Study Guide: Chemical Bonding and Molecular Geometry (Chapters 8 & 9)

Chapter 8: Basic Concepts of Chemical Bonding

Lewis Symbols and the Octet Rule

The Lewis symbol represents the valence electrons of an atom as dots around the chemical symbol. The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, resembling the electron configuration of a noble gas.

• Drawing Lewis Symbols: Place one dot for each valence electron around the element symbol.

• Determining Electron Loss or Gain: Atoms lose or gain electrons to complete their octet (8 electrons).

• Example: For K (potassium), the Lewis symbol is K• (one valence electron); for N3−, the symbol is N with eight dots and a 3− charge.

Stability of Ionic Compounds

Ionic compounds form when electrons are transferred from a metal to a nonmetal, resulting in oppositely charged ions that attract each other.

• Electron Transfer: Metals lose electrons to form cations; nonmetals gain electrons to form anions.

• Example: Mg (magnesium) loses two electrons to O (oxygen), which gains two electrons, forming MgO.

• Number of Electrons Transferred: Equal to the number needed for both atoms to achieve noble gas configurations.

• Predicting Formulas: Combine ions in ratios that balance total positive and negative charges (e.g., AlF3, K2S, Y2O3, Mg3N2).

Lattice Energy

Lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions. It is a measure of the strength of the ionic bonds in a solid.

• Trends:

  • Increases as the charges of the ions increase.

  • Decreases as the sizes (radii) of the ions increase.

• Ordering Compounds: For example, lattice energy increases in the order: KI < LiBr < MgS < GaN.

Preferred Charges on Ions and the Octet Rule

Atoms form ions with charges that allow them to achieve a noble gas electron configuration, typically following the octet rule.

• Metals: Lose electrons to form positive ions (cations).

• Nonmetals: Gain electrons to form negative ions (anions).

Covalent vs. Ionic Bonding

Covalent bonds involve the sharing of electron pairs between atoms, typically between nonmetals. Ionic bonds involve the transfer of electrons from one atom to another, typically between metals and nonmetals.

• Examples:

  • Iron: Metallic bonding

  • Sodium chloride: Ionic bonding

  • Water, oxygen: Covalent bonding

  • Argon: No bonding (noble gas)

• Physical Properties: Covalent substances often have lower boiling points than ionic substances.

Electronegativity and Periodic Trends

Electronegativity is the ability of an atom to attract shared electrons in a chemical bond. It increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Most Electronegative Elements: In each set, the element furthest to the right and top is most electronegative (e.g., F, O, N, Cl).

Bond Polarity and Electronegativity Difference

Bonds are classified based on the difference in electronegativity between the two atoms:

• Nonpolar Covalent: Electronegativity difference < 0.5

• Polar Covalent: 0.5 < difference < 2.0

• Ionic: Difference > 2.0

• Examples: B—F and Se—O are polar; Cl—Cl is nonpolar; H—I is polar (I is more electronegative).

Lewis Structures and Formal Charge

Lewis structures show the arrangement of atoms, bonds, and lone pairs in a molecule. Formal charge is a bookkeeping tool to determine the most stable Lewis structure.

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

• Application: Draw Lewis structures for CF4, NO+, SO32−, HCN, BF4−, HOCl, etc., and calculate formal charges for specific atoms.

Chapter 9: Molecular Geometry and Bonding Theories

VSEPR Model and Molecular Shapes

The Valence Shell Electron Pair Repulsion (VSEPR) model predicts the three-dimensional shapes of molecules based on the repulsion between electron domains (bonding and lone pairs) around a central atom.

• ABn Molecules: 'A' is the central atom, 'B' are surrounding atoms, 'n' is the number of B atoms.

• Electron-Domain Geometry: Determined by the total number of electron domains (bonding + lone pairs).

• Molecular Geometry: Determined by the arrangement of atoms (ignoring lone pairs).

• Example: AB2 with only single bonds and no lone pairs is linear; not all AB2 molecules are linear if lone pairs are present.

Characteristic Bond Angles

Each molecular geometry has characteristic bond angles:

• Trigonal Planar: 120°

• Tetrahedral: 109.5°

• Octahedral: 90°

• Linear: 180°

Deriving Geometries from Tetrahedral

Removing B atoms or replacing them with lone pairs in a tetrahedral AB4 molecule leads to different shapes:

• Trigonal Pyramidal: One lone pair (e.g., NH3)

• Bent: Two lone pairs (e.g., H2O)

Electron-Domain and Molecular Geometries (VSEPR)

Electron-domain geometry considers all electron domains; molecular geometry considers only the positions of atoms.

• Examples:

  • AB3 with 109° bond angles: Tetrahedral electron-domain, trigonal pyramidal molecular geometry if one lone pair is present.

  • AB6 with no lone pairs: Octahedral geometry.

  • AB4 with two lone pairs: Electron-domain geometry is octahedral; molecular geometry is square planar.

Table: Electron-Domain and Molecular Geometries

Electron Domains	Bonding Domains	Lone Pairs	Electron-Domain Geometry	Molecular Geometry
2	2	0	Linear	Linear
3	3	0	Trigonal Planar	Trigonal Planar
3	2	1	Trigonal Planar	Bent
4	4	0	Tetrahedral	Tetrahedral
4	3	1	Tetrahedral	Trigonal Pyramidal
4	2	2	Tetrahedral	Bent
5	5	0	Trigonal Bipyramidal	Trigonal Bipyramidal
6	6	0	Octahedral	Octahedral
6	4	2	Octahedral	Square Planar

Predicting Molecular Polarity

A molecule is polar if it has a net dipole moment due to an uneven distribution of electron density. Molecular geometry and bond polarity both affect overall polarity.

• Examples:

  • BF3: Nonpolar (trigonal planar, dipoles cancel)

  • SO3: Nonpolar (trigonal planar, dipoles cancel)

  • PCl3: Polar (trigonal pyramidal, dipoles do not cancel)

  • SF6: Nonpolar (octahedral, dipoles cancel)

  • IF5: Polar (square pyramidal, dipoles do not cancel)

Hybridization of Atomic Orbitals

Hybrid orbitals are formed by the combination of atomic orbitals on the same atom to explain molecular shapes.

• sp Hybridization: Linear geometry (180° bond angle)

• sp2 Hybridization: Trigonal planar geometry (120° bond angle)

• sp3 Hybridization: Tetrahedral geometry (109.5° bond angle)

• Examples: BCl3 (sp2), AlCl4− (sp3), CS2 (sp), GeH4 (sp3)

Sigma (σ) and Pi (π) Bonds

Sigma (σ) bonds are formed by head-on overlap of orbitals, while pi (π) bonds are formed by side-on overlap of p orbitals.

• σ Bond: Stronger, allows free rotation, present in all single bonds.

• π Bond: Weaker, restricts rotation, present in double and triple bonds (along with σ bonds).

• Formation: σ bonds can form from s-s, s-p, or p-p overlap; π bonds form only from p-p overlap.

Table: Comparison of σ and π Bonds

Bond Type	Orbital Overlap	Strength	Rotation
σ (Sigma)	Head-on (s-s, s-p, p-p)	Stronger	Free rotation
π (Pi)	Side-on (p-p)	Weaker	No free rotation

• True Statements:

  • σ bond is generally stronger than a π bond.

  • π bond is made from sideways overlap of two p orbitals.

  • π bond has two regions of overlap on opposite sides of the internuclear axis.

  • Two s orbitals cannot make a π bond.

Additional info: The VSEPR model and hybridization theory are foundational for predicting and explaining molecular shapes, bond angles, and reactivity in general chemistry. Mastery of Lewis structures, formal charge, and polarity is essential for understanding chemical behavior and properties.