Chloroform (\(\text{CHCl}_3\)), or trichloromethane, is a common organic compound often used as a solvent in laboratory and industrial settings. Determining the nature of its chemical bonds is a fundamental step in understanding its properties and behavior. The classification of \(\text{CHCl}_3\) is overwhelmingly covalent, meaning its atoms share electrons rather than transferring them. This conclusion is based on a close examination of the differences in electron-attracting power between the atoms involved.
Defining the Ionic and Covalent Spectrum
Chemical bonds exist along a continuous spectrum, ranging from purely ionic to purely covalent, based on how electrons are distributed. An ionic bond represents one extreme, where electrons are transferred from one atom to another, typically occurring between a metal and a non-metal. This transfer results in the formation of full positive and negative ions that are held together by strong electrostatic attraction.
The other end of the spectrum is the nonpolar covalent bond, where electrons are shared equally between two identical non-metal atoms. Between these two extremes lies the polar covalent bond, where electrons are shared unequally because one atom attracts the electron pair more strongly. This unequal sharing creates partial charges, denoted as delta positive (\(\delta+\)) and delta negative (\(\delta-\)), on the respective atoms.
The deciding factor for where a bond falls on this spectrum is the difference in electronegativity (\(\Delta\text{EN}\)) between the two bonded atoms. A difference greater than 1.7 indicates a bond with significant ionic character, suggesting electron transfer. Conversely, bonds between two non-metals, like those found in \(\text{CHCl}_3\), have a \(\Delta\text{EN}\) far below this threshold, classifying them as covalent bonds.
Applying Electronegativity to \(\text{CHCl}_3\) Bonds
The chloroform molecule contains a central carbon atom bonded to one hydrogen atom and three chlorine atoms, resulting in two distinct connections: Carbon-Hydrogen (C-H) and Carbon-Chlorine (C-Cl). The established electronegativity values are: Carbon (2.5), Hydrogen (2.1), and Chlorine (3.0).
The C-H bond has an electronegativity difference of 0.4 (2.5 minus 2.1). This small difference places the C-H bond near the nonpolar covalent range, meaning the electron sharing is nearly equal and the resulting bond dipole is weak. The C-Cl bond, in contrast, has a difference of 0.5 (3.0 minus 2.5).
Chlorine is the more electronegative atom, pulling shared electrons more strongly. This gives the chlorine atom a partial negative charge (\(\delta-\)) and the carbon atom a partial positive charge (\(\delta+\)). The largest \(\Delta\text{EN}\) in \(\text{CHCl}_3\) is 0.5, which is significantly lower than the 1.7 threshold required for ionic classification. Since all constituent atoms are non-metals and the \(\Delta\text{EN}\) values are small, the connections are confirmed to be polar covalent bonds.
Molecular Structure and Polarity
A molecule’s behavior depends not just on the individual bonds, but on its three-dimensional shape. Chloroform exhibits a tetrahedral molecular geometry. The central carbon atom is positioned at the center, with the four surrounding atoms (\(\text{H}\) and three \(\text{Cl}\)) directed toward the corners of the tetrahedron.
This shape is geometrically asymmetric because the three polar C-Cl bonds are not balanced by a fourth identical bond. The three strong C-Cl bond dipoles pull electron density toward the chlorine atoms and do not cancel each other out. This effect is compounded by the much weaker, oppositely directed dipole from the C-H bond.
The combination of these individual bond dipoles is calculated as a vector sum. Due to the molecular asymmetry, the sum results in a net molecular dipole moment of 1.04 Debye. This non-zero dipole moment confirms that \(\text{CHCl}_3\) is a polar molecule, cementing its classification as a polar covalent compound.