Question

Determine the kinds of intermolecular forces that are present in each element or compound. a. \(\mathrm{Kr}\) b. \(\mathrm{NCl}_{3}\) c. \(\mathrm{SiH}_{4}\) d. HF

Short answer

a. Kr exhibits London dispersion forces. b. NCl3 has dipole-dipole interactions and London dispersion forces. c. SiH4 exhibits London dispersion forces. d. HF shows hydrogen bonding and dipole-dipole interactions.

Step-by-step solution

Identify Intermolecular Forces for Kr

Kr is a noble gas with a closed valence shell, which suggests that its primary intermolecular force is London dispersion forces (LDF), also known as induced dipole-induced dipole interactions. These are the weakest type of intermolecular forces and exist between all molecules, particularly in noble gases and nonpolar molecules.

Assess Intermolecular Forces in NCl3

Nitrogen trichloride ((NCl)_3) has a pyramid-shaped molecular geometry, where the nitrogen atom is at the apex, which leads to a non-uniform distribution of charge. This causes NCl3 to have a dipole moment, making the dominant intermolecular force dipole-dipole interactions. There is also the presence of London dispersion forces, but these are weaker compared to the dipole-dipole interactions.

Examine Intermolecular Forces in SiH4

Silane ((SiH)_4) has a tetrahedral shape with hydrogen atoms symmetrically distributed around the silicon atom. This symmetry leads to a non-polar molecule with no permanent dipole moment. Therefore, the primary intermolecular force is the London dispersion forces.

Determine Intermolecular Forces for HF

Hydrogen fluoride (HF) is a polar molecule due to the high electronegativity difference between hydrogen and fluorine, which results in a significant dipole moment. Besides strong dipole-dipole interactions, HF can also exhibit hydrogen bonding, which is a special, stronger type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like fluorine.

Key concepts

Understanding London Dispersion Forces

Imagine a group of noble gases or nonpolar molecules gently tumbling and influencing each other in a subtle dance. This is the essence of London dispersion forces (LDF), the weakest yet most pervasive form of intermolecular attractions. Despite their fleeting nature, these forces are critical in maintaining the structural integrity of various substances.

LDFs arise from momentary shifts in electron density, leading to fleeting dipole moments which, in turn, induce a corresponding dipole in neighboring atoms or molecules. This temporary alignment of charge creates an attraction that, while feeble when considered between two particles, can accumulate to hold masses of atoms or molecules together. Think of it as the collective strength of a crowd each gently holding hands, as seen in noble gases like krypton (Kr) and nonpolar molecules, where they are the primary form of cohesion.

Dipole-Dipole Interactions

Now, let's delve into the world of molecular magnetism with dipole-dipole interactions. These interactions are a notch above LDFs in terms of strength and occur in molecules that are polar—molecules where electrons are not shared equally between atoms, leading to a separation of charge, or a dipole moment.

The classic image is of tiny magnets within a substance, each having a positive and a negative pole, aligning themselves in a head-to-tail fashion. This alignment ensures that the positive end of one molecule is close to the negative end of another, creating an attractive force. This type of interaction is like a team pulling on a rope in a tug of war, stable and directional, exemplified by nitrogen trichloride (NCl3), with its pyramid molecular shape creating a distinct polar character.

The Special Case of Hydrogen Bonding

Hydrogen bonding is the star performer among intermolecular forces. It's a special kind of dipole-dipole interaction that exhibits exceptional strength and plays a vital role in many biological processes. This bond is no ordinary connection; it's formed specifically when hydrogen is directly attached to one of the profoundly electronegative elements like fluorine, nitrogen, or oxygen.

In essence, this is a love story between hydrogen and an electron-rich atom, with hydrogen (the proton) carrying a naked positive charge that is irresistibly attracted to the lone pair of electrons on the electronegative atom. Such bonds are strong, directional, and have a profound influence on physical properties like boiling and melting points. An exemplary molecule showcasing hydrogen bonding is hydrogen fluoride (HF), where hydrogen is bound to the greedy electron-loving fluorine, contributing to HF's higher boiling point and unique properties.

Molecular Geometry and Its Effects

The architecture of a molecule, known as molecular geometry, is a determining factor in its behavior and interaction with other molecules. It's not just about aesthetics; molecular geometry defines the distribution of electrons and therefore the overall polarity of the molecule. Imagine miniature sculptures, with angles and arrangements determining how the molecules fit together in space.

Geometric shapes can vary from linear to trigonal planar, tetrahedral to complex three-dimensional structures. For instance, the pyramid shape of NCl3 creates a polar molecule, while the tetrahedral layout of silane (SiH4) keeps it nonpolar, dictating the type and strength of intermolecular forces that will prevail. This symmetry—or lack thereof—can dramatically influence the molecule's physical and chemical properties.

Deciphering Molecular Polarity

When understanding the polarity of molecules, think of a tug-of-war on a molecular level. In the battleground of chemical bonding, atoms 'pull' on shared electrons, and the strength of their 'pull' or electronegativity determines the distribution of electron density. A molecule is considered polar if there is an uneven distribution of electron density, leading to a molecule with a positive end and a negative end.

This polarity has real-world effects. For example, it affects how a substance dissolves in solvents, its boiling and melting points, and how it interacts with other substances. The polarity of molecules is like the social preferences of people: some are exuberant and interact with everyone (polar molecules in solvents), while others are reserved and keep to themselves (nonpolar molecules forming LDFs). By examining a molecule's shape and the electronegativity of its atoms, we can predict its polarity and the resulting intermolecular forces.