Summary of Formula Writing and Nomenclature of Inorganic Compounds

I. Oxidation Numbers

  • Definition: oxidation number = electrons lost, gained, or shared by an atom in bonding. Values assigned to atoms in a compound.
  • Rules ( Essentials )
    • Elements that lose electrons or have electrons shared with electron-deficient elements are assigned positive oxidation numbers; elements that gain electrons are assigned negative oxidation numbers.
    • In ionic compounds: NaCl formation example: $2 Na + Cl_2 \rightarrow 2 Na^+ Cl^-$; Na is +1, Cl is -1.
    • In covalent bonding (e.g., H$_2$O): O more electronegative, so O is -2 and H is +1 each.
    • Polyatomic groups behave as if they were single atoms (polyatomic ions). Example: NO$3^-$ and NH$4^+$ transfer intact in reactions.
    • Common oxidation numbers to memorize:
    • H: +1 (except metal hydrides where H is -1)
    • Alkali metals (Group 1): +1
    • Alkaline earth metals (Group 2): +2
    • B, Al: +3
    • O: -2 (peroxides: -1)
    • F: -1
    • Cl, Br, I: usually -1 in simple compounds
    • Some elements have multiple oxidation numbers (e.g., Cu, Fe, Sn, Pb, Mn). When this occurs, the metal name is followed by a Roman numeral in parentheses indicating oxidation state (Table 2 reference).
  • Key concepts
    • The oxidation numbers in a neutral compound sum to zero: ∑oxidation numbers=0\sum\text{oxidation numbers} = 0
    • First element in a compound usually takes its most common positive oxidation state.
  • Example highlights
    • KMnO$_4$: K = +1, O = -2 each; Mn must be +7 to balance: (+1)+4(−2)+x=0⇒x=+7(+1) + 4(-2) + x = 0 \Rightarrow x = +7
    • CaS$2$O$6$: Ca = +2; O = -2 (6 O => -12); two S must contribute +10 total, so each S = +5.
  • Polyatomic ions to know (summary): common ions have fixed charges (e.g., NO$3^-$, SO$4^{2-}$, NH$_4^+$).

II. Writing Formulas of Compounds

  • Core rule: the total oxidation number of the positive part must balance the total oxidation number of the negative part (sum = 0).
  • Steps to write formulas
    • Identify oxidation numbers from Table 1 (or from known polyatomic ion charges).
    • Balance to zero by choosing subscripts.
    • Place the positive ion first in the formula.
    • For polyatomic ions with subscripts ≥ 2, enclose the polyatomic ion in parentheses so the subscript applies to the whole ion.
    • Subscripts apply to the element immediately to their right.
  • Examples
    • Potassium bromide: KBrKBr (K = +1, Br = -1; sum 0).
    • Iron(II) bromide: FeBr$_2$ (Fe$^{+2}$ and Br$^{-1}$; need 2 Br$^-$).
    • Calcium acetate: Ca(C<em>2H</em>3O<em>2)</em>2Ca(C<em>2H</em>3O<em>2)</em>2 (Ca$^{+2}$ with two acetate ions, each acetate = -1).
    • Aluminum sulfate: Al<em>2(SO</em>4)<em>3Al<em>2(SO</em>4)<em>3 (Al$^{+3}$ and sulfate SO$4^{2-}$; balances to zero).
  • Important notes
    • Subscripts apply only to the immediately following element (or to the entire polyatomic ion if in parentheses).
    • A polyatomic ion with a subscript 2 or more must be placed in parentheses when balancing.

III. Determining the Oxidation Number of an Element in a Compound

  • Objective: find oxidation states that are not the element's most common state using the zero-sum rule and reference states.
  • Reference states (single oxidation state elements):
    • H: +1 (except hydrides: -1)
    • Group 1 (H, Li, Na, K, Rb, Cs): +1
    • Group 2 (Be, Mg, Ca, Sr, Ba): +2
    • B, Al: +3
    • O: -2 (peroxides: -1)
    • F: -1
    • Cl, Br, I: usually -1 unless in oxyacids or with fluorine or other halogens of higher electronegativity
  • Approach
    • Use known oxidation states and the equation for the sum to zero to solve for the unknown oxidation number.
  • Examples
    • Mn in KMnO$_4$:
    • K$^+$ = +1, O$^{-2}$ × 4 = -8, total from known atoms = -7; Mn must be +7 to balance: +1+4(−2)+x=0⇒x=+7.+1 + 4(-2) + x = 0 \Rightarrow x = +7.
    • S in CaS$2$O$6$:
    • Ca$^{+2}$, O$^{-2}$ × 6 = -12; sum from Ca + O = -10; two S must supply +10 total → each S = +5.
  • Additional examples (from polyatomic ions)
    • Cu in CuSO$_4$: sulfate = -2; Cu = +2.
    • Fe in Fe$3$(PO$4$)$2$: total PO$4$ = 2×(-3) = -6; Fe must contribute +6 total; each Fe = +2.

IV. Nomenclature of Compounds

  • A. Binary compounds (two elements, no polyatomic ions)
    • Metal + non-metal (ionic): Name metal first, then stem of non-metal with -ide ending.
    • If metal has multiple oxidation states, include Roman numeral after the metal name.
    • Examples:
    • NaCl → sodium chloride
    • AgI → silver iodide
    • CaO → calcium oxide
    • HgBr → mercury(I) bromide; HgBr$_2$ → mercury(II) bromide
    • CoCl$2$ → cobalt(II) chloride; CoCl$3$ → cobalt(III) chloride
  • B. Binary compounds (two non-metals)
    • Name elements in order of increasing electronegativity (less electronegative first).
    • Use prefixes to indicate numbers: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca- (mono- is often omitted for the first element).
    • Second element ends with -ide.
    • Examples:
    • CO → carbon monoxide
    • CO$_2$ → carbon dioxide
    • NO → nitrogen monoxide
    • PCl$_3$ → phosphorus trichloride
    • Alternative naming (difficulty-based): use oxidation numbers in parentheses after the first element when convenient (e.g., CO → carbon(II) oxide; CO$_2$ → carbon(IV) oxide).
  • C. Naming of Bases
    • Metal hydroxide: name metal (with oxidation state if needed) + hydroxide.
    • Examples: NaOH → sodium hydroxide; Fe(OH)$2$ → iron(II) hydroxide; Fe(OH)$3$ → iron(III) hydroxide.
    • Note: No molecular form of NH4OH; aqueous ammonia is NH3; in solution NH4$^+$ and OH$^-$ are present.
  • D. Acids
    • Binary acids (hydrogen + non-metal): hydro- + stem of non-metal + -ic + acid. Examples: HCl → hydrochloric acid; HBr → hydrobromic acid.
    • Oxygen-containing (oxy) acids: stem name of the non-metal + suffix -ous (lower oxidation state) or -ic (higher oxidation state) + acid.
    • HNO$_2$ → nitrous acid (N in +3)
    • HNO$_3$ → nitric acid (N in +5)
    • H$2$SO$3$ → sulfurous acid (S in +4)
    • H$2$SO$4$ → sulfuric acid (S in +6)
    • H$3$PO$3$ → phosphorous acid (P in +3)
    • H$3$PO$4$ → phosphoric acid (P in +5)
    • Carbonic acid H$2$CO$3$ does not exist as a molecular acid; CO$2$ in solution forms HCO$3^-$ and H$^+$ in equilibrium.
    • Hypo-/per- series for halogen oxyacids (Cl, Br, I): HClO (hypochlorous), HClO$2$ (chlorous), HClO$3$ (chloric), HClO$_4$ (perchloric).
    • E. Salts of Oxygen Acids
    • Name metal (with oxidation state if needed) + name of the oxoanion (from the acid) with -ite or -ate as appropriate.
    • If from a given hydrogen-containing acid with partial replacement, use hydrogen-containing salts: e.g., NaHCO$3$ (sodium hydrogen carbonate or sodium bicarbonate), NaH$2$PO$_4$ (sodium dihydrogen phosphate).
    • Examples: KClO (hypochlorite), FeSO$3$ (iron(II) sulfite), Fe$2$(SO$4$)$3$ (iron(III) sulfate).
    • F. Organic Acids (common/common IUPAC forms)
    • Formic acid: HCOOH (common formic) → systematic: methanoic acid.
    • Acetic acid: CH$_3$COOH → ethanoic acid.
    • Oxalic acid: H$2$C$2$O$_4$ → ethane-1,2-dioic acid (ethanedioic acid).
    • Tartaric acid: H$2$C$4$H$4$O$6$ → 2,3-dihydroxybutanedioic acid.
  • G. Prefix table (common Latin/Greek prefixes)
    • mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-
    • Note: mono- is often omitted for the first element in binary compounds.
  • H. Quick naming flow (summary from the end of the chapter)
    • Is it binary? Yes → name as binary; is it ionic (metal + nonmetal)? If yes → metal name first, non-metal stem + -ide (with Roman numeral if needed).
    • For two non-metals: use prefixes to indicate numbers, or use oxidation-number parentheses for the first element as an alternative.

V. Organic Acids (brief reference)

  • Common/trivial names with systematic IUPAC equivalents:

    • Formic acid: HCHO2 — methanoic acid
    • Acetic acid: HC2H3O2 — ethanoic acid
    • Oxalic acid: H2C2O4 — ethane-1,2-dioic acid
    • Tartareic acid: H2C4H4O6 — 2,3-dihydroxybutanedioic acid

  • Notes

    • Organic acids follow non-systematic/common names in general chemistry; systematic names exist in organic chemistry.

  • For quick review:

    • The algebraic sum of oxidation numbers in a compound equals zero.
    • The first element typically has its most common oxidation state unless otherwise indicated.
    • Binary compounds use -ide endings for the second element; prefixes indicate numbers when using two non-metals.
    • Acids are named with hydro- for binary acids and with -ous/-ic endings for oxyacids depending on oxidation state.
    • Polyatomic ions are treated as units with fixed charges; parentheses are used when needed to group ions in formulas.