four regions of electron density; tetrahedral
- carbon ( C O 2 )—three regions of electron density; trigonal planar
- oxygen ( O H)—four regions of electron density; tetrahedral
The local structures:
- nitrogen––three bonds, one lone pair; trigonal pyramidal
- carbon ( C H 2 )—four bonds, no lone pairs; tetrahedral
- carbon ( C O 2 )—three bonds (double bond counts as one bond), no lone pairs; trigonal planar
- oxygen ( O H)—two bonds, two lone pairs; bent (109°)
Exercise \(\PageIndex{5}\)
Another amino acid is alanine, which has the Lewis structure shown here. Predict the electron-pair geometry and local structure of the nitrogen atom, the three carbon atoms, and the oxygen atom with hydrogen attached:
- Answer
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electron-pair geometries: nitrogen––tetrahedral; carbon ( C H)—tetrahedral; carbon ( C H 3 )—tetrahedral; carbon ( C O 2 )—trigonal planar; oxygen ( O H)—tetrahedral; local structures: nitrogen—trigonal pyramidal; carbon ( C H)—tetrahedral; carbon ( C H 3 )—tetrahedral; carbon ( C O 2 )—trigonal planar; oxygen ( O H)—bent (109°)
Example \(\PageIndex{6}\): Molecular Simulation
Using this molecular shape simulator allows us to control whether bond angles and/or lone pairs are displayed by checking or unchecking the boxes under “Options” on the right. We can also use the “Name” checkboxes at bottom-left to display or hide the electron pair geometry (called “electron geometry” in the simulator) and/or molecular structure (called “molecular shape” in the simulator).
Build the molecule HCN in the simulator based on the following Lewis structure:
Click on each bond type or lone pair at right to add that group to the central atom. Once you have the complete molecule, rotate it to examine the predicted molecular structure. What molecular structure is this?
Solution
The molecular structure is linear.
Exercise \(\PageIndex{6}\)
Build a more complex molecule in the simulator. Identify the electron-group geometry, molecular structure, and bond angles. Then try to find a chemical formula that would match the structure you have drawn.
- Answer
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Answers will vary. For example, an atom with four single bonds, a double bond, and a lone pair has an octahedral electron-group geometry and a square pyramidal molecular structure. XeOF 4 is a molecule that adopts this structure.
Molecular Polarity and Dipole Moment
As discussed previously, polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge (δ+) and the other atom with a partial negative charge (δ–), as the electrons are pulled toward the more electronegative atom. This separation of charge gives rise to a bond dipole moment. This bond moment can be represented as a vector, a quantity having both direction and magnitude (Figure \(\PageIndex{12}\)). Dipole vectors are shown as arrows pointing along the bond from the less electronegative atom toward the more electronegative atom. A small plus sign is drawn on the less electronegative end to indicate the partially positive end of the bond. The length of the arrow is proportional to the magnitude of the electronegativity difference between the two atoms.
A whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. If such a charge separation exists, the molecule is said to be a polar molecule (or dipole); otherwise the molecule is said to be nonpolar. The dipole moment measures the extent of net charge separation in the molecule as a whole. We determine the dipole moment by adding the bond moments in three-dimensional space, taking into account the molecular structure.
For diatomic molecules, there is only one bond, so its bond dipole moment determines the molecular polarity. Homonuclear diatomic molecules such as Br 2 and N 2 have no difference in electronegativity, so their dipole moment is zero. For heteronuclear molecules such as CO, there is a small dipole moment. For HF , there is a larger dipole moment because there is a larger difference in electronegativity.
When a molecule contains more than one bond, the geometry must be taken into account. If the bonds in a molecule are arranged such that their bond moments cancel (vector sum equals zero), then the molecule is nonpolar. This is the situation in CO 2 (Figure \(\PageIndex{13A}\)). Each of the bonds is polar, but the molecule as a whole is nonpolar. From the Lewis structure, and using VSEPR theory, we determine that the CO 2 molecule is linear with polar C=O bonds on opposite sides of the carbon atom. The bond moments cancel because they are pointed in opposite directions. In the case of the water molecule (Figure \(\PageIndex{13B}\)), the Lewis structure again shows that there are two bonds to a central atom, and the electronegativity difference again shows that each of these bonds has a nonzero bond moment. In this case, however, the molecular structure is bent because of the lone pairs on O, and the two bond moments do not cancel. Therefore, water does have a net dipole moment and is a polar molecule (dipole).
The OCS molecule has a structure similar to CO 2 , but a sulfur atom has replaced one of the oxygen atoms. To determine if this molecule is polar, we draw the molecular structure. VSEPR theory predicts a linear molecule:
The C–O bond is considerably polar. Although C and S have very similar electronegativity values, S is slightly more electronegative than C, and so the C-S bond is just slightly polar. Because oxygen is more electronegative than sulfur, the oxygen end of the molecule is the negative end.
Chloromethane, CH 3 Cl, is another example of a polar molecule. Although the polar C–Cl and C–H bonds are arranged in a tetrahedral geometry, the C–Cl bonds have a larger bond moment than the C–H bond, and the bond moments do not completely cancel each other. All of the dipoles have a upward component in the orientation shown, since carbon is more electronegative than hydrogen and less electronegative than chlorine:
When we examine the highly symmetrical molecules BF 3 (trigonal planar), CH 4 (tetrahedral), PF 5 (trigonal bipyramidal), and SF 6 (octahedral), in which all the polar bonds are identical, the molecules are nonpolar. The bonds in these molecules are arranged such that their dipoles cancel. However, just because a molecule contains identical bonds does not mean that the dipoles will always cancel. Many molecules that have identical bonds and lone pairs on the central atoms have bond dipoles that do not cancel. Examples include H 2 S and NH 3 . A hydrogen atom is at the positive end and a nitrogen or sulfur atom is at the negative end of the polar bonds in these molecules:
To summarize, to be polar, a molecule must:
- Contain at least one polar covalent bond.
- Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.
Properties of Polar Molecules
Polar molecules tend to align when placed in an electric field with the positive end of the molecule oriented toward the negative plate and the negative end toward the positive plate (Figure \(\PageIndex{14}\)). We can use an electrically charged object to attract polar molecules, but nonpolar molecules are not attracted. Also, polar solvents are better at dissolving polar substances, and nonpolar solvents are better at dissolving nonpolar substances.
Example \(\PageIndex{7}\): Polarity Simulations
Open the molecule polarity simulation and select the “Three Atoms” tab at the top. This should display a molecule ABC with three electronegativity adjustors. You can display or hide the bond moments, molecular dipoles, and partial charges at the right. Turning on the Electric Field will show whether the molecule moves when exposed to a field, similar to Figure \(\PageIndex{14}\).
Use the electronegativity controls to determine how the molecular dipole will look for the starting bent molecule if:
- A and C are very electronegative and B is in the middle of the range.
- A is very electronegative, and B and C are not.
Solution
- Molecular dipole moment points immediately between A and C.
- Molecular dipole moment points along the A–B bond, toward A.
Exercise \(\PageIndex{7}\)
Determine the partial charges that will give the largest possible bond dipoles.
- Answer
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The largest bond moments will occur with the largest partial charges. The two solutions above represent how unevenly the electrons are shared in the bond. The bond moments will be maximized when the electronegativity difference is greatest. The controls for A and C should be set to one extreme, and B should be set to the opposite extreme. Although the magnitude of the bond moment will not change based on whether B is the most electronegative or the least, the direction of the bond moment will.
Summary
VSEPR theory predicts the three-dimensional arrangement of atoms in a molecule. It states that valence electrons will assume an electron-pair geometry that minimizes repulsions between areas of high electron density (bonds and/or lone pairs). Molecular structure, which refers only to the placement of atoms in a molecule and not the electrons, is equivalent to electron-pair geometry only when there are no lone electron pairs around the central atom. A dipole moment measures a separation of charge. For one bond, the bond dipole moment is determined by the difference in electronegativity between the two atoms. For a molecule, the overall dipole moment is determined by both the individual bond moments and how these dipoles are arranged in the molecular structure. Polar molecules (those with an appreciable dipole moment) interact with electric fields, whereas nonpolar molecules do not.
Glossary
- axial position
- location in a trigonal bipyramidal geometry in which there is another atom at a 180° angle and the equatorial positions are at a 90° angle
- bond angle
- angle between any two covalent bonds that share a common atom
- bond distance
- (also, bond length) distance between the nuclei of two bonded atoms
- bond dipole moment
- separation of charge in a bond that depends on the difference in electronegativity and the bond distance represented by partial charges or a vector
- dipole moment
- property of a molecule that describes the separation of charge determined by the sum of the individual bond moments based on the molecular structure
- electron-pair geometry
- arrangement around a central atom of all regions of electron density (bonds, lone pairs, or unpaired electrons)
- equatorial position
- one of the three positions in a trigonal bipyramidal geometry with 120° angles between them; the axial positions are located at a 90° angle
- linear
- shape in which two outside groups are placed on opposite sides of a central atom
- molecular structure
- structure that includes only the placement of the atoms in the molecule
- octahedral
- shape in which six outside groups are placed around a central atom such that a three-dimensional shape is generated with four groups forming a square and the other two forming the apex of two pyramids, one above and one below the square plane
- polar molecule
- (also, dipole) molecule with an overall dipole moment
- tetrahedral
- shape in which four outside groups are placed around a central atom such that a three-dimensional shape is generated with four corners and 109.5° angles between each pair and the central atom
- trigonal bipyramidal
- shape in which five outside groups are placed around a central atom such that three form a flat triangle with 120° angles between each pair and the central atom, and the other two form the apex of two pyramids, one above and one below the triangular plane
- trigonal planar
- shape in which three outside groups are placed in a flat triangle around a central atom with 120° angles between each pair and the central atom
- valence shell electron-pair repulsion theory (VSEPR)
- theory used to predict the bond angles in a molecule based on positioning regions of high electron density as far apart as possible to minimize electrostatic repulsion
- vector
- quantity having magnitude and direction