Chapter 8: Basic Concepts of Chemical Bonding – Study Notes

Basic Concepts of Chemical Bonding

Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. There are three primary types of chemical bonds, each with distinct characteristics and mechanisms.

• Ionic Bonds: Formed by electrostatic attraction between oppositely charged ions (typically between metals and nonmetals).

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

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

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

8.1 Lewis Symbols and the Octet Rule

Lewis symbols are a simple way to represent valence electrons around an atom. The octet rule states that atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons, achieving a noble gas configuration.

• Lewis Symbol: Dots around an element symbol represent valence electrons.

• Octet Rule: Atoms form compounds to achieve eight valence electrons.

Example: Oxygen (O) has six valence electrons, so its Lewis symbol is O with six dots.

8.2 Ionic Bonding

Ionic bonding occurs between metals and nonmetals (except group 8A), involving the transfer of electrons from one atom to another. This process is highly exothermic.

• One atom (metal) loses electrons (low ionization energy).

• Another atom (nonmetal) gains electrons (high electron affinity).

• Electron transfer is indicated by arrows in reaction equations.

Example: $\mathrm{Na} + \mathrm{Cl} \rightarrow \mathrm{Na}^+ + [\cdot\cdot\cdot\mathrm{Cl}\cdot\cdot\cdot]^-$

Properties of Ionic Substances

• Brittle

• High melting points

• Crystalline structure

• Cleave along smooth lines

Example: NaCl forms a cubic crystal lattice.

Energetics of Ionic Bonding

The formation of ionic compounds involves several energy changes, summarized by the Born–Haber cycle.

• Energy required to convert elements to atoms (endothermic)

• Energy required to create a cation (endothermic)

• Energy released to form an anion (exothermic)

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

Example: Formation of NaCl from Na(s) and Cl2(g) is overall exothermic.

Lattice Energy

Lattice energy is a measure of the stability of an ionic solid, defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.

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

• Calculated using the Born–Haber cycle:

$\Delta H_f^\circ(s) = \Delta H_f^\circ[\mathrm{M}(g)] + \Delta H_f^\circ[\mathrm{NM}(g)] + IE(\mathrm{M}) + EA(\mathrm{NM}) - \Delta H_{\text{lattice}}(s)$

Where M = metal, NM = nonmetal.

Table: Lattice Energies for Some Ionic Compounds

Compound	Lattice Energy (kJ/mol)	Compound	Lattice Energy (kJ/mol)
LiF	1030	MgCl2	2526
LiCl	834	SrCl2	2127
LiI	730	MgO	3795
NaF	910	CaO	3414
NaCl	788	SrO	3217
NaBr	732	ScN	7547
NaI	682
KF	808
KCl	701
KBr	671
CsCl	657
CsI	600

Trends in Lattice Energy

• Lattice energy increases with increasing charge on the cations: $+1 < +2 < +3$

• Lattice energy increases with decreasing size of the ions: $r(\mathrm{Li}^+) < r(\mathrm{Cs}^+)$

$E_{el} = \dfrac{\lambda Q_1 Q_2}{d}$

Electron Configuration of Ions

The electron configuration of ions depends on whether the atom is a main group metal, nonmetal, or transition metal.

• 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 lose valence electrons first (highest n), then d-electrons as needed for the ion charge.

8.3 Covalent Bonding

Covalent bonding involves the sharing of electrons between atoms, typically 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: $\mathrm{H}-\mathrm{H}$ (hydrogen molecule), $\cdot\cdot\cdot\mathrm{Cl}-\mathrm{Cl}\cdot\cdot\cdot$ (chlorine molecule)

Number of Bonds for Nonmetals

The number of bonds a nonmetal forms is determined by the number of electrons needed to complete its octet.

• Group number = number of valence electrons

• Number of bonds = electrons needed to complete octet

Example:

• F (Group 7A, 7 valence electrons) needs 1 bond

• O (Group 6A, 6 valence electrons) needs 2 bonds

• N (Group 5A, 5 valence electrons) needs 3 bonds

• C (Group 4A, 4 valence electrons) needs 4 bonds

Electrons Using Lewis Structures

• Lone pairs: Unshared pairs of electrons located on one atom

• Bonding pairs: Shared electrons between two atoms (represented by two dots or a line)

Multiple Bonds

• Single bond: One pair of shared electrons

• Double bond: Two pairs of shared electrons

• Triple bond: Three pairs of shared electrons

Example: $\mathrm{O}=\mathrm{C}=\mathrm{O}$ (CO2), $\mathrm{N}\equiv\mathrm{N}$ (N2)

8.4 Polarity of Bonds and Electronegativity

Electrons in covalent bonds are not always shared equally. Bond polarity measures the equality of electron sharing.

• Nonpolar covalent bond: Electrons shared equally

• Polar covalent bond: One atom attracts electrons more strongly

Example: In HF, fluorine attracts electrons more strongly than hydrogen, making the bond polar.

Electronegativity

Electronegativity is the ability of an atom in a molecule to attract electrons to itself. On the periodic table, electronegativity increases from left to right across a period and from bottom to top within a group.

Electronegativity and Polar Covalent Bonds

• Unequal sharing of electrons leads to a polar covalent bond.

• The more electronegative atom gains a partial negative charge ($\delta^-$), while the other atom becomes partially positive ($\delta^+$).

Example: $\delta^+\mathrm{H}-\delta^-\mathrm{F}$

Table: Bond Lengths, Electronegativity Differences, and Dipole Moments of Hydrogen Halides

Compound	Bond Length (Å)	Electronegativity Difference	Dipole Moment (D)
HF	0.92	1.9	1.82
HCl	1.27	0.9	1.06
HBr	1.41	0.7	0.82
HI	1.61	0.4	0.44

Dipoles

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

$\mu = Qr$

where $Q$ is the charge and $r$ is the separation distance. Dipole moments are measured in debyes (D).

Comparing Ionic and Covalent Bonding

• Ionic: Complete electron transfer (metal + nonmetal)

• Covalent: Electron pair sharing (two nonmetals)

• Electronegativity difference > 2.0 is often ionic

• Exceptions: High oxidation numbers can lead to covalent character

• Physical properties (e.g., melting points) can help distinguish bond type

8.5 Drawing Lewis Structures

Lewis structures are drawn using a systematic approach:

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

2. Write the symbols for the atoms, show connections with single bonds.

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

4. Place remaining electrons on the central atom.

5. If the central atom lacks an octet, form multiple bonds as needed.

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

Formal charge formula:

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

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

The Best Lewis Structure and Resonance

Some molecules, like ozone (O3), cannot be accurately depicted by a single Lewis structure. Both O–O bond lengths are the same, indicating resonance.

8.6 Resonance Structures

Resonance structures are multiple valid Lewis structures used to describe molecules where electron delocalization occurs. The true structure is a hybrid of all resonance forms.

• Ozone (O3): Two resonance structures, bond lengths intermediate between single and double bonds.

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

Localized electrons: On one atom or shared between two atoms. Delocalized electrons: Shared by multiple atoms.

*Additional info: Resonance and delocalization are further explored in Chapter 9.*