Problem 22 - Chemistry and Chemical Reactivity

Question

The following molecules or ions all have three oxygen atoms attached to a central atom. Draw a Lewis structure for each one and then describe the electron-pair geometry and the molecular geometry around the central atom. Comment on similarities and differences in the series. (a) \(\mathrm{CO}_{3}^{2-}\) (b) \(\mathrm{NO}_{3}^{-}\) (c) \(\mathrm{SO}_{3}^{2-}\) (d) \(\mathrm{ClO}_{3}^{-}\)

Step-by-Step Solution

Answer

All molecules have three oxygens, but \(\mathrm{CO}_{3}^{2-}\) and \(\mathrm{NO}_{3}^{-}\) are trigonal planar, while \(\mathrm{SO}_{3}^{2-}\) and \(\mathrm{ClO}_{3}^{-}\) are trigonal pyramidal due to lone pairs.

Step 1: Drawing Lewis Structure for \(\mathrm{CO}_{3}^{2-}\)

To draw the Lewis structure for the carbonate ion \(\mathrm{CO}_{3}^{2-}\):- Carbon is the central atom. Attach three oxygen atoms to carbon.- Carbon has 4 valence electrons, each oxygen has 6, and there are 2 extra electrons due to the charge, yielding \(4 + 3\times6 + 2 = 24\) valence electrons.- Form single bonds between carbon and each oxygen, using 6 electrons.- Distribute the remaining 18 electrons among the oxygen atoms as lone pairs to complete their octet.- Assign formal charges to keep as low as possible: one oxygen forms a double bond with carbon (no charge on that oxygen and carbon), each of the other two oxygen atoms have a single bond (negative formal charge) to maintain the correct ionic charge.

Step 2: Determine Geometry for \(\mathrm{CO}_{3}^{2-}\)

In the \(\mathrm{CO}_{3}^{2-}\) structure, count the number of electron pairs around the central carbon atom:- There are 3 bonds (one of them is a double bond acting as one electron domain).- This yields 3 electron domains, which leads to a trigonal planar electron-pair geometry.- As all bonds are equivalent, the molecular geometry is also trigonal planar.

Step 3: Drawing Lewis Structure for \(\mathrm{NO}_{3}^{-}\)

To draw the Lewis structure for the nitrate ion \(\mathrm{NO}_{3}^{-}\):- Nitrogen is the central atom. Attach three oxygen atoms to nitrogen.- Nitrogen has 5 valence electrons, each oxygen has 6, and one extra for the charge, totaling \(5 + 3\times6 + 1 = 24\) valence electrons.- Form single bonds from nitrogen to each oxygen using 6 electrons.- Distribute the remaining 18 electrons as lone pairs to complete oxygen's octet.- Minimize formal charges by making one N=O double bond. The nitrogen and one oxygen have no formal charge, while the other two oxygens have a formal charge of -1 (each).

Step 4: Determine Geometry for \(\mathrm{NO}_{3}^{-}\)

In the \(\mathrm{NO}_{3}^{-}\) structure, determine the geometry:- Count 3 bonds around the central nitrogen atom (including the N=O double bond counted as a single electron domain).- This also results in 3 electron domains, with a trigonal planar electron-pair geometry.- Due to symmetrical bonding, the molecular geometry is trigonal planar as well.

Step 5: Drawing Lewis Structure for \(\mathrm{SO}_{3}^{2-}\)

To draw the Lewis structure for the sulfite ion \(\mathrm{SO}_{3}^{2-}\):- Sulfur is the central atom. Attach the three oxygen atoms.- Sulfur has 6 valence electrons, plus 3×6 from oxygens, and 2 more for charge, totaling \(6 + 3\times6 + 2 = 26\) valence electrons.- Form single bonds between sulfur and each oxygen with 6 electrons.- Use the remaining 20 electrons to complete the octets of each oxygen atom as lone pairs.- Minimal formal charges are managed with one S=O double bond, balancing charges between sulfur and one oxygen.

Step 6: Determine Geometry for \(\mathrm{SO}_{3}^{2-}\)

Count the domains around sulfur in the \(\mathrm{SO}_{3}^{2-}\):- Includes 3 bonds (2 singles, 1 double) from sulfur.- Leads to 4 electron domains due to a lone pair and a trigonal pyramidal electron-pair geometry.- The molecular geometry is trigonal pyramidal as a result of the lone electron pair on sulfur.

Step 7: Drawing Lewis Structure for \(\mathrm{ClO}_{3}^{-}\)

To achieve the Lewis structure for the chlorate ion \(\mathrm{ClO}_{3}^{-}\):- Chlorine is the central atom and attaches to three oxygen atoms.- Chlorine has 7 valence electrons, each oxygen has 6, and 1 additional electron from the charge totals \(7 + 3\times6 + 1 = 26\).- Use 6 electrons for °O-Cl bonds.- Place the remaining 20 electrons around oxygen to meet octet.- Create one Cl=O double bond to maintain low formal charges.

Step 8: Determine Geometry for \(\mathrm{ClO}_{3}^{-}\)

In \(\mathrm{ClO}_{3}^{-}\), calculate the electron domains around chlorine:- 3 oxygen bonds (single and double counted once) plus a lone pair.- Results in 4 domains, corresponding to a trigonal pyramidal electron-pair geometry.- The molecular geometry, just like \(\mathrm{SO}_{3}^{2-}\), is trigonal pyramidal.

Step 9: Compare Structures and Geometries

Compare the structures and geometries of all ions:- All ions have 3 oxygen atoms attached to a central atom.- \(\mathrm{CO}_{3}^{2-}\) and \(\mathrm{NO}_{3}^{-}\) are symmetrical with trigonal planar molecular geometry due to no lone pairs on central atoms.- \(\mathrm{SO}_{3}^{2-}\) and \(\mathrm{ClO}_{3}^{-}\) are trigonal pyramidal because of one lone pair on the central atom, differentiating their geometries from the trigonal planar.

Key Concepts

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. Understanding this concept aids in predicting the shape and function of molecules. For example, in the exercises provided, the carbonate ion (\(\mathrm{CO}_{3}^{2-}\)) and nitrate ion (\(\mathrm{NO}_{3}^{-}\)) both exhibit trigonal planar molecular geometry. This shape indicates that all atoms are in the same plane, around a central atom, and spaced at 120-degree angles.

In contrast, the sulfite ion (\(\mathrm{SO}_{3}^{2-}\)) and chlorate ion (\(\mathrm{ClO}_{3}^{-}\)) have trigonal pyramidal molecular geometries. This occurs because of the presence of lone pairs on the central atoms, which affect the overall shape. In these cases, the atoms form a pyramid shape with a central atom at the apex and a lone pair on top.

It's crucial to consider factors like lone pairs and bonding pairs when predicting molecular geometry. These shapes not only influence the molecule's appearance but also its chemical reactivity and interaction with other molecules.

Electron-Pair Geometry

Electron-pair geometry takes into account both bonding and non-bonding (lone) electron pairs to describe the 3D arrangement of electrons around the central atom. This is a critical concept that predicts the spatial orientation affecting molecular shapes.

In the case of \(\mathrm{CO}_{3}^{2-}\) and \(\mathrm{NO}_{3}^{-}\), the electron-pair geometry is trigonal planar because they involve three bonding pairs without any lone pairs on the central atom. Every electronic domain is mutually equidistant, forming a flat triangular shape.

For \(\mathrm{SO}_{3}^{2-}\) and \(\mathrm{ClO}_{3}^{-}\), the presence of one lone pair results in trigonal pyramidal geometry, deviating from the planar form. Here, the electron domains including the lone pair are geometrically aligned in a way that forms a three-sided pyramid with the central atom at the peak.

• Each identical atom arrangement leads to varying shapes based on lone pair presence.
• Lone pairs exert more repulsion than bonding pairs, resulting in geometry adjustments.

Formal Charge

Formal charge is an essential concept used to determine the most stable Lewis structure. It is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms. To calculate the formal charge, use the formula:
\[\text{Formal Charge} = \text{Valence electrons} - (\text{Lone pair electrons} + \frac{1}{2}\times \text{Bonding electrons})\]

For instance, in the carbonate ion \(\mathrm{CO}_{3}^{2-}\), one of the oxygen atoms forms a double bond with carbon. This keeps the formal charge of oxygen and carbon at zero, while the other two oxygens, which are single-bonded, carry a charge of -1 each to match the overall ion charge.

Proper formal charge allocation ensures minimized charges, contributing to the stability and accuracy of Lewis structures. In every molecular example provided, reducing formal charges to the lowest possible value has been an integral part of elucidating the structure correctly.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a pivotal role in chemical bonding and the formation of molecules. They are the electrons primarily involved in forming bonds.

In drawing Lewis structures, knowing the exact count of valence electrons is the starting point. The central atoms in each of the molecules considered—carbon in \(\mathrm{CO}_{3}^{2-}\), nitrogen in \(\mathrm{NO}_{3}^{-}\), sulfur in \(\mathrm{SO}_{3}^{2-}\), and chlorine in \(\mathrm{ClO}_{3}^{-}\)—each have a specific number of valence electrons. For example:

• Carbon typically has 4 valence electrons
• Nitrogen possesses 5 valence electrons
• Sulfur and chlorine both have 6 and 7 respectively

Understanding the distribution and count of valence electrons allows chemists to predict how atoms will bond, the nature of molecular structures, and the development of different molecular geometries.