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Chapter 9: Problem 16

What property of the electron causes electron domains to have an effect on molecular shapes?

Short Answer

Expert verified
The property of electrons that causes electron domains to have an effect on molecular shapes is their negative charge. This leads to repulsion between electron domains, governed by the Valence Shell Electron Pair Repulsion (VSEPR) theory, which in turn dictates the overall geometry of molecules to minimize these repulsive forces.

Step by step solution

01

Understand electron domains

Electron domains are regions around an atom where valence electrons are most likely to be found. They include both bonding and non-bonding electron pairs. The electron pairs occupy space around the central atom, and they repel each other in an attempt to minimize the repulsive forces between them. This repulsion between electron domains determines the overall geometry of the molecule.
02

Electron properties responsible for domain repulsion

The property of electron responsible for the repulsion between electron domains is their negative charge. Electrons carry a negative charge, and as a result, they repel other negatively charged particles, such as other electrons. This repulsion is governed by the principle known as the Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that electron pairs in the valence shell of an atom repel one another, and they arrange themselves in a way to minimize these repulsive forces.
03

How electron repulsion affects molecular shape

The repulsion between electron domains plays a significant role in determining the molecular shape. In an attempt to minimize the repulsive forces, the electron domains occupy positions that maximize the angles between them and dictate the molecular geometry. For example, if a central atom has two electron domains around it, they will arrange themselves linearly with an angle of 180 degrees. If there are three electron domains, they will arrange in a trigonal planar shape with angles of 120 degrees, and so on. In conclusion, the property of electrons that causes electron domains to have an effect on molecular shapes is their negative charge. This leads to repulsion between electron domains, which in turn dictates the overall geometry of molecules to minimize these repulsive forces.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Electron Domains
Electron domains are essentially areas around a central atom where electrons are most likely to be found. These regions can consist of different types of electron pairs, such as bonding pairs that form covalent bonds between atoms, and non-bonding pairs, also known as lone pairs. Each domain takes up space and needs to be considered when predicting a molecule's shape.
One might wonder why these regions are important. They are crucial because they reflect the behavior of electrons due to their negative charge. When electrons are grouped into domains around an atom, they are not static; they occupy specific regions due to forces that keep them repelled from each other.
  • Bonding pairs: Electrons shared between atoms, forming a chemical bond.
  • Lone pairs: Electrons that are not used for bonding but occupy space.
In essence, understanding electron domains gives you insight into the structure of a molecule since they actively define its geometrical layout.
Molecular Geometry
The molecular geometry of a molecule is determined largely by the electron domains surrounding its central atom. Each arrangement seeks to minimize the repulsion between electron domains, thus dictating the shape of the molecule. Angles between these domains help establish this geometry.
Think about a simple example: water. This molecule has two hydrogen atoms bonded to an oxygen atom and two lone pairs also influencing its shape. While one might predict a linear shape based only on two bonds, the influence of lone pairs results in a bent molecular shape instead. This bending arises because lone pairs take up more space and push the bonding pairs closer together.
Molecular shapes you might encounter include:
  • Linear - at 180 degrees (two domains).
  • Trigonal planar - at 120 degrees (three domains).
  • Tetrahedral - at 109.5 degrees (four domains).
Understanding molecular geometry is essential as it influences molecular polarity, reactivity, phase of matter, and more.
Electron Repulsion
At the heart of electron domains and molecular geometry is the concept of electron repulsion. Electrons with their negative charges naturally repel each other. This repulsion forms the basis for the Valence Shell Electron Pair Repulsion (VSEPR) theory. According to this principle, electron pairs arrange themselves as far apart as possible around a central atom to minimize repulsive forces.
Consider electron repulsion like traffic flow: each car (or electron domain) needs its space to move smoothly and avoid collisions. This results in an organizational pattern that ensures minimal conflict among domains.
Several factors influence how electrons repel each other:
  • Electron charge: All electrons inherently repel due to their inherent negative charge.
  • Type of electrons: Lone pairs repel more strongly than bonding pairs.
  • Electronegativity: Atoms with high electronegativity pull electron domains closer, affecting repulsion.
Understanding electron repulsion is key to predicting how molecules behave in chemical reactions and their physical properties.

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Most popular questions from this chapter

Methyl isocyanate, \(\mathrm{CH}_{3} \mathrm{NCO},\) was made infamous in 1984 when an accidental leakage of this compound from a storage tank in Bhopal, India, resulted in the deaths of about 3,800 people and severe and lasting injury to many thousands more. (a) Draw a Lewis structure for methyl isocyanate. (b) Draw a ball-and-stick model of the structure, including estimates of all the bond angles in the compound. (c) Predict all the bond distances in the molecule. (d) Do you predict that the molecule will have a dipole moment? Explain.

An \(\mathrm{AB}_{3}\) molecule is described as having a trigonal-bipyramidal electron-domain geometry. How many nonbonding domains are on atom A? Explain.

(a) What is the difference between a localized \(\pi\) bond and a delocalized one? (b) How can you determine whether a molecule or ion will exhibit delocalized \(\pi\) bonding? (c) Is the \(\pi\) bond in \(\mathrm{NO}_{2}^{-}\) localized or delocalized?

Azo dyes are organic dyes that are used for many applications, such as the coloring of fabrics. Many azo dyes are derivatives of the organic substance azobenzene, \(\mathrm{C}_{12} \mathrm{H}_{10} \mathrm{~N}_{2}\). A closely related substance is hydrazobenzene, \(\mathrm{C}_{12} \mathrm{H}_{12} \mathrm{~N}_{2} .\) The Lewis structures of these two substances are (Recall the shorthand notation used for benzene.) (a) What is the hybridization at the \(\mathrm{N}\) atom in each of the substances? (b) How many unhybridized atomic orbitals are there on the \(\mathrm{N}\) and the \(C\) atoms in each of the substances? (c) Predict the \(\mathrm{N}-\mathrm{N}-\mathrm{C}\) angles in each of the substances. (d) Azobenzene is said to have greater delocalization of its \(\pi\) electrons than hydrazobenzene. Discuss this statement in light of your answers to (a) and (b). (e) All the atoms of azobenzene lie in one plane, whereas those of hydrazobenzene do not. Is this observation consistent with the statement in part (d)? (f) Azobenzene is an intense red-orange color, whereas hydrazobenzene is nearly colorless. Which molecule would be a better one to use in a solar energy conversion device? (See the "Chemistry Put to Work" box for more information about solar cells.)

The azide ion, \(\mathrm{N}_{3}^{-}\), is linear with two \(\mathrm{N}-\mathrm{N}\) bonds of equal length, \(1.16 \AA\). (a) Draw a Lewis structure for the azide ion. (b) With reference to Table \(8.5,\) is the observed \(\mathrm{N}-\mathrm{N}\) bond length consistent with your Lewis structure? (c) What hybridization scheme would you expect at each of the nitrogen atoms in \(\mathrm{N}_{3}^{-}\) ? (d) Show which hybridized and unhybridized orbitals are involved in the formation of \(\sigma\) and \(\pi\) bonds in \(\mathrm{N}_{3}^{-}\). (e) It is often observed that \(\sigma\) bonds that involve an \(s p\) hybrid orbital are shorter than those that involve only \(s p^{2}\) or \(s p^{3}\) hybrid orbitals. Can you propose a reason for this? Is this observation applicable to the observed bond lengths in \(\mathrm{N}_{3}^{-}\) ?
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