Question

Write a Lewis formula, including the resonance forms, for each of the following molecules or ions. (a) \(\mathrm{NO}_{2}^{-}\) (c) \(\mathrm{CO}_{2}\) (e) \(\mathrm{HNO}_{3}^{-}\) (b) \(\mathrm{SO}_{3}\) (d) \(\mathrm{CO}_{3}{ }^{2-}\)

Short answer

The Lewis structures include: \(\mathrm{NO}_2^{-}\) with two resonance forms, each has a double and a single bond with oxygen; \(\mathrm{CO}_2\) with no resonance has double bonds with each oxygen; \(\mathrm{HNO}_3^{-}\) with multiple resonance forms with an oxygens double bond and two oxygens single bond, along with an hydrogen fused with Nitrogen; \(\mathrm{SO}_3\) has three resonance structures with single bonds with three Oxygen atoms; \(\mathrm{CO}_3^{2-}\) has three resonance forms each having one double and two single bonds with oxygen.

Step-by-step solution

Determine the Lewis structure for \(\mathrm{NO}_{2}^{-}\)

To draw the Lewis structure, we first add up the total number of valence electrons. Nitrogen has 5 valence electrons and each Oxygen has 6 valence electrons. The negative charge indicates an extra electron, so our total is 18 electrons. Nitrogen is the central atom. Now, if we place a double bond between Nitrogen and one Oxygen, and a single bond with the other Oxygen, we will use up (2*5) 10 electrons, leaving eight electrons which can be placed on the oxygens as lone pairs. This is one possible Lewis structure for \(\mathrm{NO}_{2}^{-}\). However, resonance can occur since the second oxygen atom also can form a double bond with Nitrogen. Thus, two resonance forms exist for \(\mathrm{NO}_{2}^{-}\).

Determine the Lewis structure for \(\mathrm{CO}_{2}\)

Carbon is the central atom surrounded by two Oxygen atoms. Carbon has 4 valence electrons and each Oxygen has 6 valence electrons. We thus have 16 electrons in total. Carbon will form double bonds with each Oxygen, using up (2*4) 8 electrons. The remaining 8 electrons will be placed on the oxygens as lone pairs. Hence, the Lewis structure for \(\mathrm{CO}_{2}\) involves no resonance as all atoms have full octets.

Determine the Lewis structure for \(\mathrm{HNO}_{3}^{-}\)

Nitrogen is the central atom, surrounded by one Hydrogen and three Oxygen atoms. We observe that \(\mathrm{HNO}_{3}^{-}\) would have a total of 24 valence electrons. Nitrogen will form a single bond with Hydrogen and two single bonds with two Oxygen atoms, while forming a double bond with the third Oxygen atom. This uses up (2*8) 16 electrons, leaving 8 electrons which will be placed on the Oxygen atoms as lone pairs. Here, resonance can also take place as the double bond with Oxygen can move about through the molecule. Therefore, \(\mathrm{HNO}_3^{-}\) has multiple Lewis structures.

Determine the Lewis structure for \(\mathrm{SO}_{3}\)

Sulphur is at the center, surrounded by three Oxygen atoms. \(\mathrm{SO}_3\) has a total of 24 valence electrons. Sulphur forms a double bond with each Oxygen atom, using 18 electrons (2*9). The remaining 6 electrons are placed as lone pairs on the Oxygen atoms. Since each Oxygen atom can form double bonds with Sulphur, it can form resonance structures.

Determine the Lewis structure for \(\mathrm{CO}_3^{2-}\)

Carbon is the central atom, surrounded by three Oxygen atoms. There are in total 24 valence electrons for \(\mathrm{CO}_3^{2-}\). Carbon will form a single bond with two Oxygen atoms and a double bond with one Oxygen atom, using up 16 electrons. Eight remaining electrons will be placed on the Oxygen atoms as lone pairs. As the double bond can move between the three Oxygen atoms, the molecule exhibits resonance, and has several Lewis structures.

Key concepts

Valence Electrons

Valence electrons are the electrons that reside in the outermost shell of an atom and are crucial in determining an atom's chemical properties. These electrons play a significant role in chemical bonding, as they are the ones that get involved in the formation of bonds between atoms.

When constructing Lewis structures, identifying the number of valence electrons is the first and foremost step. For instance, in the textbook exercise provided, the calculation of the total valence electrons is a critical step for determining the correct Lewis structure. Nitrogen has 5 valence electrons, oxygen has 6, carbon has 4, and sulfur has 6. Moreover, when charged species like \(\mathrm{NO}_{2}^{-}\) are considered, it is essential to add or subtract electrons based on the charge to get the correct count of valence electrons available for bonding.

Understanding the concept of valence electrons allows students to grasp why certain molecules form the number and type of bonds that they do. This lays the foundation for predicting the behavior of atoms in chemical reactions and is integral for mastering the Lewis structure representation of molecules.

Resonance Forms

Resonance forms are multiple Lewis structures that collectively represent the electronic structure of a molecule. These forms are a way to visualize the delocalization of electrons within a molecule, which cannot be represented by a single Lewis structure. Essentially, resonance forms are like individual frames of a movie that, when combined, give a full picture of the electron distribution in a molecule.

In the exercise examples, species like \(\mathrm{NO}_{2}^{-}\) and \(\mathrm{CO}_3^{2-}\) show resonance. This is because their double bonds can be placed in different positions without changing the arrangement of the atoms. It's important to note that resonance structures are not real, alternating states, but rather a tool used to describe a molecule where the actual electronic structure is a hybrid of the possible resonance forms.

To improve comprehension of this concept, students should practice drawing all possible resonance structures for a given molecule and understand that the actual distribution of electrons is an average of these structures. Knowing when and why resonance occurs strengthens a student's ability to predict molecular behavior and stability.

Chemical Bonding

Chemical bonding is the force that holds atoms together in molecules. The Lewis structures in the given exercise illustrate how atoms bond to achieve stable electron configurations, often resembling that of noble gases. Bonds can be single, double, or triple, with single bonds sharing two electrons, double bonds four, and triple bonds six.

For example, in the Lewis structure for \(\mathrm{CO}_{2}\), carbon forms double bonds with each oxygen atom. This indicates that carbon prefers to share four electrons with each oxygen to fill its valence shell, aligning with the octet rule which states that atoms tend to surround themselves with eight electrons to achieve a noble gas configuration.

Understanding bonding patterns, including the octet rule and exceptions to it, allows students to predict how atoms might combine. Bonding theories not only explain structural arrangements in molecular geometry but also allow insights into the properties and reactivity of substances, making it an essential concept in chemistry education.