Chapter 8: Basic Concepts of Chemical Bonding

Basic Concepts of Chemical Bonding

Introduction to Chemical Bonds

Chemical bonds are the attractive forces that hold atoms or ions together in compounds. The nature of these bonds determines the structure and properties of substances.

• Ionic Bonds: Electrostatic attraction between oppositely charged ions. Example: NaCl (sodium chloride)

• Covalent Bonds: Sharing of electron pairs between atoms. Example: H2O (water)

• Metallic Bonds: Free electrons are shared among a lattice of metal atoms. Example: Fe (iron metal)

8.1. Lewis Symbols

Valence Electrons and Lewis Symbols

Valence electrons are the electrons in the outermost shell of an atom and are involved in chemical bonding. Lewis symbols represent these electrons as dots around the chemical symbol of an element.

• Lewis Symbol: Chemical symbol plus a dot for each valence electron.

• Dot Arrangement: Dots are placed on four sides (top, bottom, left, right) of the symbol, with up to two electrons per side.

• Example: Sulfur (S) with six valence electrons: :S:

Octet Rule

Stability and Electron Configuration

Atoms tend to gain, lose, or share electrons to achieve a stable configuration with eight valence electrons, similar to noble gases. This is known as the octet rule.

• Noble Gases: Have eight valence electrons (except He, which has two) and are chemically stable.

• Application: The octet rule mainly applies to s- and p-block elements; transition metals often do not follow this rule due to d-electrons.

8.2. Ionic Bonding

Formation and Properties of Ionic Compounds

Ionic bonding occurs between metals and nonmetals (excluding group 8A). It involves the transfer of electrons, resulting in the formation of cations and anions.

• Electron Transfer: Metals lose electrons (form cations), nonmetals gain electrons (form anions).

• Example Reaction: $\text{Na}(s) + \frac{1}{2}\text{Cl}_2(g) \rightarrow \text{NaCl}(s)$

• Energy Changes: Formation of ionic compounds is highly exothermic due to lattice energy release.

Predicting Ionic Charges

The charge of ions formed by main group elements can be predicted based on their group number in the periodic table.

Group	Common Ion Charge
1A	+1
2A	+2
7A	-1
6A	-2
3A	+3

Properties of Ionic Substances

• Crystalline structure

• Brittle and high melting points

• Cleave along smooth lines

Energetics of Ionic Bonding

• Ionization energy: Energy required to remove electrons from atoms (endothermic).

• Electron affinity: Energy released when atoms gain electrons (exothermic).

• Lattice energy: Energy released when gaseous ions form an ionic solid.

Lattice Energy Equation:

$\text{NaCl}(s) \rightarrow \text{Na}^+(g) + \text{Cl}^-(g) \qquad \Delta H_{\text{lattice}} = +788\ \text{kJ/mol}$

Lattice energy increases with increasing ionic charge and decreasing ionic size:

$E_{el} = \kappa \frac{Q_1 Q_2}{d}$

Electron Configuration of Ions

• Main group metals lose electrons to achieve the configuration of the previous noble gas.

• Nonmetals gain electrons to achieve the configuration of the nearest noble gas.

• Transition metals may not follow the octet rule due to d-electrons.

8.3. Covalent Bonding

Nature of Covalent Bonds

Covalent bonds involve the sharing of electron pairs between atoms. The stability of these bonds arises from the balance of attractive and repulsive electrostatic forces.

• Attractions between electrons and nuclei

• Repulsions between electrons

• Repulsions between nuclei

For a bond to form, attractions must outweigh repulsions.

Lewis Structures

• Represent shared electron pairs as lines and lone pairs as dots.

• Each atom aims to achieve a noble gas configuration (octet).

• Example: H2: H–H; Cl2: Cl–Cl

Number of Bonds for Nonmetals

• The number of covalent bonds needed is equal to the number of electrons required to complete the octet.

• Example: Nitrogen (N) forms three bonds to complete its octet.

Writing Lewis Structures

1. Sum valence electrons from all atoms (adjust for charge).

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

3. Connect atoms with single bonds.

4. Complete octets for outer atoms, then central atom.

5. If electrons remain, place them on the central atom; if octet is incomplete, form multiple bonds.

Bond Types

• Single bond: One pair of shared electrons

• Double bond: Two pairs of shared electrons

• Triple bond: Three pairs of shared electrons

8.4. Polarity of Bonds

Bond Polarity and Electronegativity

Bond polarity describes the distribution of electron density between two atoms in a bond.

• Nonpolar covalent bond: Electrons shared equally (e.g., Cl2).

• Polar covalent bond: Electrons shared unequally due to differences in electronegativity (e.g., HF).

• Electronegativity: The ability of an atom to attract electrons in a bond. Increases across a period and up a group.

Electronegativity Difference	Bond Type
> 2.0	Ionic
0.5 – 2.0	Polar Covalent
< 0.5	Nonpolar Covalent

8.5. Writing Lewis Structures (Covalent Molecules)

Steps for Drawing Lewis Structures

1. Sum valence electrons (adjust for charge).

2. Choose the central atom (least electronegative).

3. Connect atoms with single bonds.

4. Complete octets for outer atoms, then central atom.

5. Use multiple bonds if necessary to complete octets.

Formal Charge

• Formal charge helps determine the most stable Lewis structure.

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

• The dominant structure has formal charges closest to zero and places negative charges on the most electronegative atoms.

8.6. Resonance

Resonance Structures

Some molecules cannot be represented by a single Lewis structure. Resonance structures are used to depict delocalized electrons within molecules (e.g., ozone, O3).

Exceptions to the Octet Rule

• Molecules with an odd number of electrons (e.g., NO)

• Molecules with less than an octet (e.g., BF3)

• Molecules with more than eight electrons (expanded octet, e.g., SF6)

Bond Enthalpy and Bond Length

• Bond enthalpy: Energy required to break a bond (always positive, endothermic).

• Multiple bonds are stronger and shorter than single bonds.

• Bond length decreases as the number of shared electron pairs increases.