Fundamentals of Chemistry: Part 2

Study Notes

Learning Outcomes:
- Name key types of atomic bonding.
- Briefly describe key types of atomic bonding relevant to organic molecules.
- Relate the Octet Rule to chemical bonds.
- Understand what is meant by a Lewis structure.
- Draw basic Lewis structures.

Pauling Scale: Measures electronegativity, the ability of an atom to attract electrons in a chemical bond.
Trend across a period: Electronegativity increases across a period as the number of protons increases, increasing the attraction for electrons.
Trend down a group: Electronegativity decreases down a group as electron shells increase, shielding the nucleus's attraction for electrons.
Electronegativity and Polarity: Higher electronegativity differences lead to more polar covalent bonds, or ionic bonds.

Ionic Bond: Occurs between a metal and a non-metal. Electrons are transferred between atoms, creating ions (+ and -). This results in a strong electrostatic attraction. Examples include salts like NaCl.
Covalent Bond: Occurs between non-metal atoms. Electrons are shared between atoms to achieve a full outer electron shell (octet rule). This results in a strong binding force. Covalent bonds can be single, double, or triple bonds.
Metallic Bond: Occurs between metal atoms. The valence electrons are delocalized and shared among all atoms in the metal. This creates a strong binding force and various metallic properties.
Examples: Sodium Chloride (NaCl) is an ionic bond, water (H₂O) is a covalent bond, and copper (Cu) is a metallic bond.

Description: Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration called an octet. This configuration usually fills an outer electron shell with eight electrons.
Significance: This principle helps predict the types of bonds atoms will form.

Ionic Bond: Atoms gain or lose electrons, forming positive (cations) or negative (anions). The opposite charges attract.
Covalent Bond: Atoms share electrons to achieve a full outer electron shell. The shared electrons form a bonding pair. This results in a molecule.

Formation: Non-metal atoms share electrons to complete their outer electron shells.
Types:
- Single bond: One shared electron pair.
- Double bond: Two shared electron pairs.
- Triple bond: Three shared electron pairs.

Description: In certain molecules, more than one valid Lewis structure is possible. These structures are resonance structures; the actual structure is a hybrid of the possible Lewis structures. This suggests that electrons are delocalized in multiple bonds and over multiple atoms, not fixed.

Description: Formal charge helps determine the most likely Lewis structure by assigning electrons to each atom in a molecule. The structure with the lowest formal charges is the most stable.
Calculation: Difference between the number of valence electrons in the free atom and the number of electrons assigned to that atom in the Lewis structure.
Importance: Lower formal charges often correspond to more stable structures. Negative formal charges are preferably on more electronegative atoms, and zero is preferred overall.

Exception to the Octet Rule: Some molecules have more or fewer than eight electrons in their outer shells, particularly with elements in later periods or involving elements that have multiple bonding possibilities.

The difference between ionic bonds and covalent bonds and the atoms that form them.
Predicting types of bonds based on electron configuration.
Representing simple molecular structures with Lewis structures.