Question

Draw correct Lewis structures for ozone \(\left(\mathrm{O}_{3}\right)\), azide ion, \(\left(\mathrm{N}_{3}^{-}\right),\) and bicarbonate ion, \(\mathrm{HCO}_{3}^{-}\). Include lone pair electrons and formal charges, and use your General Chemistry textbook to review VSEPR theory, which will enable you to draw correct bond geometries.

Short answer

Answer: Ozone (O3) has a bent molecular geometry with formal charges of -1 on the outer oxygen atoms and +1 on the central oxygen atom. The azide ion (N3-) has a linear molecular geometry with a -1 formal charge on the central nitrogen atom and 0 formal charges on the terminal nitrogen atoms. The bicarbonate ion (HCO3-) has a trigonal planar molecular geometry with a -1 formal charge on the doubly bonded oxygen atom and 0 formal charges on the carbon and other oxygen atoms.

Step-by-step solution

1. Ozone (O3)

To draw the Lewis structure for ozone, first count the total number of valence electrons: 6 * 3 = 18 electrons. Place one oxygen atom in the center and connect the other two around it with single bonds. Distribute the remaining 12 electrons as lone pairs to ensure each oxygen atom has 8 electrons around it. The central oxygen has one lone pair and a single bond with each of the other two oxygen atoms. Next, add formal charges: -1 on the outer oxygen atoms, +1 on the central oxygen atom. Finally, apply VSEPR theory for bond geometry, which results in a bent molecular structure.

2. Azide ion (N3-)

To draw the Lewis structure for an azide ion, first count the total number of valence electrons: 5 * 3 + 1 = 16 electrons. Place one nitrogen in the center and connect the other two with single bonds. Distribute the remaining 10 electrons as follows: two lone pairs on the terminal nitrogen atoms, one lone pair and a negative charge on the central nitrogen atom. Next, add formal charges: -1 on the central nitrogen atom and 0 on the terminal nitrogen atoms. Applying VSEPR theory for bond geometry, we get a linear molecular structure for the azide ion.

3. Bicarbonate ion (HCO3-)

To draw the Lewis structure for a bicarbonate ion, first count the total number of valence electrons: 4 + 6 + 1 + 1 = 12 electrons. Place the carbon atom in the center and arrange one oxygen atom and the hydrogen atom around it. Using the formal charges, distribute the remaining electrons as follows: add single bonds between the carbon and the hydrogen atom and on one of the oxygen atoms, and place a double bond to the third oxygen atom. Add two lone pairs on the bonded oxygen atom and three lone pairs on the other two oxygen atoms. Next, add formal charges: -1 on the doubly bonded oxygen atom, 0 on the carbon and other oxygen atoms. Applying VSEPR theory for bond geometry, we get a trigonal planar molecular structure for the bicarbonate ion.

Key concepts

Understanding VSEPR Theory

VSEPR theory, or Valence Shell Electron Pair Repulsion theory, is a model used in chemistry to predict the geometry of individual molecules based on the number of electron pairs surrounding their central atoms. To put it simply, this theory suggests that electron pairs will arrange themselves as far apart as possible around a central atom to minimize repulsion, which directly influences the shape of a molecule.

The basic premise is that pairs of valence electrons, both bonding and nonbonding (lone pairs), repel each other. As a result, they will adopt an orientation that reduces this repulsion, leading to distinctive molecular geometries. For instance, if there are two electron pairs, they will be opposite from each other (linear configuration). As we add more electron pairs, the geometry changes to accommodate these pairs at maximum distances from each other, which explains the bent shape of ozone (O3), the linear structure of azide ion (N3-), and the trigonal planar shape of the bicarbonate ion (HCO3-).

Formal Charges Explained

Formal charges are a tool used by chemists to keep track of the charge distribution throughout a molecule. The formal charge is the hypothetical charge assigned to an atom within a molecule, assuming that electrons in chemical bonds are shared equally between atoms, regardless of relative electronegativity.

To calculate the formal charge on an atom, you can use the formula:
\[ FC = V - (N_b + \frac{1}{2}B_e) \]
where \( FC \) is the formal charge, \( V \) represents the number of valence electrons in the free atom, \( N_b \) is the number of non-bonding electrons, and \( B_e \) is the number of electrons in bonds (bonding electrons). It's essential to estimate the formal charge to predict the distribution of electrons in a molecule and thereby its stability, as seen in the solutions for ozone, azide, and bicarbonate ions.

Molecular Geometry essentials

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. It is crucial for determining various physical and chemical properties such as reactivity, polarity, color, phase of matter, and biological activity. While VSEPR theory provides a qualitative understanding of molecular shape, molecular geometry can be quantitatively assessed using bond angles and distances acquired through experimentation.

It's essential to differentiate between electronic geometry, which considers both bonding and lone pair electrons, and molecular geometry, which only looks at the arrangement of atoms (disregarding lone pairs). For example, in the case of the bicarbonate ion, its electronic geometry is tetrahedral, but with one vertex occupied by a lone pair, the molecular geometry is trigonal planar. Understanding these geometries aids in visualizing molecules and predicting how they will interact with other substances.