Question

Draw the Lewis structures for each of the following ions or molecules. Identify those that do not obey the octet rule, and explain why they do not. (a) \(\mathrm{SO}_{3}{ }^{2-}\), (b) \(\mathrm{AlH}_{3}\), (c) \(\mathrm{N}_{3}^{-}\), (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) (e) \(\mathrm{SbF}_{5}\).

Short answer

(a) \(\mathrm{SO}_{3}{ }^{2-}\) does not obey the octet rule, as sulfur has an expanded octet due to its empty d-orbitals. (b) \(\mathrm{AlH}_{3}\) does not obey the octet rule, as aluminum has a sextet due to its low electronegativity and available empty orbitals. (c) \(\mathrm{N}_{3}^{-}\) obeys the octet rule, as each nitrogen atom has a full octet. (d) \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\) obeys the octet rule, as each atom has a full octet (or duet for hydrogen). (e) \(\mathrm{SbF}_{5}\) does not obey the octet rule, as antimony has 10 electrons in its valence shell due to its empty d-orbitals.

Step-by-step solution

(a) \(\mathrm{SO}_{3}{ }^{2-}\)#

1. Sulfur (S) has 6 valence electrons, oxygen (O) has 6 valence electrons, and the ion has 2 extra electrons due to the 2- charge. Therefore, the total number of valence electrons is 6+3*(6)+2= 26. 2. Sulfur is the central atom being surrounded by three oxygen atoms. 3. We distribute the remaining electrons, starting with oxygen atoms to complete their octets, and then completing the sulfur atom octets as well. 4. Since this is an ion, we have considered the extra 2 electrons. 5. In the final structure, sulfur atom has more than 8 electrons (an expanded octet). So, it does not obey the octet rule. Sulfur can have more than 8 electrons due to its empty d-orbitals that can accommodate extra electrons.

(b) \(\mathrm{AlH}_{3}\)#

1. Aluminum (Al) has 3 valence electrons and hydrogen (H) has 1 valence electron. The total number of valence electrons is 3+3*(1) = 6. 2. Aluminum is the central atom surrounded by three hydrogen atoms. 3. We distribute the remaining electrons, starting with hydrogen atoms, and then completing any remaining electrons to the aluminum atom. 4. There is no charge in this molecule, so no extra electrons are needed. 5. In the final structure, the aluminum atom has only 6 electrons (a sextet) instead of 8 electrons. So, it does not obey the octet rule. Aluminum can do this due to a low electronegativity and the availability of empty orbitals for bonding.

(c) \(\mathrm{N}_{3}^{-}\)#

1. Nitrogen (N) has 5 valence electrons, and the ion has 1 extra electron due to the 1- charge. Therefore, the total number of valence electrons is 3*(5)+1 = 16. 2. There is only one type of atom in this ion, so the central atom is nitrogen, and it is surrounded by two other nitrogen atoms. 3. We distribute the remaining electrons to complete the octets for each nitrogen atom. 4. Since this is an ion, we have considered the extra electron. 5. In the final structure, each nitrogen atom has 8 electrons, obeying the octet rule.

(d) \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\)#

1. Carbon (C) has 4 valence electrons, hydrogen (H) has 1 valence electron, and chlorine (Cl) has 7 valence electrons. The total number of valence electrons is 4+2*(1)+2*(7) = 20. 2. Carbon is the central atom, surrounded by two hydrogen atoms and two chlorine atoms. 3. We distribute the remaining electrons to complete the octets for each hydrogen and chlorine atom and then add any remaining electrons to complete the octet for the carbon atom. 4. There is no charge in this molecule, so no extra electrons are needed. 5. In the final structure, each atom has a full octet (or duet). So, it obeys the octet rule.

(e) \(\mathrm{SbF}_{5}\)#

1. Antimony (Sb) has 5 valence electrons, and fluorine (F) has 7 valence electrons. The total number of valence electrons is 5+5*(7) = 40. 2. Antimony is the central atom surrounded by five fluorine atoms. 3. We distribute the remaining electrons first, completing the octets for each fluorine atom and then adding any remaining electrons to complete the octet for the antimony atom. 4. There is no charge in this molecule, so no extra electrons are needed. 5. In the final structure, the antimony atom has 10 electrons (exceeding the octet). So, it does not obey the octet rule. Antimony can have an expanded octet due to its empty d-orbitals that can accommodate extra electrons.

Key concepts

The Octet Rule Explained

The octet rule is a foundational concept when it comes to understanding chemical bonding, particularly in covalent molecules. It states that most elements achieve a stable electronic configuration when they have eight electrons in their valence shell, which is reminiscent of the electron configuration of noble gases.

This rule guides the arrangement of electrons for elements in the main groups of the periodic table and is crucial for drawing Lewis structures. However, there are exceptions to the octet rule. For instance, molecules or ions like \(\mathrm{SO}_{3}{ }^{2-}\) and \(\mathrm{SbF}_{5}\) showcase that elements from period 3 or higher can accommodate more than eight electrons, known as an expanded octet. This occurs because of the presence of available d-orbitals in these elements that can be used for bonding.

Conversely, species like \(\mathrm{AlH}_{3}\) have less than eight electrons around the central atom and still are stable, a situation that is explained by their capacity to form incomplete octets.

Valence Electrons and Chemical Bonding

Valence electrons are the outermost electrons of an atom and are crucial in determining how the atom can bond with others. They are responsible for forming chemical bonds, whether they are ionic or covalent. The number of valence electrons an atom has influences the shape and type of Lewis structure it can form.

In Lewis structures, valence electrons are represented by dots surrounding the atomic symbols. Counting these electrons is the first step in predicting how atoms will bond. For example, oxygen has six valence electrons, while hydrogen has one. In exercises like drawing the Lewis structure for \(\mathrm{CH}_{2} \(\mathrm{Cl}\)_{2}\), carbon forms four bonds — two with hydrogen and two with chlorine — using its four valence electrons to achieve a full octet.

Counting Valence Electrons

It’s crucial for students to learn how to correctly count the valence electrons for each element involved in a molecule or ion. This is because each step in the creation of the Lewis structure, from selecting the central atom to distributing the electrons, relies on the accurate count of valence electrons.

Molecular Geometry: Shapes and Structures

Once the Lewis structures are drawn, understanding the molecular geometry of the molecule becomes key. Molecular geometry, also known as molecular shape, is the three-dimensional arrangement of atoms within a molecule. The geometry is significant because it influences the molecule's properties, including reactivity, polarity, and phase of matter under standard conditions.

VSEPR (Valence Shell Electron Pair Repulsion) theory is often used to predict molecular geometry. It dictates that electron pairs around a central atom will arrange themselves in a way that minimizes repulsion between them, leading to the observation of various geometric shapes such as linear, trigonal planar, tetrahedral, and more.

Practice Through Examples

The question on determining the shapes and explaining the molecular geometry for ions like \(\mathrm{N}_{3}^{-}\) can be an excellent practical application for students to visualize these shapes. It is essential for them to understand that the molecular geometry not only depends on the bonding pairs of electrons but also on the lone pairs, as they affect the spatial arrangement of the atoms involved.