Nomenclature of Chemical Compounds

Naming Chemical Compounds

Section 5.1: Naming Binary Compounds

• Binary Compounds: Compounds composed of two different elements.
• Classification: Divided into two broad classes:
  • Compounds containing a metal and a nonmetal.
  • Compounds containing two nonmetals.
• Flowchart for Naming Binary Compounds:
  1. Is it a Binary Compound? If no, follow other naming strategies.
  2. Is a Metal Present?
     • No Metal Present (Type III): Uses prefixes to denote the number of atoms for each nonmetal.
     • Metal Present: Proceed to the next question.
  3. Does the Metal Form More Than One Cation?
     • No (Type I): Uses the simple element name for the cation (metals with fixed charges, typically Group 1, Group 2, Aluminum, Zinc, Silver, Cadmium).
     • Yes (Type II): The charge of the cation must be determined and specified using a Roman numeral after the element name for the cation.

Section 5.2: Naming Binary Compounds That Contain a Metal and a Nonmetal (Types I and II)

Binary Ionic Compounds (Type II)

• Characteristics:
  • Metals in these compounds can form more than one type of cation (variable charges).
  • The charge on the metal ion must be specified.
  • A Roman numeral in parentheses indicates the charge of the metal cation.
  • Transition metal cations usually require a Roman numeral.
    • Example: $Fe$ (Iron) is a transition metal, often requiring a Roman numeral.
• Rules for Naming Type II Ionic Compounds:
  1. The cation is always named first and the anion second.
  2. The cation's charge is specified by a Roman numeral in parentheses because it can assume more than one charge.
     • Example: For $Fe_2O_3$:
       • Ions Present: $Fe^{3+}$ and $O^{2-}$.
       • Charge Balance: $2(3+) + 3(2-) = 0$ (Net charge is zero).
       • Ion Names: The cation is iron(III), and the anion is oxide.
       • Comment: Iron is a transition metal and requires a (III) to specify the charge on the cation.

Binary Ionic Compounds (Type I)

• Rules for Naming Type I Ionic Compounds:
  1. The cation is always named first and the anion second.
  2. A simple cation takes its name directly from the name of the element.
  3. A simple anion is named by taking the first part of the element name (the root) and adding the suffix -ide.
     • Example: $Cl^-$ is called chloride.
     • Example: For $AlCl_3$:
       • Ions Present: $Al^{3+}$ and $Cl^-$.
       • Ion Names: Aluminum and Chloride.
       • Comment: $Al$ (Group 3) always forms $Al^{3+}$. $Cl$ (Group 7) always forms $Cl^-$. No Roman numeral is needed for aluminum as its charge is fixed.

Alternative Nomenclature (Older System)

• For metals that form two common cations:
  • The ion with the higher charge has a name ending in -ic.
  • The ion with the lower charge has a name ending in -ous.
  • Example: $Cu^+$ (cuprous) vs $Cu^{2+}$ (cupric).

Examples of Type I and Type II Compounds

• $CuBr$ -> Copper(I) bromide
• $FeS$ -> Iron(II) sulfide
• $PbO_2$ -> Lead(IV) oxide
• $CoBr_2$ -> Cobalt(II) bromide
• $CaCl_2$ -> Calcium chloride
• Manganese(IV) oxide -> $MnO_2$
• Lead(II) chloride -> $PbCl_2$
• Chromium(III) chloride -> $CrCl_3$
• Gallium iodide -> $GaI_3$

Oxidation States

• A reference for common charges metals can form (e.g., $K$ $1+$, $Fe$ $2+,3+$, $Sn$ $2+,4+$, $V$ $5+(2+,3+,4+)$). This information is crucial for determining Roman numerals in Type II compounds and for writing formulas.

Section 5.3: Naming Binary Compounds That Contain Only Nonmetals (Type III)

• Rules for Naming Type III Binary Compounds:
  1. Formed between two nonmetals.
  2. The first element in the formula is named first, using its full element name.
  3. The second element is named as though it were an anion (root name + -ide).
  4. Prefixes are used to denote the number of atoms of each element present.
  5. The prefix mono- is never used for naming the first element.
• Prefixes:
  • mono- (1)
  • di- (2)
  • tri- (3)
  • tetra- (4)
  • penta- (5)
  • hexa- (6)
  • hepta- (7)
  • octa- (8)
• Examples:
  • $CO_2$ Carbon dioxide
  • $SF_6$ Sulfur hexafluoride
  • $N_2O_4$ Dinitrogen tetroxide
  • $PCl_5$ Phosphorus pentachloride
  • $PCl_3$ Phosphorus trichloride
  • $SO_2$ Sulfur dioxide

Section 5.4: Naming Binary Compounds: A Review

• Flow Chart: The section reiterates the initial flowchart for classifying and naming binary compounds into Type I, II, or III based on metal presence and variable charge.
• Naming Binary Compounds with Metalloids:
  • Determine Nature: Identify if the compound is ionic (metal + metalloid) or covalent (nonmetal + metalloid).
  • Ionic Compounds: The metalloid ends in -ide and uses conventional Type I & II binary nomenclature conventions.
    • Example: $Sr_2Si$ Strontium silicide
    • Example: $Mg_3As_2$ Magnesium arsenide
  • Covalent Compounds: Prefixes are used to determine the number of atoms in the molecule, as in Type III binary compounds.
    • Example: $SiCl_4$ Silicon tetrachloride
    • Example: $AsF_3$ Arsenic trifluoride

Section 5.5: Naming Compounds That Contain Polyatomic Ions

• Polyatomic Ions: Ions composed of two or more atoms covalently bonded together that have an overall positive or negative charge.
• Naming Convention: Naming ionic compounds containing polyatomic ions follows rules similar to those for binary compounds:
  1. The cation is named first, followed by the anion.
  2. The name of a polyatomic cation (e.g., $NH_4^+$) is used directly.
  3. The name of a polyatomic anion (e.g., $SO_4^{2-}$) is used directly.
  • Example: $NH_4C_2H_3O_2$ Ammonium acetate
  • Example: $NaOH$ Sodium hydroxide
  • Example: $Mg(NO_3)_2$ Magnesium nitrate
  • Example: $Fe_3(PO_4)_2$ Iron(II) phosphate

Naming Oxyanions (Polyatomic Anions Containing Oxygen)

• Two Members in Series:
  • The member with the smaller number of $O$ atoms ends with -ite.
  • The member with the larger number of $O$ atoms ends with -ate.
  • Example: $NO_2^-$ (nitrite) and $NO_3^-$ (nitrate).
• More Than Two Members in Series:
  • Uses the prefix hypo- (less than) to name members with the fewest $O$ atoms.
  • Uses the prefix per- (more than) to name members with the most $O$ atoms.
  • Pattern: hypo- (fewest O) <-ite < -ate < per- (most O).

Common Polyatomic Ions

• $NH_4^+$ - ammonium
• $C_2H_3O_2^-$ - acetate
• $BrO_3^-$ - bromate
• $BrO_2^-$ - bromite
• $CN^-$ - cyanide
• $CO_3^{2-}$ - carbonate
• $HCO_3^-$ - hydrogen carbonate (bicarbonate)
• $ClO_3^-$ - chlorate
• $ClO_2^-$ - chlorite
• $ClO^-$ - hypochlorite
• $ClO_4^-$ - perchlorate
• $CrO_4^{2-}$ - chromate
• $Cr_2O_7^{2-}$ - dichromate
• $OH^-$ - hydroxide
• $NO_3^-$ - nitrate
• $NO_2^-$ - nitrite
• $C_2O_4^{2-}$ - oxalate
• $MnO_4^-$ - permanganate
• $O_2^{2-}$ - peroxide
• $PO_4^{3-}$ - phosphate
• $HPO_4^{2-}$ - hydrogen phosphate
• $H_2PO_4^-$ - dihydrogen phosphate
• $PO_3^{3-}$ - phosphite
• $SO_4^{2-}$ - sulfate
• $HSO_4^-$ - hydrogen sulfate (bisulfate)
• $SO_3^{2-}$ - sulfite
• $HSO_3^-$ - hydrogen sulfite (bisulfite)
• $CNS^-$ - thiocyanate
• $IO_3^-$ - iodate
• $S_2O_3^{2-}$ - thiosulfate
• $SiO_3^{2-}$ - silicate
• $AsO_4^{3-}$ - arsenate
• $BO_3^{3-}$ - borate
• $Fe(CN)_6^{3-}$ - ferricyanide
• $H_3O^+$ - hydronium

Section 5.6: Naming Acids and Bases

Acids

• Definition: Molecules that produce $H^+$ ions when dissolved in water. $H^+$ is essentially a bare proton (hydrogen nucleus without electrons).
• Recognition: Acids can be recognized by the hydrogen ($H$) that appears first in the formula (e.g., $HCl$). They are molecules with one or more $H^+$ ions attached to an anion.
• Nomenclature depends on whether the anion contains oxygen.

Rules for Naming Acids (Anions Without Oxygen)

• The acid is named with the prefix hydro- and the suffix -ic attached to the root name for the element.
• Examples:
  • $HF$ - hydrofluoric acid
  • $HCl$ - hydrochloric acid
  • $HBr$ - hydrobromic acid
  • $HI$ - hydroiodic acid
  • $HCN$ - hydrocyanic acid
  • $H_2S$ - hydrosulfuric acid

Rules for Naming Acids (Anions With Oxygen - Polyatomic based)

• The acid name is formed from the root name of the central element of the anion or the anion name, with a suffix of -ic or -ous.
  • If the anion name ends in -ate, the suffix -ic is added to the root name (or ate is replaced by ic).
  • If the anion has an -ite ending, the -ite is replaced by -ous.
• Examples:
  • $H_2SO_4$ - Sulfuric acid (from $SO_4^{2-}$ sulfate)
  • $H_3PO_4$ - Phosphoric acid (from $PO_4^{3-}$ phosphate)
  • $HC_2H_3O_2$ - Acetic acid (from $C_2H_3O_2^-$ acetate)
  • $HNO_3$ - Nitric acid (from $NO_3^-$ nitrate)
  • $HNO_2$ - Nitrous acid (from $NO_2^-$ nitrite)
  • $H_2SO_3$ - Sulfurous acid (from $SO_3^{2-}$ sulfite)
• Example: $HBrO_3$ - Bromic acid (from $BrO_3^-$ bromate)
• Example: $HFO$ - Hypofluorous acid (from $FO^-$ hypofluorite)

Bases

• Definition:
  • A substance that releases a hydroxide ion ($OH^-$) in an aqueous solution (Arrhenius definition).
    • Example: $NaOH$ $ightarrow$ $Na^+$ + $OH^-$ in water.
  • A chemical that accepts a proton ($H^+$) (Brønsted-Lowry definition).
    • Example: $NH_3$ (Ammonia) $NH_3 + H^+ ightleftharpoons NH_4^+$
    • Example: $HCO_3^-$ (Bicarbonate) $HCO_3^- + H^+ ightleftharpoons H_2CO_3$
• Strong Bases:
  • $LiOH$ - Lithium hydroxide
  • $NaOH$ - Sodium hydroxide
  • $KOH$ - Potassium hydroxide
  • $RbOH$ - Rubidium hydroxide
  • $CsOH$ - Cesium hydroxide
  • $Ba(OH)_2$ - Barium hydroxide
  • $Ca(OH)_2$ - Calcium hydroxide
  • $Sr(OH)_2$ - Strontium hydroxide

Acid-Base Strength: pH Scale

• pH: Stands for