Problem 53 - General Chemistry: Principles and Modern Applications

Question

Write plausible Lewis structures for the following odd-electron species: (a) \(\mathrm{CH}_{3} ;\) (b) \(\mathrm{ClO}_{2} ;\) (c) \(\mathrm{NO}_{3}\).

Step-by-Step Solution

Answer

The Lewis structures for the odd-electron species CH3, ClO2 and NO3 have been drawn. For CH3, the odd electron is located on the Carbon atom. For ClO2, the odd electron is on the central Oxygen. For NO3, the Nitrogen atom carries the unpaired electron.

Step 1: Determine the total valence electrons for each molecule

For CH3, Carbon (C) has 4 valence electrons and Hydrogen (H) has 1, multiplied by 3 gives 3 electrons. Total is 4 + 3 = 7 valence electrons. For ClO2, Chlorine (Cl) has 7 valence electrons and Oxygen (O) has 6, multiplied by 2 gives 12 electrons. Total is 7 + 12 = 19 valence electrons. For NO3, Nitrogen (N) has 5 valence electrons and Oxygen (O) has 6, multiplied by 3 gives 18 electrons. Total is 5 + 18 = 23 valence electrons.

Step 2: Assign the central atom and distribute the electrons

CH3: Carbon is the central atom with Hydrogen atoms around it. We put 2 electrons (a single bond) between the C and each H. This uses up 6 of our 7 electrons. The single remaining electron is placed on the central Carbon, making it an odd-electron species. ClO2: Oxygen is the central atom with Chlorine and another Oxygen around it. Single bond is made between Cl-O and O-O using 4 electrons. The rest are placed as lone pairs. The remaining electron (making it an odd-electron species) will be added to the central Oxygen. NO3: Nitrogen is the central atom with Oxygen atoms around it. Single bond is made among N-Ox3, using 6 electrons. Remaining ones are placed as lone pairs. The total gives 24 electrons, but only 23 are available so Nitrogen will carry one unpaired electron.

Step 3: Draw the Lewis structure

Draw the molecules according to the previous steps, making sure to appropriately place the odd (unpaired) electron.

Key Concepts

Odd-Electron Species

When dealing with chemical species that have an unpaired electron, we refer to them as odd-electron species or radicals. These types of molecules are quite special because most stable compounds want to have all their electrons paired up.

However, in some cases, due to the nature and the availability of electrons, species end up with an odd number of valence electrons. This makes them more reactive than their even-electron counterparts.

• For example, in oindent CH extsubscript{3} , each H contributes 1 electron and C contributes 4 electrons, =7 total, causing an unpaired electron.
• oindent ClO extsubscript{2} has 7 from Cl and 6 from each O, =19 electrons, so Cl gets an odd number.
• For NO extsubscript{3}, Nitrogen ends up with an unpaired electron.

It's this lone electron that often leads them to participate actively in reactions, as they seek to pair up or share their odd electron for stability.

Valence Electrons

Understanding valence electrons is crucial to grasping chemical bonding and drawing Lewis structures. Valence electrons are the electrons found in the outermost shell of an atom and determine how atoms interact during a chemical bond.

When counting valence electrons, we consider the outer shell's electron count. For example:

• Carbon (C) has 4 valence electrons, Hydrogen (H) has 1, Chlorine (Cl) has 7, and Oxygen (O) has 6.

These electrons are used in bonds, whether they're shared (covalent bonding) or transferred (ionic bonding). In Lewis structures, valence electrons are shown as dots around a symbol, or as lines representing shared pairs in bonds. Having a clear understanding of these electrons allows us to confidently predict and describe molecular structure and behavior.

Chemical Bonding

Chemical bonding is the process by which atoms combine to form molecules or compounds. Bonds can form as atoms share or transfer electrons to achieve a more stable electron configuration.

In molecules like CH extsubscript{3} , ClO extsubscript{2} , and NO extsubscript{3}, the structure is simplified and represented using Lewis structures, showcasing the bonding and non-bonding electron pairs. In a Lewis structure:

• A single line represents a bond, showing two shared electrons.
• Dots signify unpaired/valence electrons not involved in bonds.
• Electrons are typically shared in homonuclear diatomic molecules for covalent bonds. Absolute sharing is more prevalent in polar covalent and ionic bonds.

The goal for atoms forming bonds is generally to achieve a pleasant state similar to noble gases with full valence shells. This often leads to stable, lower-energy states, making chemical bonding vital for molecule stability and function.