Ionic Formulas, Polyatomic Ions & Exam Reminders

Periodic-Table Charge Patterns

  • Moving across main-group columns the common ionic charges often change in steps:
    • Representative metals tend to show +1,+2+1, +2 (e.g. Groups 1 & 2).
    • P-block metals (Sn, Pb, Bi, etc.) can exhibit two charges differing by +2+2 (e.g. Sn2+^{2+} and Sn4+^{4+}).
    • Non-metals commonly form anions in a pattern of −1,−2,−3,−4-1, -2, -3, -4 as you move rightward.
  • Transition-metal charges are not fixed; Roman numerals in the name identify the actual oxidation state.

Systematic Procedure for Writing an Ionic Formula

  1. Identify each ion and its charge.
    • The metal (left of the staircase) is the cation ( + ).
    • The non-metal or polyatomic group is the anion ( – ).
  2. List charge multiples until a total positive charge equals a total negative charge (Least Common Multiple approach).
    • Ex.: Sn4+^{4+} → +4,+8,+12,+16+4, +8, +12, +16
    • P3−^{3-} → −3,−6,−9,−12-3, -6, -9, -12
    • First match: ±12\pm12
  3. Derive subscripts from the count of each ion needed to reach equality.
    • Above example gives 3 Sn and 4 P ⟹ Sn<em>3P</em>4\text{Sn}<em>3\text{P}</em>4.
  4. Write the formula without charges.
    • Omit a subscript of 1.
    • Keep element order: cation first, anion second.
  5. Use parentheses only when a polyatomic ion’s subscript is ≥2\ge 2
    • Pb(OH)<em>4\text{Pb(OH)}<em>4 not PbOH</em>4\text{PbOH}</em>4.
    • Never place parentheses if the polyatomic ion appears once.
Shortcut (Swap-and-Reduce Method)
  • Mathematically, if cation charge is +a+a and anion charge is −b-b, rudimentary swapping gives subscripts bb and aa; then divide by the greatest common divisor if necessary.

Polyatomic Ions to Memorize

Cation (only 1 polyatomic)
  • NH4+\text{NH}_4^{+} Ammonium ⇢ polyatomic
Common Anions (polyatomic unless noted)
  • C<em>2H</em>3O2−\text{C}<em>2\text{H}</em>3\text{O}_2^{-} Acetate
  • HCO3−\text{HCO}_3^{-} Bicarbonate / Hydrogen carbonate
  • HSO3−\text{HSO}_3^{-} Bisulfite
  • CO32−\text{CO}_3^{2-} Carbonate
  • CN−\text{CN}^{-} Cyanide
  • OH−\text{OH}^{-} Hydroxide
  • NO3−\text{NO}_3^{-} Nitrate
  • PO43−\text{PO}_4^{3-} Phosphate
  • SO42−\text{SO}_4^{2-} Sulfate
  • C<em>8H</em>4O42−\text{C}<em>8\text{H}</em>4\text{O}_4^{2-} Phthalate (sometimes written "gophthalate")
  • Monatomic anions to remember: F−, Cl−, Br−, I−, S2−, O2−, N3−, P3−\text{F}^{-},\,\text{Cl}^{-},\,\text{Br}^{-},\,\text{I}^{-},\,\text{S}^{2-},\,\text{O}^{2-},\,\text{N}^{3-},\,\text{P}^{3-}

Worked Examples

1. Tin(IV) Phosphide
  • Ions: Sn4+^{4+} and P3−^{3-}.
  • LCM(4,3)=12 → 3 Sn & 4 P.
  • Formula: Sn<em>3P</em>4\boxed{\text{Sn}<em>3\text{P}</em>4}.
2. Aluminum Phosphate
  • Al3+^{3+} and PO43−\text{PO}_4^{3-} already balance (charges cancel 1 : 1).
  • Formula: AlPO4\boxed{\text{AlPO}_4}.
3. Lead(IV) Hydroxide
  • Pb4+^{4+} with OH−^{-}.
  • Need 4 hydroxides ⟹ Pb(OH)4\text{Pb(OH)}_4 (parentheses because 4 > 1).
4. Iron(III) Carbonate
  • Fe3+^{3+} ; CO32−\text{CO}_3^{2-}.
  • LCM(3,2)=6 → 2 Fe (total +6), 3 CO$_3$ (total –6).
  • Fe<em>2(CO</em>3)3\boxed{\text{Fe}<em>2(\text{CO}</em>3)_3}.
5. Naming from a Formula
FormulaCharge LogicName
CaCl2\text{CaCl}_21 Ca 2+^{2+} vs 2 Cl −^{-}Calcium chloride
Li2S\text{Li}_2\text{S}2 Li +^{+} vs 1 S 2−^{2-}Lithium sulfide
Na<em>3PO</em>4\text{Na}<em>3\text{PO}</em>43 Na +^{+} vs PO43−\text{PO}_4^{3-}Sodium phosphate
Cu(OH)2\text{Cu(OH)}_22 OH−^{-} ⇒ –2, so Cu must be +2Copper(II) hydroxide
CuOH\text{CuOH}1 OH−^{-} ⇒ –1, so Cu must be +1Copper(I) hydroxide
SnI4\text{SnI}_44 I−^{-} (–4) ⇒ Sn is +4Tin(IV) iodide

Parentheses Rule Recap

  • Required when: polyatomic ion count ≥2\ge 2.
    Example Al<em>2(SO</em>4)3\text{Al}<em>2(\text{SO}</em>4)_3.
  • Omitted when:
    • Ion is monatomic (Cl−^{-}, S2−^{2-}, O2−^{2-}, etc.).
    • Subscript equals 1.

Recognizing Ionic vs Covalent Compounds

  • Metal + non-metal (or a polyatomic ion) ⇒ ionic.
  • Two non-metals ⇒ covalent/molecular (Greek prefixes used instead of charge balancing).

Back-Calculating the Roman Numeral

  1. Sum the negative charge from all anions.
  2. Divide by the number of metal cations in the formula.
  3. Result = oxidation state ⇒ Roman numeral in the name.

Exam & Course Logistics (mentioned in lecture)

  • First exam: tomorrow; covers Chapters 1-4.
    • Lewis structures/Chapter 6 material will NOT be on this exam.
  • Supplies required: pencil, non-programmable calculator, Scantron #882-E (50 questions, A–E or 1–5 choices).
    • Phones may NOT be used as calculators; instructor owns only two spare calculators (first-come, first-served).
  • Practice exam: posted on Canvas; hard copies available in class.
  • Chapter 6 worksheet is online and due Saturday; question numbers dealing exclusively with ions may be omitted from grading (#6 and parts of #7, as announced).

Practical Study Tips Highlighted by Instructor

  • Memorize the one polyatomic cation (ammonium) and the listed common polyatomic anions.
  • Drill charge balancing by writing multiples ( +4, +8, … / –3, –6, … ) until they match.
  • Do not write the charge or a subscript 1 in the final formula.
  • Always double-check parentheses usage with polyatomic ions.