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Chapter 10: Problem 3

Specify the electron-pair and molecular geometry for each of the following. Describe the hybrid orbital set used by the underlined atom in each molecule or ion. (a) \(\underline{\mathrm{BBr}}_{3}\) (b) \(\underline{\mathrm{CO}}_{2}\) (c) \(\underline{\mathrm{CH}}_{2} \mathrm{Cl}_{2}\) (d) \(\underline{\mathrm{CO}}_{3}^{2-}\)

Short Answer

Expert verified
BBr鈧�: Trigonal Planar, sp虏; CO鈧�: Linear, sp; CH鈧侰l鈧�: Tetrahedral, sp鲁; CO鈧兟测伝: Trigonal Planar, sp虏.

Step by step solution

01

Determine BBr鈧� Geometry and Hybridization

For BBr鈧�, Boron (B) is the central atom. Boron has three valence electrons and forms three sigma bonds with Bromine. There are no lone pairs on Boron. - **Electron-Pair Geometry**: Trigonal Planar - **Molecular Geometry**: Trigonal Planar - **Hybridization of B**: sp虏 (Boron uses three sp虏 hybrid orbitals to bond with three Bromine atoms)
02

Determine CO鈧� Geometry and Hybridization

For CO鈧�, Carbon (C) is the central atom. Carbon forms two double bonds with Oxygen atoms, and there are no lone pairs on Carbon. - **Electron-Pair Geometry**: Linear - **Molecular Geometry**: Linear - **Hybridization of C**: sp (Carbon uses two sp hybrid orbitals for the double bonds with each Oxygen)
03

Determine CH鈧侰l鈧� Geometry and Hybridization

For CH鈧侰l鈧�, Carbon (C) is the central atom. Carbon forms two sigma bonds with Hydrogen and two sigma bonds with Chlorine, with no lone pairs on Carbon. - **Electron-Pair Geometry**: Tetrahedral - **Molecular Geometry**: Tetrahedral - **Hybridization of C**: sp鲁 (Carbon uses four sp鲁 hybrid orbitals to form bonds with two Hydrogens and two Chlorines)
04

Determine CO鈧兟测伝 Geometry and Hybridization

For CO鈧兟测伝, Carbon (C) is the central atom. Carbon forms three sigma bonds and has a resonance structure with double bonds distributed among three Oxygen atoms. There are no lone pairs on Carbon. - **Electron-Pair Geometry**: Trigonal Planar - **Molecular Geometry**: Trigonal Planar - **Hybridization of C**: sp虏 (Carbon uses three sp虏 hybrid orbitals to bond with Oxygen atoms)

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Hybridization
The concept of hybridization is essential in understanding molecular geometry. It refers to the mixing of atomic orbitals in an atom to form a new set of hybrid orbitals. This occurs during the formation of chemical bonds to optimize the arrangement of electrons.
  • In BBr鈧�, boron undergoes sp虏 hybridization. Here, one s and two p orbitals mix, resulting in three equivalent sp虏 hybrid orbitals. These orbitals form sigma bonds with the bromine atoms.
  • In CO鈧�, carbon uses sp hybridization. One s orbital and one p orbital mix to form two sp hybrid orbitals. Each orbital forms a sigma bond with an oxygen atom, aligning linearly.
  • CH鈧侰l鈧� involves carbon undergoing sp鲁 hybridization. All four hybridized orbitals are used to bond with hydrogen and chlorine atoms, leading to a tetrahedral shape.
  • For CO鈧兟测伝, carbon is sp虏 hybridized. Similarly to BBr鈧�, this leads to three hybrid orbitals, but with resonance structures affecting electron distribution.
Understanding hybridization helps explain how different molecules achieve their specific shapes and bond types.
Electron-Pair Geometry
Electron-pair geometry considers the spatial arrangement of all electron pairs around a central atom, including both bonding and lone pairs. This helps in predicting the molecular structure.
  • For BBr鈧�, the electron-pair geometry is trigonal planar. This symmetrical arrangement comes from the three shared electron pairs and zero lone pairs around boron.
  • In CO鈧�, with only two regions of electron density, the geometry is linear, resulting from its two double bonds.
  • CH鈧侰l鈧� has a tetrahedral electron-pair geometry because the four areas of electron density鈥攂onds to two hydrogens and two chlorines鈥攕pread out evenly in three dimensions.
  • CO鈧兟测伝 also shares a trigonal planar electron-pair geometry, with three equivalent bonding regions around carbon reflecting shared and resonance-affected double bonds.
Understanding electron-pair geometry is crucial to predicting how molecules interact and react in a three-dimensional space.
Sigma Bonds
Sigma bonds (\(\sigma\) bonds) are the first and strongest type of covalent bonds formed via the head-on overlap of atomic orbitals. They exist in all single bonds and are necessary for determining a molecule's structure and stability.
  • In BBr鈧�, three sigma bonds form between boron and each bromine atom, securing a planar shape.
  • CO鈧� forms sigma bonds from sp hybridized orbitals of carbon bonded to oxygen atoms. The linear molecular shape features these bonds.
  • CH鈧侰l鈧� contains four sigma bonds, with each sp鲁 hybridized orbital of carbon forming a bond with hydrogen or chlorine.
  • For CO鈧兟测伝, three sigma bonds exist between carbon and oxygen atoms. The resonance structure alters bond characteristics but maintains sigma bond presence.
Sigma bonds are fundamental to the rigid framework in molecular structures.
Resonance Structure
Resonance structures are different ways of representing the electron distribution within a molecule that can't be captured by a single Lewis structure. They are critical for understanding the electron delocalization seen in many molecules.
  • In CO鈧兟测伝, resonance plays a crucial role. While drawing one structure suggests a single double bond and two single bonds, resonance structures indicate that all C-O bonds are equivalent by sharing electron density.
  • Unlike CO鈧兟测伝, BBr鈧�, CO鈧�, and CH鈧侰l鈧� do not exhibit resonance structures since they are adequately described by one traditional Lewis structure.
Recognizing the concept of resonance is pivotal for comprehending molecules' true electron arrangements and their resultant chemical properties.

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Most popular questions from this chapter

Consider the three fluorides \(\mathrm{BF}_{4}^{-}, \mathrm{SiF}_{4},\) and \(\mathrm{SF}_{4}\) (a) Identify a molecule that is isoelectronic with \(\mathrm{BF}_{4}^{-}\). (b) Are \(\mathrm{SiF}_{4}\) and \(\mathrm{SF}_{4}\) isoelectronic? (c) What is the hybridization of the central atom in each of these species?

Draw the Lewis structure for chloroform, CHCl \(_{3}\). What are its electron- pair and molecular geometries? What orbitals on \(\mathrm{C}, \mathrm{H},\) and $$\mathrm{Cl}$ overlap to form bonds involving these elements?

Draw the Lewis structure for CIF \(_{3}\). What are its electronpair and molecular geometries? What is the hybridization of the chlorine atom? What orbitals on Cl and F overlap to form bonds between these elements?

What is the hybridization of the carbon atom in phosgene, \(\mathrm{Cl}_{2}\) CO? Give a complete description of the \(\sigma\) and \(\pi\) bonding in this molecule.

Draw the Lewis structure for \(\mathrm{NF}_{3} .\) What are its electronpair and molecular geometries? What is the hybridization of the nitrogen atom? What orbitals on \(\mathrm{N}\) and \(\mathrm{F}\) overlap to form bonds between these elements?
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