Question

\(\mathrm{FClO}_{2}\) and \(\mathrm{F}_{3} \mathrm{ClO}\) can both gain a fluoride ion to form stable anions. \(\mathrm{F}_{3} \mathrm{ClO}\) and \(\mathrm{F}_{3} \mathrm{ClO}_{2}\) will both lose a fluoride ion to form stable cations. Draw the Lewis structures and describe the hybrid orbitals used by chlorine in these ions.

Short answer

The Lewis structures for the given molecules and their corresponding ions are: \(\mathrm{FClO}_{2}\): ``` F=Cl–O || O ``` \(\mathrm{F}_{3} \mathrm{ClO}\): ``` O || F–Cl–F | F ``` \(\mathrm{F_{2}ClO_{2}}^{-}\): ``` F=Cl–O | O | F ``` \(\mathrm{F_{4}ClO}^{-}\): ``` O || F–Cl–F | F | F ``` \(\mathrm{ClO_{2}}^{+}\): ``` Cl–O || O ``` \(\mathrm{F_{2}ClO}^{+}\): ``` O || F–Cl–F ``` The hybrid orbitals used by chlorine in these ions are: - \(\mathrm{FClO}_{2}\): sp3 - \(\mathrm{F}_{3} \mathrm{ClO}\): sp3 - \(\mathrm{F_{2}ClO_{2}}^{-}\): sp3d - \(\mathrm{F_{4}ClO}^{-}\): sp3d - \(\mathrm{ClO_{2}}^{+}\): sp2 - \(\mathrm{F_{2}ClO}^{+}\): sp2

Step-by-step solution

Draw the Lewis structures for the given molecules

We will start by drawing the Lewis structures for the given molecules (\(\mathrm{FClO}_{2}\) and \(\mathrm{F}_{3} \mathrm{ClO}\)). Remember that each atom should have an octet of electrons (except hydrogen which can only have 2). For \(\mathrm{FClO}_{2}\): Chlorine is the central atom having a valence of 7, and it is bonded with two oxygen atoms having a valence of 6 each and one fluorine atom having a valence of 7. Chlorine shares one electron with each oxygen and fluorine atom, and the remaining electrons can form double bonds (resulting in six electrons on each oxygen and fluorine atom). Lewis structure of \(\mathrm{FClO}_{2}\): ``` F=Cl–O || O ``` For \(\mathrm{F}_{3} \mathrm{ClO}\): Chlorine is the central atom having a valence of 7, and it is connected to three fluorine atoms having a valence of 7 each and one oxygen atom having a valence of 6. Chlorine shares one electron with each fluorine atom and the remaining two electrons can form a double bond with oxygen. Lewis structure of \(\mathrm{F}_{3} \mathrm{ClO}\): ``` O || F–Cl–F | F ```

Draw the Lewis structures for the anions formed by gaining a fluoride ion

When each molecule gains a fluoride ion, they become negatively charged and form stable anions. For \(\mathrm{FClO}_{2}\), when a fluoride ion is gained, we get the ion \(\mathrm{F_{2}ClO_{2}}^{-}\). The Lewis structure changes by adding a fluoride ion in a single bond with the central chlorine atom. Lewis structure of \(\mathrm{F_{2}ClO_{2}}^{-}\): ``` F=Cl–O | O | F ``` For \(\mathrm{F}_{3} \mathrm{ClO}\), when a fluoride ion is gained, we get the ion \({\mathrm{F_{4}ClO}}^{-}\). The Lewis structure changes by adding a fluoride ion in a single bond with the central chlorine atom. Lewis structure of \(\mathrm{F_{4}ClO}^{-}\): ``` O || F–Cl–F | F | F ```

Draw the Lewis structures for the cations formed by losing a fluoride ion

When each molecule loses a fluoride ion, they become positively charged and form stable cations. For \(\mathrm{FClO}_{2}\), when a fluoride atom is lost, we get the ion \(\mathrm{ClO_{2}}^{+}\). The Lewis structure changes with the removal of the fluorine atom and its bond with the central chlorine atom. Lewis structure of \(\mathrm{ClO_{2}}^{+}\): ``` Cl–O || O ``` For \(\mathrm{F}_{3} \mathrm{ClO}\), when a fluoride atom is lost, we get the ion \(\mathrm{F_{2}ClO}^{+}\). The Lewis structure changes with the removal of one fluorine atom and its bond with the central chlorine atom. Lewis structure of \(\mathrm{F_{2}ClO}^{+}\): ``` O || F–Cl–F ```

Identify the hybrid orbitals used by the chlorine atoms in these ions

Now, we will identify the hybrid orbitals used by the chlorine atoms in these ions. For \(\mathrm{FClO}_{2}\): Chlorine is surrounded by four electron domains (1 F-atom, 2 O-atoms, and 1 lone pair). Therefore, it will have sp3 hybridization. For \(\mathrm{F}_{3} \mathrm{ClO}\): Chlorine is surrounded by four electron domains (3 F-atoms, and 1 O-atom). Therefore, it will have sp3 hybridization. For \(\mathrm{F_{2}ClO_{2}}^{-}\): Chlorine is surrounded by five electron domains (2 F-atoms, 2 O-atoms, and 1 lone pair). Therefore, it will have sp3d hybridization. For \(\mathrm{F_{4}ClO}^{-}\): Chlorine is surrounded by five electron domains (4 F-atoms, and 1 O-atom). Therefore, it will have sp3d hybridization. For \(\mathrm{ClO_{2}}^{+}\): Chlorine is surrounded by three electron domains (2 O-atoms and 1 lone pair). Therefore, it will have sp2 hybridization. For \(\mathrm{F_{2}ClO}^{+}\): Chlorine is surrounded by three electron domains (2 F-atoms and 1 O-atom). Therefore, it will have sp2 hybridization.

Key concepts

Hybridization

When atoms form bonds, their valence orbitals often mix to create new hybrid orbitals, a process known as hybridization. These orbitals provide the shared space for bonding and ensure the stability of the molecule. The type of hybridization is directly related to the number of electron domains around the central atom.
Let's look at a few key scenarios.

• sp³ hybridization: When there are four electron domains, such as in \({\mathrm{FClO}_{2}}\) and \({\mathrm{F}_{3} \mathrm{ClO}}\), the chlorine atom undergoes sp³ hybridization. This means one s and three p orbitals mix, forming four equivalent orbitals that shape the tetrahedral arrangement essential for accommodating bonds and lone pairs.
• sp³d hybridization: With five electron domains, like in \({\mathrm{F_{2}ClO_{2}}^{-}}\) and \({\mathrm{F_{4}ClO}^{-}}\), the chlorine atom mixes one s, three p, and one d orbital, resulting in a trigonal bipyramidal arrangement that allows electrons to spread efficiently.
• sp² hybridization: In molecules such as \({\mathrm{ClO_{2}}^{+}}\) and \({\mathrm{F_{2}ClO}^{+}}\), containing three electron domains, one s and two p orbitals hybridize to form three sp² orbitals. They are arranged in a trigonal planar geometry, facilitating flexible bonding angles.

Electron Domains

Electron domains are regions where electrons are most likely to be found. They include bonds (single, double, or triple) and lone pairs around a central atom. Understanding electron domains helps predict both the geometry and hybridization of molecules.
Consider each domain like a force field that dictates molecular shape:

• In \({\mathrm{FClO}_{2}}\), the chlorine atom is surrounded by four domains, including bonds to one fluorine and two oxygens plus a lone pair. This arrangement leads to an overall tetrahedral electron domain geometry, determining its sp³ hybridization.
• \({\mathrm{F_{3}ClO}}\) also has four electron domains comprising three fluorines and one oxygen, promoting tetrahedral geometry.
• For \({\mathrm{F_{2}ClO_{2}}^{-}}\) and \({\mathrm{F_{4}ClO}^{-}}\), five electron domains form, moving hybridization to sp³d, which supports a trigonal bipyramidal geometry due to additional atomic interactions.
• With \({\mathrm{ClO_{2}}^{+}}\) and \({\mathrm{F_{2}ClO}^{+}}\), three domains result in a planar shape, simplified to sp² hybridization and supporting flexibility in molecular shape dynamics.

Central Atom

The central atom in a molecule is the anchor point, typically possessing the most available bonding sites. In both \({\mathrm{FClO}_{2}}\) and \({\mathrm{F}_{3} \mathrm{ClO}}\), chlorine acts as the central atom. This decision arises mainly due to chlorine's valence electron availability, which allows it to form multiple bonds necessary to maintain octets in surrounding atoms.
Why is chlorine the central atom?

• Chlorine's capacity to form bonds with multiple other atoms stems from its seven valence electrons, making it ideal for drawing electrons from surrounding fluorine and oxygen to complete compound structures.
• Another factor is the electron negativity. Chlorine can facilitate the distribution of electrons needed for stable configurations due to its moderate size and ability to balance the electron density around it.

These properties make chlorine an effective central atom in both compounds, ensuring optimal electron domain arrangement and overall molecular stability.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms relative to each other in a molecule. This configuration is important because it influences physical properties such as polarity and reactivity.
Here's a look at how geometry is determined and why it matters:

• In \({\mathrm{FClO}_{2}}\), the tetrahedral electron domain geometry rearranges to a bent molecular geometry due to the lone pair's repulsion, affecting angles and electron distribution.
• For \({\mathrm{F_{3}ClO}}\), while the electron domain is originally tetrahedral, the actual molecular shape ends up being seesaw-like due to occupied lone pairs and bond repulsions, causing deviations from perfect tetrahedral angles.
• In the ions \({\mathrm{F_{2}ClO_{2}}^{-}}\) and \({\mathrm{F_{4}ClO}^{-}}\), trigonal bipyramidal geometry is observed due to five electron domains. The shape balances minimizing repulsion and maximizing spacing among domains around the central chlorine.
• Molecules like \({\mathrm{ClO_{2}}^{+}}\) and \({\mathrm{F_{2}ClO}^{+}}\) demonstrate linear molecular geometry, derived from a planar electron domain arrangement. Their shape is crucial for predicting molecular interactions and potential reactivity patterns.