Question

Write a Lewis structure that obeys the octet rule for each molecule or ion. Include resonance structures if necessary and assign formal charges to each atom. \begin{equation} \mathrm { a. } {SeO}_{2} \quad \text { b. } \mathrm{CO}_{3}^{2-} \quad \text { c. } \mathrm{ClO}^{-} \quad \text { d. } \mathrm{NO}_{2}^{-} \end{equation}

Short answer

SeO2: Se central atom, double bonds with each O, no formal charges. CO3^2-: C central atom, single bonds with each O, formal charge of -1 on each O and +1 on C, with resonance. ClO^-: Cl central atom, single bond with O, formal charge of -1 on O, no charge on Cl. NO2^-: N central atom, one double bond and one single bond with the Os, formal charge of -1 on single-bonded O, with resonance.

Step-by-step solution

Determine the total number of valence electrons for SeO2

Selenium (Se) has 6 valence electrons as a member of group 16. Each Oxygen (O) has 6 valence electrons. Summing these up: Se: 6e-, O: 6e- * 2 = 12e-, Total: 6e- + 12e- = 18 valence electrons.

Draw the skeleton structure for SeO2

Place Selenium in the center since it is less electronegative. Connect it with two oxygens using single bonds. This accounts for 4 electrons, leaving 14 more to be placed.

Complete the octets for Oxygen in SeO2

Place six electrons (three lone pairs) around each Oxygen to satisfy the octet rule, using 12 of the remaining electrons. This leaves 2 electrons.

Place remaining electrons and assign formal charges for SeO2

Place the remaining 2 electrons on Selenium as a lone pair. Calculate formal charges: Oxygen has no formal charge (6 valence electrons - 6 nonbonding - 1 bonding), and Selenium has no formal charge (6 valence electrons - 2 nonbonding - 2 bonding).

Determine the total number of valence electrons for CO3^2-

Carbon (C) has 4 valence electrons, each Oxygen (O) has 6 valence electrons, and the 2- charge adds 2 more electrons. Summing these up: C: 4e-, O: 6e- * 3 = 18e-, charge: 2e-, Total: 4e- + 18e- + 2e- = 24 valence electrons.

Draw the skeleton structure for CO3^2-

Place Carbon in the center and connect it with three Oxygens using single bonds. This accounts for 6 electrons, leaving 18 more to be placed.

Complete the octets on Oxygen in CO3^2-

Place six electrons (three lone pairs) around each Oxygen to satisfy the octet rule, using 18 of the remaining electrons. All valence electrons are now used.

Place any remaining electrons, assign formal charges, and show resonance for CO3^2-

No remaining electrons. Calculate formal charges and include resonance structures. Each Oxygen will have a -1 formal charge (6 valence electrons - 6 nonbonding - 1 bonding), and Carbon will have a +1 formal charge (4 valence electrons - 3 bonding). For resonance, the negative charge should rotate among the three Oxygens.

Determine the total number of valence electrons for ClO-

Chlorine (Cl) has 7 valence electrons and Oxygen (O) has 6 valence electrons. The 1- charge adds 1 more electron. Summing these up: Cl: 7e-, O: 6e-, charge: 1e-, Total: 7e- + 6e- + 1e- = 14 valence electrons.

Draw the skeleton structure for ClO-

Place Chlorine in the center and connect it with Oxygen using a single bond. This uses 2 electrons, leaving 12 more to be placed.

Complete the octets and assign formal charges for ClO-

Place six electrons (three lone pairs) on each atom to satisfy the octet rule, using all the remaining electrons. Calculate formal charges: Oxygen will have a -1 formal charge (6 valence electrons - 6 nonbonding - 1 bonding), and Chlorine will have no formal charge (7 valence electrons - 6 nonbonding - 2 bonding).

Determine the total number of valence electrons for NO2^-

Nitrogen (N) has 5 valence electrons, each Oxygen (O) has 6 valence electrons, and the 1- charge adds 1 more electron. Summing these up: N: 5e-, O: 6e- * 2 = 12e-, charge: 1e-, Total: 5e- + 12e- + 1e- = 18 valence electrons.

Draw the skeleton structure for NO2^-

Place Nitrogen in the center and connect it with two Oxygens using single bonds. This uses 4 electrons, leaving 14 more to be placed.

Complete the octets and assign formal charges, and show resonance for NO2^-

Place six electrons (three lone pairs) on one Oxygen and five electrons (two lone pairs and one single electron) on Nitrogen, leaving three electrons. Use a double bond to connect Nitrogen with the other Oxygen, satisfying the octet rule for Nitrogen and the second Oxygen. Calculate formal charges and show the resonance structure where the double bond switches between the two Oxygens.

Key concepts

Valence Electrons

In chemistry, valence electrons are the electrons in the outermost shell of an atom that can participate in forming chemical bonds with other atoms. They are crucial for understanding how atoms interact and bond in a molecule.

Regarding the example of SeO2, each oxygen atom contributes six valence electrons while selenium brings in another six. By summing these values, we find that SeO2 has a total of 18 valence electrons to work with when creating a Lewis structure. Similarly, when determining the Lewis structure for polyatomic ions such as CO3^2-, we must consider the charge of the ion which affects the total count of valence electrons. In this cases the extra electrons due to the 2- charge are added to the total valence electron count, resulting in 24 valence electrons for CO3^2-.

The concept of valence electrons is deeply intertwined with the octet rule, which states that atoms are generally more stable when they have eight electrons in their valence shell.

Formal Charges

Formal charge is a bookkeeping tool that helps chemists keep track of the electrons. It is used to predict the most stable structure among possible Lewis structures. It's calculated based on the number of valence electrons an atom owns compared to the number it would have in a molecule or polyatomic ion.

The formula to calculate formal charge is: \[ \text{Formal charge} = (\text{Valence electrons}) - (\text{Non-bonding electrons}) - \frac{1}{2}(\text{Bonding electrons}) \] For the SeO2 molecule, both oxygen atoms and the selenium atom have a formal charge of zero, which means electrons have been distributed in the molecule in such a way that all atoms have the same number of electrons as they would have in most stable molecules or ions. However, in species like CO3^2-, where resonance structures are present, the negative charges rotate among the oxygen atoms, and each structure predicts a different distribution of formal charges.

Understanding formal charges is essential for predicting the most likely arrangement of atoms in a given molecule, especially when multiple valid structures can exist. It also provides a deeper insight into the distribution of electron density within a molecule.

Resonance Structures

Resonance structures come into play when a single Lewis structure is not sufficient to describe the electron distribution within a molecule or ion. These are not real structures but are hypothetical intermediates between real, possible structures. They help predict properties like shape, polarity, and reactivity.

Take CO3^2- for instance. This ion has three resonance structures that distribute the double bond and negative charges among the oxygen atoms differently in each structure. The real structure is a hybrid that stabilizes the electron distribution across the molecule, known as a resonance hybrid. Similarly, in the NO2^- ion, the double bond switches positions between the two oxygen atoms in different resonance structures. Resonance is crucial in understanding molecules and ions where electron delocalization contributes to stability, often leading to lower energy configurations and affecting reactivity and other chemical properties.