Question

Explain why \(C O\) has a dipole moment but \(C O_{2}\) does not.

Short answer

CO has a dipole moment due to its polar bond, but CO₂ does not because its dipoles cancel out.

Step-by-step solution

Understanding Molecular Polarity

A molecule has a dipole moment when there is an uneven distribution of electron density, causing a separation of charge. This occurs when there are polar bonds in the molecule, and their dipoles do not cancel out.

Analyze the Polarity of CO

In carbon monoxide (CO), there is a polar bond between carbon and oxygen. Oxygen is more electronegative than carbon, resulting in a partial negative charge on oxygen and a partial positive charge on carbon. The molecule is linear with this single bond, so the dipole moment does not cancel, and thus, CO has a dipole moment.

Analyze the Polarity of CO₂

Carbon dioxide (CO₂) also has polar bonds between the carbon and each oxygen. However, CO₂ is a linear molecule with two equal but opposite CO bond dipoles. These dipoles are 180° apart and cancel each other out, resulting in no net dipole moment.

Conclusion

CO has a dipole moment because it is a linear molecule with a single polar bond. CO₂ does not have a dipole moment because its linear geometry causes the bond dipoles to cancel out.

Key concepts

Dipole Moment

A dipole moment exists in a molecule when there is an uneven distribution of electrons between atoms involved in a bond. This uneven distribution leads to a separation of electric charge within the molecule. Dipole moments occur when there are polar bonds, meaning one atom draws electrons towards itself more strongly due to higher electronegativity. This results in a partial negative charge on one end and a partial positive charge on the other.
Understanding dipole moments is important in molecular chemistry because they affect how molecules interact with electric fields and other molecules. A good indication that a molecule has a dipole moment is if the molecule has a non-symmetrical shape along with polar bonds. This non-symmetrical arrangement ensures that the dipole moments from different bonds do not cancel each other out, leading to a net dipole moment across the molecule.
Key points to consider:

• Dipole moments arise from polar bonds
• The structure of a molecule (linear, bent, etc.) plays a crucial role
• Non-canceling dipoles lead to an overall dipole moment

Carbon Monoxide

Carbon monoxide (CO) is a simple linear molecule with a single carbon atom bonded to an oxygen atom. The bond between carbon and oxygen is polar due to oxygen's greater electronegativity. This means oxygen attracts the shared electrons in the bond more strongly, causing a partial negative charge (\(-\)) on the oxygen and a partial positive charge (\(+\)) on the carbon.
Because CO has only one C-O bond, and the molecule itself is linear, there is no way for this dipole to cancel out. This results in CO having a net dipole moment, making it a polar molecule. It's important to note that the presence of a dipole moment affects the physical and chemical properties of CO, including its reactivity and interaction with external electric fields.
Highlights about CO:

• Linear structure
• Single polar C-O bond
• Resulting dipole moment makes it a polar molecule

Carbon Dioxide

Carbon dioxide (CO₂) also consists of carbon and oxygen atoms connected through C-O bonds. Importantly, it has a linear molecular geometry. Despite the fact that each C-O bond in CO₂ is polar, the molecule as a whole does not have a dipole moment.
This is because the two C-O bonds in CO₂ are arranged 180° apart from each other, leading them to be opposite but equal in magnitude. As a result, their dipoles cancel each other out completely. Therefore, there is no net separation of charge across the molecule, making CO₂ a nonpolar molecule.
Key aspects about CO₂:

• Linear geometry with two equal and opposite C-O bonds
• Polar bonds cancel each other
• Results in no dipole moment, making CO₂ nonpolar

Polar Bonds

Polar bonds occur when two atoms in a molecule have different electronegativities, causing electrons to be shared unevenly. This results in one atom carrying a partial negative charge and the other a partial positive charge. Polar bonds are key to creating molecular dipoles.
The degree of polarity in a bond is determined by the difference in electronegativity between the two atoms. Higher differences result in stronger dipoles. Such bonds significantly influence the overall polarity of the molecule, as seen with CO and CO₂.
Key considerations about polar bonds:

• Arise from differences in electronegativity
• Cause partial charges on atoms involved
• Important for determining molecular dipole moments