Classification, Naming, and Writing Formulas of Inorganic Compounds

Classification and Naming of Binary Ionic Compounds

• Binary compounds are chemical compounds composed of exactly two different elements.

  • Examples of binary compounds include $NaCl$, $CO_2$, $NH_3$, and $P_4O_{10}$.

• A binary ionic compound is an ionic compound in which one element present is a metal and the other element present is a non-metal.

Binary Ionic Compounds with Fixed Oxidation States

• Metals in these compounds exhibit fixed-charge behavior and form only one type of cation.

• Naming Rule: Name of the metal + stem of the non-metal + suffix -ide.

  • General Formula: Metal name + non-metal stem + -ide.
  • Example ($NaF$): Start with the full name of the metal (sodium), follow it with the stem of the non-metal (fluor-), and add the suffix -ide to yield sodium fluoride.
  • Example ($NaCl$): Sodium chloride.
  • Example ($KBr$): Potassium bromide.

• Names, stems, and ionic formulas of selected common non-metallic elements:

  • Bromine: Stem = brom-, Name of ion = bromide, Formula = $Br^-$
  • Carbon: Stem = carb-, Name of ion = carbide, Formula = $C^{4-}$
  • Chlorine: Stem = chlor-, Name of ion = chloride, Formula = $Cl^-$
  • Fluorine: Stem = fluor-, Name of ion = fluoride, Formula = $F^-$
  • Hydrogen: Stem = hydr-, Name of ion = hydride, Formula = $H^-$
  • Iodine: Stem = iod-, Name of ion = iodide, Formula = $I^-$
  • Nitrogen: Stem = nitr-, Name of ion = nitride, Formula = $N^{3-}$
  • Oxygen: Stem = ox-, Name of ion = oxide, Formula = $O^{2-}$
  • Phosphorus: Stem = phosph-, Name of ion = phosphide, Formula = $P^{3-}$
  • Sulfur: Stem = sulf-, Name of ion = sulfide, Formula = $S^{2-}$

• Practice Compound Naming:

  1. $MgO$: Magnesium oxide
  2. $Al_2S_3$: Aluminum sulfide
  3. $K_3N$: Potassium nitride
  4. $CaCl_2$: Calcium chloride

• Metals that exhibit fixed-charge behavior:

  • Monovalent ($+1$): $Li^+$, $Na^+$, $K^+$, $Rb^+$, $Cs^+$, $Ag^+$
  • Divalent ($+2$): $Be^{2+}$, $Mg^{2+}$, $Ca^{2+}$, $Sr^{2+}$, $Ba^{2+}$, $Zn^{2+}$, $Cd^{2+}$
  • Trivalent ($+3$): $Al^{3+}$, $Ga^{3+}$

Binary Ionic Compounds with Variable Oxidation States

• Metals with variable ionic charges exhibit less predictable behavior because they form more than one type of positive ion.
  • Example: Iron forms both $Fe^{2+}$ and $Fe^{3+}$ ions depending on chemical conditions.
• When naming compounds containing metals with variable charges, the charge on the metal ion is specified in the name using Roman numerals in parentheses immediately following the metal's name (Stock system) or using traditional/classical suffixes (-ous for lower charge, -ic for higher charge).
  • Chlorides of $Fe^{2+}$ and $Fe^{3+}$ ($FeCl_2$ and $FeCl_3$):
  • $FeCl_2$: Iron (II) chloride or ferrous chloride
  • $FeCl_3$: Iron (III) chloride or ferric chloride
  • Example ($CuO$): Copper (II) oxide or cupric oxide
• Determination of charge: The fixed charge on the non-metal ion is used to calculate the unknown charge on the metal ion.
• Summary of metals with variable charge characteristics:
  • Iron ($Fe$): Charge = $+2, +3$; Names = Ferrous ($+2$), Ferric ($+3$) or Iron (II), Iron (III)
  • Chromium ($Cr$): Charge = $+2, +3$; Names = Chromium (II), Chromium (III)
  • Manganese ($Mn$): Charge = $+2, +3$; Names = Manganese (II), Manganese (III)
  • Cobalt ($Co$): Charge = $+2, +3$; Names = Cobaltous ($+2$), Cobaltic ($+3$) or Cobalt (II), Cobalt (III)
  • Nickel ($Ni$): Charge = $+2, +3$; Names = Nickel (II), Nickel (III)
  • Copper ($Cu$): Charge = $+1, +2$; Names = Cuprous ($+1$), Cupric ($+2$) or Copper (I), Copper (II)
  • Gold ($Au$): Charge = $+1, +3$; Names = Gold (I), Gold (III)
  • Tin ($Sn$): Charge = $+2, +4$; Names = Stannous ($+2$), Stannic ($+4$) or Tin (II), Tin (IV)
  • Lead ($Pb$): Charge = $+2, +4$; Names = Plumbous ($+2$), Plumbic ($+4$) or Lead (II), Lead (IV)
  • Mercury ($Hg$): Charge = $+1, +2$; Names = Mercurous ($Hg_2^{2+}$), Mercuric ($Hg^{2+}$)

Ionic Compounds Containing Polyatomic Ions

Categories and Properties of Ions

• Monoatomic Ion: An ion formed from a single atom through the loss or gain of electrons.
  • Examples: $Cl^-$, $Na^+$, $Ca^{2+}$, $Mg^{2+}$, $N^{3-}$.
• Polyatomic Ion: An ion formed from a group of atoms held together by covalent bonds that carries a net positive or negative charge through loss or gain of electrons.
  • Example: Sulfate ion ($SO_4^{2-}$).
  • Polyatomic ions are highly stable and maintain their fundamental identity during chemical reactions.
  • Polyatomic ions are charged pieces of compounds, not independent molecules; they never exist in isolation and are always associated with counter-ions of opposite charge.

Comprehensive List of Polyatomic Ions

• Positive Polyatomic Ions (Cations):

  • $NH_4^+$: Ammonium
  • $H_3O^+$: Hydronium

• Negative Polyatomic Ions (Anions):

  • $NO_3^-$: Nitrate
  • $NO_2^-$: Nitrite
  • $N_3^-$: Azide
  • $SO_4^{2-}$: Sulfate
  • $HSO_4^-$: Bisulfate (Hydrogen sulfate)
  • $SO_3^{2-}$: Sulfite
  • $HSO_3^-$: Bisulfite (Hydrogen sulfite)
  • $S_2O_3^{2-}$: Thiosulfate
  • $PO_4^{3-}$: Phosphate
  • $HPO_4^{2-}$: Biphosphate (Hydrogen phosphate)
  • $H_2PO_4^-$: Dihydrogen phosphate
  • $PO_3^{3-}$: Phosphite
  • $ClO_4^-$: Perchlorate
  • $ClO_3^-$: Chlorate
  • $ClO_2^-$: Chlorite
  • $ClO^-$: Hypochlorite
  • $CO_3^{2-}$: Carbonate
  • $HCO_3^-$: Bicarbonate (Hydrogen carbonate)
  • $C_2O_4^{2-}$: Oxalate
  • $C_2H_3O_2^-$: Acetate
  • $CN^-$: Cyanide
  • $MnO_4^-$: Permanganate
  • $CrO_4^{2-}$: Chromate
  • $Cr_2O_7^{2-}$: Dichromate
  • $OH^-$: Hydroxide

Structural Rules for Polyatomic Ions

• Charge Distribution: Most polyatomic ions carry negative charges ranging from $-1$ to $-3$. Only two positive polyatomic ions are commonly listed: $NH_4^+$ and $H_3O^+$.
• Suffix Rules:
  • Only two negative polyatomic ions end in -ide: hydroxide ($OH^-$) and cyanide ($CN^-$).
  • The remaining negative polyatomic ions end in -ate or -ite.
• The -ate and -ite Comparison:
  • In an -ate/-ite pair, the -ate ion always possesses one more oxygen atom than the -ite ion (e.g., $SO_4^{2-}$ sulfate vs. $SO_3^{2-}$ sulfite).
  • Both ions in an -ate/-ite pair carry the exact same electrical charge.
• Hydrogen Addition:
  • When a hydrogen atom is added to an anion (such as $CO_3^{2-}$ forming $HCO_3^-$), the charge on the resulting hydrogenated ion is increased by $+1$ (becoming 1 unit less negative).

Formula Writing and Naming Rules for Polyatomic Compounds

• Formula Writing Conventions:
  1. Parentheses Usage: When more than one polyatomic ion of a given type is needed to balance charge, the polyatomic ion is enclosed in parentheses, and a subscript is placed outside the right parenthesis. Example: $Fe(OH)_3$.
  2. Repeated Symbols: Elemental symbols may appear more than once in a formula to preserve the distinct structural identity of polyatomic ions. Example: $NH_4NO_3$ contains nitrogen ($N$) in two separate locations corresponding to $NH_4^+$ and $NO_3^-$.
• Examples of Determining Chemical Formulas from Ion Pairs:
  1. $Na^+$ and $SO_4^{2-}$: $Na_2SO_4$
  2. $Mg^{2+}$ and $NO_3^-$: $Mg(NO_3)_2$
  3. $NH_4^+$ and $CN^-$: $NH_4CN$
• Naming Modifications for Polyatomic Ionic Compounds:
  • Rule A: If a positive polyatomic ion is present, its specific name is substituted for the metal name.
  • Rule B: If a negative polyatomic ion is present, its specific name is substituted for the non-metal stem + -ide ending.
  • Rule C: If both positive and negative polyatomic ions are present, a dual substitution is performed using both polyatomic ion names.
• Examples of Polyatomic Compound Naming:
  1. $Ca_3(PO_4)_2$: Calcium phosphate
  2. $Fe_2(SO_4)_3$: Iron (III) sulfate or ferric sulfate
  3. $(NH_4)_2CO_3$: Ammonium carbonate

Binary Molecular Compounds

Structure and Naming Rules

• A binary molecular compound is a molecular compound composed of only two non-metallic elements.
• Naming Convention: Name of non-metal of lower electronegativity + stem of the more electronegative non-metal + suffix -ide.
• Numerical prefixes indicating the exact number of atoms of each element present precede the names of both non-metals.

Numerical Prefixes

• Numerical Prefix List:

  • 1: mono-
  • 2: di-
  • 3: tri-
  • 4: tetra-
  • 5: penta-
  • 6: hexa-
  • 7: hepta-
  • 8: octa-
  • 9: nona-
  • 10: deca-

• Table of Prefix Usage and Molecular Examples:

  • 1 (mono-): $CO$ — Carbon monoxide
  • 2 (di-): $H_2F_2$ — Dihydrogen difluoride
  • 3 (tri-): $NCl_3$ — Nitrogen trichloride
  • 4 (tetra-): $S_4N_4$ — Tetrasulfur tetranitride
  • 5 (penta-): $ClF_5$ — Chlorine pentafluoride
  • 6 (hexa-): $I_2F_6$ — Diiodine hexafluoride
  • 7 (hepta-): $IF_7$ — Iodine heptafluoride
  • 8 (octa-): $P_4O_8$ — Tetraphosphorus octoxide
  • 9 (nona-): $P_4S_9$ — Tetraphosphorus nonasulfide
  • 10 (deca-): $P_4Se_{10}$ — Tetraphosphorus decaselenide

• Specific Rule Constraints:

  • The prefix mono- is never used for the first element in a compound name, but it must be used for the second element if only one atom is present (e.g., $CO$ is carbon monoxide, not monocarbon monoxide).
  • Hydrogen Exception: Compounds in which hydrogen is listed as the first element in the formula are named without numerical prefixes.

• Additional Examples:

  • $NF_3$: Nitrogen trifluoride
  • $I_2F_6$: Diiodine hexafluoride

Common Name Exemptions for Binary Molecular Compounds

• Certain binary molecular compounds possess traditional common names completely unrelated to systematic prefix naming rules:
  • $H_2O$: Water
  • $H_2O_2$: Hydrogen peroxide
  • $NH_3$: Ammonia
  • $N_2H_4$: Hydrazine
  • $CH_4$: Methane
  • $C_2H_6$: Ethane
  • $PH_3$: Phosphine
  • $AsH_3$: Arsine

Acids, Bases, and Hydrates

Naming Acids

• Binary Acids (contain hydrogen and one other non-metallic element):

  • Formula Rule: hydro- + root word + -ic acid
  • Example: $H_2S$ — hydrosulfuric acid
  • Example: $HCl$ — hydrochloric acid

• Polyatomic Acids (Oxyacids):

  • $HClO_4$: Perchloric acid
  • $HClO_3$: Chloric acid
  • $HClO_2$: Chlorous acid
  • $HClO$: Hypochlorous acid
  • $HNO_3$: Nitric acid
  • $HNO_2$: Nitrous acid
  • $H_3PO_3$: Phosphorous acid
  • $H_2CO_3$: Carbonic acid

Naming Bases

• Systematic Rule: Name of metal + hydroxide
• Example: $Mg(OH)_2$ — magnesium hydroxide
• Exception to the rule: $NH_4OH$ — ammonium hydroxide

Naming Hydrates

• Hydrates: Inorganic salts that capture and incorporate a specific number of water molecules into their crystalline matrix.
• Naming Convention: Systematic salt name + Greek prefix indicating water molecule count + hydrate.
• Examples:
  • $CuSO_4 \cdot 5H_2O$: Cupric sulfate pentahydrate
  • $MgSO_4 \cdot 7H_2O$: Magnesium sulfate heptahydrate
• Anhydrous / Anhydrates: Compounds containing no water molecules of crystallization.
• Reference for Greek Prefixes:
  • 1: mono
  • 2: di
  • 3: tri
  • 4: tetra
  • 5: penta
  • 6: hexa
  • 7: hepta
  • 8: oct
  • 9: non
  • 10: dec
  • 11: undec
  • 12: dodec