Chem chapter 5 study guide

5.1 Morphine: A Molecular Imposter

Implications of Molecular Shape

The shape of molecules plays a crucial role in their biological activity. For example, the molecular shape of morphine allows it to bind to pain receptors in the body, mimicking the action of natural endorphins.

• VSEPR theory and the Lewis model are essential tools for predicting molecular shape.

• Molecular shape determines how molecules interact with biological targets.

• Example: Morphine's shape enables it to act as a molecular imposter, fitting into endorphin receptors.

5.2 Electronegativity and Bond Polarity

Bond Polarity and Electronegativity

Electrons in a bond are not always shared equally between two atoms. The unequal sharing leads to polar covalent bonds.

• Electronegativity (EN): A measure of an atom's ability to attract electrons in a bond. It can be found on the periodic table.

• Bonds can range from nonpolar covalent (equal sharing) to polar covalent (unequal sharing) to ionic (complete transfer).

• Dipole moment: A quantitative measure of bond polarity, calculated as $\mu = q \times r$, where $q$ is the charge and $r$ is the distance between charges.

• Example: The H–Cl bond is polar because Cl is more electronegative than H.Percent ionic character: The extent to which a bond is ionic versus covalent.

5.3 Writing Lewis Structures for Molecular Compounds and Polyatomic Ions

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule.

• Draw Lewis structures for both molecular compounds and polyatomic ions.

• Determine the central atom (usually the least electronegative).

• Account for all valence electrons.

• Example: The Lewis structure for CO2 shows double bonds between C and each O, with lone pairs on the oxygens.

5.4 Resonance and Formal Charge

Resonance Structures and Formal Charge

Some molecules cannot be represented by a single Lewis structure. Instead, they have resonance structures—multiple valid Lewis structures that differ only in the placement of electrons.

• Resonance: Delocalization of electrons across multiple atoms, stabilizing the molecule.

• Formal charge: Calculated for each atom in a Lewis structure as:

$\text{Formal charge} = \text{(valence electrons)} - \text{(nonbonding electrons)} - \frac{1}{2}\text{(bonding electrons)}$

• Helps identify the most stable resonance structure (lowest formal charges preferred).

• Example: Ozone (O3) has two resonance structures with different placements of double bonds.

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

Octet Rule Exceptions

While many molecules follow the octet rule (eight electrons around each atom), there are important exceptions:

• Odd-electron species: Molecules with an odd number of electrons (e.g., NO).

• Incomplete octets: Molecules where some atoms (e.g., B, Be) have fewer than eight electrons.

• Expanded octets: Elements in period 3 or higher can have more than eight electrons (e.g., SF6).

• Draw Lewis structures for each type of exception.

• Example: BF3 (boron trifluoride) has only six electrons around boron.

5.6 Bond Energies and Bond Lengths

Bond Energy and Bond Length

Bond energy is the energy required to break one mole of a bond in a molecule in the gas phase. Bond length is the average distance between the nuclei of two bonded atoms.

• Bond energies can be used to estimate reaction enthalpies using the formula:

$\Delta H_{\text{rxn}} = \sum \text{(bond energies of bonds broken)} - \sum \text{(bond energies of bonds formed)}$

• There is an inverse relationship: as bond length decreases, bond strength (energy) increases.

• Example: Triple bonds are shorter and stronger than double or single bonds.

5.7 VSEPR Theory: The Effect of Electron Groups

Electron Groups and Molecular Shape

VSEPR theory (Valence Shell Electron Pair Repulsion) predicts the shape of molecules based on the repulsion between electron groups (bonding and lone pairs) around a central atom.

• Electron groups repel each other and arrange themselves as far apart as possible.

• Five basic shapes are predicted based on the number of electron groups:

Electron Groups	Base Shape	Bond Angles
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	90°, 120°
6	Octahedral	90°

• Recognize molecules in their core shapes based on the number of electron groups.

5.8 VSEPR Theory: The Effect of Lone Pairs

Lone Pairs and Molecular Geometry

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

• Electron geometry considers all electron groups; molecular geometry considers only the arrangement of atoms.

• Lone pairs can compress bond angles, leading to shapes such as bent or trigonal pyramidal.

• Example: In NH3, the presence of a lone pair results in a trigonal pyramidal shape rather than a perfect tetrahedron.

5.9 VSEPR Theory: Predicting Molecular Geometries

Procedures for Predicting and Drawing Molecular Geometries

To predict molecular geometry:

1. Draw the Lewis structure.

2. Count the number of electron groups around the central atom.

3. Determine the electron geometry (based on total electron groups).

4. Determine the molecular geometry (based on bonding groups only).

5. Draw the molecule, showing the correct geometry.

• Apply these steps to molecules with more than one central atom as well.

5.10 Molecular Shape and Polarity

Polarity and Dipole Moments

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

• Polar bonds arise from differences in electronegativity ($\Delta EN$).

• Vector addition of bond dipoles determines the overall molecular dipole moment.

• Nonpolar molecules have dipoles that cancel; polar molecules have a net dipole moment.

• Example: CO2 has polar bonds but is nonpolar overall due to its linear shape.

• Molecular polarity affects macroscopic properties, such as solubility (e.g., water and oil do not mix due to polarity differences).