Question

Write a Lewis formula, including the resonance forms, for each of the following molecules or ions. (a) \(\mathrm{OCN}^{-}\) (c) \(\mathrm{ClO}_{2}{ }^{-}\) (e) \(\mathrm{H}_{2} \mathrm{CO}_{3}\) (b) \(\mathrm{N}_{3}{ }^{-}\) (d) \(\mathrm{PO}_{4}{ }^{3-}\)

Short answer

Lewis structures for each molecule or ion include: \(\mathrm{OCN}^{-}\) with carbon in the center with a single bond to oxygen and nitrogen. \(\mathrm{ClO}_{2}{ }^{-}\) with chlorine in the middle and single bonds with either oxygen atom. \(\mathrm{H}_{2} \mathrm{CO}_{3}\) with carbon in the center with a double bond to one oxygen and single bonds to other oxygen and hydrogen atoms. \(\mathrm{N}_{3}{ }^{-}\) with central nitrogen atom triple bonded to the two other nitrogen atoms. \(\mathrm{PO}_{4}{ }^{3-}\) with phosphorus in the center forming a double bond with one oxygen and single bonds with the remaining three oxygen atoms.

Step-by-step solution

Drawing Lewis structure for \(\mathrm{OCN}^{-}\)

Count the total number of valence electrons, which is 16. Place the least electronegative atom in the center (C). Now, form a single bond between O and C and between C and N. After forming bonds, distribute the remaining electrons as lone pairs. The left over electron gives it the negative charge.

Drawing Lewis structure for \(\mathrm{ClO}_{2}{ }^{-}\)

Total valence electrons are 19. Oxygens are more electronegative than chlorine, so put chlorine (Cl) in the middle. Each oxygen atom forms a single bond with the Cl atom. The remaining valence electrons are used to complete the octets of each atom. One more electron will be added since it carries a negative (-1) charge.

Drawing Lewis structure for \(\mathrm{H}_{2} \mathrm{CO}_{3}\)

Start with 24 valence electrons. Carbon (C) is the least electronegative, hence it goes in the center. Now, form a double bond between C and one O and a single bond with other O and two H atoms respectively. Distribute the remaining electrons as lone pairs.

Drawing Lewis structure for \(\mathrm{N}_{3}{ }^{-}\)

There are 18 valence electrons. The two outer N atoms form a triple bond with the central N atom. The remaining electron gives the negative charge.

Drawing Lewis structure for \(\mathrm{PO}_{4}{ }^{3-}\)

Total valence electrons are 32. P is the least electronegative atom so it's in the center forming a bond with each O. One of the Oxygen atoms forms a double bond with P to complete its octet. Distribute the remaining electrons as lone pairs. Add 3 more lone pairs due to the 3- charge.

Key concepts

Resonance Structures

Resonance structures are an important aspect of chemical bonding.
They provide a way to represent molecules that cannot be described with a single Lewis structure.
This often occurs in molecules or ions where electrons are delocalized, meaning they are spread out across multiple atoms rather than being localized between two atoms in a single bond.

Resonance structures show the different possible ways the electron pairs can be arranged in the molecule to highlight this delocalization.

• These structures have the same placement of atoms but different arrangements of electrons.
• The true structure of the molecule is a hybrid of all possible resonance forms.

The resonance structures for molecules like \( ext{OCN}^{-}\) show different possible configurations for its electrons.
This allows for greater stability and explains the molecule's actual state, which is an average or hybrid of these different forms.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding.
They determine how an atom interacts with others and dictates the type of bonds that can form.

For example, in the Lewis structures, these are the electrons involved in forming covalent bonds or remaining as lone pairs.

• In \( ext{OCN}^{-}\), the total number of valence electrons is 16; this helps in forming single bonds and lone pairs.
• Valence electrons are essential in achieving stable electron configurations, often by filling or emptying these outer electron shells.

Determining the number of valence electrons is the first step in drawing any Lewis structure.
It helps predict the bonding capacity and arrangements of atoms in the molecule or ion.

Chemical Bonding

Chemical bonding is the process that links atoms together to form molecules and compounds.
There are several types of chemical bonds, but covalent bonds often featured in Lewis structures involve the sharing of electrons between atoms.

Understanding the type of bond helps in constructing accurate molecular representations.

• In \( ext{N}_3^{-}\), a sequence of triple bonds helps stabilize the molecule with 18 valence electrons involved.
• The choice of forming single, double, or triple bonds often depends on achieving full octets for stability.

In addition to covalent bonds, one also often considers electron sharing in the context of the molecule's geometry and reactivity.
Different arrangements of electrons lead to different molecular shapes and properties.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms in a molecule.
It's based on the idea that electron pairs around a central atom will arrange themselves to minimize repulsion.

This concept helps understand physical and chemical properties such as polarity, reactivity, and phase of matter.

• It uses the Valence Shell Electron Pair Repulsion (VSEPR) theory to predict shapes around a central atom.
• For example, \( ext{PO}_4^{3-}\) might adopt a tetrahedral shape due to the arrangement of its 32 valence electrons around phosphorus.

By predicting the molecular geometry, you can infer potential interactions with other molecules.
This is crucial for understanding reactions and properties of substances in chemistry.