Question

For each of the following bonds, indicate the positive end with \(\delta^{+}\) and the negative end with \(\delta\). Draw an arrow to show the dipole for each. a. \(\mathrm{N}-\mathrm{F}\) b. \(\mathrm{Si}-\mathrm{Br}\) c. \(\mathrm{C}-\mathrm{O}\) d. \(\mathrm{P}-\mathrm{Br}\) e. \(\mathrm{N}-\mathrm{P}\)

Short answer

a. \mathrm{N}-\mathrm{F}: \delta^{+}\mathrm{N} \delta^{-}\mathrm{F}\ (points \rightarrow \mathrm{F})\ b. \mathrm{Si}-\mathrm{Br}: \delta^{+}\mathrm{Si} \delta^{-}\mathrm{Br}\ (points \rightarrow \mathrm{Br})\ c. \mathrm{C}-\mathrm{O}: \delta^{+}\mathrm{C} \delta^{-}\mathrm{O}\ (points \rightarrow \mathrm{O})\ d. \mathrm{P}-\mathrm{Br}: \delta^{+}\mathrm{P} \delta^{-}\mathrm{Br}\ (points \rightarrow \mathrm{Br})\ e. \mathrm{N}-\mathrm{P}: \delta^{+}\mathrm{P} \delta^{-}\mathrm{N}\ (points \rightarrow \mathrm{N})

Step-by-step solution

Understand Bond Dipoles

The bond dipole is the result of differences in electronegativity between the two atoms in a covalent bond. The more electronegative atom attracts the shared electrons more strongly, becoming the negative end \(\delta^{-}\), while the less electronegative atom becomes the positive end \(\delta^{+}\).

Determine Electronegativity Differences

Find the electronegativity values of the atoms in each bond. The atom with the higher electronegativity will be \(\delta^{-}\) and the atom with the lower electronegativity will be \(\delta^{+}\).

Identify the Dipole for \(\mathrm{N}-\mathrm{F}\)

\(\mathrm{N}\) has an electronegativity of 3.04 and \(\mathrm{F}\) has an electronegativity of 3.98. Since \(\mathrm{F}\) is more electronegative, it will be \(\delta^{-}\) and \(\mathrm{N}\) will be \(\delta^{+}\). Draw the arrow pointing towards \(\mathrm{F}\).

Identify the Dipole for \(\mathrm{Si}-\mathrm{Br}\)

\(\mathrm{Si}\) has an electronegativity of 1.90 and \(\mathrm{Br}\) has an electronegativity of 2.96. Since \(\mathrm{Br}\) is more electronegative, it will be \(\delta^{-}\) and \(\mathrm{Si}\) will be \(\delta^{+}\). Draw the arrow pointing towards \(\mathrm{Br}\).

Identify the Dipole for \(\mathrm{C}-\mathrm{O}\)

\(\mathrm{C}\) has an electronegativity of 2.55 and \(\mathrm{O}\) has an electronegativity of 3.44. Since \(\mathrm{O}\) is more electronegative, it will be \(\delta^{-}\) and \(\mathrm{C}\) will be \(\delta^{+}\). Draw the arrow pointing towards \(\mathrm{O}\).

Identify the Dipole for \(\mathrm{P}-\mathrm{Br}\)

\(\mathrm{P}\) has an electronegativity of 2.19 and \(\mathrm{Br}\) has an electronegativity of 2.96. Since \(\mathrm{Br}\) is more electronegative, it will be \(\delta^{-}\) and \(\mathrm{P}\) will be \(\delta^{+}\). Draw the arrow pointing towards \(\mathrm{Br}\).

Identify the Dipole for \(\mathrm{N}-\mathrm{P}\)

\(\mathrm{N}\) has an electronegativity of 3.04 and \(\mathrm{P}\) has an electronegativity of 2.19. Since \(\mathrm{N}\) is more electronegative, it will be \(\delta^{-}\) and \(\mathrm{P}\) will be \(\delta^{+}\). Draw the arrow pointing towards \(\mathrm{N}\).

Key concepts

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons. It's represented by a numerical value.
The higher the value, the more an atom attracts electrons.
For instance, fluorine has the highest electronegativity value of 3.98. This means it strongly pulls electrons towards itself.
By comparing electronegativity values, we can predict the direction of the dipole in a bond.
Generally, moving left to right across a period in the periodic table, electronegativity increases.
Moving down a group, electronegativity decreases. Understanding these trends helps us determine which atom in a bond will be \(\delta^{-}\) and which will be \(\delta^{+}\).

Covalent Bonds

Covalent bonds occur when two atoms share electrons. It's the sharing of these electrons that bond the atoms together.
The shared electrons allow each atom to attain a stable electron configuration, similar to noble gases.
Not all covalent bonds are equal. When atoms with different electronegativities form a bond, the shared electrons spend more time closer to one atom.
This creates a partial charge on both atoms, making one end \(\delta^{-}\) and the other \(\delta^{+}\). Thus, a bond dipole is established. The strength and direction of this dipole depend on the electronegativity difference between the two atoms involved.

Dipole Moment

The dipole moment is a measure of the separation of positive and negative charges in a molecule.
It's a vector quantity, which means it has both a direction and a magnitude.
The direction of the dipole moment is from the less electronegative atom to the more electronegative atom. It's often represented by an arrow.
The bigger the difference in electronegativity between the two atoms, the larger the dipole moment.
Dipole moments are crucial for predicting the behavior of molecules in electric fields and understanding molecular interactions.

Molecular Polarity

Molecular polarity depends on both the individual bond dipoles and the geometry of the molecule.
Even if a molecule has polar bonds, if the shape of the molecule is symmetrical, the bond dipoles can cancel out.
This makes the molecule nonpolar overall.
For example, carbon dioxide (CO2) has polar C=O bonds, but it's a linear molecule, making the entire molecule nonpolar.
However, if the molecule has an asymmetrical shape, the bond dipoles do not cancel out. This results in a polar molecule, like water (H2O).
Polarity affects properties such as solubility, boiling point, and interactions with other molecules.