Rucete ✏ SAT Chemistry In a Nutshell
4. Chemical Formulas
This chapter focuses on writing and naming chemical compounds, including ionic and covalent compounds, polyatomic ions, acids, and bases. It also covers empirical formulas, percent composition, and how to write and balance chemical equations—key skills for SAT Chemistry.
Binary Ionic Compounds: Category I
- Metals from Groups 1 and 2 form only one type of cation.
- Name = cation + anion (with “-ide” ending), e.g., NaCl = sodium chloride.
- Use crisscross rule to determine subscripts, e.g., CaCl₂.
- Reduce subscripts to lowest whole-number ratio.
Binary Ionic Compounds: Category II
- Transition metals with multiple charges use Roman numerals (e.g., Fe²⁺ = iron(II)).
- Older names: “-ic” for higher charge, “-ous” for lower (ferric vs. ferrous).
- Ag⁺, Zn²⁺, Cd²⁺ always have fixed charges—no numeral needed.
Polyatomic Ions and Ternary Compounds
- Polyatomic ions act as single charged units, e.g., NH₄⁺, SO₄²⁻.
- Ternary compounds contain 3 elements—name cation first.
- Use parentheses when needed: Fe₂(SO₄)₃.
- “-ate” = more O, “-ite” = fewer O (e.g., nitrate vs. nitrite).
- Prefixes: “per-” = most O, “hypo-” = least O (e.g., ClO₄⁻ = perchlorate).
Binary Covalent Compounds: Category III
- Between two nonmetals; no ions.
- First element uses full name; second ends in “-ide.”
- Use prefixes: mono-, di-, tri-, tetra-, penta-, etc.
- Never use “mono-” on first element; drop vowels when needed.
- Examples: N₂O₄ = dinitrogen tetroxide, SF₆ = sulfur hexafluoride
Naming Acids and Bases
- Binary acids: hydro- + root + -ic (e.g., HCl = hydrochloric acid)
- Ternary acids:
- “-ate” → “-ic” (e.g., HNO₃ = nitric acid)
- “-ite” → “-ous” (e.g., HNO₂ = nitrous acid)
- “per-ate” → “per-ic” (e.g., HClO₄)
- “hypo-ite” → “hypo-ous” (e.g., HClO)
- Bases = hydroxides (e.g., NaOH = sodium hydroxide)
- Salt naming:
- -ic acid → -ate salt
- -ous acid → -ite salt
- hydro-...-ic → -ide salt
Formula Mass and Percentage Composition
- Formula shows element types and ratios.
- Formula mass = sum of atomic masses × subscripts
- Molecular mass = total mass of actual molecule
- Empirical = simplest ratio; Molecular = actual atoms
- Example: CH₂ = empirical; C₂H₄ = true formula
- Percent composition = (mass of element ÷ total mass) × 100
Determining Empirical Formulas
- Assume 100 g → % = grams
- Convert grams to moles
- Divide by smallest mol value → whole number ratio
- Example: 60% Mg, 40% O → MgO
- If molecular mass is given, divide it by empirical formula mass to get multiplier
Laws of Composition
- Definite Composition: A compound always has the same percent composition.
- Multiple Proportions: Mass ratios of elements in different compounds are small whole numbers.
Writing and Balancing Chemical Equations
- Use correct symbols and formulas.
- Balance with coefficients only.
- Example: H₂ + O₂ → H₂O → 2H₂ + O₂ → 2H₂O
- Example: 2C₄H₁₀ + 13O₂ → 8CO₂ + 10H₂O
Showing Phases and Ionic Equations
- Phase notations:
- (s) = solid
- (l) = liquid
- (g) = gas
- (aq) = aqueous
- Example: 2HCl(aq) + Zn(s) → ZnCl₂(aq) + H₂(g)
- Net ionic equations show only the reacting species:
- Remove spectator ions
- Example: 2H⁺(aq) + Zn(s) → Zn²⁺(aq) + H₂(g)
In a Nutshell
This chapter explains how to name and write chemical formulas for ionic, covalent, and acid-base compounds. It also introduces polyatomic ions, the Stock system, and naming conventions. Calculations like formula mass, percent composition, and empirical formula help understand compound makeup. Lastly, balancing chemical equations and writing net ionic equations are essential tools in representing chemical changes quantitatively.