Chapter 10: Basic Concepts of Chemical Bonding – Study Notes

Basic Concepts of Chemical Bonding

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in chemical bonding and reactivity. Their arrangement determines how an element interacts with other atoms.

• Valence electrons are those in the highest principal quantum number (n) and any in partially filled d and f orbitals.

• Core electrons are inner electrons that are not involved in bonding.

• Valence electrons are more available for reactions because they are farther from the nucleus.

Types of Chemical Bonds

Atoms are held together by chemical bonds, which can be classified into two main types:

• Ionic Bond: Attraction between a cation (positively charged ion) and an anion (negatively charged ion). Electrons are transferred from one atom to another.

• Covalent Bond: Sharing of a pair of electrons, with one electron contributed by each atom.

Example: In sodium chloride (NaCl), Na+ and Cl- are held together by ionic bonds. In H2, two hydrogen atoms share electrons via a covalent bond.

Bonding Electrons and Lewis Symbols

Lewis symbols are a simple way to represent valence electrons and predict bonding behavior.

• Each dot around an element symbol represents a valence electron.

• Lewis symbols help visualize how atoms bond in molecules.

Lewis Symbols of Atoms

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

The Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full set of eight valence electrons, resulting in greater stability.

• Most stable electron configurations have eight outer shell electrons (an octet).

• Lewis symbols are designed to accommodate up to eight electrons (four pairs).

• There are exceptions to the octet rule (see below).

Ionic Bonding

Ionic bonding involves the transfer of electrons from one atom to another, forming ions that are held together by electrostatic attraction.

• Formation of ions requires electron transfer.

• Oppositely charged ions attract to form an ionic bond.

• Lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions.

Formula for Lattice Energy (Coulomb's Law):

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

• $k$ = Coulomb's constant ($8.99 \times 10^9$ J·m/C2)

• $Q_1$, $Q_2$ = charges on the ions

• $d$ = distance between nuclei

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	SO	3217
NaBr	732	SeN	7547
NaI	682
KF	808
KCl	701
KI	649
CaCl2	657
CsI	600

Trends: Lattice energy increases with higher ionic charges and decreases with larger ionic radii.

The Born-Haber Cycle

The Born-Haber cycle is a thermodynamic cycle used to analyze the steps in the formation of an ionic compound and to calculate lattice energy.

• It involves ionization energy, electron affinity, sublimation, and bond dissociation energies.

• Helps relate the lattice energy to measurable quantities.

The Covalent Bond

Covalent bonds form when two atoms share one or more pairs of electrons, resulting in a stable balance of attractive and repulsive forces between atoms.

• Electron density is concentrated between the nuclei.

• Bond strength and length depend on the number of shared electron pairs.

Covalent Bond Enthalpies

Bond enthalpy is the energy required to break one mole of a bond in a molecule in the gas phase.

Bond	Bond Enthalpy (kJ/mol)	Bond	Bond Enthalpy (kJ/mol)
C–H	413	O–H	463
C–C	348	N–H	391
C–N	305	O–O	146
C–O	358	N–N	163
C–F	485	O=O	498
C=O	799	N≡N	941

Covalent Bond Lengths

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

Bond	Bond Length (Å)	Bond	Bond Length (Å)
C–C	1.54	N–N	1.47
C=C	1.34	N=N	1.24
C≡C	1.20	N≡N	1.10
C–N	1.47	N–O	1.36
C=N	1.30	N=O	1.22
C≡N	1.16	O=O	1.21
C–O	1.43	C=O	1.23
C≡O	1.13

Electronegativity

Electronegativity (χ) is the ability of an atom in a molecule to attract electrons to itself.

• Higher electronegativity means stronger attraction for electrons.

• Lower electronegativity means a greater tendency to lose electrons.

• The difference in electronegativity between two atoms helps determine bond type.

Ionic or Covalent?

Bonds exist on a continuum from non-polar covalent to ionic, depending on the electronegativity difference (Δχ):

• Δχ ≈ 0: Non-polar covalent

• 0 < Δχ < 2: Polar covalent

• Δχ ≥ 2: Ionic

Formula:

$\Delta \chi = |\chi_1 - \chi_2|$

Electronegativity Trends

• Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

Different Types of Bonds

The type of bond (ionic, polar covalent, non-polar covalent) depends on the electronegativity difference between the atoms involved.

• HF (Δχ = 1.9): More ionic character

• HCl (Δχ = 0.9): Polar covalent

• HBr (Δχ = 0.7): Less polar covalent

• HI (Δχ = 0.4): More like non-polar covalent

Dipole Moments

Dipole moments arise in molecules with polar covalent bonds due to unequal sharing of electrons.

• One atom attracts electrons more strongly, resulting in partial charges (δ- and δ+).

• Dipole moment (μ) is a measure of bond polarity:

$\mu = Qr$

• Q = magnitude of charge, r = distance between charges

Trends for the Hydrogen Halides

Compound	Bond Length (Å)	Electronegativity Difference	Dipole Moment (D)
HF	0.92	1.9	1.82
HCl	1.27	0.9	1.08
HBr	1.41	0.7	0.82
HI	1.61	0.4	0.44

Lewis Structures

Lewis structures represent the arrangement of electrons in molecules and ions, showing both bonding and non-bonding (lone pair) electrons.

• Steps to draw Lewis structures:

  1. Add up all valence electrons from the atoms (add for anions, subtract for cations).

  2. Arrange the atoms and connect them with single bonds.

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

  4. Place any leftover electrons on the central atom.

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

• Use formal charge to help determine the most stable structure.

Formal Charge

Formal charge helps identify the most stable Lewis structure by minimizing charges on atoms.

• Formal charge = (valence electrons) – (non-bonding electrons + ½ bonding electrons)

• Sum of formal charges should equal the overall charge of the molecule or ion.

• Central atom can only have four bonds or lone pairs.

• For ions, place the charge outside square brackets.

Examples: Nitrate (NO3-), Acetone ((CH3)2CO), Thiocyanate (NCS-), Bromate (BrO3-), Ozone (O3).

Resonance Structures

Resonance structures are different Lewis structures for the same molecule or ion that have the same arrangement of atoms but different arrangements of electrons.

• Resonance adds stability to the molecule or ion.

• Actual structure is a hybrid of all resonance forms.

• Example: Ozone (O3) has two resonance structures, and both O–O bonds are the same length (127.8 pm).

Exceptions to the Octet Rule

Some molecules and ions do not follow the octet rule:

• Odd number of electrons: e.g., ClO2, NO, NO2

• Less than an octet: e.g., BF3

• More than an octet: e.g., PF5, PO43-, ICl4-, SF4, AsF6-

• Atoms with more than an octet must be large enough to accommodate extra electron pairs.

Additional info: The above notes expand on the provided slides and text, adding definitions, formulas, and examples for clarity and completeness.