CHEM051 - Nomenclature: Part 2 – Binary Compounds Polyatomic & Oxyacid-Derived Ion

Overview & Context

• Part Two of the video series “Inorganic Chemical Nomenclature” focuses on compounds that contain polyatomic ions.

• Assumes you have already mastered Part One (binary compounds).

• Central skills:

  • Memorize key oxyacids and the polyatomic ions they form.

  • Apply the Stock (oxidation-number) system.

  • Correctly place parentheses, subscripts, and Roman numerals when writing formulas and names.

• Real-world relevance: oxyacids and their salts appear in batteries, rust removal (steel pickling), beverages, cleaning products, titrations, medicinal/industrial oxidizers.

Key Definitions

• Oxyacid: contains hydrogen, oxygen, and another non-metal; yields hydrogen ions in water.

• Polyatomic ion: charged species composed of ≥2 atoms covalently bonded, acting as a single ion in salts.

• Monoprotic acid: releases one $H^+$ per molecule (e.g., $HNO_3$).

• Polyprotic acid: releases >1 $H^+$; subdivided into diprotic, triprotic, etc.

• “Big brother / normal / younger / baby” terminology describes oxyacids with +1, 0, −1, or −2 O atoms relative to the main “-ic” acid.

Core Oxyacids ("-ic" Acids)

Mnemonic: "A Car Never Stays Perfectly Clean, (Bring Ice)"
(Acetic, Carbonic, Nitric, Sulfuric, Phosphoric, Chloric, Bromic, Iodic)

• $CH_3COOH$ – Acetic acid (1 acidic H) – vinegar (5 %)

• $H<em>2CO</em>3$ – Carbonic acid – forms when $CO_2$ dissolves in water; carbonation.

• $HNO_3$ – Nitric acid – monoprotic, strong oxidizer.

• $H<em>2SO</em>4$ – Sulfuric acid – battery acid, diprotic.

• $H<em>3PO</em>4$ – Phosphoric acid – flavoring in cola; triprotic.

• $HClO<em>3$ – Chloric acid – halogen congener templates. Analogues: $HBrO</em>3$ (bromic), $HIO_3$ (iodic).

Derived Oxyacid Variants

• "-ous" acids: 1 less O than the corresponding "-ic".
  Examples: $HNO<em>2$ (nitrous), $H</em>2SO<em>3$ (sulfurous), $H</em>3PO<em>3$ (phosphorus), $HClO</em>2$ (chlorous).

• Hypo-"-ous": 2 fewer O than "-ic".
  $HClO$ (hypochlorous), $HBrO$, $HIO$; $H<em>3PO</em>2$ (hypophosphorus).

• Per-"-ic": 1 more O than "-ic".
  $HClO<em>4$ (perchloric), $HBrO</em>4$ (perbromic), $HIO_4$ (periodic).

Polyatomic Anions from Oxyacids

General rule: remove all acidic H → suffix switches:

• "-ic" → "-ate"

• "-ous" → "-ite"

• Hypo-"-ous" → hypo-"-ite"

• Per-"-ic" → per-"-ate"

Main "-ate" anions (charge determined by lost H):

• Acetate $CH_3COO^-$

• Carbonate $CO_3^{2-}$

• Nitrate $NO_3^-$

• Sulfate $SO_4^{2-}$

• Phosphate $PO_4^{3-}$

• Chlorate $ClO<em>3^-$; likewise $BrO</em>3^-$ (bromate), $IO_3^-$ (iodate)

"-ite" series examples:

• Nitrite $NO<em>2^-$; Sulfite $SO</em>3^{2-}$; Chlorite $ClO_2^-$ etc.

Hypo-"-ite" examples:

• Hypochlorite $ClO^-$ – active ingredient in bleach.

• Hypophosphite $H<em>2PO</em>2^-$ (special, see below).

Per-"-ate" examples: $ClO<em>4^-$ (perchlorate), $BrO</em>4^-$ (perbromate), $IO_4^-$ (periodate).

Special Case: Phosphorus Oxyacids & Their Ions

Phosphorus breaks the “# of O changes, H stays” rule.

• $H<em>3PO</em>4$ (phosphoric): 3 acidic H → phosphate $PO_4^{3-}$.

• $H<em>3PO</em>3$ (phosphorus): only 2 acidic H (one H bonded to P).
  • Dihydrogen phosphite $H<em>2PO</em>3^-$ (lose 1 H)
  • Monohydrogen phosphite / phosphite $HPO_3^{2-}$ (lose 2 H).

• $H<em>3PO</em>2$ (hypophosphorus): only 1 acidic H.
  • Hypophosphite $H<em>2PO</em>2^-$ after full ionization.
  Structural rationale: acidic H must be bonded to O, not directly to P.

Monoprotic vs Polyprotic & Stepwise Ionization

• Polyprotic acids ionize one $H^+$ at a time, forming intermediate anions that include remaining H.
  E.g. Carbonic sequence:
  $H<em>2CO</em>3 \rightarrow HCO<em>3^- + H^+$ (bicarbonate) $HCO</em>3^- \rightarrow CO_3^{2-} + H^+$ (carbonate)

• Naming conventions for partial salts:
  • “Bicarbonate” or “bisulfate” means half-neutralized (−1 charge).
  • Prefix “monohydrogen-/dihydrogen-” explicitly states remaining H.
  • Use ‘bi-’ only for diprotic acids; never for phosphate series.

Additional Polyatomic Ions (Non-oxyacid Derived)

• Cyanide $CN^-$ ← hydrocyanic (hydrogen cyanide) acid; toxic, almond odor.

• Hydroxide $OH^-$ ← deprotonated water.

• Ammonium $NH_4^+$ ← protonated ammonia; ONLY common polyatomic cation here; “-ium” signals cation.

Rules for Naming Compounds Containing Polyatomic Ions

1. Treat like binary salts: cation name first, anion second.

2. Use Stock Roman numerals for metals with variable oxidation states.

3. Never state oxidation number for fixed-valence metals (Group 1, 2, Al, Zn, Cd, Ag).

4. Parentheses:

   • Omit around monatomic ions (e.g., $CaCl_2$).

   • Use around polyatomic ions when more than ONE unit appears: $Ca(OH)_2$.

5. Preserve identity of each polyatomic ion—do not merge formulas (e.g., $NH<em>4CN$ not $N</em>2H_4C$).

Naming Partially Neutralized Salts (Hydrogen- or “Bi-” Forms)

• With monovalent cations (Li⁺, Na⁺, K⁺, NH₄⁺): multiple correct names allowed.
  Example $K<em>2HPO</em>4$:
  • Dipotassium monohydrogen phosphate
  • Potassium monohydrogen phosphate
  • (Dipotassium) phosphate

• With polyvalent cations (Ca²⁺, Al³⁺, Fe³⁺):
  • Always state number of hydrogens (mono-, di-)
  • Do NOT number the metal.
  Example $CaHPO_4$ → calcium monohydrogen phosphate (NOT monocalcium…).

Parentheses & Formula Writing Conventions

• Cross-over (inverse) rule to balance charges.
  Example: Chromium(VI) nitrate → $Cr(NO<em>3)</em>6$.

• Check sum of charges = 0.
  $Zn^{2+} + 2 CH<em>3COO^- → Zn(CH</em>3COO)_2$.

Practice/Example Highlights

• $Fe^{3+} + 3 NO<em>2^- → Fe(NO</em>2)_3$ – iron(III) nitrite.

• $Mn(ClO<em>2)</em>7$ – deduce $Mn^{7+}$; name manganese(VII) chlorite.

• $Cd(IO<em>4)</em>2$ – cadmium periodate (cadmium fixed +2).

• $Sn(ClO<em>3)</em>2$ → tin(II) chlorate (Sn determined +2 from two –1 anions).

Special Transition-Metal Oxyanions (Strong Oxidizers)

Permanganate series

• Permanganic acid $HMnO_4$ (monoprotic).

• Anion $MnO<em>4^-$ (permanganate) – deep purple; $KMnO</em>4$ common titrant; Mn oxidation state $+7$.

Chromate / Dichromate series

• Chromic acid $H<em>2CrO</em>4$ (diprotic) → $CrO<em>4^{2-}$ chromate; bright yellow $K</em>2CrO_4$ indicator.

• Dichromic acid $H<em>2Cr</em>2O<em>7$ → $Cr</em>2O<em>7^{2-}$ dichromate; orange $K</em>2Cr<em>2O</em>7$; Cr oxidation state $+6$.

Oxidation Numbers & Stock System Reminders

• Oxidation state inferred from total anion charge and formula subscripts.
  Example: $Mo(MnO<em>4)</em>6$ → anions total −6 → Mo +6 → molybdenum(VI) permanganate.

• Always verify both mass and charge balance in equations and formulas.

Common Pitfalls & Tips

• Do NOT use di-/tri- prefixes on simple binary ionic salts (avoid “calcium dichloride”).

• “Bi-” acceptable only for $HSO<em>4^-$, $HCO</em>3^-$; NEVER for phosphate derivatives.

• Remember $NH_4^+$ is the lone common polyatomic cation—use Roman numerals only if cation is a variable-valence metal, not for ammonium.

• Parentheses around polyatomic ions only when the subscript > 1.

Ethical, Practical, & Laboratory Notes

• $H<em>2SO</em>4$ pickling protects steel but poses burn hazards.

• $KMnO<em>4$ and $K</em>2Cr<em>2O</em>7$ are powerful oxidizers; strict safety protocols, especially environmental disposal (Cr(VI) carcinogenic).

• $HCN$ highly toxic; almond odor sometimes undetectable to genetically insensitive individuals.

Study Strategies & Mnemonics

• Master the “-ic” acids first; derive others by ±O.

• Use flashcards: front = formula, back = name & charge.

• Write stepwise ionization trees for each polyprotic acid (visual memory).

• Practice cross-over rule daily—it automates formula balancing.

Quick Reference Equations & Charges

\begin{aligned}
HNO3 &\rightarrow H^+ + NO3^-\
H2SO4 &\xrightarrow[-H^+]{} HSO4^- \xrightarrow[-H^+]{} SO4^{2-}\
H3PO4 &\rightarrow H2PO4^- \rightarrow HPO4^{2-} \rightarrow PO4^{3-}\
H2CO3 &\rightarrow HCO3^- \rightarrow CO3^{2-}
\end{aligned}