Problem 27 - Chemistry: Principles and Reactions

Question

Give the Lewis structure of (a) \(\mathrm{NO}_{2}\) (b) \(\mathrm{NO}\) (c) \(\mathrm{SO}_{2}\) (d) \(\mathrm{SO}_{3}\)

Step-by-Step Solution

Answer

Answer: The final Lewis structures for the molecules are as follows: (a) [O=N-O]⁻, (b) O-N (with N having an unpaired electron), (c) O=S-O, and (d) O=S(=O)-O.

Step 1: Count the valence electrons

For each molecule, count the total number of valence electrons from the atoms. Nitrogen (N) has 5 valence electrons, Oxygen (O) has 6, and Sulfur (S) has 6. (a) NO₂: N has 5 valence electrons and each O has 6, so total electrons = 5 + 6 + 6 = 17 (b) NO: N has 5 valence electrons and O has 6, so total electrons = 5 + 6 = 11 (c) SO₂: S has 6 valence electrons and each O has 6, so total electrons = 6 + 6 + 6 = 18 (d) SO₃: S has 6 valence electrons and each O has 6, so total electrons = 6 + 6 + 6 + 6 = 24

Step 2: Arrange the atoms

For each molecule, place the central atom (the one with the lowest electronegativity) and connect the surrounding atoms with a single bond. (a) NO₂: N is the central atom, with two O atoms connected by single bonds. (b) NO: N and O are connected by a single bond. (c) SO₂: S is the central atom, with two O atoms connected by single bonds. (d) SO₃: S is the central atom, with three O atoms connected by single bonds.

Step 3: Distribute the remaining electrons

For each molecule, place the remaining electrons on the surrounding atoms to satisfy the octet rule (except for N in NO and NO₂, which will have an odd number of electrons). (a) NO₂: Assign 2 lone pairs (4 electrons) to each O atom, and 1 non-bonding electron to N. Total electrons: 17. (b) NO: Assign 3 lone pairs (6 electrons) to O and 2 non-bonding electrons to N. Total electrons: 11. (c) SO₂: Assign 2 lone pairs (4 electrons) to each O atom, and 1 lone pair (2 electrons) to S. Total electrons: 18. (d) SO₃: Assign 2 lone pairs (4 electrons) to each O atom. No lone pairs are left for S. Total electrons: 24.

Step 4: Check the octet rule

Check if each atom satisfies the octet rule (8 electrons). If not, form double or triple bonds by sharing additional electron pairs. (a) NO₂: N has 7 electrons, O₂ has 8. Form a double bond between N and one of the O atoms. Final structure: [O=N-O]⁻ (b) NO: N has 7 electrons, O has 8. We cannot form any double or triple bond, so the structure remains: O-N (with N having an unpaired electron) (c) SO₂: S has 6 electrons, each O has 8. Form a double bond between S and one of the O atoms. Final structure: O=S-O (d) SO₃: S and each O has 8 electrons. Final structure: O=S(=O)-O

Key Concepts

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom, which can participate in the formation of chemical bonds. These electrons determine how atoms interact with each other, as they are the ones that are exchanged or shared when atoms bond chemically. Understanding valence electrons is essential to determining how many bonds an atom can form.

For example, in the exercise provided, the number of valence electrons was counted for molecules like \(\mathrm{NO}_{2}\), \(\mathrm{NO}\), \(\mathrm{SO}_{2}\), and \(\mathrm{SO}_{3}\). Nitrogen (N) has 5 valence electrons while oxygen (O) and sulfur (S) both have 6. Calculating the total count of valence electrons is the first step to creating the Lewis structures, which visually represent the arrangement of electrons in molecules.

Octet Rule

The octet rule is a chemical rule of thumb that states atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. The rule applies to main group elements, with notable exceptions for hydrogen (which seeks a duet of electrons), helium, and transition metals.

Atoms will share, donate or receive electrons to achieve a full octet. This can result in the formation of ionic or covalent bonds. The steps provided in the exercise include verification of the octet rule for each atom within the molecules. If an atom does not have an octet, adjustments like forming double or triple bonds between atoms are made. Such an adjustment is seen in \(\mathrm{NO}_{2}\) and \(\mathrm{SO}_{2}\), where double bonds were necessary to comply with the octet rule.

Lewis Dot Structures

Lewis dot structures, also known as Lewis structures, are diagrams that represent the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. The Lewis structure can predict the number of bonds that an atom can form, the arrangement of atoms, and whether or not additional electrons need to be added as dots to represent lone pairs.

A crucial aspect of drawing Lewis structures is to ensure that the valence shell is filled appropriately according to the octet rule. In the exercise, each step of the process involves placing dots for lone pairs and connecting atoms with lines to indicate chemical bonds. For instance, \(\mathrm{NO}_{2}\) was depicted with a double bond to fulfill nitrogen's octet.

Chemical Bonding

Chemical bonding is the physical process that allows for the formation of molecules and compounds through the sharing, exchanging, or transferring of valence electrons. There are several types of bonds including ionic, covalent, and metallic bonds.

In the provided exercise, covalent bonding was the main focus, where atoms share pairs of electrons. For each molecule, such as \(\mathrm{SO}_{3}\), it is imperative to assess whether the atoms are sharing electrons in a manner that allows for a stable structure in accordance with the octet rule. The final Lewis structures signify the full representation of the chemical bonding in the molecules.