Notes on Molecular vs Ionic Compounds and Nomenclature

Types of Compounds and Naming

  • Compounds come in two broad types: molecular (covalent) compounds and ionic compounds.
  • Molecular compounds involve sharing electrons (covalent bonds) between nonmetal atoms.
  • Ionic compounds involve transfer of electrons between a metal (or ammonium) and a nonmetal (or polyatomic ion), forming cations and anions that result in a neutral overall charge.

Molecular (Covalent) Compounds

  • Key idea: atoms share electrons to fill outer (valence) shells, forming covalent bonds.
  • Definitions:
    • Covalent: sharing of electrons between atoms.
    • Valence electrons: electrons in the outermost shell; involved in bonding.
  • Common examples discussed:
    • Water: ext{H}_2 ext{O} (H2O)
    • Carbon monoxide: ext{CO} (CO)
    • Ammonia: ext{NH}_3 (NH3)
    • Hydrogen sulfide: ext{H}_2 ext{S} (H2S) [note on naming: dihydrogen sulfide is the systematic form; commonly hydrogen sulfide is used]
    • Chloroform: ext{CHCl}_3
    • Phosphorus pentafluoride: ext{PF}_5
    • Dinitrogen pentoxide: ext{N}2 ext{O}5
    • Tetraphosphorus hexoxide: ext{P}4 ext{O}6 ext{ or } ext{P}4 ext{O}6 ext{ (tetraphosphorus hexoxide)}
  • Prefix system for naming molecular compounds (to show the number of atoms of each element):
    • mono-, di-, tri-, tetra-, penta-, hex-, hepta-, octa-, nona-, deca-
    • Rules:
    • Use prefixes for both elements to indicate atom counts.
    • Do not use the prefix mono for the first element when naming the first element (e.g.,
      HCl is not monohydrogen chloride; it’s simply hydrogen chloride).
    • For the second element, replace the ending with -ide (when appropriate) to form a name like oxide, sulfide, chloride, etc.
  • Examples illustrating naming rules:
    • ext{H}_2 ext{S}
      ightarrow ext{dihydrogen sulfide} (commonly hydrogen sulfide)
    • ext{N}2 ext{O}5
      ightarrow ext{dinitrogen pentoxide}
    • ext{P}4 ext{O}6
      ightarrow ext{tetraphosphorus hexoxide}
    • ext{HCl}
      ightarrow ext{hydrogen chloride} ext{ (aqueous form: } ext{HCl (aq)}
      ightarrow ext{ hydrochloric acid)}
  • Important nuance: for a naming exercise, the first element’s name is not altered by a prefix when it’s “one” in number; the second element’s name is altered to end in -ide (e.g., chlorine → chloride) and prefixes indicate quantity.

Ionic Compounds

  • Ionic compounds form from a metal (or ammonium, NH4+) and a nonmetal (or a polyatomic ion).
  • Polyatomic ions you may encounter include:
    • Acetate: ext{C}2 ext{H}3 ext{O}_2^-
    • Bromate: ext{BrO}_3^-
    • Nitrate: ext{NO}_3^-
    • Nitrite: ext{NO}_2^-
    • Phosphate: ext{PO}_4^{3-}
    • Carbonate: ext{CO}_3^{2-}
  • The periodic table helps distinguish metals (left) from nonmetals (right). Metals tend to form cations; nonmetals form anions.
  • Ammonium (NH4+) is an exception: it is not a metal, but it forms ionic compounds as a cation.
  • Examples of ionic compounds:
    • Sodium chloride: ext{NaCl} (Na+ and Cl−)
    • Calcium sulfate: ext{CaSO}_4 (Ca^{2+} and SO4^{2-})
    • Sodium carbonate: ext{Na}2 ext{CO}3 (Na+ and CO3^{2-})
    • Magnesium sulfate: ext{MgSO}_4
    • Ammonium nitrate: ext{NH}4 ext{NO}3 (NH4+ and NO3−)
  • The key rule: ionic compounds are built from positive ions (cations) and negative ions (anions) and are electrically neutral overall.
  • Cations vs anions:
    • Cation: positive ion (e.g., Na+, Fe^{3+}, NH4+).
    • Anion: negative ion (e.g., Cl−, NO3−, SO4^{2-}).
  • Group trends for charges (helpful starting point):
    • Group 1 metals: +1 charge (e.g., Na+, K+).
    • Group 2 metals: +2 charge (e.g., Mg^{2+}, Ca^{2+}).
    • Nonmetals typically gain electrons to form anions with charges corresponding to their common oxidation states (e.g., O^{2-}, N^{3-}).
  • Special cases:
    • Transition metals can have more than one possible oxidation state (e.g., Fe can be Fe^{2+} or Fe^{3+}). In naming, the charge is indicated with a Roman numeral in the name: e.g., iron(III) chloride → ext{FeCl}3; iron(II) chloride → ext{FeCl}2.
    • Polyatomic ions keep their own names; when forming ionic compounds with polyatomic anions, use the polyatomic ion name directly (e.g., calcium sulfate, sulfate is the anion name).
  • Ionic compound naming convention:
    • Name the metal (or ammonium) cation first, unchanged.
    • Name the anion second:
    • If the anion is monoatomic, change the end of the root to -ide (e.g., chloride, sulfide).
    • If the anion is polyatomic, use the polyatomic ion name (e.g., sulfate, phosphate) without -ide.
    • Do not use prefixes in ionic compound names.
  • Examples of naming from formulas:
    • ext{NaCl}
      ightarrow ext{sodium chloride}
    • ext{CaSO}_4
      ightarrow ext{calcium sulfate}
    • ext{FeCl}_3
      ightarrow ext{iron(III) chloride}
    • ext{Cu(OH)}_2
      ightarrow ext{copper(II) hydroxide}
    • ext{TiCl}_4
      ightarrow ext{titanium(IV) chloride}
    • ext{Zn}3( ext{PO}4)_2
      ightarrow ext{zinc phosphate} (polyatomic group PO4^{3-} requires parentheses in the formula and in the name the - phosphate remains as is)
  • Writing formulas from names (ionic):
    • Step 1: Identify the ions and their charges from the name (including for polyatomic ions and ammonium).
    • Step 2: Cross the charges to become subscripts (the “criss-cross” method): the charge of one ion becomes the subscript of the other ion, and vice versa.
    • Step 3: If the resulting subscripts share a common factor, reduce to the smallest whole-number ratio.
    • Step 4: Use parentheses around polyatomic ions if the polyatomic ion appears more than once in the formula.
    • Example 1: Aluminum oxide from name Aluminum oxide
    • Ions: Al^{3+} and O^{2-}.
    • Cross: Al2O3. There is no simplification needed.
    • Final formula: ext{Al}2 ext{O}3
    • Example 2: Iron(III) chloride from name Iron(III) chloride
    • Ions: Fe^{3+} and Cl^{-}.
    • Cross: FeCl3. Final formula: ext{FeCl}3
    • Example 3: Calcium phosphate from name Calcium phosphate
    • Ions: Ca^{2+} and PO_4^{3-}.
    • Cross: Ca3(PO4)_2. The phosphate is polyatomic and appears more than once, so parentheses are used.
    • Final formula: ext{Ca}3( ext{PO}4)_2
    • Example 4: Ammonium phosphate from name Ammonium phosphate
    • Ions: NH4^{+} and PO4^{3-}.
    • Cross: (NH4)3PO_4. Ammonium is a polyatomic cation that is written as a unit; parentheses are used around NH4 if needed.
    • Final formula: ( ext{NH}4)^3 ext{PO}4 ext{ or } ext{(NH}4)3 ext{PO}_4
    • Example 5: Copper(II) hydroxide from name Copper(II) hydroxide
    • Ions: Cu^{2+} and OH^{-}.
    • Cross: Cu(OH)2. Final formula: ext{Cu(OH)}2
  • Summary of rules for ionic vs molecular naming and formulas:
    • Ionic compounds: metal (or ammonium) + nonmetal (or polyatomic ion); use criss-cross to balance charges; Roman numerals for transition metal charges; no prefixes in naming.
    • Molecular compounds: nonmetals only; use prefixes to show atom counts; second element ends with -ide; first element may omit the prefix if its count is 1.
    • Overall charge of ionic compounds must be neutral; this balancing is what guides the subscripts in the formula.
    • In formulas with polyatomic ions, use parentheses when the polyatomic unit appears more than once (e.g., Ca3(PO4)_2).

Quick Reference: Common Ions and Examples

  • Monatomic anions (typical endings):
    • Chloride: ext{Cl}^-
    • Oxide: ext{O}^{2-}
    • Nitride: ext{N}^{3-}
    • Sulfide: ext{S}^{2-}
  • Polyatomic anions you should know:
    • Nitrate: ext{NO}_3^{-}
    • Nitrite: ext{NO}_2^{-}
    • Carbonate: ext{CO}_3^{2-}
    • Sulfate: ext{SO}_4^{2-}
    • Phosphate: ext{PO}_4^{3-}
    • Acetate: ext{C}2 ext{H}3 ext{O}_2^{-}
    • Hydroxide: ext{OH}^{-}
    • Chlorate: ext{ClO}_3^{-}
  • Common cations:
    • Sodium: ext{Na}^+
    • Potassium: ext{K}^+
    • Calcium: ext{Ca}^{2+}
    • Magnesium: ext{Mg}^{2+}
    • Ammonium: ext{NH}_4^{+}
    • Iron can have multiple charges (Fe^{2+}, Fe^{3+}); named with Roman numerals in compounds:
    • Iron(III) chloride: ext{FeCl}_3
  • Useful cautions:
    • Ammonium salts are ionic, even though ammonium is not a metal.
    • A neutral formula results from equal total positive and negative charges.
    • For transition metals, the Roman numeral in parentheses in the name indicates the metal's charge in that compound (oxidation state concept).

Worked Practice Highlights (from the session’s examples)

  • Potassium iodide: ext{KI} (K^+ and I^-)
  • Sodium chloride: ext{NaCl} (Na^+ and Cl^-)
  • Hydrogen chloride gas: ext{HCl}; aqueous form: ext{HCl (aq)}
    ightarrow ext{hydrochloric acid}
  • Nitrogen dioxide: ext{NO}_2 (molecular)
  • Phosphorus pentafluoride: ext{PF}_5 (molecular)
  • Calcium phosphate: ext{Ca}3( ext{PO}4)_2
  • Zinc phosphate: ext{Zn}3( ext{PO}4)_2
  • Titanium tetrachloride: ext{TiCl}4 (Ti^{IV}Cl4)
  • Aluminum oxide: ext{Al}2 ext{O}3
  • Copper(II) hydroxide: ext{Cu(OH)}_2
  • Ammonium nitrate: ext{NH}4 ext{NO}3
  • Calcium phosphate vs calcium sulfate distinction based on the polyatomic ion name:
    • Phosphate: ext{PO}4^{3-} → calcium phosphate: ext{Ca}3( ext{PO}4)2
    • Sulfate: ext{SO}4^{2-} → calcium sulfate: ext{CaSO}4

Practical Takeaways for Exam Prep

  • Be able to classify a formula as molecular or ionic by identifying cations and anions (metals or ammonium vs nonmetals/polyatomic ions).
  • For molecular compounds, use prefixes to name and recall that the second element ends with -ide; the first element’s prefix may be omitted if its count is 1.
  • For ionic compounds, name the cation first (including NH4^+ as ammonium), then the anion. Use -ide for monoatomic anions and the standard polyatomic ion names for polyatomic anions. Do not use prefixes for ionic compounds.
  • For transition metals, expect multiple charges; the charge is shown by a Roman numeral in the name (e.g., iron(III) chloride → FeCl3).
  • Practice cross-mistry (criss-cross) to write formulas from names and practice writing names from formulas, including using parentheses for polyatomic ions when needed.
  • Always ensure the overall formula is electrically neutral; adjust subscripts to balance total charges.
  • Keep a handy reference sheet of common polyatomic ions and a simple periodic-table cheat sheet to speed up problem solving during exams.