General and Analytical Chemistry I Lecture 11

Covalent Bond: A bond that results from the sharing of electrons between atoms.
Formation of bonds can occur from a head-to-head overlap of atomic orbitals on neighboring atoms. For example, hydrogen (H) and chlorine (Cl) form covalent bonds through the overlap of H (1s) and Cl (2p) orbitals.
A Molecule: The unit of matter held together by covalent bonds, characterized by each atom having a single, unpaired electron.

Molecular Compounds: Compounds whose particles are molecules that consist only of nonmetals.
Ionic Compounds: Compounds formed by particles that are cations (positively charged) and anions (negatively charged).
Atomic Elements: Elements whose particles consist of single atoms.
Molecular Elements: Elements whose particles are multi-atom molecules.

Covalent bonds can form between both unlike and like atoms, leading to a broad variety of molecular compounds.
Most covalent molecules achieve a noble gas configuration by sharing enough electrons.
Water (H₂O) is an example where two hydrogen atoms and one oxygen atom are covalently bonded.

Valence Electrons and Bonding:
- 3A: 3 bonds (e.g., BH₃)
- 4A: 4 bonds (e.g., CH₄)
- 5A: 3 bonds (e.g., NH₃)
- 6A: 2 bonds (e.g., H₂O)
- 7A: 1 bond (e.g., HF)
- 8A: 0 bonds (Ne)
For elements in the third period and below (P, S, Cl), the number of covalent bonds may vary and can include expanded octets.

Single Bond: Formed by sharing two electrons or one pair; represented by a single line between atoms.
Double Bond: Formed by sharing four electrons or two pairs; represented by two lines between atoms.
Triple Bond: Formed by sharing six electrons or three pairs; represented by three lines between atoms.

Ionic Compounds: Ions are arranged in a crystal lattice.
Table of Physical Properties:
- NaCl:
- Physical appearance: White solid
- Type of bond: Ionic
- Melting point: 801 °C
- Boiling point: 1465 °C
- HCl:
- Physical appearance: Colorless gas
- Type of bond: Covalent
- Melting point: -115 °C
- Boiling point: -84.9 °C

The bond formation involves attractive and repulsive forces:
- Attraction: The nucleus-electron attractions outweigh the nucleus-nucleus and electron-electron repulsions, creating a net attractive force.
- Bond Length: The optimum distance between nuclei in a covalent bond, indicated as 74 picometers (pm) for H-H bonds.
If atoms are too close, strong repulsions occur; too far apart leads to weak attractions and no bonding.

Bond Energy: The amount of energy needed to break one mole of a bond in a compound. It is measured homolytically in the gas state.
Reaction Enthalpy (rxn): Calculated by comparing the energy required to break bonds against the energy released when new bonds are formed:
- \Delta H_{rxn} = (\text{energy required to break bonds}) - (\text{energy evolved when bonds are made})
Trends in Bond Energies:
- More electrons shared leads to stronger bonds:
- For example, C \equiv C (837 kJ) > C = C (611 kJ) > C - C (347 kJ)
- Shorter covalent bonds are stronger: e.g., Br - F (237 kJ) > Br - Cl (218 kJ) > Br - Br (193 kJ)

Table of Average Bond Dissociation Energies (kJ/mol):
- H-H: 436
- C-H: 410
- N-H: 390
- O-F: 180
- I-I: 151
- H-F: 570
- C≡C: 728
- C=O: 732

Covalent bonding between unlike atoms leads to unequal electron sharing, creating a polar covalent bond.
In a polar bond, one end has higher electron density (partially negative charge) and the other has lower electron density (partially positive charge)
- Example: H-Cl ext{ has } d+ ( ext{H}) ext{ and } d- ( ext{Cl}),
Electronegativity: A measure of an atom's ability to attract bonding electrons, proposed by Linus Pauling:
- Increases across periods (left to right) and decreases down groups (top to bottom).
- Most electronegative: Fluorine
- Least electronegative: Francium
Polar bonds are influenced by the difference in electronegativity:
- Pure covalent (0), nonpolar covalent (0.1 - 0.4), polar covalent (0.5 - 1.9), ionic (≥ 2.0).

Dipole Moment: A quantitative measure of bond polarity, dependent on the size of partial charges and the distance between them. Measured in Debyes (D).
The percent ionic character indicates the bond's dipole moment relative to that of full ions.
Examples of dipole moments for various molecules in gas phase:
- Cl₂: 0
- HF: 1.9
- LiF: 3.0

Naming binary molecular compounds involves:
- The combination of elements from groups 4A–7A with each other or with hydrogen.
- Utilize prefix rules:
1. Mono- (note: not used on the first nonmetal)
2. Di-
3. Tri-
4. Tetra-
5. Penta-
6. Hexa-
7. Hepta-
8. Octa-
9. Nona-
10. Deca-
Example - Naming Binary Molecular Compounds:
- BF₃:
- Name first element: Boron.
- Name second element with -ide suffix: Fluorine ➔ Fluoride.
- Prefix: Monoboron, Trifluoride (drop mono from the first element).
More examples include:
- NO₂ ➔ Nitrogen Dioxide,
- PCl₅ ➔ Phosphorus Pentachloride,
- I₂F₇ ➔ Diiodine Heptafluoride.

Representation of valence electrons as shared (bond pairs) or unshared (lone pairs).
Octet Rule: Most elements binding to achieve 4 pairs (8 electrons) maximum.
Common bonding patterns include:
- Carbon: 4 bonds, 0 lone pairs.
- Nitrogen: 3 bonds, 1 lone pair.
- Oxygen: 2 bonds, 2 lone pairs.

Identify the Central Atom: Usually the least electronegative excluding hydrogen.
Count Valence Electrons: Total up considering all atoms and charges.
Skeletal Structure: Form bonds between central and surrounding atoms.
Complete Octets: Distribute remaining electrons to fulfill octet rule as needed.

Examples of Lewis Structures:
- Ammonia (NH₃), Carbon Dioxide (CO₂): Both demonstrate the bonding patterns for proper structure representation.
From this guidance, students should be able to articulate the electron distribution and bond properties, predict molecular structures, and determine the physical and chemical properties based on bonding interactions.

Questions: Building structures for various molecules, including CO₂, SeOF₂, H₃PO₄, and SO₃, to ensure understanding of bonding principles and octet completion.