Chapter 8: Basic Concepts of Chemical Bonding – Study Notes

Basic Concepts of Chemical Bonding

Introduction to Chemical Bonds

Chemical bonding is a fundamental concept in chemistry that explains how atoms combine to form compounds. There are three primary types of chemical bonds, each with distinct characteristics and mechanisms.

• Ionic Bonds: Formed by electrostatic attraction between ions, typically between metals and nonmetals.

• Covalent Bonds: Involve the sharing of electrons between atoms, usually nonmetals.

• Metallic Bonds: Characterized by free electrons that hold metal atoms together in a 'sea of electrons.'

Example: Table salt (NaCl) is held together by ionic bonds, while water (H2O) is held together by covalent bonds.

Lewis Symbols and the Octet Rule

Valence Electrons and Lewis Symbols

G. N. Lewis introduced a method to represent valence electrons using dots around the element symbol. This notation helps visualize potential bonding electrons.

• Valence Electrons: Electrons in the outermost shell, involved in chemical bonding.

• Lewis Symbol: Element symbol surrounded by dots representing valence electrons.

Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons, similar to the electron configuration of noble gases.

Group	Element	Electron Configuration	Lewis Symbol
1A	Li	[He]2s1	Li•
2A	Be	[He]2s2	Be••
7A	Cl	[Ne]3s23p5	Cl•••••••
8A	Ne	[He]2s22p6	Ne••••••••

Ionic Bonding

Formation and Characteristics

Ionic bonding occurs between metals and nonmetals (except group 8A elements) and involves the transfer of electrons. This process is highly exothermic and results in the formation of ions with opposite charges.

• Electron Transfer: One atom (metal) loses electrons (low ionization energy), while another (nonmetal) gains electrons (high electron affinity).

• Exothermic Reaction: The formation of ionic compounds releases energy.

Example: Sodium (Na) reacts with chlorine (Cl) to form sodium chloride (NaCl):

$\mathrm{Na} + \mathrm{Cl} \rightarrow \mathrm{Na}^+ + \mathrm{Cl}^-$

Arrows in Lewis structures indicate the transfer of electrons from Na to Cl.

Properties of Ionic Compounds

• Crystalline Structure: Ionic compounds form well-defined three-dimensional lattices.

• Brittleness: Ionic solids are brittle and cleave along smooth lines.

• High Melting Points: Due to strong electrostatic forces between ions.

Example: In NaCl, each Na+ ion is surrounded by six Cl- ions and vice versa.

Energetics of Ionic Bonding

The formation of ionic bonds involves several energy changes:

• Energy is required to convert elements to atoms (endothermic).

• Energy is required to create cations (endothermic).

• Energy is released to form anions (exothermic).

• Formation of the solid lattice releases a large amount of energy (exothermic).

Born-Haber Cycle: A method to calculate the overall energy change in the formation of ionic compounds.

Lattice Energy

Lattice energy is a measure of the stability gained by arranging oppositely charged ions in an ionic solid. It is defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.

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

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

Compound	Lattice Energy (kJ/mol)
NaCl	788
MgO	3923
LiF	1036
CsI	600
Additional info: Values inferred from typical textbook data.

Electron Configuration of Ions

Formation of Ions

When metals lose electrons, they achieve the electron configuration of the previous noble gas. Nonmetals gain electrons to reach the configuration of the nearest noble gas. Transition metals may not follow the octet rule and lose valence electrons first, then d-electrons as needed for their ion charge.

• Example: Na: [Ne] after losing one electron; Cl: [Ar] after gaining one electron.

Covalent Bonding

Formation and Characteristics

Covalent bonds involve the sharing of electrons between atoms, primarily nonmetals. Several electrostatic interactions occur in covalent bonds:

• Attractions between electrons and nuclei

• Repulsions between electrons

• Repulsions between nuclei

For a bond to form, attractions must outweigh repulsions.

Lewis Structures

Lewis structures visually represent the sharing of electrons in covalent bonds. Each atom aims to achieve a noble gas configuration by sharing electrons.

• Example: H2: $\mathrm{H}-\mathrm{H}$; Cl2: $\mathrm{Cl}-\mathrm{Cl}$

Number of Bonds for Nonmetals

The group number indicates the number of valence electrons. To achieve an octet, nonmetals form a number of bonds equal to the electrons needed to complete the octet.

Element	Valence Electrons	Bonds Needed
H	1	1
O	6	2
N	5	3
C	4	4

Lone Pairs and Bonding Pairs

• Lone Pairs: Unshared pairs of electrons located on a single atom.

• Bonding Pairs: Shared electrons between two atoms, represented by a line or two dots.

Multiple Bonds

• Single Bond: One pair of shared electrons.

• Double Bond: Two pairs of shared electrons.

• Triple Bond: Three pairs of shared electrons.

Example: CO: $\mathrm{O}=\mathrm{C}=\mathrm{O}$; N2: $\mathrm{N}\equiv\mathrm{N}$

Polarity of Bonds and Electronegativity

Bond Polarity

Bond polarity measures how equally electrons are shared in a covalent bond. In nonpolar covalent bonds, electrons are shared equally. In polar covalent bonds, one atom attracts electrons more strongly.

• Nonpolar Example: F2 (fluorine molecule)

• Polar Example: HF (hydrogen fluoride), where F attracts electrons more strongly than H

Electronegativity

Electronegativity is the ability of an atom in a molecule to attract electrons to itself. It generally increases from left to right across a period and from bottom to top within a group.

• High Electronegativity: Fluorine (most electronegative element)

• Low Electronegativity: Cesium, Francium

Polar Covalent Bonds

When two atoms share electrons unequally, a polar covalent bond results. The more electronegative atom acquires a partial negative charge ($\delta^-$), while the other becomes partially positive ($\delta^+$).

Example: HF: $\mathrm{H}-\mathrm{F}$, with F being $\delta^-$ and H being $\delta^+$

Dipole Moments

A dipole moment occurs when two electrical charges of equal magnitude but opposite sign are separated by a distance. The dipole moment ($\mu$) is calculated as:

$\mu = Q \times r$

Measured in debyes (D).

Comparing Ionic and Covalent Bonding

• Ionic Bonding: Complete electron transfer (metal + nonmetal)

• Covalent Bonding: Electron pair sharing (two nonmetals)

• Electronegativity Difference: >2.0 is often ionic; exceptions exist

Drawing Lewis Structures

Steps for Drawing Lewis Structures

1. Sum the valence electrons from all atoms, adjusting for overall charge.

2. Write the symbols for the atoms, connect them with single bonds.

3. Complete the octets around all atoms bonded to the central atom.

4. Place any remaining electrons on the central atom.

5. If not enough electrons for the central atom's octet, try multiple bonds.

6. Assign formal charges to atoms to determine the most stable structure.

Formal Charge Formula:

$\text{Formal charge} = \text{valence electrons} - (\text{bonding electrons}/2 + \text{nonbonding electrons})$

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

Resonance Structures

Resonance in Molecules

Some molecules cannot be accurately depicted by a single Lewis structure. Resonance structures are used to represent delocalized electrons shared among multiple atoms.

• Example: Ozone (O3) has two resonance structures, with bond lengths intermediate between single and double bonds.

• Benzene (C6H6): Six equivalent C–C bonds, depicted as a hexagon with a circle.

Delocalized electrons are shared among several atoms, not just between two.

Exceptions to the Octet Rule

Types of Exceptions

• Ions or molecules with an odd number of electrons (e.g., NO)

• Ions or molecules with less than an octet (e.g., BF3)

• Ions or molecules with more than eight valence electrons (expanded octet, e.g., PF5, phosphate ion)

Example: NO has 11 valence electrons, resulting in an odd-electron molecule.

Strengths and Lengths of Single and Multiple Bonds

Bond Enthalpy and Bond Length

Bond enthalpy is the energy required to break a bond. Multiple bonds are stronger and shorter than single bonds.

• Single Bond: Longest and weakest

• Double Bond: Intermediate length and strength

• Triple Bond: Shortest and strongest

Bond Enthalpy Calculation:

$\Delta H_{\text{rxn}} = \sum \text{(bond enthalpies of bonds broken)} - \sum \text{(bond enthalpies of bonds formed)}$

Example: For the reaction $\mathrm{CH}_4 + \mathrm{Cl}_2 \rightarrow \mathrm{CH}_3\mathrm{Cl} + \mathrm{HCl}$:

$\Delta H = [D(\mathrm{C}-\mathrm{H}) + D(\mathrm{Cl}-\mathrm{Cl})] - [D(\mathrm{C}-\mathrm{Cl}) + D(\mathrm{H}-\mathrm{Cl})]$

Numerical example: $\Delta H = (413 + 242) - (328 + 431) = 655 - 759 = -104\ \mathrm{kJ}$

Bond Type	Average Bond Enthalpy (kJ/mol)	Bond Length (pm)
C–C (single)	348	154
C=C (double)	614	134
C≡C (triple)	839	120
Additional info: Values inferred from standard textbook tables.

Summary: Understanding chemical bonding is essential for predicting molecular structure, properties, and reactivity. Mastery of these concepts is foundational for further study in chemistry.