How do I determine which molecule of each pair has the largest dipole moment?

case one: CO and NO

case two: SF4 and CF4

I have to use their electronegativities, right?

For linear molecules, such as CO and NO, yes, use electronegativity.
For non-linear molecules, such as SF4 and CF4, EN PLUS shape. Sometimes the shape cancels the individual internal dipoles.

How do I determine if they cancel or not?

Determine the shape of SF4 and CF4. CF4, for rexample, is tetrahedral (therefore, symmetrical). Each C and F has an EN such that you would expect each C-F bond to be polar (and it is) but four C-F bonds in a tetrahedral arrangement is symmetrical and the molecule as a whole has no dipole moment. So you know the answer now for CF4.

Explain Bot Explain Bot answered
9 months ago

To determine if the dipoles cancel or not in a molecule, you need to consider its molecular geometry. In the case of SF4, it has a seesaw shape with four fluorine atoms bonded to a central sulfur atom. The sulfur atom is more electronegative than fluorine, so the bonds between them are polar with the fluorine atoms being slightly negative and the sulfur atom slightly positive.

However, the seesaw shape of SF4 means that the dipole moments of the individual S-F bonds do not cancel each other out. This asymmetry results in an overall dipole moment for the molecule, meaning SF4 does have a dipole moment.

On the other hand, CF4 has a tetrahedral shape with four fluorine atoms bonded to a central carbon atom. Each carbon-fluorine bond in CF4 is polar with carbon being slightly positive and fluorine slightly negative due to the difference in electronegativity. However, the tetrahedral arrangement of the bonds in CF4 leads to an overall cancellation of the dipole moments, resulting in a molecule with no net dipole moment.

Therefore, in this case, SF4 has a larger dipole moment compared to CF4.