Question

What is a dipole moment? Give four examples of molecules that possess dipole moments, and draw the direction of the dipole as shown in Section 12.3 .

Short answer

A dipole moment is a measure of the separation of positive and negative charges within a molecule caused by the presence of polar bonds, which are formed when electrons are shared unequally between atoms with different electronegativities. Examples of molecules that possess dipole moments include: 1. Water (H2O): O-H bonds are polar, and the dipole moment points towards the more negative oxygen atom. 2. Hydrogen chloride (HCl): H-Cl bond is polar, and the dipole moment points towards the more negative chlorine atom. 3. Ammonia (NH3): N-H bonds are polar, and the dipole moment points towards the more negative nitrogen atom. 4. Carbon dioxide (CO2): Although it has polar C=O bonds, the net dipole moment is zero due to its symmetrical structure.

Step-by-step solution

(Define Dipole Moment)

A dipole moment is a measure of the separation of positive and negative charges within a molecule or an atom. It is a vector quantity and is represented by an arrow with its tail at the more positive end and its head at the more negative end. The direction of the arrow indicates the flow of electrons. In molecules, a dipole moment occurs due to the presence of polar bonds, which are formed when electrons are shared unequally between atoms with different electronegativities.

(First Example: Water)

Water (H2O) is a molecule that possesses a dipole moment due to the polar O-H bonds. The lone pairs of electrons on the oxygen atom make it more electronegative, causing the dipole moment to point towards the oxygen atom. Here is a representation of the dipole in water: O / \ δ+ H-H δ+ ↑ The arrow points towards the more negative oxygen atom, and the δ+ symbol represents the partial positive charge on the hydrogen atoms.

(Second Example: Hydrogen Chloride)

Hydrogen chloride (HCl) is another example of a molecule with a dipole moment. In this case, the more electronegative chlorine atom attracts the shared electron pair leading to polarity within the molecule. The dipole moment points towards the chlorine atom as shown below: H-Cl δ+↗ δ- The arrow points towards the more negative chlorine atom, and the δ+ symbol represents the partial positive charge on the hydrogen atom.

(Third Example: Ammonia)

Ammonia (NH3) has a dipole moment due to the polar N-H bonds. Nitrogen is more electronegative than hydrogen, resulting in a dipole moment that points towards the nitrogen atom. The structure of ammonia with the dipole moment is represented as follows: N / | \ H H H ↘ δ+ ↙ The arrow points towards the more negative nitrogen atom, and the δ+ symbol represents the partial positive charge on the hydrogen atoms.

(Fourth Example: Carbon Dioxide)

Although carbon dioxide (CO2) has two polar C=O bonds, the symmetrical linear geometry of the molecule causes the two dipole moments to cancel each other, resulting in a net dipole moment of zero. Therefore, CO2 is a nonpolar molecule and does not possess a dipole moment. In conclusion, the dipole moment is a measure of the separation of positive and negative charges within molecules due to the presence of polar bonds. Examples of molecules that possess dipole moments include H2O, HCl, and NH3, while CO2, despite having polar bonds, has a net dipole moment of zero due to its symmetrical structure.

Key concepts

Polar Bonds

Understanding polar bonds is essential when exploring why certain molecules have dipole moments. Polar bonds are formed between two atoms of different electronegativities—the measure of an atom's ability to attract and hold onto electrons. When atoms with dissimilar electronegativities form a covalent bond, the shared electrons are pulled more towards the atom with the higher electronegativity, causing a partial charge distribution.

Take for instance the bond between hydrogen and oxygen in a water molecule. Oxygen's higher electronegativity means it will have a partial negative charge (designated as \( \delta- \) ), while hydrogen exhibits a partial positive charge ( \( \delta+ \) ). This unequal sharing of electrons is the fundamental cause of a polar bond. Whether in organic compounds or simple diatomic molecules, polar bonds lead to the characteristic uneven distribution of electrical charges within a molecule that can result in molecular polarity.

For students working on understanding polar bonds, visualizing electron distribution and recognizing electronegativity differences is key. By looking at the periodic table, you can predict bond polarity by noting the position of the elements involved—the farther apart they are, generally, the more polar their bond will be.

Molecular Polarity

Molecular polarity arises from the vector summation of all individual bond dipoles within a molecule. It is a holistic property that considers both the molecular geometry and bond polarity. A molecule with polar bonds may not always be polar; it depends on the spatial arrangement of these bonds.

For instance, water is a polar molecule because its bent shape means the dipole moments of the O-H bonds don't cancel out. Ammonia, with its trigonal pyramidal geometry, also exhibits polarity. Contrast this with carbon dioxide, a molecule that despite having polar bonds (C=O), is nonpolar due to its linear shape that leads to the cancellation of dipole moments.

Thus, the determination of molecular polarity is a crucial skill in chemistry. It involves evaluating both the molecular shape and the electronegativity of atoms to understand how it affects the overall dipole moment. Usually, molecules with symmetrical shapes, like CO2, are nonpolar even when containing polar bonds, while those with asymmetrical shapes, like H2O and NH3, display a net dipole moment. Intuitively, if the shape of a molecule prevents the vectorial 'tug-of-war' between dipoles from canceling out, the molecule will be polar.

Electronegativity

Electronegativity is a central concept when discussing dipole moments and molecular polarity. It quantitatively represents an atom's tendency to attract a shared pair of electrons towards itself in a chemical bond. The Pauling scale is commonly used to assign electronegativity values to the elements, with fluorine being the most electronegative element at 4.0.

The differences in electronegativity between two bonding atoms will dictate the polarity of the bond. A larger difference implies a more polar bond. For instance, chlorine is more electronegative than hydrogen, making the H-Cl bond polar. In compounds with multiple atoms, such as water, electronegativity helps explain the directional nature of dipole moments and their magnitudes.

Students should grasp the impact of electronegativity differences on bond character: small differences generally lead to nonpolar covalent bonds, moderate differences result in polar covalent bonds, and large differences (typically greater than 2.0 on the Pauling scale) often lead to ionic bonds. In a molecular context, understanding these electronegativity concepts is indispensable for predicting molecular behavior, such as dipole moments and the interactions molecules will have with one another in different environments.