Question

(a) How does a polar molecule differ from a nonpolar one? (b) Atoms X and Y have different electronegativities. Will the diatomic molecule \(X-Y\) necessarily be polar? Explain. (c) What factors affect the size of the dipole moment of a diatomic molecule?

Short answer

A polar molecule has a net dipole moment due to the presence of polar bonds with a difference in electronegativity, while a nonpolar molecule does not have a net dipole moment. The diatomic molecule \(X-Y\) with different electronegativities will necessarily be polar, as it will have an uneven electron distribution and a net dipole moment. Factors affecting the dipole moment of a diatomic molecule include the charge difference (electronegativity) and bond length; dipole moment can be calculated using the formula: \[\text{Dipole moment} = (\text{Charge difference}) \times (\text{Bond length})\]

Step-by-step solution

Definition of Polar and Nonpolar Molecules

A polar molecule is a molecule that has a net dipole moment due to the presence of polar bonds in the molecule. These molecules have a difference in electronegativity between the atoms that results in an uneven distribution of electrons, creating a positive and negative end (or pole) within the molecule. Nonpolar molecules, on the other hand, have an even distribution of electrons within the molecule and do not have a net dipole moment.

Determine if the \(X-Y\) Molecule is Polar

If atoms X and Y have different electronegativities, it means that there will be an uneven distribution of electrons between the two atoms in the \(X-Y\) diatomic molecule. This will create a polar bond, and since the molecule is diatomic, there will be a net dipole moment. Therefore, the \(X-Y\) diatomic molecule will necessarily be polar.

Factors Affecting the Dipole Moment of a Diatomic Molecule

The dipole moment of a diatomic molecule depends on two factors: the charge difference (or the difference in electronegativity) between the two atoms and the bond length (or distance) between the two atoms. The dipole moment can be calculated using the formula: \[ \text{Dipole moment} = (\text{Charge difference}) \times (\text{Bond length}) \] - A larger difference in electronegativity between the two atoms in a diatomic molecule will result in a larger dipole moment, as the electrons will be more unevenly distributed and create a stronger charge difference. - A larger bond length between the two atoms in a diatomic molecule will also result in a larger dipole moment. If the bond length increases, it means that the charge difference is spread over a larger distance, which increases the dipole moment.

Key concepts

Electronegativity Difference

Electronegativity is a measure of how strongly an atom can attract electrons in a chemical bond. When two atoms form a bond, the difference in their electronegativities determines how the electrons are distributed.

• If the electronegativity difference is large, one atom will attract electrons more strongly, leading to an uneven distribution. This results in a polar bond with distinct positive and negative poles.
• If the electronegativity difference is small or nonexistent, electrons are shared more equally, resulting in a nonpolar bond.

Understanding electronegativity is crucial when determining whether a molecule is polar or nonpolar. In diatomic molecules like \(X-Y\), differing electronegativities between X and Y will create a polar molecule because of the resulting uneven electron distribution.

Dipole Moment

The dipole moment is a vector quantity that measures the separation of charge within a molecule. It indicates how polar a molecule is. In simple terms, the dipole moment can be thought of as a vector pointing from the positive to the negative charge center. Several factors influence the dipole moment:

• Electronegativity Difference: As the difference in electronegativity between two bonded atoms increases, so does the dipole moment. This is because the charge separation becomes more pronounced.
• Bond Length: The greater the distance between two atoms, the larger the dipole moment. Since the charge is spread over a longer distance, the dipole moment increases. The dipole moment can be calculated using the formula: \[ \text{Dipole moment} = (\text{Charge difference}) \times (\text{Bond length}) \]

A molecule with a significant dipole moment is said to be polar, with examples of such molecules including water \((H_2O)\), which has a strong dipole due to its angular shape and electronegativity differences.

Diatomic Molecules

Diatomic molecules consist of two atoms, which may either be the same or different elements. These molecules are a great way to explore fundamental concepts like bond polarity and dipole moments because they are simple and small.

• Homonuclear Diatomic Molecules: When both atoms are the same (e.g., \(N_2\), \(O_2\)), they share electrons equally, resulting in nonpolar molecules.
• Heteronuclear Diatomic Molecules: When the atoms are different (e.g., \(HCl\), \(CO\)), their differing electronegativities can result in polar bonds and thus polar molecules with dipole moments.

The simplicity of diatomic molecules makes them ideal for studying the basic principles of molecular polarity, as the absence of complex geometry helps focus solely on electronegativity and its effects on the bond and molecule as a whole.