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Molecular Geometry and Dipole Moments

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VSEPR theory helps to determine electron-pair geometries and molecular geometries.

A series of steps is used to predict the geometry and bond angles of molecules, such as phosphorus trichloride.

The first step is to draw the Lewis structure of the molecule.

Next, count the total number of electron groups on the central atom. Around phosphorus, there are four electron groups: three bonding pairs and one lone pair.

Now, determine the electron-pair geometry. The electron-pair geometry is tetrahedral. However, because of the lone pair, the molecular geometry is trigonal pyramidal. The lone pair reduces the bond angle to less than 109.5°.

Bonding electron pairs are not always shared equally between the two bonding atoms.

In a covalent bond like that of hydrofluoric acid, the electrons are pulled toward the more electronegative atom, indicated by a partial charge. Such bonds are called polar bonds.

The charge separation creates a vector called the bond dipole moment, which is indicated by the Greek letter µ. Its value is the product of the magnitude of the partial charges and the distance between them.

Dipole moments are commonly expressed in debyes. One debye is equal to 3.336 × 10⁻³⁰ coulomb-meters.

The vector points from the less to the more electronegative atom and indicates the bond dipole moment. Its length is proportional to the magnitude of the electronegativity difference between the two atoms. Most diatomic molecules containing atoms of different elements have dipole moments and therefore are polar molecules.

In polyatomic compounds, the net dipole moment is determined by the individual bond dipole moments and geometry of the compound.

Consider a water molecule with two polar bonds. It has a bent shape and is a polar molecule.

In contrast, a carbon dioxide molecule is linear. The two carbon–oxygen bonds are polar but are oriented in opposite directions, canceling out each other's dipole moment and making the overall molecule nonpolar.

Overview

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The VSEPR theory can be used to determine the electron pair geometries and molecular structures as follows:

Write the Lewis structure of the molecule or polyatomic ion.
Count the number of electron groups (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 electron groups.
Use the number of lone pairs to determine the molecular structure. If more than one arrangement of lone pairs and chemical bonds is possible, choose the one that will minimize repulsions.

Dipole Moment of a Molecule

When atoms with different electronegativities form a bond, the electrons are pulled toward the more electronegative atom, leaving one atom with a partial positive charge (δ+) and the other atom with a partial negative charge (δ–). Such bonds are called polar covalent bonds, and the separation of charge gives rise to a bond dipole moment. The magnitude of a bond dipole moment is represented by the Greek letter µ and is given by:

μ = Qr

where Q is the magnitude of the partial charges (determined by the electronegativity difference), and r is the distance between them. Dipole moments are commonly expressed in debyes, where one debye is equal to 3.336 × 10⁻³⁰ C·m.

The bond dipole moment is a vector represented by an arrow pointing along the bond from the less electronegative toward the more electronegative atom, with a small plus sign on the less electronegative end.

A whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. Such molecules are said to be polar. The dipole moment measures the extent of net charge separation in the molecule as a whole. In diatomic molecules, the bond dipole moment determines the molecular polarity.

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 the vector sum of their bond moments equals zero, then the molecule is nonpolar (e.g., CO₂). The water molecule has a bent molecular structure, and the two bond moments do not cancel. Therefore, water is a polar molecule with a net dipole moment.

This text has been adapted fromOpenstax, Chemistry 2e, Section 7.6 Molecular Structure and Polarity.

Transcript

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$$\rightleftharpoonup{xx}$$ $$\longleftharp{xx}$$, $$\longrightharp{xx}$$,

VSEPR theory helps to determine electron-pair geometries and molecular geometries.

A series of steps is used to predict the geometry and bond angles of molecules, such as phosphorus trichloride.

The first step is to draw the Lewis structure of the molecule.

Next, count the total number of electron groups on the central atom. Around phosphorus, there are four electron groups: three bonding pairs and one lone pair.

Now, determine the electron-pair geometry. The electron-pair geometry is tetrahedral. However, because of the lone pair, the molecular geometry is trigonal pyramidal. The lone pair reduces the bond angle to less than 109.5°.

Bonding electron pairs are not always shared equally between the two bonding atoms.

In a covalent bond like that of hydrofluoric acid, the electrons are pulled toward the more electronegative atom, indicated by a partial charge. Such bonds are called polar bonds.

The charge separation creates a vector called the bond dipole moment, which is indicated by the Greek letter µ. Its value is the product of the magnitude of the partial charges and the distance between them.

Dipole moments are commonly expressed in debyes. One debye is equal to 3.336 × 10⁻³⁰ coulomb-meters.

The vector points from the less to the more electronegative atom and indicates the bond dipole moment. Its length is proportional to the magnitude of the electronegativity difference between the two atoms. Most diatomic molecules containing atoms of different elements have dipole moments and therefore are polar molecules.

In polyatomic compounds, the net dipole moment is determined by the individual bond dipole moments and geometry of the compound.

Consider a water molecule with two polar bonds. It has a bent shape and is a polar molecule.

In contrast, a carbon dioxide molecule is linear. The two carbon–oxygen bonds are polar but are oriented in opposite directions, canceling out each other's dipole moment and making the overall molecule nonpolar.

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