Chapter 8: Basic Concepts of Chemical Bonding

Types of Chemical Bonds:
  1. Ionic Bonds: Electrostatic forces that hold ions together (e.g., NaCl).
  2. Covalent Bonds: Result from sharing electrons between atoms (e.g., Cl2).
  3. Metallic Bonds: Refers to metal nuclei floating in a sea of electrons (e.g., Na).

Lewis Symbols: Represent valence electrons as dots around the element's symbol, typically arranged on four sides of a square.
Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons, enhancing stability.

Enthalpy (H):
  - Defined as the internal energy (E) plus the product of the pressure (P) and volume (V) of the system:
   $H = E + PV$
  - Change in enthalpy (ΔH) can be expressed as:
   $ΔH = H_{final} - H_{initial} = ΔE + PΔV$
  - Typically, for many reactions, $PΔV$ is small, thus:
   $ΔH ext{ (approx)} = ΔE$
Heat Transfer:
- If heat is transferred from surroundings to the system, ΔH > 0 (endothermic reaction).
- If heat is transferred from the system to the surroundings, ΔH < 0 (exothermic reaction).

Energetics of Ionic Bond Formation:
  - The formation reaction for NaCl:
   $Na(s) + rac{1}{2}Cl_2(g) <br>ightarrow NaCl(s) ext{ with } ΔH_{reaction} = -410.9 ext{ kJ/mol (exothermic)}$
  - Dissociation of NaCl is endothermic:
   $NaCl(s) <br>ightarrow Na^+(g) + Cl^-(g) ext{ with } ΔH = 788 ext{ kJ/mol}$
Lattice Energy:
  - Defined as the energy required to separate one mole of a solid ionic compound into gaseous ions.
  - Depends on the charges of the ions and their sizes.
  - Formula:
   $E = rac{k imes Q_1 imes Q_2}{d}$
  - Where:
    - E = potential energy of interacting charged particles
    - $Q_1, Q_2$ = charges on the particles
    - $d$ = distance between their centers
    - $k$ = constant (8.99 × 10^9 J.m/C²)
Lattice Energies for Selected Ionic Compounds:
  - Table:
    | Compound | Lattice Energy (kJ/mol) |
    |----------|-------------------------|
    | LiF | 1030 |
    | NaCl | 788 |
    | MgO | 3795 |
    | CaO | 3414 |
    | … | … |
  - General observation: As $Q_1$ and $Q_2$ increase, E increases, whereas as $d$ increases, E decreases.

Definition: A chemical bond formed by sharing a pair of electrons (e.g., H-H).
Characteristics:
- Bonds show electrostatic interactions including attractions between electron and nuclei, and repulsion between electrons and nuclei.
- Representation: Shared electron pairs are shown as lines (single bonds), while unshared pairs are shown as dots.
Multiple Bonds:
  - Covalent bonds can involve sharing more than one pair of electrons:
    - Single Bond: H-H
    - Double Bond: O=O
    - Triple Bond: N≡N

Bond Polarity: Refers to the distribution of electron density in a bond.
  - Types:
    1. Nonpolar Covalent Bond: Electrons shared equally (e.g., F2).
    2. Polar Covalent Bond: Electrons shared unequally (e.g., HF).
    3. Ionic Bond: When electron sharing difference is significant (e.g., NaBr).
Electronegativity:
  - Range from 0.7 (Cs) to 4.0 (F) on the Pauling scale; increases across a period and decreases down a group.
  - Example for Polar Covalent Bond: In HF:
   $ext{Electronegativity Difference} = 4 - 2.1 = 1.9 ext{ (polar covalent bond)}$
Differences in Electronegativity and Polarity:
  - Nonpolar covalent bonds: difference close to zero.
  - Polar covalent bonds: difference ranging between 0.5 and 1.9.
  - Ionic bonds: difference greater than 2.0.

Guidelines for Drawing Lewis Structures:
  1. Sum the valence electrons from all atoms (adding or subtracting for any charge).
  2. Write symbols for atoms and connect them with single bonds.
  3. Identify the central atom, often the least electronegative one.
  4. Complete octets for all atoms (except H, which requires only 2 electrons).
  5. Place leftover electrons on the central atom.
  6. Use lone pairs on terminal atoms to form multiple bonds when additional electrons for octet.
  7. Calculate formal charges to ensure the most stable structure.
Formal Charge Calculation:
ext{Formal Charge} = ext{# of valence e-} - ( ext{# of lone pair e-} + rac{1}{2} imes ext{# of bonding e-})
Stability Criteria: Aim for structures with formal charges close to zero, with negative charges on the most electronegative atoms.

Definition: Molecules can exist in multiple forms called resonance structures.
Example: Ozone (O3) can be represented as having one single bond and one double bond, but experimentally both bonds are identical. Hence, the actual structure is an average of its resonance structures.
Other examples include NO3– and SO3.

Categories:
  1. Odd Number of Electrons: Molecules such as ClO2, NO, and NO2 cannot complete octet on all atoms.
  2. Less than Octet: Common in compounds with elements from Groups 1A, 2A, and 3A (e.g., BF3).
  3. More than Octet: Atoms from period 3 or greater can have expanded octets (e.g., PF5, SF4).

Bond Entropy (D): The energy required to break a particular covalent bond in one mole of a substance:
   $D(C–H) = rac{ΔH}{4}$
  - For example, for CH4:
   $ΔH_{reaction} = 1660 ext{ kJ for all bonds broken}$
  - Therefore, bond strengths can be calculated using average bond enthalpies.
Average Bond Enthalpies:
  - Represent the average amount of energy required to break bonds across different compounds, reflecting bond strengths.
  - Example:
    | Single Bonds | Value (kJ/mol) |
    |----------------------|-----------------|
    | C-H | 413 |
    | N-H | 391 |
    | O-H | 463 |
    | … | … |
  - A more extensive list of average bond enthalpies for various bonds can be referenced.
Bond Length and Bond Strength:
- Generally: The more bonds between two atoms, the stronger the bond and lesser the distance (bond length decreases as bond order increases).