Study Notes

Atoms bond together to achieve greater stability, typically by filling their outermost electron shell (octet rule — 8 valence electrons, except H which wants 2).

Ionic Bonds: form between a METAL and a NONMETAL (large difference in electronegativity)

• The metal TRANSFERS electrons to the nonmetal

• Metal becomes a cation (+), nonmetal becomes an anion (−)

• Strong electrostatic attraction between oppositely charged ions

• Forms crystal lattice structures (not individual molecules)

• Example: NaCl — Na⁺ and Cl⁻ held together in a lattice

Covalent Bonds: form between NONMETALS (small difference in electronegativity)

• Atoms SHARE electrons

• Both atoms benefit from the shared electrons counting toward their octets

• Form discrete molecules

• Can be polar (unequal sharing) or nonpolar (equal sharing)

• Example: H₂O, CO₂, CH₄

Metallic Bonds: form between METALS

• Metal atoms release electrons into a "sea of electrons"

• The positive metal ions are surrounded by freely moving electrons

• Explains why metals conduct electricity and are malleable/ductile

The ability to conduct electricity depends on whether charged particles (ions or electrons) can move freely.

Ionic compounds (e.g., NaCl):

• Solid state: DOES NOT conduct — ions are locked in crystal lattice, cannot move

• Molten (liquid) state: DOES conduct — ions can move freely when lattice breaks down

• Aqueous (dissolved in water): DOES conduct — ions separate and move in solution

Covalent compounds (e.g., sugar, water):

• Solid, liquid, aqueous states: generally DOES NOT conduct — no ions, no free electrons

• Exception: some covalent acids ionize in water and conduct

Metallic compounds:

• Solid AND liquid state: DOES conduct — electrons are always free to move

This property is a key way to identify bond type in a lab setting.

Ionic compounds form between metals (or polyatomic cations) and nonmetals (or polyatomic anions).

Binary ionic compounds (metal + nonmetal):

• Name the metal first (unchanged, or with Roman numeral for variable charge)

• Then name the nonmetal with "-ide" ending

• Example: NaCl → sodium chloride; MgO → magnesium oxide

Transition metals with variable charges: use Roman numerals

• Fe²⁺ = iron(II); Fe³⁺ = iron(III)

• Example: FeCl₃ → iron(III) chloride; FeCl₂ → iron(II) chloride

Polyatomic ions: know the name (given on reference sheet)

• Examples: SO₄²⁻ (sulfate), NO₃⁻ (nitrate), OH⁻ (hydroxide), NH₄⁺ (ammonium)

Formula from name:

• Use ion charges to determine ratio for electrical neutrality (total charge = 0)

• Example: Calcium fluoride: Ca²⁺ and F⁻ → need 2 fluoride ions → CaF₂

• Simplify to lowest whole number ratio when needed

Covalent compounds form between two nonmetals. Use Greek numerical prefixes to indicate the number of each atom.

Prefixes:

1 = mono- (only on second element, or if ambiguous)

2 = di-

3 = tri-

4 = tetra-

5 = penta-

6 = hexa-

7 = hepta-

8 = octa-

Rules:

• First element: use prefix only if more than one (usually omit "mono" for first)

• Second element: always use prefix + "-ide" ending

• Drop the last vowel of prefix if element name starts with a vowel

Examples:

• CO = carbon monoxide

• CO₂ = carbon dioxide

• N₂O₄ = dinitrogen tetroxide

• PCl₅ = phosphorus pentachloride

Formula from name:

• Use the prefixes to determine the subscripts directly

• Example: dinitrogen trioxide → N₂O₃

Lewis structures show how covalent molecules share electrons. Each atom should have 8 electrons around it (octet), except H (2).

Drawing Lewis structures:

1. Count total valence electrons (add all atoms' valence electrons; add 1 per negative charge)

2. Arrange atoms — least electronegative atom is usually central

3. Place a single bond between each atom pair (uses 2 electrons per bond)

4. Complete octets on outer atoms with lone pairs

5. Give remaining electrons to the central atom

6. If central atom lacks octet, form double or triple bonds by converting lone pairs

Examples:

• H₂O: O has 6 valence electrons; 2 bonds to H, 2 lone pairs on O → octets satisfied

• CO₂: C forms 2 double bonds with each O → O=C=O

• NH₃: N forms 3 bonds to H with 1 lone pair on N

When to simplify ionic formulas:

• Ionic compounds should be written in the lowest whole-number ratio

• Example: Ca₂O₂ simplifies to CaO

• But H₂O₂ (hydrogen peroxide) is NOT simplified — it's a covalent molecule where the structure matters