3.1 Electron Arrangements and the Octet Rule

Introduction to Electron Shells and the Octet Rule

Understanding how electrons are arranged in atoms is fundamental to predicting chemical behavior. The octet rule helps explain why atoms form certain ions and compounds.

• Electron Shells: Electrons occupy energy levels (shells) around the nucleus. The first shell holds up to 2 electrons, the second and third up to 8 each.

• Valence Electrons: The electrons in the outermost shell are called valence electrons. They determine an element's chemical properties.

• Predicting Valence Electrons: For main-group elements (Groups 1A-8A) in the first four periods, the group number often equals the number of valence electrons.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a stable configuration of 8 valence electrons (an octet), similar to noble gases.

Example: Oxygen (Group 6A) has 6 valence electrons and tends to gain 2 electrons to achieve an octet.

3.2 In Search of an Octet, Part 1: Ion Formation

Ion Formation and Polyatomic Ions

Atoms form ions to achieve a stable electron configuration. The periodic table helps predict ionic charges.

• Ion Formation: Cations are formed when atoms lose electrons; anions are formed when atoms gain electrons.

• Predicting Ionic Charge: Main-group elements form ions with charges based on their group number (e.g., Group 1A forms +1, Group 7A forms -1).

• Naming Ions: Cations are named after the element (e.g., Na+ is sodium ion); anions often end in '-ide' (e.g., Cl- is chloride ion).

• Symbols for Ions: The symbol includes the element and its charge (e.g., Ca2+).

• Polyatomic Ions: These are ions composed of two or more atoms covalently bonded, carrying a net charge (e.g., NH4+, SO42-).

Example: Magnesium (Mg) loses two electrons to form Mg2+.

3.3 Ionic Compounds—Electron Give and Take

Formation and Naming of Ionic Compounds

Ionic compounds are formed by the electrostatic attraction between cations and anions. Their formulas and names follow specific rules.

• Formation: Ionic compounds are formed when metals transfer electrons to nonmetals, resulting in oppositely charged ions.

• Naming Ionic Compounds: The cation is named first, followed by the anion (e.g., NaCl is sodium chloride).

• Writing Formulas: The formula reflects the ratio of ions needed to balance charges (e.g., CaCl2 for calcium chloride).

• Predicting Ionic Charges: Use the periodic table to determine the charge of ions in a compound.

• Transition Metals: These can have multiple possible charges; the charge is indicated in the name (e.g., iron(III) chloride for FeCl3).

Example: Al3+ and O2- combine to form Al2O3 (aluminum oxide).

3.4 In Search of an Octet, Part 2: Covalent Bond

Covalent Compounds and Lewis Structures

Covalent compounds are formed when atoms share electrons to achieve an octet. Their structures and naming follow distinct conventions.

• Covalent Bond: A chemical bond formed by sharing electrons between nonmetal atoms.

• Ionic vs. Covalent: Ionic compounds involve electron transfer; covalent compounds involve electron sharing.

• Valence Electrons and Bonding: The number of bonds an atom forms is related to its number of valence electrons (e.g., oxygen typically forms 2 bonds).

• Lewis Structures: Diagrams showing the arrangement of electrons in a molecule. Dots represent valence electrons.

• Binary Covalent Compounds: Named using prefixes to indicate the number of atoms (e.g., CO2 is carbon dioxide).

Example: Water (H2O) has two single covalent bonds between oxygen and hydrogen.

3.5 The Mole: Counting Atoms and Compounds

Mole Concept and Calculations

The mole is a fundamental unit for counting atoms and molecules in chemistry. It relates mass, number of particles, and chemical quantities.

• The Mole: One mole contains particles (Avogadro's number).

• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).

• Conversions:

  • Particles to moles:

  • Mass to moles:

Example: 18 g of H2O is 1 mole, containing molecules.

3.6 Getting Covalent Compounds into Shape

Molecular Geometry and VSEPR Theory

The shape of covalent molecules is determined by the arrangement of atoms and electron pairs, predicted by VSEPR theory.

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular shapes based on electron pair repulsion.

• Common Shapes: Linear, bent, trigonal planar, tetrahedral, etc.

• Wedge-and-Dash Notation: Used to represent three-dimensional arrangement of atoms in molecules.

• Lewis Structures: Used to determine the number and arrangement of atoms and electron pairs.

Example: Methane (CH4) is tetrahedral; water (H2O) is bent.

3.7 Electronegativity and Molecular Polarity

Bond Polarity and Molecular Polarity

Electronegativity differences between atoms lead to polar bonds and affect the overall polarity of molecules.

• Electronegativity: A measure of an atom's ability to attract shared electrons. Fluorine is the most electronegative element.

• Bond Polarity: A bond is polar if the atoms have different electronegativities; electrons are shared unequally.

• Molecular Polarity: Determined by bond polarities and molecular shape. Symmetrical molecules may be nonpolar even if they contain polar bonds.

Example: Water (H2O) is a polar molecule due to its bent shape and polar O-H bonds.

Additional info: Academic context and examples have been expanded for clarity and completeness. All equations are provided in LaTeX format as required.