Question

Write a Lewis structure for each of the following polyatomic ions. Show all bonding valence electron pairs as lines and all nonbonding valence electron pairs as dots. For those ions that exhibit resonance, draw the various possible resonance forms. a. carbonate ion, ${\mathrm{CO}}_3^{2-}$ b. ammonium ion, ${\mathrm{NH}}_4^{+}$ c. hypochlorite ion, ${\mathrm{ClO}}^{-}$

Short answer

The Lewis structures for the polyatomic ions are: a. Carbonate ion, ${\mathrm{CO}}_3^{2-}$: Resonance structures $$\begin{array}{ccc}\text{O}-\text{C}-\text{O}& \text{O}\equiv \text{C}-\text{O}& \text{O}-\text{C}\equiv \text{O}\\ \text{|}& & \text{ensuremath downarrow}\\ \text{O}& & \text{O}\end{array}$$ b. Ammonium ion, ${\mathrm{NH}}_4^{+}$: $$\begin{gathered} \text{H}-\text{N}-\text{H} \\ \text{|} \\ \text{H} \end{gathered}$$ c. Hypochlorite ion, ${\mathrm{ClO}}^{-}$: $$\text{O}-\text{Cl}:$$

Step-by-step solution

Count the total number of valence electrons.

For each ion, we will first count the total number of valence electrons. Recall that the valence electrons are the electrons located in the outermost shell of an atom. a. Carbonate ion, ${\mathrm{CO}}_3^{2-}$: Carbon (C) has 4 valence electrons, oxygen (O) has 6 valence electrons, and there are also 2 additional electrons due to the 2- charge of the ion. So, we have $4+3\times 6+2=24$ valence electrons. b. Ammonium ion, ${\mathrm{NH}}_4^{+}$: Nitrogen (N) has 5 valence electrons, hydrogen (H) has 1 valence electron, and there is one electron less due to the 1+ charge of the ion. So, we have $5+4\times 1-1=8$ valence electrons. c. Hypochlorite ion, ${\mathrm{ClO}}^{-}$: Chlorine (Cl) has 7 valence electrons, oxygen (O) has 6 valence electrons, and there is one additional electron due to the 1- charge of the ion. So, we have $7+6+1=14$ valence electrons.

Drawing the Lewis structures.

Now that we know the total number of valence electrons for each ion, we can begin to draw the Lewis structures. a. Carbonate ion, ${\mathrm{CO}}_3^{2-}$: We start with a single bond between C and each of the three O atoms, which accounts for six valence electrons used. Then, we complete the octet for each oxygen atom, which requires 6 more electrons per oxygen atom. The total number of electrons used so far is $6+3\times 6=24$, which is equal to the total number of valence electrons. We can draw the different resonance forms: O - C - O | O O O \ // C - O O - C - O \ O These three structures are resonance structures. b. Ammonium ion, ${\mathrm{NH}}_4^{+}$: In this case, we have a single bond between N and each of the four H atoms, which accounts for eight valence electrons used. This ion does not exhibit resonance. H - N - H | H c. Hypochlorite ion, ${\mathrm{ClO}}^{-}$: We start with a single bond between Cl and O, which accounts for two valence electrons. Then, we complete the octet for both atoms, requiring 6 electrons for O and 6 electrons for Cl. The total number of electrons used is $2+6+6=14$, which is equal to the total number of valence electrons. The hypochlorite ion does not exhibit resonance. O // Cl | :

Key concepts

Valence Electrons

In the realm of chemistry, valence electrons play a critical role as they are the electrons found in the outermost shell of an atom. These electrons are essential because they are involved in bonding between atoms, particularly in the formation of molecules and ions. Valence electrons help determine how atoms interact with each other and therefore dictate the chemical properties of an element.
To find the number of valence electrons in an ion, start by identifying the valence electrons in each of the atoms that constitute the ion. For neutral atoms, the valence electrons are generally equal to the group number of the element on the periodic table. However, ions can carry extra or fewer electrons due to their charges:

• Negative ions (anions) gain extra electrons, which adds to the total valence count.
• Positive ions (cations) lose electrons, reducing the total valence electron count.

By counting the valence electrons accurately, as demonstrated with the carbonate, ammonium, and hypochlorite ions, you can correctly set the stage for constructing accurate Lewis structures.

Resonance Structures

Resonance structures are different ways of representing a molecule or ion that cannot be described by a single Lewis structure. This occurs when there are multiple valid arrangements of electrons within the molecule. These structures help us understand the delocalization of electrons across different parts of the molecule, providing a more comprehensive picture of bonding.
For example, the carbonate ion ${\mathrm{CO}}_3^{2-}$ offers an excellent illustration of resonance. It has three Lewis structures because the double bond between carbon and oxygen can be placed in three different positions. Despite the different placements, none of these structures accurately represents the full nature of the ion alone. Instead, the true structure is a hybrid, sharing characteristics of all the forms:

• Each resonance structure maintains the same arrangement of atoms, only the positions of the electrons differ.
• These structures are not isolated, static entities; they mix in real molecular systems.
• Resonance structures contribute to a lower energy state and therefore more stability.

The concept of resonance is vital because it allows chemists to conceive of molecular structures as more dynamic and complex than a single snapshot.

Polyatomic Ions

Polyatomic ions are composed of two or more atoms covalently bonded together, that carry a net charge, either positive or negative. Understanding polyatomic ions is essential for grasping the nature of ionic compounds and the roles these ions play in chemical reactions.
A key aspect of polyatomic ions is the fact that they do not exist as a mere collection of individual atoms but as a single, charged entity that acts as a unit during chemical processes. Some important characteristics of polyatomic ions include:

• The entire ion carries a charge, not just any single atom within it.
• They generally contain an atom that functions as the central atom, bonded to surrounding atoms or groups.
• The overall charge of a polyatomic ion arises from the loss or gain of electrons, making them anions or cations.

To draw Lewis structures for polyatomic ions, it is crucial to account for both the shape of the ion and its charge by adjusting the total number of valence electrons. The ammonium ion ${\mathrm{NH}}_4^{+}$, for instance, loses one electron because it carries a positive charge, ensuring proper formation and representation in its structure.