Chapter 8: Basic Concepts of Chemical Bonding – Study Notes

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 properties of substances.

• Metallic bonds: Formed by electrons that are relatively free to move throughout the metal lattice (e.g., Fe in a body-centered cubic lattice).

• Ionic bonds: Result from the electrostatic attraction between oppositely charged ions (e.g., NaCl).

• Covalent bonds: Involve the sharing of electron pairs between atoms (e.g., H2O).

8.2 Ionic Bonding

Formation of Ionic Bonds

Ionic bonds are formed when one atom transfers electrons to another, resulting in the formation of oppositely charged ions. This typically occurs between metals (which lose electrons easily) and nonmetals (which gain electrons easily).

• Electron transfer: One atom with low ionization energy loses electrons, and another with high electron affinity gains them.

• Example: $\text{Na}(s) + \frac{1}{2}\text{Cl}_2(g) \rightarrow \text{NaCl}(s)$ $\Delta H_f = -410.9\ \text{kJ}$

• Each ion achieves a stable octet configuration (Na+: [Ne], Cl-: [Ar]).

Characteristics of Ionic Substances

• Brittle

• High melting points

• Usually crystalline

• Can be cleaved along planes

Energetics of Ionic Bond Formation

• Lattice Energy (U): The energy required to completely separate one mole of a solid ionic compound into its gaseous ions.

• When ions come together, energy is released and a solid forms.

• Lattice energy is a measure of the stability of the ionic solid.

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

Lattice Energies for Some Ionic Compounds

Compound	Lattice Energy (kJ/mol)	Compound	Lattice Energy (kJ/mol)
LiF	1030	MgCl2	2326
LiCl	834	SrCl2	2127
NaF	910	MgO	3795
NaCl	788	CaO	3414
KF	808	SeN	7547
KCl	701	CaI2	600

Qualitatively Determining Relative Lattice Energies

The magnitude of lattice energy depends on the charges of the ions, their sizes, and their arrangement in the solid. The lattice energy can be estimated by:

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

• Lattice energy increases as the charges on the ions increase.

• Lattice energy increases as the ionic radii decrease.

Example: Arrange NaF, CsI, and CaO in order of increasing lattice energy. (Hint: Consider charge and ionic radius.)

8.3 Covalent Bonding

Formation of Covalent Bonds

A covalent bond is formed by the sharing of a pair of electrons between two atoms. The shared electron density is concentrated between the nuclei, holding the atoms together.

• Lewis symbols can be used to represent the formation of covalent bonds.

• Hydrogen is an exception to the octet rule, achieving a stable two-electron configuration like helium.

Lewis Symbols and the Octet Rule

The electrons involved in chemical bonding are the valence electrons, those in the outermost shell. Lewis electron-dot symbols represent valence electrons as dots around the element symbol.

• Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (full s and p subshells).

• Only main-group (s- and p-block) elements are represented by Lewis-dot symbols.

Group	Element	Electron Configuration	Lewis Symbol
1A	Li	[He]2s1	Li·
2A	Be	[He]2s2	Be··
3A	B	[He]2s22p1	B···
4A	C	[He]2s22p2	C····
5A	N	[He]2s22p3	N·····
6A	O	[He]2s22p4	O······
7A	F	[He]2s22p5	F·······
8A	Ne	[He]2s22p6	Ne········

Multiple Bonds and Bond Strength

• Atoms can share more than one pair of electrons to achieve octets.

• Double bonds involve two shared pairs; triple bonds involve three.

• Bond length decreases and bond strength increases as the number of shared pairs increases.

Bond	Bond Length (Å)
N—N (single)	1.47
N=N (double)	1.24
N≡N (triple)	1.10

8.4 Bond Polarity and Electronegativity

Bond Polarity

Bond polarity describes how equally electrons are shared in a covalent bond.

• Nonpolar covalent bond: Electrons are shared equally (e.g., Cl2, N2).

• Polar covalent bond: Electrons are shared unequally; one atom attracts electrons more strongly (e.g., HF).

Electronegativity

Electronegativity is the ability of an atom in a molecule to attract electrons to itself. It increases across a period and decreases down a group in the periodic table.

• Electronegativity values are used to predict bond type and polarity.

Using Electronegativity Differences

ΔEN	Prediction
< 0.5	Non-polar covalent
0.5 – 1.6	Polar covalent
1.6 – 2.0	Polar covalent (if both nonmetals); Ionic (if a metal is present)
> 2.0	Ionic

• Example: C–O bond in CO2: ΔEN = 1.0 (polar covalent); Mg–Cl bond in MgCl2: ΔEN = 1.8 (ionic).

Partial Charges

• In polar bonds, the more electronegative atom acquires a partial negative charge (δ–), and the less electronegative atom a partial positive charge (δ+).

• Example: In HF, F is δ– and H is δ+.

8.5 Drawing Lewis Structures

Steps for Drawing Lewis Structures

1. Sum the valence electrons for all atoms.

2. Write the symbols for the atoms, showing which are attached to which, and connect them with single bonds.

3. Complete the octets around all outer atoms (except hydrogen).

4. Place any leftover electrons on the central atom.

5. If the central atom does not have an octet, try forming multiple bonds.

• Practice: Draw Lewis structures for PCl3, CH2Cl2, HCN, NO+, C2H4, BrO3-, ClO2-, PO43-, H2CO, H2O2, C2F6, AsO33-, H2SO3, C2H2.

Formal Charge

• Formal charge helps determine the most stable Lewis structure when multiple are possible.

• Formal charge = (Valence electrons) – (Nonbonding electrons + 1/2 Bonding electrons)

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

• The sum of formal charges equals the overall charge of the molecule or ion.

Example: For CO2, the structure with formal charges of zero on all atoms is dominant.

8.6 Resonance Structures

Resonance

Some molecules cannot be adequately represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures that differ only in the placement of electrons, not atoms.

• The actual structure is a resonance hybrid, a blend of all resonance forms.

• Resonance stabilizes the molecule and predicts equal bond lengths where resonance occurs.

• Example: SO3 and the carbonate ion (CO32-) have multiple resonance structures.

8.7 Exceptions to the Octet Rule

Types of Exceptions

1. Molecules and ions with an odd number of electrons: Complete pairing is impossible (e.g., NO, NO2).

2. Molecules and ions with fewer than an octet: Common for compounds of boron (B) and beryllium (Be), which can be stable with less than eight electrons (e.g., BF3, BeF2).

3. Molecules and ions with more than an octet (hypervalent): Central atoms from period 3 or below can accommodate more than eight electrons (e.g., SF6, XeF4).

Summary Table: Types of Octet Rule Exceptions

Type	Example	Description
Odd number of electrons	NO	Cannot pair all electrons
Fewer than octet	BF3	Boron stable with 6 electrons
More than octet	SF6	Central atom has 12 electrons

Additional info: The ability to expand the octet is due to the availability of d orbitals in elements from period 3 and beyond.