Compounds: How Elements Combine – Study Notes

Chapter 3: Compounds – How Elements Combine

3.1 Electrons and Energy Levels

Atoms form compounds to achieve greater stability, which is often accomplished by achieving a full outer shell of electrons. The arrangement and behavior of electrons play a crucial role in chemical bonding.

• Electron Cloud: Electrons move rapidly in a region around the nucleus called the electron cloud. Their exact location cannot be precisely determined.

• Energy Levels: Electrons occupy specific energy levels (shells) around the nucleus. Lower energy levels are filled first.

• Maximum Electrons per Level: The maximum number of electrons in an energy level is given by the formula: $2n^2$ where n is the principal energy level (1, 2, 3, ...).

• Examples:

  • First energy level (n = 1): 2 electrons

  • Second energy level (n = 2): 8 electrons

  • Third energy level (n = 3): 18 electrons

3.1 Valence Electrons and the Octet Rule

The electrons in the outermost energy level (valence shell) are called valence electrons. These electrons are responsible for chemical bonding.

• Valence Shell: The highest energy level containing electrons.

• Groups and Periods: In the periodic table, groups (columns) indicate the number of valence electrons for main-group elements; periods (rows) indicate the outermost energy level.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell (a stable octet).

• Ion Formation:

  • Gaining electrons forms an anion (negative ion).

  • Losing electrons forms a cation (positive ion).

3.2 Ion Formation and Naming

Atoms form ions to achieve stable electron configurations. The naming of ions follows specific conventions.

• Metal Ions: Name of the metal + "ion" (e.g., Na+ is sodium ion).

• Transition Metals: Use a Roman numeral to indicate the charge (e.g., Fe2+ is iron(II) ion).

• Nonmetal Ions: Replace the ending with "-ide" (e.g., fluoride).

• Polyatomic Ions: Most end in "-ate"; "-ite" is used for related ions with one fewer oxygen. Exceptions include hydroxide (OH−), hydronium (H3O+), and cyanide (CN−).

3.3 Ionic Bonds and Compounds

Ionic bonds form when electrons are transferred from a metal to a nonmetal, resulting in oppositely charged ions that attract each other.

• Ionic Bond: The electrostatic attraction between cations and anions.

• Formation Steps:

  1. Determine the charge of each ion.

  2. Combine ions so the total charge is zero.

  3. Check the formula for correct ratios.

• Naming Ionic Compounds:

  • Combine the names of the ions (cation first), omitting the word "ion" (e.g., NaCl is sodium chloride).

  • For transition metals, include the Roman numeral (e.g., CuO is copper(II) oxide).

  • If a polyatomic ion is present, its name remains unchanged (e.g., Ca3(PO4)2 is calcium phosphate).

3.4 Covalent Bonds and Molecular Compounds

Covalent bonds form when nonmetals share valence electrons to achieve an octet. The resulting compounds are called molecules.

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

• Lewis Structures: Visual representations showing shared pairs of electrons as lines (bonds) and lone pairs as dots.

• Bonding Patterns:

  • Nitrogen forms 3 covalent bonds.

  • Oxygen forms 2 covalent bonds.

  • Fluorine and chlorine form 1 covalent bond.

  • Carbon forms 4 covalent bonds (single, double, or triple bonds).

• Molecular Formulas: Show the actual number of each atom in a molecule (e.g., glucose is C6H12O6).

• Naming Covalent Compounds:

  1. Name the first element.

  2. Name the second element, changing the ending to "-ide".

  3. Use Greek prefixes to indicate the number of each atom (e.g., CO2 is carbon dioxide).

  Note: Some compounds have traditional names (e.g., H2O is water).

3.5 The Mole: Counting Atoms and Compounds

The mole is a fundamental unit in chemistry for counting atoms, molecules, or ions. It relates the mass of a substance to the number of particles it contains.

• Definition: One mole contains $6.02 \times 10^{23}$ particles (Avogadro’s number).

• Molar Mass: The mass of one mole of a substance (in grams) is numerically equal to its atomic or molecular mass.

• Conversions:

  • Atoms ↔ Moles: Use Avogadro’s number as a conversion factor.

  • Moles ↔ Mass: Use molar mass as a conversion factor.

3.6 Molecular Shapes and VSEPR Theory

The three-dimensional shape of a molecule is determined by the arrangement of electron pairs around the central atom, as described by the Valence Shell Electron Pair Repulsion (VSEPR) model.

• VSEPR Model: Electron pairs (bonding and nonbonding) repel each other and arrange themselves as far apart as possible.

• Types of Electron Clouds:

  • Bonding clouds (B): Shared pairs of electrons (bonds).

  • Nonbonding clouds (N): Lone pairs of electrons.

• Common Shapes:

  • Tetrahedral: 4 bonds (e.g., methane, CH4).

  • Trigonal planar: 3 bonds (e.g., formaldehyde, CH2O).

  • Linear: 2 bonds (e.g., carbon dioxide, CO2).

• Effect of Lone Pairs: Nonbonded electrons occupy more space and can alter bond angles, making molecules like water (H2O) bent rather than linear.

3.7 Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract shared electrons in a covalent bond. Differences in electronegativity determine bond type and polarity.

• Electronegativity: Increases across a period and decreases down a group. Fluorine is the most electronegative element.

• Bond Types:

  • Nonpolar Covalent: Electrons shared equally (identical atoms).

  • Polar Covalent: Electrons shared unequally (different atoms; partial charges result).

  • Ionic: Electrons transferred (large electronegativity difference, ≥ 1.8).

• Partial Charges: Indicated by δ+ (partial positive) and δ− (partial negative), or by a dipole arrow pointing toward the more electronegative atom.

• Polarity and Molecular Shape: The overall polarity of a molecule depends on both the polarity of its bonds and its shape (e.g., CO2 is nonpolar due to its linear shape, while H2O is polar due to its bent shape).

Table: Summary of Bond Types by Electronegativity Difference

Electronegativity Difference	Bond Type	Example
0	Nonpolar Covalent	Cl2
0 < ΔEN < 1.8	Polar Covalent	H2O
≥ 1.8	Ionic	NaCl

Example: In water (H2O), the O–H bonds are polar, and the molecule is bent, resulting in a polar molecule. In carbon dioxide (CO2), the C=O bonds are polar, but the molecule is linear, so the dipoles cancel and the molecule is nonpolar.

Additional info: The VSEPR model is essential for predicting molecular geometry, which in turn affects molecular polarity and properties such as solubility and boiling point.