Question

Draw Lewis structures for these ions, and show which atom in each bears the formal charge. (a) \(\mathrm{CH}_{3} \mathrm{NH}_{3}^{+}\) (b) \(\mathrm{CO}_{3}{ }^{2-}\) (c) \(\mathrm{OH}^{-}\)

Short answer

(a) \(\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}\) (b) \(\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{OH}^{-}\) Answer: (a) In \(\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}\), the Nitrogen atom bears the formal charge. (b) In \(\mathrm{CO}_{3}^{2-}\), one of the Oxygen atoms bears the formal charge. (c) In \(\mathrm{OH}^{-}\), the Oxygen atom bears the formal charge.

Step-by-step solution

Count total electrons

Count the total number of valence electrons in the molecule, adding or subtracting one electron per positive or negative charge. Here, we have 3 Hydrogen atoms (3 x 1 e-), 1 Carbon atom (4 e-), 1 Nitrogen atom (5 e-) and a +1 charge (-1 e-). Thus, we have a total of 3 + 4 + 5 - 1 = 11 electrons.

Arrange atoms and place electrons

Start with the least electronegative atom (Carbon) in the center and arrange other atoms around it, resulting in the structure H - C - N - H, with a hydrogen atom attached to both C and N. Distribute the electrons around the atoms in pairs, starting with the atoms directly bonded together.

Complete octets and formal charges

Fill up the remaining electrons in the outer shell of each atom, making sure not to exceed the octet rule. Calculate the formal charge of each atom in the ion using this formula: Formal Charge = Valence electrons - Non-bonding electrons - 1/2 Bonding electrons The \(\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}\) Lewis structure becomes: H H H | | | C―N | H Since Nitrogen has a lower formal charge than the other atoms, it will bear the formal charge in the ion. (b) For \(\mathrm{CO}_{3}{ }^{2-}\)

Count total electrons

Count the total number of valence electrons in the ion, adding or subtracting one electron per positive or negative charge. Here, we have 1 Carbon atom (4 e-), 3 Oxygen atoms (3 x 6 e-) and a -2 charge (+2 e-). Thus, we have a total of 4 + 18 + 2 = 24 electrons.

Arrange atoms and place electrons

Start with the least electronegative atom (Carbon) in the center and arrange other atoms (Oxygen) around it. Connect Carbon to each Oxygen with a single bond, and distribute the electrons around the atoms in pairs, starting with the atoms directly bonded together.

Complete octets and formal charges

Fill up the remaining electrons in the outer shell of each atom, deciding on single, double, or triple bonds as necessary to satisfy the octet rule. Calculate the formal charge of each atom in the ion. The \(\mathrm{CO}_{3}{ }^{2-}\) Lewis structure becomes: O // C \ O- \ O Each of the three Oxygen atoms has a formal charge of -1. One Oxygen atom bears a formal charge in the ion. (c) For \(\mathrm{OH}^{-}\)

Count total electrons

Count the total number of valence electrons in the ion, adding or subtracting one electron per positive or negative charge. Here, we have 1 Hydrogen atom (1 e-), 1 Oxygen atom (6 e-), and a -1 charge (+1 e-). Thus, we have a total of 1 + 6 + 1 = 8 electrons.

Arrange atoms and place electrons

Start with the least electronegative atom (Oxygen) and connect with bonding electrons. In this ion, there are only two atoms: Oxygen and Hydrogen, which will be connected by a single bond.

Complete octets and formal charges

Fill up the remaining electrons in the outer shell of each atom, making sure not to exceed the octet rule. Calculate the formal charge of each atom in the ion. The \(\mathrm{OH}^{-}\) Lewis structure becomes: O // H The Oxygen has a formal charge of -1 and bears the formal charge in the ion.

Key concepts

Formal Charge

Formal charge is a concept used to determine the actual load an atom carries in a molecule or ion. To calculate the formal charge, a simple formula is used: Formal Charge = Valence electrons - Non-bonding electrons - 1/2 Bonding electrons.

• **Valence Electrons**: These are the electrons present in the outer shell of an atom. Different elements have a different count of valence electrons
• **Non-bonding Electrons**: Also known as lone pairs, these electrons are not involved in forming bonds.
• **Bonding Electrons**: Electrons that participate in bonding between atoms. If an atom forms two bonds, it has four bonding electrons.

By applying this formula, you can check if the charges in a Lewis structure are balanced and identify the atom with a higher tendency to hold a charge. In Lewis structures, it's crucial to minimize formal charges to stabilize the structure and predict the behavior of the molecule.

Valence Electrons

Valence electrons play a vital role in forming chemical bonds and determining the Lewis structure of a molecule or ion. The entities participate in the first instance of bonding from the atoms' outer shell.

For example, carbon has four valence electrons, while oxygen has six. Knowing the number of valence electrons helps in counting the total electrons needed for Lewis structure.

To find the total electrons in a neutral molecule:

• Count the valence electrons from each atom.
• Add or subtract electrons equal to the ion's charge (add for negative charge, subtract for positive charge).

Counting valence electrons accurately ensures the correct distribution of electrons in the molecule's structure. This fundamental step helps locate non-bonding and bonding electrons easily and complete the understanding of the molecular shape.

Octet Rule

The octet rule, a fundamental principle in chemistry, posits that atoms tend to establish bonds until they hold eight electrons in their valence shell to achieve a stable electronic configuration. Most Lewis structures obey this rule, ensuring the stability and neutrality of the molecule.

While drawing a Lewis structure, use the octet rule to:

• Ensure that most atoms (like carbon, nitrogen, and oxygen) form bonds that complete their outer shells to eight electrons.
• Recognize exceptions: Elements like hydrogen are stable with only two electrons, while others like sulfur might expand their octet.

When correctly achieved, a structure that follows the octet rule generally translates to a stable and balanced molecular or ionic entity. Identifying these completed octets can also help in predicting potential reactions and the strength of bonds within the molecule.