General Chemistry: Atomic Structure, Bonding, and Molecular Geometry – Exam 4 Review Key Takeaways

Periodic Trends and Atomic Structure

Trends in Atomic Properties

Understanding periodic trends is essential for predicting the properties of elements. These trends are observed as you move across periods (rows) and down groups (columns) in the periodic table.

• Effective Nuclear Charge (Zeff): Increases from left to right across a period due to increasing number of protons, but remains relatively constant down a group.

• Atomic Radius: Decreases across a period (left to right) and increases down a group. This is because increased Zeff pulls electrons closer, while adding shells increases size.

• Ionization Energy: Increases across a period and decreases down a group. Higher Zeff makes it harder to remove electrons.

• Electron Affinity: Generally becomes more negative (more exothermic) across a period, but there are exceptions.

• Example: Sodium (Na) has a larger atomic radius than chlorine (Cl), but Cl has a higher ionization energy.

Electronic Configurations and Ions

Writing Electron Configurations for Ions

When writing electron configurations for cations, electrons are removed from the highest principal quantum number (n) orbital first, not necessarily the last filled in the Aufbau order.

• For transition metals: Remove electrons from the s orbital before the d orbital.

• Example: Fe: [Ar] 4s2 3d6; Fe2+: [Ar] 3d6

Bonding and Molecular Geometry

Types of Bonds and Hybridization

The type of hybridization and the number of bonds determine molecular geometry and properties.

• sp Hybridization: Linear geometry, 180° bond angle.

• sp2 Hybridization: Trigonal planar geometry, 120° bond angle.

• sp3 Hybridization: Tetrahedral geometry, 109.5° bond angle.

• Pi (π) Bonds: Formed from unhybridized p orbitals; must involve unhybridized orbitals.

• Example: Ethene (C2H4) has sp2 hybridization and a double bond (one σ and one π bond).

Molecular Geometry and VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron domain repulsions.

• Common Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Bond Angles: Tetrahedral (109.5°), trigonal planar (120°), linear (180°).

• Example: NH3 is tetrahedral in electron geometry but trigonal pyramidal in molecular geometry due to one lone pair.

Ionic Compounds and Lattice Energy

Lattice Energy

Lattice energy is the energy released when gaseous ions form an ionic solid. It is influenced by the charges and radii of the ions.

• Magnitude: Increases with higher ionic charges and smaller ionic radii.

• Formula: $U = k \frac{Q_1 Q_2}{r}$ (where $Q_1$ and $Q_2$ are the charges, $r$ is the distance between ions, and $k$ is a proportionality constant).

• Example: MgO has a higher lattice energy than NaCl due to higher charges on Mg2+ and O2−.

Lewis Structures and Formal Charge

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules. They help predict molecular shape, reactivity, and formal charge.

• Steps:

  1. Count total valence electrons.

  2. Draw skeletal structure with least electronegative atom in the center.

  3. Distribute electrons to satisfy octets (or duets for H).

  4. Assign remaining electrons as lone pairs.

  5. Check for multiple bonds if octet is not satisfied.

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

• Best Structure: Minimizes formal charges and places negative charges on more electronegative atoms.

Bond Order, Bond Length, and Bond Strength

Relationship Between Bonds

The number of bonds between two atoms (bond order) affects bond length and strength.

• Bond Order: Single < Double < Triple (increasing order)

• Bond Length: Decreases as bond order increases.

• Bond Strength: Increases as bond order increases.

• Example: N≡N (triple bond) is shorter and stronger than O=O (double bond).

Resonance and Delocalization

Resonance Structures

Some molecules have multiple valid Lewis structures (resonance forms). The true structure is a hybrid of these forms.

• Delocalized π Bonding: Occurs when electrons are shared over more than two atoms, leading to resonance stabilization.

• Localized π Bonding: Electrons are shared between only two atoms.

• Example: Ozone (O3) has two resonance structures with delocalized π electrons.

Summary Table: Hybridization, Geometry, and Bond Angles

Hybridization	Geometry	Bond Angle	Example
sp	Linear	180°	BeCl2
sp2	Trigonal planar	120°	BF3
sp3	Tetrahedral	109.5°	CH4
sp3d	Trigonal bipyramidal	90°, 120°	PCl5
sp3d2	Octahedral	90°	SF6

Additional info:

• Some context and explanations have been expanded for clarity and completeness.

• Where the original notes referenced "see question X," the relevant content has been integrated into the appropriate sections above.