Atoms, Elements, Molecules, Compounds, and Inorganic Nomenclature

Subatomic Particles and Atomic Structure

• All atoms are composed of three primary subatomic particles: protons, neutrons, and electrons.
• The physical properties, mass, and charges of subatomic particles are defined as follows:
  • Proton:
  • Mass in kilograms: $1.67262 \times 10^{-27}\,\text{kg}$
  • Mass in atomic mass units (amu): $1.00727\,\text{amu}$
  • Relative charge: $+1$
  • Charge in Coulombs (C): $+1.60218 \times 10^{-19}\,\text{C}$
  • Neutron:
  • Mass in kilograms: $1.67493 \times 10^{-27}\,\text{kg}$
  • Mass in atomic mass units (amu): $1.00866\,\text{amu}$
  • Relative charge: $0$
  • Charge in Coulombs (C): $0$
  • Electron:
  • Mass in kilograms: $0.00091 \times 10^{-27}\,\text{kg}$
  • Mass in atomic mass units (amu): $0.00055\,\text{amu}$
  • Relative charge: $-1$
  • Charge in Coulombs (C): $-1.60218 \times 10^{-19}\,\text{C}$

The Periodic Table and Element Classifications

• Elements differ from one another based on the number of subatomic particles contained within their atoms.
• The Periodic Table displays three primary pieces of information for each element:
  • Chemical Symbol: A unique representation for each element, consisting of one capital letter or one capital letter followed by a lower-case letter (e.g., $\text{C}$ for carbon, $\text{Na}$ for sodium).
  • Atomic Number ($Z$): The total number of protons present in the nucleus of an atom of the element. Elements are arranged sequentially by increasing atomic number across horizontal rows (periods).
  • Average Atomic Mass: The weighted average mass of all naturally occurring isotopes of an element, reflecting the combined mass of protons and neutrons.

• Major categories and group designations on the Periodic Table include:
  • Alkali Metals: Group 1A (Group 1, excluding hydrogen). Reactive metals forming $+1$ cations.
  • Alkaline Earth Metals: Group 2A (Group 2). Reactive metals forming $+2$ cations.
  • Transition Metals: Groups 3B through 2B (Groups 3 to 12). Metals capable of forming variable cation charges.
  • Halogens: Group 7A (Group 17). Reactive nonmetals forming $-1$ anions.
  • Noble Gases: Group 8A (Group 18). Unreactive monatomic gases.
  • Lanthanides: Inner transition elements with atomic numbers 58 through 71 (from Cerium $\text{Ce}$ to Lutetium $\text{Lu}$).
  • Actinides: Inner transition elements with atomic numbers 90 through 103 (from Thorium $\text{Th}$ to Lawrencium $\text{Lr}$).

Formation of Ions and Predictable Ion Charges

• Elements form ions by gaining or losing electrons to achieve stable electron configurations.
• Predictable charges based on periodic table group placement:
  • Group 1A: Forms $+1$ cations ($\text{Li}^+$, $\text{Na}^+$, $\text{K}^+$, $\text{Rb}^+$, $\text{Cs}^+$). Hydrogen forms $\text{H}^+$ (cation) or hydride $\text{H}^-$ (anion).
  • Group 2A: Forms $+2$ cations ($\text{Mg}^{2+}$, $\text{Ca}^{2+}$, $\text{Sr}^{2+}$, $\text{Ba}^{2+}$).
  • Group 3A: Aluminum forms $\text{Al}^{3+}$.
  • Group 5A: Nitrogen forms nitride $\text{N}^{3-}$.
  • Group 6A: Forms $-2$ anions ($\text{O}^{2-}$, $\text{S}^{2-}$, $\text{Se}^{2-}$, $\text{Te}^{2-}$).
  • Group 7A: Forms $-1$ anions ($\text{F}^-$, $\text{Cl}^-$, $\text{Br}^-$, $\text{I}^-$).
  • Group 8A: Noble gases remain uncharged.

Classification of Matter: Atomic-Level View

• Pure substances are categorized into elements and compounds based on their structural particles.

• Elements:
  • Atomic Elements: Elements whose fundamental structural particles are single, individual atoms (e.g., Neon $\text{Ne}$, Xenon $\text{Xe}$).
  • Molecular Elements: Elements whose fundamental particles are multi-atom molecules composed of two or more bound atoms of the same element.
  • Diatomic Molecular Elements: Hydrogen ($\text{H}_2$), Nitrogen ($\text{N}_2$), Oxygen ($\text{O}_2$), Fluorine ($\text{F}_2$), Chlorine ($\text{Cl}_2$), Bromine ($\text{Br}_2$), Iodine ($\text{I}_2$).
  • Polyatomic Molecular Elements: Phosphorus ($\text{P}_4$), Sulfur ($\text{S}_8$), Selenium ($\text{Se}$).

• Compounds:
  • Molecular Compounds: Compounds whose fundamental particles are molecules composed strictly of nonmetal elements covalently bonded (e.g., Water $\text{H}_2\text{O}$, Nitrogen dioxide $\text{NO}_2$).
  • Ionic Compounds: Compounds composed of cations (metals or polyatomic cations) and anions (nonmetals or polyatomic anions) arranged in a 3D crystalline lattice array of formula units (e.g., Sodium chloride $\text{NaCl}$, Nickel(II) chloride $\text{NiCl}_2$).

Worked Examples: Classifying Pure Substances

• Classification Analysis:
  • Xenon ($\text{Xe}$): Xenon is an element that exists as isolated single atoms, not multi-atom molecules. It is classified as an atomic element.
  • Nickel(II) Chloride ($\text{NiCl}_2$): Composed of a metal (nickel) and a nonmetal (chlorine). It is classified as an ionic compound.
  • Bromine ($\text{Br}_2$): Bromine is an element that exists as a diatomic molecule. It is classified as a molecular element.
  • Nitrogen Dioxide ($\text{NO}_2$): Composed entirely of nonmetals (nitrogen and oxygen). It is classified as a molecular compound.
  • Sodium Nitrate ($\text{NaNO}_3$): Composed of a metal cation ($\text{Na}^+$) and a polyatomic anion ($\text{NO}_3^-$). It is classified as an ionic compound.

Nomenclature of Ionic Compounds

• Ionic compounds are formed from metals (cations) and nonmetals or polyatomic ions (anions) arranged in a formula unit matrix.

Type I Ionic Compounds (Invariant Cation Charge)

• Type I ionic compounds contain a metal whose charge is invariant (constant across all compounds).
• Metals with invariant charges:
  • Group 1A ($+1$
  • Group 2A ($+2$
  • Scandium ($\text{Sc}^{3+}$), Aluminum ($\text{Al}^{3+}$), Zinc ($\text{Zn}^{2+}$), Silver ($\text{Ag}^+$)

• Naming Rule for Type I Binary Ionic Compounds:

• Formula: [Name of Cation (Metal)] + [Base Name of Anion (Nonmetal) + -ide]

• Common Monoatomic Anions:

• Fluorine ($\text{F}^-$) -> Base: fluor -> Fluoride

• Chlorine ($\text{Cl}^-$) -> Base: chlor -> Chloride

• Bromine ($\text{Br}^-$) -> Base: brom -> Bromide

• Iodine ($\text{I}^-$) -> Base: iod -> Iodide

• Oxygen ($\text{O}^{2-}$) -> Base: ox -> Oxide

• Sulfur ($\text{S}^{2-}$) -> Base: sulf -> Sulfide

• Nitrogen ($\text{N}^{3-}$) -> Base: nitr -> Nitride

• Phosphorus ($\text{P}^{3-}$) -> Base: phosph -> Phosphide

  • Type I Practice Examples:

• $\text{KCl}$: Potassium chloride

• $\text{CaO}$: Calcium oxide

• $\text{MgBr}_2$: Magnesium bromide

• $\text{Al}_2\text{S}_3$: Aluminum sulfide

Type II Ionic Compounds (Variable Cation Charge)

• Type II ionic compounds contain a metal that can form cations with different charges.
• Common in transition metals and selected main group metals such as Lead ($\text{Pb}$), Tin ($\text{Sn}$), and Titanium ($\text{Ti}$).

• Naming Rule for Type II Ionic Compounds:

  • Formula: [Name of Cation (Metal)] + (Charge of Cation in Roman Numerals in Parentheses) + [Base Name of Anion + -ide]

• Common Variable-Charge Metals:

• Chromium:

  • $\text{Cr}^{2+}$ -> Chromium(II) [Older name: Chromous]
  • $\text{Cr}^{3+}$ -> Chromium(III) [Older name: Chromic]

• Iron:

  • $\text{Fe}^{2+}$ -> Iron(II) [Older name: Ferrous]
  • $\text{Fe}^{3+}$ -> Iron(III) [Older name: Ferric]

• Cobalt:

  • $\text{Co}^{2+}$ -> Cobalt(II) [Older name: Cobaltous]
  • $\text{Co}^{3+}$ -> Cobalt(III) [Older name: Cobaltic]

• Copper:

  • $\text{Cu}^+$ -> Copper(I) [Older name: Cuprous]
  • $\text{Cu}^{2+}$ -> Copper(II) [Older name: Cupric]

• Tin:

  • $\text{Sn}^{2+}$ -> Tin(II) [Older name: Stannous]
  • $\text{Sn}^{4+}$ -> Tin(IV) [Older name: Stannic]

• Mercury:

  • $\text{Hg}_2^{2+}$ -> Mercury(I) [Older name: Mercurous] (Note: $\text{Hg}_2^{2+}$ is a diatomic polyatomic cation)
  • $\text{Hg}^{2+}$ -> Mercury(II) [Older name: Mercuric]

• Lead:

  • $\text{Pb}^{2+}$ -> Lead(II) [Older name: Plumbous]

  • $\text{Pb}^{4+}$ -> Lead(IV) [Older name: Plumbic]

  • Type II Practice Examples:

• $\text{FeS}$: Iron(II) sulfide (contains $\text{Fe}^{2+}$)

• $\text{Fe}_2\text{S}_3$: Iron(III) sulfide (contains $\text{Fe}^{3+}$)

• $\text{Cu}_2\text{O}$: Copper(I) oxide (contains $\text{Cu}^+$)

• $\text{CuO}$: Copper(II) oxide (contains $\text{Cu}^{2+}$)

• $\text{TiCl}_4$: Titanium(IV) chloride

• $\text{PbBr}_2$: Lead(II) bromide

• Combination of Copper(II) ($\text{Cu}^{2+}$) and Nitride ($\text{N}^{3-}$): $\text{Cu}_3\text{N}_2$

• Combination of Iron(III) ($\text{Fe}^{3+}$) and Bromide ($\text{Br}^-$): $\text{FeBr}_3$

Polyatomic Ions and Oxyanions

• Polyatomic ions consist of multiple atoms held together by covalent bonds that collectively carry an overall ionic charge.

• Common Polyatomic Ions:

  • Acetate: $\text{C}_2\text{H}_3\text{O}_2^-$
  • Carbonate: $\text{CO}_3^{2-}$
  • Hydrogen carbonate (bicarbonate): $\text{HCO}_3^-$
  • Hydroxide: $\text{OH}^-$
  • Nitrite: $\text{NO}_2^-$
  • Nitrate: $\text{NO}_3^-$
  • Chromate: $\text{CrO}_4^{2-}$
  • Dichromate: $\text{Cr}_2\text{O}_7^{2-}$
  • Phosphate: $\text{PO}_4^{3-}$
  • Hydrogen phosphate: $\text{HPO}_4^{2-}$
  • Dihydrogen phosphate: $\text{H}_2\text{PO}_4^-$
  • Ammonium: $\text{NH}_4^+$
  • Hypochlorite: $\text{ClO}^-$
  • Chlorite: $\text{ClO}_2^-$
  • Chlorate: $\text{ClO}_3^-$
  • Perchlorate: $\text{ClO}_4^-$
  • Permanganate: $\text{MnO}_4^-$
  • Sulfite: $\text{SO}_3^{2-}$
  • Hydrogen sulfite (bisulfite): $\text{HSO}_3^-$
  • Sulfate: $\text{SO}_4^{2-}$
  • Hydrogen sulfate (bisulfate): $\text{HSO}_4^-$
  • Cyanide: $\text{CN}^-$
  • Peroxide: $\text{O}_2^{2-}$

• Oxyanion Nomenclature System:

  • Oxyanions are polyatomic ions containing oxygen and another element.
  • When a series contains two oxyanions:
  • The ion with more oxygen atoms ends in -ate (e.g., $\text{NO}_3^-$ nitrate, $\text{SO}_4^{2-}$ sulfate).
  • The ion with fewer oxygen atoms ends in -ite (e.g., $\text{NO}_2^-$ nitrite, $\text{SO}_3^{2-}$ sulfite).
  • When a series contains four oxyanions (e.g., halogen oxyanions):
  • Fewest oxygen atoms: hypo- + -ite (e.g., $\text{ClO}^-$ hypochlorite, $\text{BrO}^-$ hypobromite)
  • Fewer oxygen atoms: -ite (e.g., $\text{ClO}_2^-$ chlorite, $\text{BrO}_2^-$ bromite)
  • More oxygen atoms: -ate (e.g., $\text{ClO}_3^-$ chlorate, $\text{BrO}_3^-$ bromate)
  • Most oxygen atoms: per- + -ate (e.g., $\text{ClO}_4^-$ perchlorate, $\text{BrO}_4^-$ perbromate)

• Polyatomic Practice Examples:

  • $\text{NH}_4\text{Cl}$: Ammonium chloride
  • $\text{Ca}(\text{C}_2\text{H}_3\text{O}_2)_2$: Calcium acetate
  • $\text{Cu}(\text{NO}_3)_2$: Copper(II) nitrate
  • Aluminum ($\text{Al}^{3+}$) and Sulfate ($\text{SO}_4^{2-}$): $\text{Al}_2(\text{SO}_4)_3$
  • Chromium(II) ($\text{Cr}^{2+}$) and Hydrogen carbonate ($\text{HCO}_3^-$): $\text{Cr}(\text{HCO}_3)_2$

Hydrated Ionic Compounds

• Hydrates are ionic compounds containing a specific stoichiometric ratio of water molecules incorporated directly into their formula unit crystal matrix.
• Numerical Prefixes for Hydrates:
  • $\text{hemi} = \frac{1}{2}$
  • $\text{mono} = 1$
  • $\text{di} = 2$
  • $\text{tri} = 3$
  • $\text{tetra} = 4$
  • $\text{penta} = 5$
  • $\text{hexa} = 6$
  • $\text{hepta} = 7$
  • $\text{octa} = 8$
  • $\text{nona} = 9$
  • $\text{deca} = 10$
• Hydrate Examples:
  • $\text{MgSO}_4 \cdot 7\text{H}_2\text{O}$: Magnesium sulfate heptahydrate (Epsom salts)
  • $\text{CoCl}_2 \cdot 6\text{H}_2\text{O}$: Cobalt(II) chloride hexahydrate
  • $\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$: Calcium sulfate hemihydrate
  • $\text{BaCl}_2 \cdot 6\text{H}_2\text{O}$: Barium chloride hexahydrate
  • $\text{CuSO}_4 \cdot 6\text{H}_2\text{O}$: Copper(II) sulfate hexahydrate

Nomenclature of Binary Molecular Compounds

• Molecular compounds are formed between two nonmetals.
• Nonmetals can combine in multiple fixed ratios to form different unique compounds (e.g., $\text{NO}$, $\text{NO}_2$, $\text{N}_2\text{O}$, $\text{N}_2\text{O}_3$, $\text{N}_2\text{O}_4$, $\text{N}_2\text{O}_5$).
• Ordering rule: If both elements lie in the same periodic table group, list the element with the greater row (period) number first.

• Naming Pattern:
  • [Prefix] + [Name of First Element] + [Prefix] + [Base Name of Second Element + -ide]
• Rules for Prefixes:
  • Omit mono- if there is only one atom of the first element.
  • Drop the final letter a or o from a prefix when the element name starts with a vowel (e.g., monoxide, pentoxide).
• Molecular Compound Examples:
  • $\text{NI}_3$: Nitrogen triiodide
  • $\text{PCl}_5$: Phosphorus pentachloride
  • $\text{P}_4\text{S}_{10}$: Tetraphosphorus decasulfide

Nomenclature and Formula Writing for Acids

• Acids are molecular compounds that form hydrogen ions ($\text{H}^+$) when dissolved in water.
• Formulas must be followed by $\text{(aq)}$ to designate an aqueous solution. Un-dissolved compounds are not named as acids.

Binary Acids

• Binary acids consist of hydrogen and one other nonmetal element.

• Naming Formula: hydro- + [Base Name of Nonmetal + -ic] + acid
• Examples:
  • $\text{HCl}(aq)$: Hydrochloric acid
  • $\text{H}_2\text{S}(aq)$: Hydrosulfuric acid
  • $\text{HBr}(aq)$: Hydrobromic acid

Oxyacids

• Oxyacids consist of hydrogen bound to an oxyanion (a polyatomic ion containing oxygen).
• Naming Rules for Oxyacids:
  • If the oxyanion name ends in -ate: Change ending to -ic + acid.

• If the oxyanion name ends in -ite: Change ending to -ous + acid.

• Oxyacid Examples:
  • $\text{HClO}_3(aq)$ (contains chlorate $\text{ClO}_3^-$): Chloric acid
  • $\text{HC}_2\text{H}_3\text{O}_2(aq)$ (contains acetate $\text{C}_2\text{H}_3\text{O}_2^-$): Acetic acid
  • $\text{H}_3\text{PO}_4(aq)$ (contains phosphate $\text{PO}_4^{3-}$): Phosphoric acid
  • $\text{H}_2\text{SO}_3(aq)$ (contains sulfite $\text{SO}_3^{2-}$): Sulfurous acid

Writing Acid Formulas

• When given an acid name ending in acid, the formula always begins with $\text{H}^+$ followed by the anion.
• Balance charges to achieve a neutral formula as if writing an ionic compound.
• Presence of hydro- prefix denotes a binary acid.
• Absence of hydro- prefix denotes an oxyacid.
• Practice Examples for Writing Formulas:
  • Chlorous Acid: Derived from chlorite ($\text{ClO}_2^-$) + $\text{H}^+$ -> $\text{HClO}_2(aq)$
  • Phosphoric Acid: Derived from phosphate ($\text{PO}_4^{3-}$) + $\text{H}^+$ -> $\text{H}_3\text{PO}_4(aq)$
  • Hydrobromic Acid: Binary acid derived from bromide ($\text{Br}^-$) + $\text{H}^+$ -> $\text{HBr}(aq)$

Inorganic Nomenclature Flowchart

• Systematic decision tree for determining inorganic chemical formulas and names:

• Step-by-Step Classification Path:
  1. IONIC (Metal + Nonmetal / Polyatomic Ion):
  • Metal forms one type of ion only (Type I):
    • Pattern: [Name of cation] + [Base name of anion + -ide]
    • Example: $\text{CaI}_2$ -> Calcium iodide
  • Metal forms more than one type of ion (Type II):
    • Pattern: [Name of cation] + (Charge in Roman numerals) + [Base name of anion + -ide]
    • Example: $\text{FeCl}_3$ -> Iron(III) chloride
  1. MOLECULAR (Nonmetals only):
  • Pattern: [Prefix] + [1st Element] + [Prefix] + [Base name of 2nd element + -ide]
  • Example: $\text{P}_2\text{O}_5$ -> Diphosphorus pentoxide
  1. ACIDS (Formula starts with H, in aqueous solution):
  • Binary acids (Two-element, no oxygen):
    • Pattern: hydro- + [Base name of nonmetal + -ic] + acid
    • Example: $\text{HCl}(aq)$ -> Hydrochloric acid
  • Oxyacids (Contains oxygen):
    • Oxyanion ends in -ate: [Base name of oxyanion + -ic] + acid (Example: $\text{H}_3\text{PO}_4$ -> Phosphoric acid)
    • Oxyanion ends in -ite: [Base name of oxyanion + -ous] + acid (Example: $\text{H}_2\text{SO}_3$ -> Sulfurous acid)