Chapter 10: Chemical Bonding – Lewis Structures and Bonding Theories

Chemical Bonding: Introduction and Applications

Bonding Models and AIDS Drugs

Bonding theories are essential for understanding how atoms combine to form molecules, which has direct applications in drug design and disease treatment.

• HIV-protease is a protein crucial for the replication of the HIV virus, which causes AIDS.

• Researchers use bonding theories to simulate how drug molecules interact with HIV-protease, leading to the development of protease inhibitors such as Indinavir.

• Protease inhibitors, when used with other drugs, can reduce HIV viral counts to undetectable levels and are also used to treat other viral diseases like Hepatitis C.

Bonding Theories: The Lewis Model

Overview of Bonding Theories

Bonding theories predict how atoms bond to form compounds and explain molecular shapes, which determine many physical and chemical properties.

• The Lewis model is a foundational bonding theory developed by G.N. Lewis.

• It uses dot structures (Lewis structures) to represent valence electrons and predict molecular stability and structure.

Representing Valence Electrons with Dots

Lewis Structures and Valence Electrons

Lewis structures visually represent the valence electrons of main-group elements as dots around the element symbol.

• The number of valence electrons equals the group number for main-group elements (except helium).

• Dots are placed singly first, then paired, with a maximum of two per side.

• Atoms with eight valence electrons (an octet) are particularly stable; neon is an example.

• Helium is an exception, achieving stability with two electrons (a duet).

Example: Phosphorus (Group 5A) has five valence electrons, represented as five dots around the symbol P.

Lewis Model of Chemical Bonding

Ionic and Covalent Bonds

In the Lewis model, chemical bonds form by sharing or transferring electrons to achieve stable electron configurations.

• Ionic bonds: Electrons are transferred from metals to nonmetals, forming cations and anions.

• Covalent bonds: Electrons are shared between nonmetals.

• The octet rule states that atoms tend to have eight electrons in their valence shell (except H and He, which require two).

Lewis Structures for Ionic Compounds

Electron Transfer and Ionic Bond Formation

• Metals lose electrons to become cations; nonmetals gain electrons to become anions.

• The resulting electrostatic attraction forms an ionic compound.

• Lewis structures show electron transfer by moving dots from metal to nonmetal.

Example: Potassium (K) transfers its valence electron to chlorine (Cl), forming K+ and [Cl]-, which combine to make KCl.

Example: Magnesium (Mg) and oxygen (O) form MgO by Mg losing two electrons (Mg2+) and O gaining two (O2-).

Example: Sodium (Na) and sulfur (S) form Na2S, requiring two Na atoms for each S atom.

Lewis Structures for Covalent Compounds

Electron Sharing and Molecular Compounds

• Nonmetals share electrons to form molecular compounds with covalent bonds.

• Shared electrons (bonding pairs) are shown between atoms; lone pairs are nonbonding electrons on a single atom.

• Bonding pairs are often represented by dashes (e.g., H—O—H for water).

Example: In H2O, each hydrogen shares electrons with oxygen, giving H a duet and O an octet.

Diatomic Molecules and Multiple Bonds

• Halogens (e.g., Cl2) and hydrogen (H2) form diatomic molecules to achieve stable configurations.

• Double bonds (e.g., O2) and triple bonds (e.g., N2) form when atoms share two or three pairs of electrons, respectively. Multiple bonds are shorter and stronger than single bonds.

Steps for Writing Lewis Structures for Covalent Compounds

1. Write the correct skeletal structure (central atom, usually the least electronegative; H is always terminal).

2. Sum the total number of valence electrons (add/subtract for polyatomic ions).

3. Distribute electrons to give octets (or duets for H), starting with bonding pairs, then lone pairs.

4. If any atom lacks an octet, form double or triple bonds as needed.

Example: For CCl4, C is central, surrounded by four Cl atoms; total electrons = 32.

Example: For CO2, C is central, double bonds to each O; total electrons = 16.

Polyatomic ions: Add electrons for negative charge, subtract for positive, and enclose the structure in brackets with the charge.

Exceptions to the Octet Rule

• Molecules with odd numbers of electrons (e.g., NO) cannot satisfy the octet rule for all atoms.

• Boron compounds (e.g., BF3, BH3) often have only six electrons around B.

• Some molecules (e.g., SF6, PCl5) have expanded octets (more than eight electrons around the central atom).

Resonance Structures

Some molecules (e.g., SO2) can be represented by more than one valid Lewis structure. The actual molecule is an average (resonance hybrid) of these structures.

• Resonance structures are shown with a double-headed arrow between them.

• Bond lengths and strengths are intermediate between single and double bonds.

Predicting Molecular Shapes: VSEPR Theory

Electron and Molecular Geometry

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

• Electron groups (lone pairs, single, double, or triple bonds) arrange themselves to minimize repulsion.

• Common geometries: linear (180°), trigonal planar (120°), tetrahedral (109.5°).

• Lone pairs affect molecular geometry, making bond angles slightly less than ideal.

Example: CH4 is tetrahedral; NH3 is trigonal pyramidal; H2O is bent.

Electronegativity and Polarity

Bond Polarity and Molecular Polarity

• Electronegativity is an element's ability to attract electrons in a bond (scale: 0.7 to 4.0; F is most electronegative).

• If two atoms have identical electronegativity, the bond is nonpolar covalent (e.g., Cl2).

• Large differences result in ionic bonds (e.g., NaCl); intermediate differences yield polar covalent bonds (e.g., HCl).

• A dipole moment arises from unequal sharing of electrons, creating partial charges (δ+ and δ−).

Table: Effect of Electronegativity Difference on Bond Type

Electronegativity Difference (ΔEN)	Bond Type	Example
0–0.4	Nonpolar covalent	Cl2
0.5–1.9	Polar covalent	HCl
≥2.0	Ionic	NaCl

Polar Molecules

• A molecule is polar if its polar bonds do not cancel out (net dipole moment).

• CO2 is nonpolar (linear geometry, dipoles cancel); H2O is polar (bent geometry, dipoles add).

Everyday Chemistry: Soap and Molecular Polarity

• Water is polar; oil and grease are nonpolar, so they do not mix.

• Soap molecules have a polar head (attracts water) and a nonpolar tail (attracts grease), allowing water and grease to mix and be washed away.

Review and Learning Objectives

• Write Lewis structures for elements, ionic, and covalent compounds.

• Distinguish between ionic and covalent compounds.

• Write resonance structures and predict molecular shapes using VSEPR theory.

• Determine molecular polarity based on bond polarity and geometry.