Question

In which of the following are there no dipoledipole forces between molecules? Justify your answer. (a) \(\mathrm{AsH}_{3}\) (b) \(\mathrm{CO}_{2}\) (c) \(\mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{SeCl}_{2}\)

Short answer

The molecule without dipole-dipole forces between its molecules is CO₂ (option b). This is because it has a linear, symmetrical molecular geometry, making it non-polar despite having polar bonds. All other molecules (AsH₃, H₂O, and SeCl₂) experience dipole-dipole forces due to their polar nature.

Step-by-step solution

Determine the Electronegativity of Each Atom

To determine if a molecule is polar, we need to evaluate the differences in electronegativity between the atoms in each molecule. Electronegativity values can typically be found in a chemistry textbook or online.

Evaluate the Electronegativity Differences

Compare the electronegativity values of each atom within a molecule to determine if there is a significant difference. A significant difference indicates polar bonds. (a) AsH₃: As has an electronegativity of 2.18, and H has an electronegativity of 2.20. The difference is 0.02, which is not significant. (b) CO₂: C has an electronegativity of 2.55, and O has an electronegativity of 3.44. The difference is 0.89, which is significant. (c) H₂O: O has an electronegativity of 3.44, and H has an electronegativity of 2.20. The difference is 1.24, which is significant. (d) SeCl₂: Se has an electronegativity of 2.55, and Cl has an electronegativity of 3.16. The difference is 0.61, which is significant.

Examine Molecular Geometry

We must now consider the molecular geometry of each molecule to determine its polarity. A molecule can have polar bonds, but if its geometry is symmetrical, it will be non-polar overall. (a) AsH₃: Trigonal pyramidal, a polar shape. (b) CO₂: Linear, a symmetrical, non-polar shape despite the presence of polar bonds. (c) H₂O: Bent, a polar shape. (d) SeCl₂: Bent, a polar shape.

Identify the Molecule Without Dipole-Dipole Forces

Based on our evaluation of the electronegativity differences and molecular geometry, we can conclude that CO₂ (option b) is the molecule without dipole-dipole forces between molecules since it's a non-polar molecule. All other molecules (AsH₃, H₂O, and SeCl₂) have dipole-dipole forces between their molecules due to their polar nature.

Key concepts

Molecular Polarity

Understanding molecular polarity is essential for grasping why certain molecules exhibit dipole-dipole forces. A molecule is termed polar if it has a resulting net dipole moment, which is the product of the charge difference and the distance between the charges. In simpler terms, a polar molecule has a distribution of electric charge that leads to the existence of distinct positive and negative ends, or 'poles'.

When considering the exercise, we look at the molecular geometry and the difference in electronegativity to assess polarity. For instance, in the case of CO₂, even though the bonds between carbon and oxygen are polar due to a significant difference in electronegativity, the linear geometry results in an overall non-polar molecule since the dipoles cancel each other out. In contrast, H₂O has a bent shape causing the dipoles not to cancel, thus making it a polar molecule with dipole-dipole interactions.

Exercise Improvement Advice: Understanding the concept of molecular polarity can be enhanced by visualizing molecules with a 3D modeling tool to see how bond polarity and molecular shape contribute to the overall molecular polarity.

Electronegativity

Electronegativity is a chemical property that describes the ability of an atom to attract electrons towards itself. It greatly influences the polarity of a molecule. Atoms with higher electronegativity will attract shared electrons more strongly, potentially creating a dipole.

In the context of the exercise, by comparing the electronegativity values of arsenic (As) and hydrogen (H), we find a minimal difference, suggesting that AsH₃ should have non-polar bonds. However, due to the shape of the molecule, overall, it experiences dipole-dipole forces. Recognizing the trends in electronegativity across the periodic table can be a pivotal step for students to predict the type of bonding in a molecule. Generally, electronegativity increases from left to right and decreases from top to bottom in the periodic table.

Molecular Geometry

Molecular geometry, the three-dimensional arrangement of atoms in a molecule, is crucial for determining a molecule’s overall polarity. Different molecular shapes such as linear, bent, tetrahedral, or trigonal pyramidal can arise from the specific ways atoms bond and the number of electron pairs surrounding the central atom.

In the exercise, CO₂'s linear shape, despite having polar bonds, results in a non-polar molecule because the bond dipoles are equal in magnitude and opposite in direction, cancelling each other out. H₂O and SeCl₂, however, have bent shapes which do not allow for the cancellation of dipole moments, leading to polar molecules with associated dipole-dipole forces.

Exercise Improvement Advice: It could be beneficial for students to practice drawing Lewis structures and using VSEPR theory to predict the geometry of a molecule. This understanding helps in determining whether the molecule will have a net dipole moment or not.