Question

Consider the compounds \(\mathrm{CHCl}_{3}\) (chloroform, an important solvent that was once used as an anesthetic) and \(\mathrm{CHBr}_{3}\) (bromoform, which has been used as a sedative). Compare the strengths of their dipole-dipole attractions and the strengths of their London forces. Their boiling points are \(61^{\circ} \mathrm{C}\) and \(149^{\circ} \mathrm{C}\), respectively. For these compounds, which kinds of attractive forces (i.e., dipoledipole or London) are more important in determining their boiling points? Justify your answer.

Short answer

The London forces are more important in determining the boiling points of CHCl3 and CHBr3, as indicated by the higher boiling point of CHBr3 (149 degrees Celsius) compared to CHCl3 (61 degrees Celsius), despite both molecules being capable of dipole-dipole interactions.

Step-by-step solution

Understand Dipole-Dipole Attractions

Dipole-dipole attractions occur between molecules that have permanent dipole moments. These are stronger interactions compared to London dispersion forces (a type of van der Waals force), which are present in all molecules due to instantaneous dipole moments caused by electron movement. Molecules with larger dipole moments should have stronger dipole-dipole interactions.

Compare the Molecular Structure of CHCl3 and CHBr3

Both CHCl3 (chloroform) and CHBr3 (bromoform) are similar in molecular shape, and both have three halogen atoms bonded to a central carbon atom. However, bromine atoms are larger and have more electrons than chlorine atoms, which could lead to stronger London forces in CHBr3 relative to CHCl3.

Analyze Boiling Points and Infer Dominant Forces

Given the boiling points of CHCl3 (61 degrees Celsius) and CHBr3 (149 degrees Celsius), it can be inferred that the stronger the intermolecular forces within a compound, the higher the boiling point. As CHBr3 has a significantly higher boiling point, it indicates that the London forces, which are stronger due to bromine being larger and having more electrons than chlorine, play a more critical role in determining the boiling point for these compounds.

Justify the Impact of London Forces on Boiling Points

The considerable difference in boiling points suggests that the London forces have a significant impact on boiling points of these molecules. The greater electron cloud and molar mass of CHBr3, result in stronger London forces compared to CHCl3, and as such, these forces are more important than dipole-dipole interactions in determining the boiling points of these particular compounds.

Key concepts

Dipole-Dipole Attractions

Dipole-dipole attractions are forces that occur between molecules with permanent dipole moments. When a molecule has regions of partial positive and negative charge—often resulting from electronegative atoms, like chlorine or bromine, drawing electrons closer—these charges interact with the opposite charges on neighboring molecules. Imagine tiny magnets that attract or repel each other; that's akin to how these dipole interactions work. Greater differences in electronegativity between atoms within a molecule can lead to stronger dipole-dipole attractions. Since both chloroform ((CHCl3)) and bromoform ((CHBr3)) have sizable permanent dipoles due to their molecular structures, they experience significant dipole-dipole attractions. However, the strength of these attractions varies slightly between the two because of differences in electronegativity and molecular shape.

London Dispersion Forces

London dispersion forces, or induced dipole-induced dipole attractions, are another type of van der Waals force. These forces are universally present in all molecules, regardless of whether they have a permanent dipole. London forces arise from the momentary uneven distribution of electrons around an atom or molecule, creating an instantaneous dipole that can induce a complementary dipole in neighboring atoms or molecules. Here's a fun analogy: think of London forces as the social networking of molecules, constantly establishing and breaking temporary connections. In molecules like chloroform and bromoform, the larger and heavier the halogen atoms, the greater the London dispersion forces. This is because larger atoms with more electrons can have a more pronounced temporary dipole. In our two compounds, bromoform, with its heavier bromine atoms in comparison to the chlorine atoms in chloroform, exhibits stronger London dispersion forces.

Boiling Point Analysis

The boiling point of a substance is a direct measure of the strength of its intermolecular forces. A higher boiling point indicates that more energy (heat) is required to overcome these forces and transition from liquid to gas. This is why boiling point analysis is a powerful tool when comparing the relative strengths of intermolecular forces within different molecules. When we observe that bromoform boils at a much higher temperature (149 degrees Celsius) than chloroform (61 degrees Celsius), it tells us that bromoform's collective intermolecular forces must be profoundly stronger. This elevation in boiling point leads to an inference—supported further by molecular analysis—that it is the London dispersion forces playing a pivotal role here, especially given the larger molecular weight and electron cloud of bromoform.

Molecular Structure Comparison

Understanding the molecular structure of substances provides valuable insight into their chemical behavior, particularly regarding intermolecular forces. Chloroform and bromoform have substantial similarities: both have a tetrahedral molecular geometry with a central carbon atom bonded to three halogen atoms and one hydrogen atom. However, the critical difference is the type of halogen atom. The larger bromine atoms in bromoform compared to the chlorine atoms in chloroform result in a bigger electron cloud and higher molar mass. These differences greatly amplify the London dispersion forces present in bromoform, as evidenced by its significantly higher boiling point compared to chloroform. This type of comparison allows chemists to predict and explain variations in physical properties, including boiling points, as a direct consequence of changes in molecular structure.