Question

Apply VSEPR theory to explain why ${\mathrm{CO}}_2$ is a nonpolar molecule even though it has two polar bonds.

Short answer

${\mathrm{CO}}_2$ is nonpolar because its linear shape allows the polar $\mathrm{C}=\mathrm{O}$ bond dipoles to cancel out.

Step-by-step solution

Identify the Molecular Formula

Begin by recognizing the molecular formula of carbon dioxide, which is ${\mathrm{CO}}_2$. It consists of one carbon ($\mathrm{C}$) atom and two oxygen ($\mathrm{O}$) atoms.

Determine the Number of Valence Electrons

Calculate the total number of valence electrons in ${\mathrm{CO}}_2$. Carbon has 4 valence electrons and each oxygen has 6, resulting in a total of $4+(6\times 2)=16$ valence electrons.

Draw the Lewis Structure

Create the Lewis structure for ${\mathrm{CO}}_2$. Place the carbon atom in the center and attach the oxygen atoms on either side. Share three pairs of electrons between carbon and each oxygen to form double bonds, using all 16 valence electrons.

Apply VSEPR Theory

Use the VSEPR (Valence Shell Electron Pair Repulsion) theory to predict the shape. With two regions of electron density around the carbon atom (due to the two double bonds), ${\mathrm{CO}}_2$ adopts a linear geometry to minimize electron pair repulsion.

Analyze Polar Bonds

Examine the individual $\mathrm{C}=\mathrm{O}$ bonds. Each $\mathrm{C}=\mathrm{O}$ bond is polar due to the difference in electronegativity between carbon and oxygen.

Evaluate Molecular Polarity

Determine the overall polarity of the molecule. Despite having polar bonds, the linear shape of ${\mathrm{CO}}_2$ means that the dipole moments of the $\mathrm{C}=\mathrm{O}$ bonds cancel out, resulting in no overall dipole moment.

Key concepts

Molecular Polarity

Molecular polarity is a concept that determines whether a molecule is polar or nonpolar. It arises from the distribution of electron density within a molecule. If a molecule has an uneven distribution of electrons, making one part more negative than the other, it is polar. This usually happens in molecules where there are significant differences in electronegativity between the bonded atoms. For example, each carbon-oxygen bond in carbon dioxide ${\mathrm{CO}}_2$ is polar due to the higher electronegativity of oxygen compared to carbon.

However, the ultimate shape of the molecule plays a crucial role in determining the overall polarity. In ${\mathrm{CO}}_2$, the linear arrangement of the atoms causes the dipoles from each polar $\mathrm{C}=\mathrm{O}$ bond to align in opposite directions. These opposing dipoles cancel each other out, leading to a nonpolar molecule overall.

• Polarity is influenced by both the individual bond dipoles and the shape of the molecule.
• A molecule can have polar bonds but still be nonpolar if the molecular geometry allows for cancellation of the dipoles.

Lewis Structure

A Lewis structure is a simple way to show the bonding between atoms in a molecule. It highlights the atoms’ valence electrons, which participate in bond formation. For ${\mathrm{CO}}_2$, constructing the Lewis structure involves organizing its 16 valence electrons (4 from carbon and 6 from each oxygen) to show how the atoms share electrons to form bonds.

Carbon is placed in the center in the ${\mathrm{CO}}_2$ Lewis structure, with each oxygen atom positioned on either side. Each pair of carbon and oxygen atoms shares two pairs of electrons, forming double bonds. This use of 16 valence electrons ensures that each atom achieves a full valence shell in line with the octet rule.

• Lewis structures help visualize molecules’ electron arrangements and predict their geometry.
• The representation shows how valence electrons are distributed in shared or lone pair configurations.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are key in forming chemical bonds. They determine an atom's chemical properties and its ability to bond with other atoms. For a molecule like ${\mathrm{CO}}_2$, knowing the number of valence electrons is crucial.

Carbon has 4 valence electrons, while each oxygen has 6. This totals to 16 valence electrons for ${\mathrm{CO}}_2$. These electrons are shared to form stable chemical bonds, adhering to the octet rule, which suggests atoms tend to form bonds until they have eight electrons in their valence shell. This rule helps predict the structure and reactivity of the molecule.

• Counting valence electrons is the first step in creating Lewis structures and understanding molecular geometry.
• Valence electrons influence how atoms connect to form molecules.