Chemical Bonding: Concepts, Types, and Lewis Structures

Chemical Bonding

Introduction to Chemical Bonding

Chemical bonding refers to the energy (or force) that holds atoms together in molecules, salts, and metals. Understanding the nature of chemical bonds is fundamental to organic chemistry, as it explains the structure and properties of molecules.

• Chemical bond: The attractive force that binds atoms together to form compounds.

• Types of chemical bonds: Ionic, covalent (polar and nonpolar), and metallic bonds.

Lewis Dot Symbols

Lewis dot symbols are a simple way to represent the valence electrons of atoms. They are used to visualize the formation of chemical bonds.

• Write the atomic symbol.

• Show the outer (valence) electrons as dots around the symbol.

• Example: Li: one dot, Be: two dots, B: three dots, C: four dots, N: five dots, O: six dots, F: seven dots, Ne: eight dots.

Types of Chemical Bonds

Ionic Bonds

Ionic bonds are formed between metals and nonmetals through the transfer of electrons. Metals lose electrons to achieve a noble gas configuration, while nonmetals gain electrons to achieve the same.

• Formation: One atom donates electrons (becomes a cation), and another atom accepts electrons (becomes an anion).

• Example:

  • $\mathrm{Na} - e^- \longrightarrow \mathrm{Na}^+$

  • $\mathrm{Cl} + e^- \longrightarrow \mathrm{Cl}^-$

  • $\mathrm{Na}^+ + \mathrm{Cl}^- \longrightarrow \mathrm{NaCl}$

• The resulting ionic bond is the electrostatic attraction between oppositely charged ions.

Covalent Bonds

Covalent bonds are formed when two atoms share one or more pairs of electrons. This allows both atoms to achieve a filled valence shell, often resembling the electron configuration of a noble gas.

• Single covalent bond: Two atoms share one pair of electrons (e.g., $\mathrm{H}_2$, $\mathrm{Cl}_2$).

• Double and triple bonds: Atoms can share two or three pairs of electrons, respectively (e.g., $\mathrm{O}_2$, $\mathrm{N}_2$).

• Example: $\mathrm{H} \cdot + \cdot \mathrm{H} \longrightarrow \mathrm{H} - \mathrm{H}$

Octet Rule

Atoms tend to form bonds in such a way that each atom has eight electrons in its valence shell, achieving a noble gas configuration. This is known as the octet rule.

• Hydrogen is an exception, aiming for two electrons (duet rule).

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It varies across the periodic table.

• Trend: Increases from left to right across a period and decreases from top to bottom down a group.

• Examples: Cesium (Cs): 0.7, Fluorine (F): 4.0 (most electronegative element).

Types of Covalent Bonds

• Nonpolar covalent bond: Electrons are shared equally between identical atoms (e.g., $\mathrm{H}_2$, $\mathrm{Cl}_2$).

• Polar covalent bond: Electrons are shared unequally between different atoms, resulting in partial charges (e.g., $\mathrm{HCl}$).

Bond Type by Electronegativity Difference

Electronegativity Difference	Bond Type
> 2.0	Ionic bond
0.3 – 2.0	Polar covalent bond
< 0.3	Nonpolar covalent bond

Polarity of Molecules

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

• If a molecule is polar, it must have at least one polar bond.

• If the molecule is not symmetric, the dipoles do not cancel, and the molecule is polar (e.g., $\mathrm{H}_2\mathrm{O}$).

• If the molecule is symmetric, the dipoles cancel, and the molecule is nonpolar (e.g., $\mathrm{CO}_2$).

Bond Length and Bond Energy

Bond Length

Bond length is the distance between the nuclei of two bonded atoms.

• Shorter bonds are generally stronger and involve more shared electrons (e.g., triple < double < single bond).

Bond Energy

Bond energy is the average energy required to break a particular type of bond between two atoms in a molecule, measured in kJ/mol.

• Higher bond order (number of shared electron pairs) generally means higher bond energy.

Lewis Structures

Steps to Draw Lewis Structures

1. Determine the total number of valence electrons.

2. Place one pair of electrons between each pair of bonded atoms.

3. Distribute the remaining electrons to satisfy the octet rule (or duet for hydrogen), starting with outer atoms.

4. If necessary, form double or triple bonds to ensure all atoms have a complete octet.

Examples of Lewis Structures

• NH3: 8 valence electrons

• H2O: 8 valence electrons

• HF: 8 valence electrons

• C2H6: 14 valence electrons

• C2H2: 10 valence electrons

Practice: Lewis Structures for Common Molecules

• H2O: $\mathrm{H}-\mathrm{O}-\mathrm{H}$, 8 e-

• CO2: $\mathrm{O}=\mathrm{C}=\mathrm{O}$, 16 e-

• CCl4: Central C atom with four Cl atoms, 32 e-

• SF6: Central S atom with six F atoms, 48 e-

• NH2CONH2: Urea, 24 e-

Summary Table: Lewis Structures of Selected Molecules

Molecule	Lewis Structure Description	Total Valence Electrons
H2O	O atom single bonded to two H atoms, two lone pairs on O	8
CO2	C atom double bonded to two O atoms, each O with two lone pairs	16
CCl4	C atom single bonded to four Cl atoms, each Cl with three lone pairs	32
SF6	S atom single bonded to six F atoms, each F with three lone pairs	48
NH2CONH2	Central C double bonded to O, single bonded to two NH2 groups	24

Additional info:

• Lewis structures are foundational for understanding molecular geometry, resonance, and reactivity in organic chemistry.

• Practice drawing Lewis structures for a variety of molecules to master the skill.