CHM111: Pre-Organic Chemistry

Oxygen vs. Sulfur
- Oxygen (Z=8) has a filled 2p orbital and lacks d orbitals in the same energy level.
- Sulfur (Z=16) has a filled 3p orbital and possesses a 3d orbital that is capable of holding electrons.
Electron Capacity of Orbitals
  - Once the 3d orbital is filled, no additional electrons will typically occupy it due to the high energy required, thus limiting the possible oxidation states of sulfur and elements below it.
  - Charge states above 6 (e.g., -10) are impractical for these elements due to high electron repulsion.
  - The capacity for holding electrons in expanded octets tends to scale with the size of the atom.

d orbitals are often introduced starting from the transition metals (row 4), not because they start there, but because that’s when electrons are typically added to those orbitals.
Elements such as phosphorus and sulfur have d orbitals but may not utilize them significantly in bonding.

Electronegativity: The ability of an atom to attract electrons when involved in a covalent bond. It quantifies the tendency of an atom to attract a bonding pair of electrons.
- Key Point: For any atom, electronegativity is defined as its ability to pull an electron toward itself when in a covalent bond.
- Most Electronegativity Atom: Fluorine (F)
General Trend: Electronegativity increases across a period (left to right) and decreases down a group in the periodic table.
- Fluorine: Positioned to gain an electron easily due to its small size and full p orbitals after gaining one more electron, achieving a noble gas configuration.
- Noble Gases: Noble gases like Ne or He do not seek additional electrons as they already have a complete octet of valence electrons.

Single Covalent Bond: Consists of two electrons shared between atoms.
Types of Bonds:
  - Covalent Bonds: between two non-metals.
  - Polar Covalent Bonds: between non-metals with significant electronegativity differences.
  - Ionic Bonds: between metals and non-metals.

Total Valence Electrons: Count the total number of valence electrons considering charge.
   - For example, in Methanol (CH₃OH):
     - Carbon (C): 4 valence electrons.
     - Hydrogen (H): 1 valence electron for each H (4 total for 4 H's).
     - Oxygen (O): 6 valence electrons.
     - Total: 4 (C) + 4 (H) + 6 (O) = 14 valence electrons.
Central Atom: Identify the central atom (usually the least electronegative).
- E.g., in CH₃OH: Carbon becomes the central atom as it can form four bonds.
Skeleton Structure: Form a skeletal representation connecting the atoms together without worrying about electron distribution yet.
Distribute Remaining Electrons: Start placing remaining electrons on outer atoms first, aiming for an octet, excluding hydrogen.
- E.g., Oxygen should be completed to have 8 electrons total by considering lone pairs.
Form Double/Triple Bonds if Necessary: If an octet is not satisfied, create multiple bonds where necessary (common with C, O, N).

Draw carbon in the center attached to three hydrogen atoms and one hydroxyl (OH) group.
Count and ensure all atoms have appropriate total electrons and octets where applicable.

Organic structures often utilize carbon's ability to form four bonds and include oxygen or nitrogen.
Functional Groups: Recognizable parts of organic molecules, e.g., carboxyl group (-COOH) in acetic acid.

Definition of Formal Charge: An alternate way to represent charge distribution across atoms in a molecule based on the arrangement of electrons.
  - The formula to calculate the formal charge involves:
    - Formal Charge = (Valence Electrons) - (Non-Bonding Electrons) - (Bonding Electrons/2)
    - Each bond counts as one electron regardless of how many atoms share it.
Steps to Calculate Formal Charge:
  - Identify the number of valence electrons for the atom.
  - Subtract the number of lone pair electrons.
  - Subtract half the number of bonding electrons contributed.

Total of valence electrons: 5 (P) + 6*4 (O) + 3 (charge) = 32 total.
Central atom: Phosphorus, with a lower electronegativity than oxygen.
Attach oxygen atoms to phosphorus and assign lone pairs to ensure octets are satisfied.

Explains how some molecules can have more than one valid Lewis structure.
Electrons in resonance structures are not particularly fixed but are distributed over atoms, giving rise to hybridization.
For example, phosphate can show resonance as the double bond can be positioned between different oxygen atoms.
Bond Strength and Length: Each bond in a resonance hybrid is characterized as one and a quarter bonds, indicative of partial double bond character due to resonance.

The concept of resonance leads to a better understanding of molecular stability and reactivity.
Recognizing formal charges assists in predicting how atoms within molecules will interact with each other, crucial for understanding reaction mechanisms and organic chemistry.

Resonance and electron distribution details add depth to molecular characterization, aiding in more comprehensive chemical understanding and reaction prediction.