Which of the following molecules is polar?

• A. H_2O
• B. CO_2
• C. CS_2
• D. CCl_4

Answer: (A)

Verified step by step guidance

1. Step 1: Understand that molecular polarity depends on both the polarity of individual bonds and the geometry of the molecule. A molecule is polar if it has a net dipole moment, meaning the bond dipoles do not cancel out.
2. Step 2: Analyze the molecular geometry of each molecule using the VSEPR theory. For example, \(\mathrm{H_2O}\) has a bent shape, \(\mathrm{CO_2}\) and \(\mathrm{CS_2}\) are linear, and \(\mathrm{CCl_4}\) is tetrahedral.
3. Step 3: Consider the electronegativity differences between the central atom and the surrounding atoms to determine if the bonds are polar. For instance, \(\mathrm{H_2O}\) has polar O-H bonds, \(\mathrm{CO_2}\) has polar C=O bonds, \(\mathrm{CS_2}\) has polar C=S bonds, and \(\mathrm{CCl_4}\) has polar C-Cl bonds.
4. Step 4: Evaluate whether the molecular geometry causes the bond dipoles to cancel out. In linear molecules like \(\mathrm{CO_2}\) and \(\mathrm{CS_2}\), the dipoles are equal and opposite, canceling each other. In tetrahedral \(\mathrm{CCl_4}\), the symmetrical arrangement also cancels dipoles.
5. Step 5: Conclude that \(\mathrm{H_2O}\) is polar because its bent shape causes the bond dipoles to add up to a net dipole moment, making the molecule polar.