Atoms, Ions, and Molecules - Chapter 2 Notes

Atomic Structure and Subatomic Particles

• Matter is composed of basic building blocks known as atoms, which consist of subatomic particles: electrons, protons, and neutrons.

• The physical properties and quantitative metrics for each subatomic particle are defined as follows:

  • Electron ($e^-$):

  • Mass: $9.10938 \times 10^{-28}\,\text{g}$

  • Charge in Coulombs: $-1.6022 \times 10^{-19}\,\text{C}$

  • Unit Charge: $-1$

  • Proton ($p$):

  • Mass: $1.67262 \times 10^{-24}\,\text{g}$

  • Charge in Coulombs: $+1.6022 \times 10^{-19}\,\text{C}$

  • Unit Charge: $+1$

  • Neutron ($n$):

  • Mass: $1.67493 \times 10^{-24}\,\text{g}$

  • Charge in Coulombs: $0\,\text{C}$

  • Unit Charge: $0$

• Mass relationships between subatomic particles:   $\text{mass } p \approx \text{mass } n \approx 1840 \times \text{mass } e^-$

Atomic Number, Mass Number, and Isotopes

• Atomic Number ($Z$): The total number of protons present in the nucleus of an atom of an element. In a neutral atom, the number of protons is equal to the number of electrons.

• Mass Number ($A$): The total combined number of protons and neutrons in the nucleus of an atom:   $\text{Mass Number } (A) = \text{Number of Protons } (Z) + \text{Number of Neutrons}$

• Isotope Notation (Symbolism): An atom is symbolically represented as ${}_Z^A\text{X}$, where:

  • $\text{X}$ = Element symbol

  • $A$ = Mass number (protons + neutrons)

  • $Z$ = Atomic number (protons)

• Isotopes: Atoms of the exact same element ($\text{X}$) possessing identical atomic numbers ($Z$) but varying mass numbers ($A$) due to differing numbers of neutrons in their nuclei.

• Isotopes of Hydrogen:

  • Protium (${}_1^1\text{H}$): $1\,\text{proton}, 0\,\text{neutrons}, 1\,\text{electron}$

  • Deuterium (${}_1^2\text{H}$): $1\,\text{proton}, 1\,\text{neutron}, 1\,\text{electron}$

  • Tritium (${}_1^3\text{H}$): $1\,\text{proton}, 2\,\text{neutrons}, 1\,\text{electron}$

• Isotopes of Uranium:

  • Uranium-235 (${}_{92}^{235}\text{U}$):

  • $A = 235$

  • $Z = 92$

  • Number of protons = $92$

  • Number of neutrons = $235 - 92 = 143$

  • Uranium-238 (${}_{92}^{238}\text{U}$):

  • $A = 238$

  • $Z = 92$

  • Number of protons = $92$

  • Number of neutrons = $238 - 92 = 146$

• Worked Example — Determining Subatomic Particle Counts:

  • (a) ${}_{11}^{20}\text{Na}$: Protons = $11$, Neutrons = $20 - 11 = 9$, Electrons = $11$

  • (b) ${}_{11}^{22}\text{Na}$: Protons = $11$, Neutrons = $22 - 11 = 11$, Electrons = $11$

  • (c) ${}_8^{17}\text{O}$: Protons = $8$, Neutrons = $17 - 8 = 9$, Electrons = $8$

  • (d) Carbon-14 (${}_6^{14}\text{C}$): Protons = $6$, Neutrons = $14 - 6 = 8$, Electrons = $6$

The Periodic Table

• The Periodic Table is a structured chart that arranges elements possessing similar chemical and physical behavior into organized groups. Elements are listed sequentially in order of increasing atomic number.

• Structural Organization:

  • Periods: Horizontal rows arranged across the periodic table.

  • Groups (Families): Vertical columns containing elements with analogous chemical properties.

• Classification of Elements:

  • Metals: Excellent conductors of heat and electricity. Located on the left and central portions of the table.

  • Nonmetals: Generally poor conductors of heat and electricity. Located on the upper right side of the table.

  • Metalloids: Elements displaying properties intermediate between metals and nonmetals.

• Named Elemental Groups:

  • Group 1A: Alkali Metals (e.g., $\text{Li}, \text{Na}, \text{K}, \text{Rb}, \text{Cs}, \text{Fr}$)

  • Group 2A: Alkaline Earth Metals (e.g., $\text{Be}, \text{Mg}, \text{Ca}, \text{Sr}, \text{Ba}, \text{Ra}$)

  • Group 7A: Halogens (e.g., $\text{F}, \text{Cl}, \text{Br}, \text{I}, \text{At}$)

  • Group 8A: Noble Gases (e.g., $\text{He}, \text{Ne}, \text{Ar}, \text{Kr}, \text{Xe}, \text{Rn}$)

  • d-Block: Transition Metals (comprising Groups 3B through 2B / Groups 3 through 12)

Molecules and Ions

• Atomic Form in Nature: Only Noble Gases (Group 8A) exist naturally as isolated single atoms, designated as monatomic species. All other matter is formed from molecules or ions.

• Molecules: A chemical combination of two or more atoms bound together in a definite arrangement by chemical forces (bonds).

• Diatomic Molecules: Contain exactly two atoms.

  • Homonuclear diatomic elements: $\text{H}_2, \text{N}_2, \text{O}_2, \text{F}_2, \text{Cl}_2, \text{Br}_2, \text{I}_2$

  • Heteronuclear diatomic compounds: $\text{HCl}, \text{CO}$

• Polyatomic Molecules: Contain more than two atoms bound together (e.g., $\text{O}_3, \text{H}_2\text{O}, \text{NH}_3, \text{CH}_4$).

• Ions: An atom or collection of bound atoms possessing a net positive or negative electric charge.

  • Cations: Positively charged species resulting from the loss of one or more electrons by a neutral atom or molecule.     $\text{Na } (11\,p, 11\,e^-) \rightarrow \text{Na}^+ (11\,p, 10\,e^-) + e^-$

  • Anions: Negatively charged species resulting from the gain of one or more electrons by a neutral atom or molecule.     $\text{Cl } (17\,p, 17\,e^-) + e^- \rightarrow \text{Cl}^- (17\,p, 18\,e^-)$

• Ionic Types by Particle Count:

  • Monatomic Ions: Formed from a single atom bearing a charge (e.g., $\text{Na}^+, \text{Cl}^-, \text{Ca}^{2+}, \text{O}^{2-}, \text{Al}^{3+}, \text{N}^{3-}$).

  • Polyatomic Ions: Groups of bound atoms bearing a collective net charge (e.g., $\text{OH}^-, \text{CN}^-, \text{NH}_4^+, \text{NO}_3^-, \text{Hg}_2^{2+}$). The species $\text{Hg}_2^{2+}$ (mercury(I)) acts specifically as a polyatomic cation.

• Common Ion Charges Based on Periodic Table Position:

• Group 1A: $+1$ charge ($\text{H}^+, \text{Li}^+, \text{Na}^+, \text{K}^+, \text{Rb}^+, \text{Cs}^+$)

• Group 2A: $+2$ charge ($\text{Mg}^{2+}, \text{Ca}^{2+}, \text{Sr}^{2+}, \text{Ba}^{2+}$)

• Group 3A: $+3$ charge ($\text{Al}^{3+}, \text{Ga}^{3+}, \text{In}^{3+}$)

• Group 4A: $-4$ charge ($\text{C}^{4-}$)

• Group 5A: $-3$ charge ($\text{N}^{3-}, \text{P}^{3-}$)

• Group 6A: $-2$ charge ($\text{O}^{2-}, \text{S}^{2-}, \text{Se}^{2-}, \text{Te}^{2-}$)

• Group 7A: $-1$ charge ($\text{F}^-, \text{Cl}^-, \text{Br}^-, \text{I}^-$)

• Transition Metals with Fixed Charges: $\text{Zn}^{2+}, \text{Cd}^{2+}, \text{Ag}^+$

• Transition/Heavy Main Group Metals with Variable Charges: $\text{Cr}^{2+}/\text{Cr}^{3+}$, $\text{Mn}^{2+}/\text{Mn}^{3+}$, $\text{Fe}^{2+}/\text{Fe}^{3+}$, $\text{Co}^{2+}/\text{Co}^{3+}$, $\text{Ni}^{2+}/\text{Ni}^{3+}$, $\text{Cu}^+/\text{Cu}^{2+}$, $\text{Au}^+/\text{Au}^{3+}$, $\text{Hg}_2^{2+}/\text{Hg}^{2+}$, $\text{Sn}^{2+}/\text{Sn}^{4+}$, $\text{Pb}^{2+}/\text{Pb}^{4+}$

Chemical Formulas

• Chemical Formula: An expression depicting the elemental composition and numerical ratios of atoms in molecules or ionic compounds using chemical symbols.

• Types of Chemical Formulas:

  • Molecular Formula: Reflects the exact number of atoms of each individual element contained within the smallest fundamental unit of a substance.

  • Empirical Formula: Represents the simplified, smallest whole-number integer ratio of atoms comprising a compound. Empirical formulas are established via experimental measurements.

• Mathematical Relation:   $\text{Molecular Formula} = n \times (\text{Empirical Formula}) \quad \text{where } n = 1, 2, 3, \dots$

• Representation Modes for Molecular Structures:

• Molecular Formula: Lists chemical symbols with subscript counts (e.g., $\text{H}_2, \text{H}_2\text{O}, \text{NH}_3, \text{CH}_4$).

• Structural Formula: Highlights connectivity and covalent bonds between atoms using dashes (e.g., $\text{H}-\text{H}$, $\text{H}-\text{O}-\text{H}$, $\text{H}-\text{N}(\text{H})-\text{H}$, $\text{H}-\text{C}(\text{H})_2-\text{H}$).

• Ball-and-Stick Model: 3D spatial view representing atoms as spheres and bonds as connecting rods.

• Space-Filling Model: 3D model illustrating truncated spheres scaled to match atomic van der Waals radii.

  • Examples of Molecular to Empirical Conversions:

• Phosphorus pentoxide: $\text{P}_4\text{O}_{10} \rightarrow \text{P}_2\text{O}_5$

• Water: $\text{H}_2\text{O} \rightarrow \text{H}_2\text{O}$

• Dinitrogen tetroxide: $\text{N}_2\text{O}_4 \rightarrow \text{NO}_2$

• Decane derivative: $\text{C}_{10}\text{H}_{22} \rightarrow \text{C}_5\text{H}_{11}$

• Organic oxide: $\text{C}_6\text{H}_{18}\text{O}_3 \rightarrow \text{C}_2\text{H}_6\text{O}$

• Pentanol isomer: $\text{C}_5\text{H}_{12}\text{O} \rightarrow \text{C}_5\text{H}_{12}\text{O}$

  • Worked Examples — Formula Determination:

• Methylamine: Based on its structural model containing 1 Carbon atom (black), 1 Nitrogen atom (blue), and 5 Hydrogen atoms (grey), its molecular formula is $\text{CH}_5\text{N}$ (or structural format $\text{CH}_3\text{NH}_2$).

• Diborane ($\text{B}_2\text{H}_6$): Divide subscripts by 2 to yield the empirical formula $\text{BH}_3$.

• Dimethyl fumarate ($\text{C}_8\text{H}_{12}\text{O}_4$): Divide subscripts by 4 to yield the empirical formula $\text{C}_2\text{H}_3\text{O}$.

• Vanillin ($\text{C}_8\text{H}_8\text{O}_3$): Subscripts share no common divisor greater than 1; the empirical formula is identical to its molecular formula, $\text{C}_8\text{H}_8\text{O}_3$.

Ionic Compounds and Crossover Method

• Broad Classification of Compounds:

  • Organic Compounds: Contain carbon in combination with elements such as hydrogen, oxygen, nitrogen, and sulfur.

  • Inorganic Compounds: All remaining chemical combinations, grouped into ionic compounds, molecular compounds, acids/bases, and hydrates.

• Composition of Ionic Compounds:

  • Formed by electrostatic attraction between cations (metals) and anions (nonmetals).

  • The formula unit represents the smallest electrically neutral ratio, matching the empirical formula.

  • The sum of positive charges from cations and negative charges from anions within a formula unit must equal zero.

• The Crossover Method for Writing Formulas:

  • The absolute magnitude of the charge on the cation becomes the subscript for the anion.

  • The absolute magnitude of the charge on the anion becomes the subscript for the cation.

  • Reduce subscripts to their lowest whole-number integer values when possible.

• Worked Example — Magnesium Nitride Formula:

  • Combining $\text{Mg}^{2+}$ and $\text{N}^{3-}$ ions:

  • Cation charge ($+2$) yields anion subscript 2.

  • Anion charge ($-3$) yields cation subscript 3.

  • Resulting formula: $\text{Mg}_3\text{N}_2$

Systematic Chemical Nomenclature

• Overview: Chemical nomenclature provides a standardized method for identifying inorganic substances based on their category: ionic compounds, molecular compounds, acids/bases, or hydrates.

1. Naming Ionic Compounds

• Ionic compounds consist of metals combined with nonmetals or polyatomic ions. Binary ionic compounds contain two distinct elements; ternary ionic compounds contain three or more elements.

Type I Binary Ionic Compounds

• Metals forming Type I compounds have a single, fixed ionic charge. These include Groups 1A and 2A metals, $\text{Al}^{3+}$, $\text{Ga}^{3+}$, $\text{In}^{3+}$, and transition metals $\text{Zn}^{2+}$, $\text{Cd}^{2+}$, and $\text{Ag}^+$.

• Naming Rule: [Name of Cation] + [Root of Anion] + -ide

• Monatomic Anions Nomenclature:

  • Group 4A: $\text{C}^{4-}$ (carbide), $\text{Si}^{4-}$ (silicide) [Note: $\text{C}_2^{2-}$ is also termed carbide]

  • Group 5A: $\text{N}^{3-}$ (nitride), $\text{P}^{3-}$ (phosphide)

  • Group 6A: $\text{O}^{2-}$ (oxide), $\text{S}^{2-}$ (sulfide), $\text{Se}^{2-}$ (selenide), $\text{Te}^{2-}$ (telluride)

  • Group 7A: $\text{F}^-$ (fluoride), $\text{Cl}^-$ (chloride), $\text{Br}^-$ (bromide), $\text{I}^-$ (iodide)

• Examples of Type I Ionic Names:

  • $\text{BaCl}_2$: barium chloride

  • $\text{K}_2\text{O}$: potassium oxide

  • $\text{Mg(OH)}_2$: magnesium hydroxide

  • $\text{KNO}_3$: potassium nitrate

  • $\text{AlCl}_3$: aluminum chloride

Type II Binary Ionic Compounds

• Metals forming Type II compounds display variable oxidation states (variable charges). Includes transition metals and heavier p-block metals like tin ($\text{Sn}$) and lead ($\text{Pb}$).

• Naming Rule: [Name of Cation] + (Charge in Roman Numerals) + [Root of Anion] + -ide

• Examples of Type II Ionic Names:

  • $\text{FeCl}_2$: iron(II) chloride

  • $\text{FeCl}_3$: iron(III) chloride

  • $\text{MnO}$: manganese(II) oxide

  • $\text{Mn}_2\text{O}_3$: manganese(III) oxide

  • $\text{MnO}_2$: manganese(IV) oxide

  • $\text{Cu(NO}_3)_2$: copper(II) nitrate

  • $\text{KH}_2\text{PO}_4$: potassium dihydrogen phosphate

  • $\text{NH}_4\text{ClO}_3$: ammonium chlorate

  • $\text{Cr}_2(\text{Cr}_2\text{O}_7)_3$: chromium(III) dichromate

• Common Polyatomic and Monatomic Ions Reference Table:

• Worked Practice Solutions:

  • Name $\text{Fe(NO}_3)_2$: iron(II) nitrate

  • Name $\text{Na}_2\text{HPO}_4$: sodium hydrogen phosphate

  • Name $\text{(NH}_4)_2\text{SO}_3$: ammonium sulfite

  • Formula for mercury(I) nitrate: $\text{Hg}_2(\text{NO}_3)_2$

  • Formula for cesium oxide: $\text{Cs}_2\text{O}$

  • Formula for strontium nitride: $\text{Sr}_3\text{N}_2$

2. Naming Molecular (Covalent) Compounds

• Binary molecular compounds (Type III) consist of nonmetals bound covalently to nonmetals.

• Greek Numerical Prefixes:

  • 1 = mono-

  • 2 = di-

  • 3 = tri-

  • 4 = tetra-

  • 5 = penta-

  • 6 = hexa-

  • 7 = hepta-

  • 8 = octa-

  • 9 = nona-

  • 10 = deca-

• Naming Rule: [Prefix] + [First Element Name] + [Prefix] + [Base Name of Second Element] + -ide

• Prefix Formatting Rules:

  • The prefix "mono-" is omitted for the first element. $\text{PCl}_3$ is phosphorus trichloride, not monophosphorus trichloride.

  • Drop the final "a" or "o" of a Greek prefix when preceding "oxide". $\text{N}_2\text{O}_4$ is dinitrogen tetroxide rather than dinitrogen tetraoxide.

• Examples of Molecular Compounds:

  • $\text{HI}$: hydrogen iodide

  • $\text{NF}_3$: nitrogen trifluoride

  • $\text{SO}_2$: sulfur dioxide

  • $\text{N}_2\text{Cl}_4$: dinitrogen tetrachloride

  • $\text{NO}_2$: nitrogen dioxide

  • $\text{N}_2\text{O}$: dinitrogen monoxide

  • $\text{PBr}_5$: phosphorus pentabromide

  • $\text{As}_2\text{O}_5$: diarsenic pentoxide

  • Bromine trifluoride: $\text{BrF}_3$

  • Diboron trioxide: $\text{B}_2\text{O}_3$

• Hydrogen-Containing Covalent Exceptions: Compounds of hydrogen with nonmetals often use common or non-prefix names:

  • $\text{B}_2\text{H}_6$: diborane

  • $\text{CH}_4$: methane

  • $\text{SiH}_4$: silane

  • $\text{NH}_3$: ammonia

  • $\text{PH}_3$: phosphine

  • $\text{H}_2\text{O}$: water

  • $\text{H}_2\text{S}$: hydrogen sulfide

3. Naming Acids and Bases

• Acid: A substance that produces hydrogen ions ($\text{H}^+$) when dissolved in water ($\text{aq}$).

• Pure gaseous state compounds differ in name from their aqueous acid forms:

  • $\text{HCl}(g)$ = hydrogen chloride (pure gas)

  • $\text{HCl}(aq)$ = hydrochloric acid (aqueous acid solution)

• Binary Acids: Composed of hydrogen and a single nonmetal element dissolved in water.

  • Naming Rule: hydro- + [Base Name of Nonmetal] + -ic acid

  • $\text{HF}(aq)$: hydrofluoric acid ($\text{F}^-$ fluoride)

  • $\text{HCl}(aq)$: hydrochloric acid ($\text{Cl}^-$ chloride)

  • $\text{HBr}(aq)$: hydrobromic acid ($\text{Br}^-$ bromide)

  • $\text{HI}(aq)$: hydroiodic acid ($\text{I}^-$ iodide)

  • $\text{HCN}(aq)$: hydrocyanic acid ($\text{CN}^-$ cyanide)

  • $\text{H}_2\text{S}(aq)$: hydrosulfuric acid ($\text{S}^{2-}$ sulfide)

• Oxoacids: Acids containing hydrogen, oxygen, and a central nonmetal element. Their formulas list $\text{H}$ first, followed by the central element and $\text{O}$.

• **Reference Oxoacids (