Question

1. which compound(s) would have dipole-dipole forces of attraction between molecules? select all that apply. cl₂ pcl₃ ccl₄ scl₂

Explanation:

Step1: Analyze \( \text{Cl}_2 \)

\( \text{Cl}_2 \) is a diatomic molecule with identical atoms, so it is non - polar. Non - polar molecules only have London dispersion forces, not dipole - dipole forces.

Step2: Analyze \( \text{PCl}_3 \)

The central atom \( \text{P} \) in \( \text{PCl}_3 \) has a lone pair of electrons. The molecular geometry of \( \text{PCl}_3 \) is trigonal pyramidal. Due to the difference in electronegativity between \( \text{P} \) and \( \text{Cl} \), and the asymmetrical shape, the molecule has a net dipole moment. So, \( \text{PCl}_3 \) is a polar molecule and will have dipole - dipole forces.

Step3: Analyze \( \text{CCl}_4 \)

The molecular geometry of \( \text{CCl}_4 \) is tetrahedral. The four \( \text{C - Cl} \) bonds are arranged symmetrically around the central \( \text{C} \) atom. The dipole moments of the individual \( \text{C - Cl} \) bonds cancel each other out, making \( \text{CCl}_4 \) a non - polar molecule. It only has London dispersion forces.

Step4: Analyze \( \text{SCl}_2 \)

The central atom \( \text{S} \) in \( \text{SCl}_2 \) has two lone pairs of electrons. The molecular geometry is bent (angular). The difference in electronegativity between \( \text{S} \) and \( \text{Cl} \) and the asymmetrical shape result in a net dipole moment. So, \( \text{SCl}_2 \) is a polar molecule and will have dipole - dipole forces.

Answer:

\( \text{PCl}_3 \), \( \text{SCl}_2 \) (The checkboxes corresponding to \( \text{PCl}_3 \) and \( \text{SCl}_2 \) should be selected)