Chem 9 Chapter 4

Predicting Molecular Shapes

• Lewis Theory and VSEPR Theory:

  • Lewis theory, combined with valence shell electron pair repulsion (VSEPR) theory, is essential for predicting molecular shapes.

  • VSEPR Theory postulates that:

    • Electron groups (lone pairs, single bonds, multiple bonds) repel each other.

    • This repulsion dictates the geometry of the molecule.

Geometry of Common Molecules

• CO2 Geometry:

  • Lewis Structure: Features two double bonds.

  • Geometry Determination: Two electron groups (double bonds) on carbon repel, leading to a bond angle of 180° and a linear geometry.

• H2CO Geometry:

  • Lewis Structure: Contains three total electron groups around carbon.

  • Outcome: The electron groups maximize distance, resulting in a bond angle of 120° and a trigonal planar geometry.

  • Bond Angle Adjustments: The C=O double bond has more electron density than C-H bonds, causing slight adjustments to bond angles: HCH is 116° and HCO is 122°.

• CH4 Geometry:

  • Electron Groups: Four electron groups around carbon signify a tetrahedral geometry with bond angles of 109.5°.

  • Shape Explanation: The tetrahedron is optimal for electron group separation, allowing maximum distance between groups.

Impact of Lone Pairs on Geometry

• NH3 Geometry:

  • Structure Review: Four electron groups (1 lone pair, 3 bonds).

  • Electron Geometry: Tetrahedral; Molecular Geometry: Trigonal pyramidal due to the lone pair's influence.

  • Bond Angle: Slightly less than 109.5°.

• H2O Geometry:

  • Structure Review: Four electron groups (2 lone pairs, 2 bonds).

  • Electron Geometry: Tetrahedral; Molecular Geometry: Bent due to lone pairs.

  • Bond Angles: Slightly less than the ideal angles due to lone pair repulsion.

Summary of Electron and Molecular Geometries

• Table 10.1 Overview:

  • Electron Groups and Configurations:

    • 2 groups: linear (180°),

    • 3 groups: trigonal planar (120°),

    • 4 groups: tetrahedral (109.5°).

Using VSEPR for Geometry Determination

1. Draw correct Lewis Structure.

2. Total electron groups around central atom.

3. Count bonding groups vs. lone pairs.

4. Refer to Table 10.1 for geometry classification.

Representation Methods

• Chemists employ specific notations to illustrate three-dimensional structures on two-dimensional surfaces.

Molecular Shape and Health Impact

• Artificial Sweeteners:

  • Example: Aspartame and saccharin have negligible calories but are sweet due to molecular shape fitting receptor sites in taste cells.

Electronegativity and Molecular Interactions

• Oil and Water Separation:

  • Oil and water do not mix due to polar and non-polar interactions.

• Electronegativity Definition:

  • Electronegativity refers to an element's ability to attract electrons in a bond.

  • Oxygen possesses higher electronegativity than hydrogen, leading to unequal electron distribution.

• Dipole Moments:

  • Unequal sharing causes the oxygen to carry a partial negative charge (δ−) and hydrogen a partial positive charge (δ+).

Polar Covalent Bonds

• Definition: Bonds exhibiting a dipole moment due to electronegativity differences.

• The strength of polarity is related to the electronegativity difference and bond length.

Identifying Bond Properties

• Identical Electronegativity:

  • Covalent bonding results in equal electron sharing, yielding a nonpolar molecule (e.g., Cl2).

• Large Electronegativity Difference:

  • Leads to ionic bonds (e.g., NaCl).

• Intermediate Difference:

  • Polar covalent bonds (e.g., HF).

Polarity in Molecules

• Criteria for Polar Molecules:

  • Presence of polar bonds does not always mean the entire molecule is polar; the configuration must also be considered.

  • Example: CO2 has polar bonds but is nonpolar due to linear cancellation of dipole moments.

  • Conversely, H2O is a polar molecule due to unbalanced dipole moments from its bent shape.

Vector Notation for Dipole Moments

• Arrows represent the direction of dipole moments, aiding in visualizing molecular polarity.

Summary of Intermolecular Forces**

• Types:

  • Dispersion Forces: Weak, present in all molecules due to temporary dipoles.

  • Dipole-Dipole Forces: Moderate strength in polar molecules; attraction between permanent dipoles.

  • Hydrogen Bonds: Strongest force, occurs in compounds with hydrogen bonded to F, O, or N.

Analysis of Water's Unique Properties**

• Water stays liquid at room temperature due to strong intermolecular hydrogen bonding despite its low molar mass.

• This significant dipole results from its bent geometry and electronegative bonds.

Chapter Summary**

• Understanding molecular and intermolecular forces is critical in chemistry for predicting chemical behavior and interactions.