CHEM LECTURE

Types of Chemical Bonds

• Ionic Bonds

  • Characterized by the exchange of electrons between atoms.

  • One atom (the lender) loses electrons to become positively charged (cation), while the other (the receiver) gains electrons to become negatively charged (anion).

  • Example: Sodium (Na) loses an electron to become Na+, and Chlorine (Cl) gains an electron to become Cl-.

• Covalent Bonds

  • Formed when two atoms cooperate by sharing electrons.

  • Can range from single bonds (one pair of shared electrons) to double bonds (two pairs) to triple bonds (three pairs).

  • Example: Hydrogen (H) and Chlorine (Cl) share one electron each to form HCl.

Lewis Structures

• A visual representation of the valence electrons in a molecule.

• Important for illustrating how atoms bond through ionic or covalent interactions.

• Formal Charge

  • A method to determine the charge on an atom in a molecule based on its bonding and lone pair electrons.

  • Calculated by:

    • Formal Charge = (Valence Electrons) - (Lone Pair Electrons) - (Bonded Pairs / 2)

Electron Configuration and Stability

• Atoms strive for a stable electron configuration, often seeking a noble gas configuration (8 electrons in the valence shell).

• Octet Rule

  • Atoms generally prefer to have eight electrons in their valence shell (with some exceptions for hydrogen and helium).

Electronegativity and Polarity

• Electronegativity

  • A measure of an atom's ability to attract and hold onto electrons.

  • Fluorine (F) is the most electronegative element, while Francium (Fr) is among the least.

• Polar Covalent Bonds

  • Occur when two atoms with different electronegativities share electrons unequally, leading to a slightly charged end of the molecule.

  • Example: HCl is polarized due to the higher electronegativity of Cl.

Ionic vs. Covalent Bonds

• Ionic bonds generally form when the electronegativity difference is 1.8 or greater.

• Covalent bonds form when the difference is 0.4 or less.

• Polar Covalent Bonds

  • Have a difference between 0.4 and 1.8, exhibiting characteristics of both ionic and covalent bonds.

Practical Applications

• When studying bonds and molecular structures, it is crucial to determine the best Lewis structure by following a systematic method:

  1. Count total valence electrons.

  2. Create a skeletal structure (least electronegative in the center).

  3. Assign electrons to terminal atoms first, then central atoms.

  4. If any electrons remain, consider them for multiple bonds if necessary to ensure stable configurations.

Summary of Bond Formation

• Cations and anions form through ionic exchanges, while covalent bonds create stable molecules by sharing electrons.

• Formal charge calculations confirm stability and appropriate representation of molecules through their Lewis structures.

• Understanding electronegativity differences plays a critical role in determining bond types and molecular behavior.