Chapter 3 Study Guide: Chemical Bonding, The Mole, and Molecular Structure

3.1 Electron Arrangements and the Octet Rule

Introduction to Electron Shells and the Octet Rule

Atoms consist of electrons arranged in shells (energy levels) surrounding the nucleus. The arrangement of these electrons determines the chemical properties of an element, especially its reactivity and bonding behavior. The octet rule is a guiding principle stating that atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons, similar to noble gases.

• Electron Shells: Electrons occupy discrete energy levels called shells, labeled n=1, n=2, etc.

• Valence Electrons: The electrons in the outermost shell are called valence electrons and are crucial for chemical bonding.

• 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 achieve eight electrons in their valence shell through chemical reactions.

Example: Oxygen (Group 6A) has 6 valence electrons and tends to form bonds to complete its octet.

Additional info: Hydrogen and helium are exceptions, aiming for two electrons (duet rule).

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

Ion Formation and the Octet Rule

Atoms can achieve an octet by gaining or losing electrons, forming ions. The periodic table helps predict the ionic charge of main-group elements.

• Cations: Atoms lose electrons to form positively charged ions (e.g., Na+).

• Anions: Atoms gain electrons to form negatively charged ions (e.g., Cl-).

• Predicting Ionic Charge: Group 1A elements form +1 ions, Group 2A form +2, Group 7A form -1, etc.

• Naming Ions: Cations are named after the element (sodium ion), anions often end in '-ide' (chloride ion).

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

• Polyatomic Ions: Ions composed of multiple atoms (e.g., SO42-).

Example: Magnesium (Group 2A) 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: Combine ions in ratios that balance total positive and negative charges.

• Naming: Name the cation first, then the anion (e.g., sodium chloride).

• Writing Formulas: Use subscripts to indicate the ratio of ions (e.g., NaCl, CaCl2).

• Transition Metals: May have multiple possible charges; use Roman numerals in names (e.g., iron(III) chloride).

• Predicting Charges: Use the periodic table and compound formula to deduce ion charges.

Example: Al3+ and O2- combine to form Al2O3.

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

Covalent Bonding and Molecular Compounds

Covalent compounds are formed when atoms share electrons to achieve an octet. These compounds differ from ionic compounds in their bonding and properties.

• Covalent Bond: A chemical bond formed by sharing electron pairs between atoms.

• Ionic vs. Covalent: Ionic involves electron transfer; covalent involves sharing.

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

• Lewis Structures: Diagrams showing the arrangement of atoms and shared electron pairs in a molecule.

• Binary Covalent Compounds: Named using prefixes (e.g., CO2 is carbon dioxide).

Example: Water (H2O) is formed by sharing electrons between hydrogen and oxygen.

3.5 The Mole: Counting Atoms and Compounds

The Mole Concept and Avogadro's Number

The mole is a fundamental unit in chemistry for counting atoms, molecules, or ions. Avogadro's number defines the number of particles in one mole.

• Mole: $1 ext{ mole} = 6.022 imes 10^{23}$ particles (Avogadro's number).

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

• Conversions:

  • Particles to moles: $n = rac{ ext{number of particles}}{6.022 imes 10^{23}}$

  • Mass to moles: $n = rac{ ext{mass (g)}}{ ext{molar mass (g/mol)}}$

Example: 18 g of H2O contains $n = rac{18}{18} = 1$ mole of water 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 using the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• VSEPR Theory: Electron pairs around a central atom repel each other, determining molecular shape.

• Common Shapes:

  • Linear

  • Trigonal planar

  • Tetrahedral

  • Trigonal pyramidal

  • Bent

• Wedge-and-Dash Notation: Used to represent 3D arrangement of atoms in Lewis structures.

Example: Methane (CH4) has a tetrahedral shape.

3.7 Electronegativity and Molecular Polarity

Bond Polarity and Molecular Polarity

Electronegativity is the tendency of an atom to attract electrons in a bond. Differences in electronegativity lead to polar bonds and, depending on molecular geometry, polar molecules.

• Electronegativity: Fluorine is the most electronegative element; values decrease down a group and increase across a period.

• Bond Polarity: A bond is polar if the atoms have different electronegativities.

• Molecular Polarity: Determined by both bond polarities and the shape of the molecule.

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