Trichlorofluoromethane, commonly known as CFC-11, has the chemical formula \(\text{CFCl}_3\). Determining whether this compound is polar or nonpolar requires examining its molecular structure and the nature of its chemical bonds. Based on these principles, \(\text{CFCl}_3\) is classified as a polar molecule, meaning it possesses an uneven distribution of electrical charge. This polarity is a direct consequence of its specific three-dimensional shape and the different elements bonded to its central atom.
The Three-Dimensional Shape of \(\text{CFCl}_3\)
The \(\text{CFCl}_3\) molecule is structured around a central carbon atom. This carbon atom forms four single covalent bonds: one with fluorine and three with chlorine. The four regions of electron density (electron domains) repel each other, forcing the atoms into a tetrahedral geometry.
In a perfectly symmetrical molecule like carbon tetrachloride (\(\text{CCl}_4\)), the tetrahedral shape ensures equal bond angles (approximately 109.5 degrees). However, because the atoms surrounding the central carbon in \(\text{CFCl}_3\) are not identical, the ideal geometry is slightly distorted. The overall shape remains tetrahedral.
Why the Bonds Are Polar
Molecular polarity begins with the polarity of individual bonds, determined by electronegativity. Electronegativity is the measure of an atom’s ability to attract shared electrons toward itself within a chemical bond. In \(\text{CFCl}_3\), the central carbon is bonded to fluorine and chlorine, both significantly more electronegative than carbon.
This difference means electrons in the \(\text{C-F}\) and \(\text{C-Cl}\) covalent bonds are not shared equally. Electron density is pulled toward the halogen atoms, creating a partial negative charge (\(\delta^-\)) on the fluorine and chlorine, and a partial positive charge (\(\delta^+\)) on the carbon. This charge separation establishes the carbon-fluorine and carbon-chlorine bonds as distinct polar bonds. The degree of polarity differs between these two bond types. Fluorine is the most electronegative element, giving the \(\text{C-F}\) bond a larger electronegativity difference than the \(\text{C-Cl}\) bonds.
How Molecular Polarity is Determined
Molecular polarity is determined by the combination of bond polarity and the molecule’s overall geometry, visualized through the net dipole moment. Each polar bond in \(\text{CFCl}_3\) has a bond dipole moment, which is a vector quantity possessing both magnitude and direction. These four individual bond dipole moments must be added together vectorially.
In a molecule with identical peripheral atoms, such as the symmetrical \(\text{CCl}_4\), all four bond vectors cancel out, resulting in a net dipole moment of zero and a nonpolar molecule. However, \(\text{CFCl}_3\) is inherently asymmetrical, having one fluorine atom and three chlorine atoms. The three \(\text{C-Cl}\) bond dipoles and the single \(\text{C-F}\) bond dipole do not have equal magnitudes or arrangement for complete cancellation.
This imbalance leads to an unequal sharing of electrons across the entire molecule, creating a region of higher electron density on one side. Since the vector sum of the individual bond dipoles is non-zero, the molecule possesses a measurable net dipole moment, approximately \(0.46 \pm 0.02\) Debye, confirming that \(\text{CFCl}_3\) is a polar molecule.
The Significance of \(\text{CFCl}_3\)
Trichlorofluoromethane (CFC-11) is most significant for its historical and environmental context as one of the original chlorofluorocarbons (CFCs). These compounds were developed in the 1930s and found widespread commercial use due to their stability, non-toxicity, and non-flammability. \(\text{CFCl}_3\) was utilized as a refrigerant, a blowing agent in foam production, and an aerosol propellant.
Its chemical stability allowed it to persist long enough to reach the stratosphere. There, high-energy ultraviolet radiation breaks the molecule down, releasing highly reactive chlorine atoms. These chlorine atoms participate in a catalytic cycle that efficiently destroys stratospheric ozone molecules.
This ozone-depleting potential led to the international phase-out of \(\text{CFCl}_3\) production under the landmark Montreal Protocol. Despite the global ban, \(\text{CFCl}_3\) remains a concern due to its long atmospheric lifetime and its status as a potent greenhouse gas. It contributes not only to ozone layer recovery delays but also to the atmospheric heat-trapping effect.