Chapter 5: Chemical Bonding I – Drawing Lewis Structures and Determining Molecular Shapes

5.1 Morphine: A Molecular Imposter

The structure and shape of molecules play a crucial role in their biological activity. Morphine, a compound derived from the opium poppy, mimics the structure of endorphins, allowing it to bind to the same pain receptors in the body. Understanding molecular shape is essential for explaining such interactions.

• Lock-and-Key Model: Describes how molecules fit together based on their shapes, similar to a key fitting into a lock.

• VSEPR Theory and Lewis Model: Useful for predicting and explaining molecular shapes.

• Example: Morphine's ability to bind to endorphin receptors is due to its similar shape to endorphins.

5.2 Electronegativity and Bond Polarity

Atoms in a molecule may share electrons unequally, leading to different types of chemical bonds. The concept of electronegativity helps predict bond polarity and the nature of chemical bonds.

• Electronegativity (EN): The ability of an atom to attract shared electrons in a chemical bond. It increases across a period and decreases down a group in the periodic table.

• Bond Polarity: Determined by the difference in electronegativity () between two atoms:

  • Nonpolar covalent: to

  • Polar covalent: to

  • Ionic:

• Dipole Moment (): A measure of bond polarity, calculated as where is the magnitude of the charge and is the distance between charges.

• Percent Ionic Character: The extent to which a bond is ionic, based on measured dipole moments compared to those expected for full ionic separation.

• Example: The bond between S (2.5) and O (3.5) has (polar covalent).

5.3 Writing Lewis Structures for Molecular Compounds and Polyatomic Ions

Lewis structures represent the arrangement of valence electrons in molecules and ions, helping to predict molecular shape and reactivity.

• Skeletal Structure: The arrangement of atoms, with the least electronegative atom (except H) typically at the center.

• Valence Electrons: Sum the valence electrons from all atoms (add or subtract electrons for ions).

• Electron Distribution: Distribute electrons to form bonds and satisfy the octet (or duet for H) rule.

• Multiple Bonds: Form double or triple bonds if necessary to complete octets.

• Example: For , C is central, total valence electrons = 16, double bonds formed to satisfy octets.

5.4 Resonance and Formal Charge

Some molecules cannot be represented by a single Lewis structure. Resonance and formal charge concepts help describe these cases and identify the most stable structures.

• Resonance Structures: Two or more valid Lewis structures for the same molecule, differing only in the placement of electrons.

• Resonance Hybrid: The actual structure is a blend of all resonance forms, leading to special stability.

• Formal Charge: Calculated as:

• Minimizing Formal Charge: The most stable Lewis structure has the smallest formal charges, with negative charges on the most electronegative atoms.

• Example: has resonance structures with formal charges distributed between N and O atoms.

5.5 Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets

Not all molecules follow the octet rule. There are three main exceptions:

• Odd-Electron Species (Radicals): Molecules with an odd number of electrons, such as .

• Incomplete Octets: Atoms (especially Be and B) with fewer than eight electrons, e.g., .

• Expanded Octets: Atoms (period 3 and beyond) with more than eight electrons, e.g., .

• Second-Period Elements: Cannot have expanded octets due to lack of available d orbitals.

5.6 Bond Energies and Bond Lengths

Bond energy and bond length are fundamental properties that influence molecular stability and reactivity.

• Bond Energy: The energy required to break one mole of a bond in the gas phase. Breaking bonds is endothermic; forming bonds is exothermic.

• Average Bond Energy: Used to estimate reaction enthalpies ():

• Bond Length: The distance between the nuclei of two bonded atoms. Multiple bonds are shorter and stronger than single bonds.

5.7 VSEPR Theory: The Five Basic Shapes

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on the repulsion between electron groups around a central atom.

• Electron Groups: Include bonds (single, double, triple) and lone pairs.

• Basic Geometries:

  • Linear (2 groups): 180° bond angle (e.g., , CO)

  • Trigonal Planar (3 groups): 120° (e.g., , )

  • Tetrahedral (4 groups): 109.5° (e.g., )

  • Trigonal Bipyramidal (5 groups): 90°, 120°, 180° (e.g., )

  • Octahedral (6 groups): 90°, 180° (e.g., )

5.8 VSEPR Theory: The Effect of Lone Pairs

Lone pairs occupy more space than bonding pairs, affecting molecular geometry and bond angles.

• Electron Geometry vs. Molecular Geometry: Electron geometry considers all electron groups; molecular geometry considers only the arrangement of atoms.

• Tetrahedral Electron Geometry:

  • Trigonal pyramidal: 1 lone pair (e.g., )

  • Bent: 2 lone pairs (e.g., )

• Trigonal Bipyramidal Electron Geometry:

  • Seesaw: 1 lone pair (e.g., )

  • T-shaped: 2 lone pairs (e.g., )

  • Linear: 3 lone pairs (e.g., )

• Octahedral Electron Geometry:

  • Square pyramidal: 1 lone pair (e.g., )

  • Square planar: 2 lone pairs (e.g., )

5.9 VSEPR Theory: Predicting Molecular Geometries

Predicting molecular geometry involves a systematic procedure based on Lewis structures and electron group counting.

1. Draw the Lewis structure.

2. Count the total number of electron groups on the central atom.

3. Determine the number of bonding groups and lone pairs.

4. Use VSEPR theory (or a reference table) to identify electron and molecular geometry.

5. For molecules with more than one central atom, determine the geometry around each central atom separately.

5.10 Molecular Shape and Polarity

The overall polarity of a molecule depends on both the polarity of its bonds and its molecular geometry.

• Polar Bonds: Identified by differences in electronegativity.

• Net Dipole Moment: Determined by vector addition of individual bond dipoles.

• Vector Addition: Used to predict whether the molecule as a whole is polar or nonpolar.

• Macroscopic Properties: Molecular polarity affects properties such as solubility and boiling point (e.g., water and oil immiscibility).

• Example: has polar bonds but is nonpolar overall due to its linear geometry.

Additional Problems and Examples

Classifying bonds and drawing Lewis structures are essential skills in understanding chemical bonding.

• Classifying Bonds by Electronegativity Difference:

• Writing Lewis Structures:

  • For : C is central, total valence electrons = 16, double bonds to S atoms.

  • For : C is central, H and Cl are terminal, total valence electrons = 14.

  • For : N is central, total valence electrons = 18 (including extra for charge), resonance structures possible.

Additional info: For more complex molecules, always check for possible resonance structures and formal charges to determine the most accurate Lewis structure. Practice with a variety of examples is essential for mastering these concepts.