Ionic and Acid Naming Quiz

Ionic Naming

• Binary Compounds: Ionic compounds composed of one metal and one nonmetal.
  • Naming Rule: Name the metal cation first, then the nonmetal anion with the -ide suffix.
  • Example: NaCl (Sodium Chloride).
• Polyatomic Ions: Ionic compounds with more than two atoms (often oxygen and another element).
  • Naming Rule: Name the metal cation first, then the polyatomic anion. Polyatomic ions usually end in -ate or -ite.
  • Example: Na₂SO₄ (Sodium Sulfate).
• Roman Numerals for Transition Metals: Many transition metals have more than one possible charge.
  • Naming Rule: Use Roman numerals in parentheses after the metal cation to indicate the charge.
  • Examples: FeCl₂ (Iron(II) Chloride), Fe₂O₃ (Iron(III) Oxide).
  • Exceptions: Silver (Ag⁺) and Zinc (Zn²⁺) always have a single charge.

Acid Naming

• Acids Ending in -ate: Replace -ate with -ic acid.
  • Example: H₂SO₄ (Sulfuric Acid).
• Acids Ending in -ite: Replace -ite with -ous acid.
  • Example: H₂SO₃ (Sulfurous Acid).
• Acids Ending in -ide: Add "hydro-" to the beginning, replace -ide with -ic acid.
  • Example: HCl (Hydrochloric Acid).

Molecular Naming

• Covalent Compounds: Compounds formed by sharing electrons between nonmetals.
  • Naming Rule: Use prefixes to denote the number of atoms of each element. Omit "mono-" for the first element.
  • Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.
  • Example: CO (Carbon Monoxide), N₂O₄ (Dinitrogen Tetraoxide), P₄O₁₀ (Tetraphosphorus Decoxide).

Ionic vs. Covalent Compounds

• Ionic Compounds: Metal and nonmetal, high melting point, conduct electricity when molten or dissolved.
• Covalent Compounds: Nonmetal and nonmetal, low melting point, generally do not conduct electricity.

Double Replacement Reactions

• Double Replacement Reaction: Two ionic compounds react, ions switch partners forming new compounds.
  • Example: AgNO₃ + NaCl → AgCl + NaNO₃

States of Matter

• Solid: Atoms/molecules are closely packed.
• Liquid: Atoms/molecules are spaced further and can move past each other.
• Gas: Atoms/molecules are widely spaced and freely move.

Diatomics

• Diatomic Elements: Elements that exist in nature as two-atom molecules.
  • List: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂.

Bonding and Octet Rule

• Ionic Bonding: Electrons are transferred.
• Covalent Bonding: Electrons are shared.
• Octet Rule: Atoms tend to complete their outer electron shell with 8 electrons (except H, which wants 2 – a "duet").
• Exceptions: Beryllium(Be) often has 4 electrons, Boron (B) often has 6, Phosphorus (P) often has 10, and Sulfur (S) often has 12 or more in its valence shell. Similarly, some elements in period 3 and onwards can exceed the octet rule.

Lewis Structures

• Ionic Bonding: Electrons transferred forming ions.
• Covalent Bonding: Electrons shared forming bonds.

Formal Charge

• Formula: valence electrons - (bonds + lone pairs)
• Purpose: To help determine the better Lewis structure – The structure with the fewest or smallest formal charges is preferred.

Charges of Ions

• Identify the element and predict the associated ion based on its position and charge tendency. -Example; Sodium (Na⁺), Oxygen (O⁻²), Nitrogen (N⁻³).

Group Names

• Group 1: Alkali Metals
• Group 2: Alkaline Earth Metals
• Group 17: Halogens
• Group 18: Noble Gases

Reactivity and Electronegativity Trends

• Reactivity Trends: Group 1 (Alkali Metals) and Group 17 (Halogens) are highly reactive.
• Electronegativity: Tends to increase across a period and decrease down a group.

Polyatomic Ions

• Polyatomic Ions: Groups of atoms that have a collective charge.
  • Examples: Nitrate (NO₃⁻), Phosphate (PO₄³⁻), Hydroxide (OH⁻).

VSEPR Theory (Molecular Shapes)

• Molecular Shapes:
  • Linear: 180° (e.g., CO₂)
  • Bent: <120° (e.g., H₂O)
  • Trigonal Planar: 120° (e.g., BF₃)
  • Tetrahedral: 109.5° (e.g., CH₄)

Dimensional Analysis

• Conversions: Use conversion factors (e.g., 1 km = 1000 m, 1 g = 1000 mg, °C to K).
• Examples: Convert km to meters, grams to milligrams, °C to Kelvin.