Question

Draw the Lewis structure for carbonyl fluoride, \(\mathrm{COF}_{2}\) What are its electron-pair geometry and molecular geometry? What is the hybridization of the carbon atom? What orbitals overlap to form the \(\sigma\) and \(\pi\) bonds between carbon and oxygen?

Short answer

Carbonyl fluoride has a trigonal planar geometry, \(sp^2\) hybridization in carbon, and uses \(sp^2\) and \(p\) orbitals for \(\sigma\) bonds, and \(p\) orbitals for \(\pi\) bonds.

Step-by-step solution

Determine the total number of valence electrons

First, count the total number of valence electrons available. Carbon (C) has 4 valence electrons, each oxygen (O) and each fluorine (F) atom has 6 and 7 valence electrons, respectively. Given the molecular formula is \(\mathrm{COF}_{2}\), sum up the valence electrons: \[4 \text{ (from C)} + 6 \text{ (from O)} + 7 \times 2 \text{ (from 2 F)} = 24\text{ valence electrons}.\]

Sketch a rough structure

Place the carbon atom in the center, as carbon typically forms the most bonds, and surround it with one oxygen and two fluorine atoms. Carbon will be the central atom with the other more electronegative atoms surrounding it.

Connect atoms with single bonds

Connect each surrounding atom to the central carbon atom using single bonds. This accounts for \(3\) bond pairs or \(6\) electrons (C-O and C-F bonds). This leaves us with \(24 - 6 = 18\) electrons for completing octets.

Complete outer atoms’ octets

Add remaining electrons to the outer atoms (Fluorine and Oxygen) to complete their octets. Each fluorine needs \(6\) more electrons, and oxygen needs \(4\) more. This uses up \(12 + 4 = 16\) of the remaining \(18\) electrons, leaving \(2\) electrons.

Use remaining electrons to form double bonds

Use the remaining electrons to form a double bond between carbon and oxygen (C=O), which completes the octet for both carbon and oxygen in the structure.

Determine geometry and hybridization

The central carbon atom is bonded to three atoms (one O and two F) with no lone pairs, which corresponds to a trigonal planar electron-pair geometry. The molecular geometry is also trigonal planar. The hybridization of the carbon atom is \(sp^2\).

Identify orbitals in bond formation

The \(\sigma\) bonds between the carbon atom and its neighbors are formed from \(sp^2\) hybrid orbitals of carbon overlapping with \(p\) orbitals of oxygen and fluorine. The \(\pi\) bond between carbon and oxygen is formed by the side-to-side overlap of unhybridized \(p\) orbitals in carbon and oxygen.

Key concepts

Chemical Bonding

Chemical bonding is fundamental in understanding how molecules are held together. In the case of carbonyl fluoride, or \(\mathrm{COF}_2\), we look at how the atoms within the molecule connect. Carbon typically takes the central position because it forms the most bonds. It connects with one oxygen and two fluorine atoms. These connections are chemical bonds formed by sharing electrons. - **Types of Bonds:** - **Single Bonds:** These are found between carbon and fluorine (C-F), where one pair of electrons is shared. - **Double Bonds:** The carbon and oxygen atoms form a double bond (C=O) involving two pairs of shared electrons. Understanding the number of shared electron pairs helps in understanding the Lewis structure of a molecule, which is a simple way to represent the elements that make up the compound and how they bond.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms in a molecule. For carbonyl fluoride, the molecular geometry is trigonal planar. This means that all the atoms bonded to carbon are in a single plane, forming a triangle shape. - **Trigonal Planar Characteristics:** - It involves three groups of electrons (three atoms bound to carbon). - The bond angles in a trigonal planar configuration are approximately \(120^\circ\). This geometry ensures that the molecule is symmetrical, which can affect properties such as polarity and reactivity.

Electron-Pair Geometry

Electron-pair geometry considers the arrangement of all electron groups around the central atom, including bonds and lone pairs, to determine shape. For \(\mathrm{COF}_2\), the electron-pair geometry is also trigonal planar. This matches the molecular geometry because the central carbon atom in \(\mathrm{COF}_2\) is surrounded by three electron groups—no lone pairs alter the basic shape here: one double bond to oxygen and two single bonds to fluorine. - **Important Points:** - Both molecular and electron-pair geometry are crucial for predicting how a molecule behaves. - Changes in electron pair arrangement can lead to changes in molecular geometry, but in \(\mathrm{COF}_2\), they align perfectly as trigonal planar.

Hybridization

Hybridization is the process by which atomic orbitals mix to form new orbitals, suited for bonding in molecules. In carbonyl fluoride, the carbon atom undergoes \(sp^2\) hybridization. - **Hybridization Steps:** - Carbon's one \(s\) orbital and two of its \(p\) orbitals mix to form three equivalent \(sp^2\) hybrid orbitals. - These \(sp^2\) orbitals form \(\sigma\) bonds with the \(p\) orbitals of the oxygen and fluorine atoms. Additionally, the carbon-oxygen double bond includes a \(\pi\) bond, formed from the side-to-side overlap of \(p\) orbitals from both atoms.By understanding hybridization, one gains insight into the molecular structure's stability and the bond strengths within the molecule.