Chemistry Atomic Structure and Bonding

Atomic Structure and Bonding

• Atoms strive for the octet rule, achieving the lowest energy state.
• Diatomic elements: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂ (Memorize the acronym "Have No Fear Of Ice Cold Beer")

Types of Chemical Bonds

• Intramolecular Bonds: Strong bonds between atoms.
• Intermolecular Bonds: Weaker bonds between molecules.

Covalent Bonds

• Formed between nonmetals by sharing electrons.
• Nonpolar Covalent: Equal sharing of electrons (similar electronegativity).
• Polar Covalent: Unequal sharing of electrons (difference in electronegativity).
  • Commonly found in all three states of matter.
  • Solubility varies; some are soluble in water.
  • Do not conduct electricity.
  • Typically have low melting points (in liquid or gaseous phases).
• Exceptions: Carbon and hydrogen always form nonpolar covalent bonds.

Ionic Bonds

• Formed between a metal and a nonmetal.
• Involve the transfer of electrons, forming cations (positive ions) and anions (negative ions).
  • The opposite charges of the ions attract each other.
• Predominantly found in the solid state.
• Often soluble in water.
• Conduct electricity when the charged particles are mobile (e.g., in solution or molten state).
• Typically have high melting points.

Factors Influencing Bond Formation

• Electronegativity: The ability of an atom to attract shared electrons in a covalent bond.
• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full outer electron shell.
• Ionization Energy (IE): Measures the energy needed to remove an electron from an atom. Corresponds with electronegativity.

Lewis Dot Structures

• Represent valence electrons of atoms.

Drawing Lewis Structures for Ions

• Depict the gain or loss of electrons outside the bracket. (Examples given: Na, Al, Cl, and N).

Drawing Lewis Structures for Covalent Bonds

• Exceptions to the Octet Rule: Hydrogen (2 valence electrons), Beryllium (4), Boron (6) Elements from period 3 and lower may have expanded octets.
• Steps:
  1. Determine the total valence electrons.
  2. Draw the skeleton structure (hydrogen is never in the center and halogens are not usually in the center)
  3. Subtract the number of electrons in bonds from the total.
  4. Add lone pairs to complete octets (except for H, B, and Be).

Energy Changes

• Endothermic Reactions: Absorb energy (+ΔH).
• Exothermic Reactions: Release energy (-ΔH).

Lattice Energy

• The energy required to separate one mole of an ionic compound into its gaseous ions.
• Higher lattice energy for compounds with higher charges or smaller ionic radii.
• Examples:
  • Comparing Lattice Energies: Identify the compound with the largest lattice energy. (Examples provided: LiF versus NaF, NaF versus MgF2,).

Energetics of Ionic Solid Formation

• The overall net energy change (sum of all energy changes) must be negative for the reaction to occur spontaneously.
• Born-Haber Cycle: A cycle that illustrates the steps in forming an ionic compound and their enthalpy changes. The cycle is used to calculate the lattice energy.
• Steps in the Cycle:
  1. Sublimation (must be gaseous form),
  2. Ionization Energy
  3. Electron Affinity
  4. Lattice Energy

Bond Lengths and Bond Order

• Bond Length: The distance between the nuclei of two bonded atoms.
• Bond Order: The number of shared electron pairs in a covalent bond.
  • Single bonds have the lowest bond order, lowest bond enthalpy, and longest bond length.
  • Triple bonds have the highest bond order, highest bond enthalpy, and shortest bond length.

Resonance Structures

• Multiple ways of drawing the same covalent bond structure.
• Often used to depict the structure that best describes a bond.

Formal Charge

• A way to determine the best resonance structure by calculating the formal charge on each atom in a possible structure
• The structure with the lowest formal charge is generally the best representation.