Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: Which one of the following pairs of molecules will have permanent dipole moments for both members?

Welcome to a fascinating journey into the microscopic world of molecules! Have you ever wondered why water forms spherical droplets, or why oil and water refuse to mix? The secret lies in a concept called molecular polarity, specifically measured by something known as the dipole moment.

Imagine a microscopic game of tug-of-war. When two different atoms form a covalent bond, they rarely share the electrons equally. The more electronegative atom pulls the electron cloud closer to itself, creating a partial negative charge, while leaving the other atom with a partial positive charge. This separation of charge creates a bond dipole, which is a vector quantity—meaning it has both magnitude and direction.

However, having polar bonds does not automatically make the entire molecule polar. The overall polarity of a molecule depends entirely on its 3D geometry. If the molecule is perfectly symmetrical, the individual bond dipoles can cancel each other out, just like two equally strong teams pulling in opposite directions in our tug-of-war. Let's apply this logic to the molecules in our problem.

Silicon Tetrafluoride (${\text{SiF}}_4$)

Let's start by examining Silicon tetrafluoride (${\text{SiF}}_4$). Silicon sits at the center, bonded to four highly electronegative fluorine atoms. Because fluorine is much more electronegative than silicon, each $\text{Si}-\text{F}$ bond is highly polar, with the dipole vector pointing towards the fluorine atom.

But what about the overall shape? According to VSEPR theory, the four bonding pairs of electrons repel each other equally, arranging themselves as far apart as possible. This results in a perfect regular tetrahedral geometry with bond angles of $109.5^{\circ }$. Because of this flawless symmetry, the four bond dipole vectors pull equally in all directions in 3D space. When you add these vectors together, they perfectly cancel each other out. Therefore, the net dipole moment is zero ($\mu =0$). ${\text{SiF}}_4$ is a non-polar molecule.

Carbon Dioxide (${\text{CO}}_2$)

Next up is Carbon dioxide (${\text{CO}}_2$). Here, a central carbon atom forms double bonds with two oxygen atoms. Oxygen is more electronegative than carbon, so the $\text{C}=\text{O}$ bonds are polar, with the dipoles pointing towards the oxygen atoms.

However, the central carbon atom has no lone pairs of electrons. The two regions of electron density (the double bonds) repel each other to the maximum extent, resulting in a perfectly linear geometry with a bond angle of $180^{\circ }$. The two bond dipoles are equal in magnitude but point in exactly opposite directions. It is a perfect tie! They cancel each other out completely, leaving ${\text{CO}}_2$ with a net dipole moment of zero ($\mu =0$).

Nitrogen Dioxide (${\text{NO}}_2$)

Now, let's look at Nitrogen dioxide (${\text{NO}}_2$). This molecule is a bit of a rebel. Nitrogen is bonded to two oxygen atoms, but it also holds onto a single, unpaired electron. This makes ${\text{NO}}_2$ an odd-electron molecule, or a radical.

This single electron occupies space and exerts a repulsive force on the two nitrogen-oxygen bonds, pushing them downward. This breaks the linear symmetry and forces the molecule into a bent or angular geometry. Because the molecule is bent, the individual bond dipoles no longer point in exactly opposite directions. They cannot cancel each other out. Instead, they add up to create a net permanent dipole moment ($\mu eq 0$). ${\text{NO}}_2$ is a polar molecule.

Ozone (${\text{O}}_3$)

Finally, let's examine Ozone (${\text{O}}_3$). You might think that a molecule made entirely of the same element must be non-polar. But ozone surprises us! The central oxygen atom is bonded to two terminal oxygen atoms and also possesses a full lone pair of electrons.

Just like the unpaired electron in ${\text{NO}}_2$, this lone pair strongly repels the bonding pairs of electrons. This repulsion forces the ozone molecule into a bent shape. Furthermore, due to resonance, the central oxygen carries a partial positive charge, while the terminal oxygens carry partial negative charges. Because of the bent geometry and this charge distribution, the individual dipoles do not cancel. They combine to form a net permanent dipole moment ($\mu eq 0$). Ozone is indeed a polar molecule.

Our mission was to find the pair where both members possess a permanent dipole moment.

We discovered that perfectly symmetrical molecules like ${\text{SiF}}_4$ (tetrahedral) and ${\text{CO}}_2$ (linear) have zero dipole moments due to vector cancellation. However, molecules with lone pairs or unpaired electrons on the central atom, like ${\text{NO}}_2$ and ${\text{O}}_3$, have bent geometries. This lack of symmetry prevents the bond dipoles from canceling, resulting in a permanent net dipole moment.

Therefore, the correct pair is ${\text{NO}}_2$ and ${\text{O}}_3$.