Recording-2025-03-04T22:02:31.251Z

Electronegativity

• Definition: Electronegativity is the tendency of an atom to attract electrons in a chemical bond.

• Trends:

  • Increases across a period from left to right.

  • Decreases down a group.

Covalent Bonds and Electronegativity

• Covalent Compounds: Atoms share electrons; however, they do not always share equally.

• Example:

  • Chlorine (Cl):

    • Chlorine gas (Cl2) has an electronegativity of 3.

    • When two chlorine atoms bond, the difference in electronegativity is 3 - 3 = 0, indicating nonpolar covalent bonds due to equal sharing of electrons.

• Polar Covalent Bonds:

  • Hydrochloric Acid (HCl):

    • Chlorine is 3, and hydrogen is 2.1.

    • Difference = 3 - 2.1 = 0.9, indicating a polar covalent bond where electrons are drawn closer to chlorine, giving it a partial negative charge and hydrogen a partial positive charge.

Ionic Bonds

• Sodium Chloride (NaCl):

  • Example of an ionic bond where electronegativity difference > 2 indicates that one atom loses an electron while the other gains it.

Dipoles

• Definition: A molecule with a positive end and a negative end due to unequal sharing of electrons in a polar bond.

• Example: In HCl, Cl is the negative end and H is the positive end.

The Octet Rule

• Basic Principle: Atoms tend to form bonds until they are surrounded by eight valence electrons, achieving a stable electron configuration.

• Exceptions to the Octet Rule:

  • Hydrogen: Forms a single bond and is stable with only 2 electrons.

  • Lithium & Beryllium: Lithium forms one bond, Beryllium typically forms two.

  • Boron: Usually makes three bonds.

  • Atoms in period 3 or higher can have expanded octets.

Drawing Lewis Structures

1. Identify the Central Atom:

   • The least electronegative element is often the central atom.

   • Note: Hydrogen is never a central atom.

2. Count Valence Electrons:

   • Add valence electrons from all atoms involved.

3. Connect Atoms:

   • Form bonds between the central atom and outer atoms, using the counted valence electrons.

4. Stabilize Outer Atoms:

   • Ensure outer atoms have full octets (2 for H).

5. Adjust for Central Atom:

   • If central atom is not stable after assigning all available electrons, form double or triple bonds as needed.

Example: Carbon Dioxide (CO2)

• Steps to Draw:

  1. Central atom: Carbon (least electronegative).

  2. Valence electrons: Carbon = 4, oxygen = 6 (x2) = 16 total.

  3. Bonds made: 2 (4 electrons used).

  4. Oxygen atoms: Each requires 4 more electrons (8 total each).

  5. Need to stabilize carbon by forming double bonds with each oxygen, giving it a satisfied octet.

Formal Charge Methodology

• Definition: Charge of an atom in a molecule as calculated from the known number of valence electrons.

• Procedure to Calculate:

  • Count valence electrons.

  • Subtract unshared electrons and half of bonded electrons.

  • Goal: Ideal formal charges for stability (ideally zero).

Conclusion

• Understanding electronegativity and bond types (covalent vs ionic, polar vs nonpolar) is crucial for predicting molecular behavior.

• The octet rule aids in predicting stability and configurations, while formal charge helps refine structures for the most stable arrangement.