A molecule’s polarity—the distribution of its electrical charge—is the primary factor that determines how molecules interact, influencing properties like solubility and boiling point. Chemical bonds within a molecule can create areas of slight positive and slight negative charge, setting up an electrical gradient. Dibromomethane, represented by the chemical formula \(\text{CH}_2\text{Br}_2\), is a simple organic compound whose physical properties are dictated by the way its atoms share electrons. Understanding the polarity of this molecule requires a step-by-step look at its internal structure, from the nature of its individual bonds to its overall three-dimensional shape.
Dibromomethane’s Polarity Status
Dibromomethane (\(\text{CH}_2\text{Br}_2\)) is classified as a polar molecule. The molecule possesses a permanent separation of charge, which creates distinct positive and negative ends. This uneven distribution of electron density dictates its physical and chemical behavior. \(\text{CH}_2\text{Br}_2\) has a measurable, non-zero dipole moment, confirming its polar nature.
The Role of Bond Polarity and Electronegativity
The first step in determining molecular polarity is analyzing the individual bonds. Electronegativity measures an atom’s ability to attract shared electrons. When atoms with different electronegativity values bond, electrons are pulled closer to the more electronegative atom, creating a polar covalent bond and resulting in partial positive (\(\delta^{+}\)) and negative (\(\delta^{-}\)) charges.
Dibromomethane contains carbon-hydrogen (\(\text{C-H}\)) and carbon-bromine (\(\text{C-Br}\)) bonds. The electronegativity difference between carbon (2.55) and hydrogen (2.20) is very small (0.35), meaning the \(\text{C-H}\) bonds are essentially nonpolar.
In contrast, the difference between carbon and bromine (2.96) is larger (0.41), making the \(\text{C-Br}\) bonds significantly polar. Bromine is more electronegative, pulling electron density toward the two bromine atoms and giving them a partial negative charge. This leaves the central carbon atom with a partial positive character.
How Molecular Geometry Creates a Net Dipole Moment
The presence of polar bonds does not guarantee the molecule is polar; the determination depends entirely on the three-dimensional shape. If the polar bonds are arranged symmetrically, their electrical pulls can perfectly cancel each other out, resulting in a nonpolar molecule.
The structure of dibromomethane is dictated by the Valence Shell Electron Pair Repulsion (VSEPR) theory. The central carbon atom in \(\text{CH}_2\text{Br}_2\) has four single bonds, forming four electron domains. This arrangement results in a tetrahedral molecular geometry, where the four peripheral atoms are positioned at the corners of a pyramid-like shape with the carbon atom at the center.
Crucially, the two hydrogen atoms and the two bromine atoms are not identical, which destroys the perfect electrical symmetry of the tetrahedral shape. In a symmetrical molecule like carbon tetrachloride (\(\text{CCl}_4\)), the four identical \(\text{C-Cl}\) dipoles pull equally in opposite directions and cancel out, making the molecule nonpolar.
However, in \(\text{CH}_2\text{Br}_2\), the two strong pulls toward the highly electronegative bromine atoms are not directly opposed by the two very weak pulls of the hydrogen atoms. The two strong bond dipoles from the \(\text{C-Br}\) bonds combine with the two weak dipoles from the \(\text{C-H}\) bonds in an asymmetrical vector addition. Because the bromine forces are much stronger, the resulting net dipole moment does not equal zero. The overall electron density is permanently shifted toward the bromine side of the molecule, creating a distinct negative pole. This permanent charge separation is the reason why dibromomethane is a polar molecule.