chapter 7

Octet Rule

  • Definition: The octet rule states that atoms tend to gain, lose, or share electrons in order to have a full outer shell of eight electrons, similar to the electron configuration of noble gases.

Valence Shell Electrons

  • Definition: Valence shell electrons are the outermost electrons in an atom that are involved in forming bonds.

  • Noble Gas Electron Configuration:

    • Noble gases in Periods 2 and above (Ne, Ar, Kr, Xe) have 8 valence electrons.

    • Helium, located in Period 1, has 2 valence electrons (1s²).

    • Full electron configuration for noble gases in Periods 2 and above: ns²np⁶.

    • Full electron configuration for helium: 1s².

Ionic and Covalent Bonding

Ionic Bonding:

  • Formation: An ionic compound forms when a metal transfers electrons to a non-metal.

  • Process:

    • The metal loses valence electrons, becoming a cation (positively charged ion).

    • The non-metal gains valence electrons, becoming an anion (negatively charged ion).

Covalent Bonding:

  • Formation: A covalent bond is formed when orbitals on one non-metal atom overlap with orbitals on another non-metal atom.

  • Process:

    • The valence electrons in these overlapping orbitals are shared between the two atoms.

Electronegativity

  • Definition: Electronegativity is the ability of a covalently bonded atom to attract electrons to itself.

  • Trends:

    • Electronegativity values vary across the periodic table:

    • Fluorine (F) is the most electronegative element with a value of 4.0.

    • Trends show that electronegativity generally increases across a period and decreases down a group.

    • Notable electronegativity values for several elements:

    • H: 2.1

    • Cl: 3.0

    • O: 3.5

    • N: 3.0

    • C: 2.5

    • Li: 1.5

    • Na: 0.91

    • K: 1.0

    • Cs: 0.7

  • Polarity of Bonds:

    • Bond polarity is determined by the difference in electronegativities.

    • A covalent bond is considered non-polar if the difference is less than 0.4.

    • A bond is polar if the difference is between 0.4 and 1.9.

    • Bonds with differences greater than 1.9 are considered ionic.

Distinguishing Non-Polar from Polar Covalent Bonds

  • Non-Polar Covalent Bond:

    • Occurs when the electronegativities of the two atoms are equal or very similar (difference < 0.4).

    • Example: F-F bond in F₂ (non-polar).

  • Polar Covalent Bond:

    • Occurs when there is a significant difference in electronegativities (difference between 0.4 and 1.9).

    • Example: H-F bond in HF where hydrogen has a partial positive charge (δ+) and fluorine has a partial negative charge (δ-).

Metallic Bonding

  • Concept: In metallic compounds, metal atoms share their valence electrons to form an electron sea that is delocalized among all metal atoms.

  • Result: This delocalization contributes to the properties of metals such as conductivity and malleability.

Lewis Dot Structures

Importance:

  • Lewis dot structures visually represent the valence electrons in a molecule or ion, showing bonding (with lines) and non-bonding (with dots) electrons.

Key Points for Drawing:

  1. Count the total number of valence electrons available.

    • For anions, add one electron for each negative charge.

    • For cations, subtract one electron for each positive charge.

  2. Identify the central atom (the least electronegative atom that isn’t hydrogen).

  3. Connect outer atoms to the central atom with single bonds (2 electrons per bond).

  4. Complete the octets of outer atoms first, then the central atom.

  5. If necessary, create multiple bonds to satisfy the octet rule for the central atom.

  6. Ensure all valence electrons are accounted for in the structure.

Molecular Geometry

VSEPR Theory:

  • Valence Shell Electron Pair Repulsion (VSEPR) Theory: Explains that electron pairs around a central atom will orient themselves as far apart as possible to minimize repulsion.

  • Electron Domains:

    • Defined as the number of lone pairs and bonded atoms surrounding a central atom.

Common Molecular Shapes and Angles:

  • Linear Geometry:

    • 2 electron domains, bond angle of 180°.

  • Trigonal Planar Geometry:

    • 3 electron domains, bond angle of 120°.

  • Tetrahedral Geometry:

    • 4 electron domains, bond angle of 109.5°.

  • Trigonal Bipyramidal Geometry:

    • 5 electron domains, bond angles of 90° and 120°.

  • Octahedral Geometry:

    • 6 electron domains, bond angle of 90°.

Covalent Bonding Strength and Length

  • Bond Strength: A greater number of shared electron pairs between two atoms results in a stronger bond.

  • Bond Length: The bond length decreases as the number of shared electron pairs increases, leading to shorter, stronger bonds.

Molecular Orbital Theory

  • Concept: Molecular orbitals form from the overlap of atomic orbitals, characterized by specific energy levels.

  • Types of Orbitals:

    • Sigma Bonds (σ): Formed by end-to-end overlap of orbitals, always the first bond formed between two atoms.

    • Pi Bonds (π): Formed by side-on overlap of p orbitals, present as the second and third bonds in multiple bonds.

Summary of Chemical Bonding Topics

  • Chemical bonds can be classified into ionic, covalent, and metallic types based on the nature of electron sharing or transfer.

  • The octet rule provides a foundational explanation for the formation of these bonds and stability of different compounds.