Question

Explain why the first molecule of each pair is polar and the second is not. (a) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CCl}_{4}\) (c) \(\mathrm{BF}_{2} \mathrm{Cl}, \mathrm{BF}_{3}\) (b) \(\mathrm{PF}_{3}, \mathrm{BF}_{3}\) (d) \(\mathrm{SO}_{2}, \mathrm{O}_{3}\)

Short answer

\(\mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{BF}_{2} \mathrm{Cl}, \mathrm{PF}_{3}, \) and \(\mathrm{SO}_{2}\) are polar due to asymmetrical shape and difference in electronegativity among atoms which causes a net dipole moment. \(\mathrm{CCl}_{4}, \mathrm{BF}_{3}, \) and \(\mathrm{O}_{3}\) are non-polar because of their symmetrical structure which cancels out the individual bond dipoles.

Step-by-step solution

Identify and evaluate the molecular geometry and electronegativity

In \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\), Hydrogen and Chlorine exhibit a dipole moment due to the difference in electronegativities, hence it is polar. In \(\mathrm{CCl}_{4}\), while there's a difference in electronegativity between Carbon and Chlorine, the symmetrical structure of the molecule cancels the dipole moments making it non-polar.

Identify and evaluate the molecular geometry and electronegativity

In \(\mathrm{BF}_{2} \mathrm{Cl}\), the differences in electronegativity between Boron, Fluorine and Chlorine lead to individual bond dipoles that, due to the molecule's geometry, do not cancel each other; hence, the molecule is polar. In \(\mathrm{BF}_{3}\), the symmetry of the molecule causes the individual dipoles to cancel each other out, making it non-polar.

Identify and evaluate the molecular geometry and electronegativity

In \(\mathrm{PF}_{3}\), Phosphorous and Fluorine have different electronegativities, and the three Fluorine atoms create a non-symmetrical structure, causing the molecule to be polar. In \(\mathrm{BF}_{3}\) (as in step 2) the symmetry of the molecule leads to the cancellation of the bond dipoles, making it non-polar.

Identify and evaluate the molecular geometry and electronegativity

In \(\mathrm{SO}_{2}\), Sulfur and Oxygen have different electronegativities and the molecule's bent structure, give it an overall dipole moment, making it polar. In \(\mathrm{O}_{3}\), or Ozone, all the atoms are the same, hence there is no difference in electronegativity, making it non-polar.

Key concepts

Molecular Geometry

Grasping the concept of molecular geometry is vital for understanding why molecules have differing polarity. To visualize this, imagine tiny magnets arranged in specific shapes. If these magnets point in opposite directions with equal force, their pull cancels out, just as the dipoles in a symmetrical molecule cancel each other.

For instance, in the molecule \textbf{CCl}\(_4\), the Chlorine atoms are spaced evenly around the central Carbon atom in a tetrahedral shape. This geometric symmetry ensures that any dipole moments due to Chlorine's higher electronegativity point in such a way that they cancel out. On the other hand, \textbf{CH}\(_2\)\textbf{Cl}\(_2\) has a different arrangement of Hydrogen and Chlorine around Carbon, leading to a shape where the dipoles do not cancel fully, rendering the molecule polar.

Understanding the three-dimensional shape of molecules helps in predicting their polarity. VSEPR (Valence Shell Electron Pair Repulsion) theory provides a useful method for determining that shape. Molecules with a non-uniform distribution of charge due to asymmetrical shape, like \textbf{SO}\(_2\) with its bent structure, typically exhibit polarity.

Electronegativity

Electronegativity refers to an atom's ability to attract and hold onto electrons within a chemical bond. Think of it as a tug-of-war between atoms: the stronger the atom, the more electrons it can pull towards itself.

The concept of electronegativity explains why some molecules have polar bonds. Atoms with significant differences in electronegativity create an uneven electron distribution, with the more electronegative atom gaining a slight negative charge and the less electronegative atom having a small positive charge. This asymmetry in charge generates a dipole. For example, the bond between Boron and Chlorine in \textbf{BF}\(_2\)\textbf{Cl} is polar because Chlorine is more electronegative than Boron, pulling electrons towards itself creating a dipole moment.

In a molecule like \textbf{O}\(_3\), or Ozone, all atoms involved have the same electronegativity, so there's no difference in electron distribution among the atoms, resulting in a non-polar molecule. It's crucial to consider both the differences in electronegativity and the molecule's geometry to fully understand the molecular polarity.

Dipole Moments

Dipole moments are vector quantities that are a measure of the separation of positive and negative charges in a molecule. It's akin to measuring the 'flow' of charge within a molecule. A molecule with a dipole moment has a sort of 'molecular north and south pole', so to speak.

The polarity of a molecule is largely influenced by its dipole moment, which arises when there is an overall imbalance in the charge distribution. A classic example is the molecule \textbf{PF}\(_3\). The Phosphorous-Fluorine bonds have significant differences in electronegativity, and because Fluorine is highly electronegative, these bonds have strong dipole moments. Since \textbf{PF}\(_3\) is not symmetrically shaped, these dipoles don't cancel out, lending a polar character to the molecule.

Conversely, a molecule like \textbf{BF}\(_3\) has no overall dipole moment because the individual dipoles from the Boron-Fluorine bonds point in opposite directions and are equidistant from the central Boron atom, balancing out the charge distribution. This highlights the importance of both molecular geometry and electronegativity when determining the dipole moments and ultimately the polarity of a molecule.