Molecules and Compounds

Chemical Bonds and Compound Classification

• Chemical bonds form because they lower the potential energy of the charged particles (protons and electrons) that compose atoms.

• Compounds possess distinct physical and chemical properties entirely different from their constituent elements (e.g., explosive $H_2(g)$ and combustible $O_2(g)$ combine to form flame-extinguishing liquid $H_2O(l)$).

• Law of Definite Proportions: Chemical compounds consist of fixed, definite ratios of elemental atoms (such as a $2:1$ ratio of $H$ to $O$ in water), whereas mixtures can contain variable proportions.

• Ionic Bonds: Result from the transfer of one or more electrons from a metal to a nonmetal, forming cations and anions held together by nondirectional electrostatic forces in a three-dimensional crystal lattice.

• Covalent Bonds: Result from the sharing of valence electron pairs between nonmetal atoms of high ionization energy, yielding discrete molecular units.

Representing Chemical Compounds

• Empirical Formula: Gives the simplest whole-number ratio of atoms of each element in a compound (e.g., $CH$ for benzene, $C_6H_6$).

• Molecular Formula: Gives the actual number of atoms of each element in a discrete molecule of a compound ($\text{molecular formula} = (\text{empirical formula})_n$).

• Structural Formula: Uses lines to represent covalent bonds and illustrates atom connectivity and geometry in space.

• Ball-and-Stick Models: Represent atoms as spheres and chemical bonds as sticks, reflecting molecular shape and bond angles.

• Space-Filling Models: Depict atoms scaled to size where spheres represent electron clouds, illustrating how atoms occupy space within the molecule.

Lewis Theory and the Octet Rule

• Lewis Electron-Dot Structures: Represent valence electrons as dots around the chemical symbol to model bonding and electron distribution.

• Octet Rule: Main-group atoms gain, lose, or share electrons to achieve a stable noble-gas configuration with eight valence electrons ($ns^2 np^6$).

• Duet Rule: Applied to light elements ($H$, $Li$, $Be$, and $B$) that achieve stability with two valence electrons, matching Helium ($He$).

• Expanded Octets: Nonmetals located in period 3 and below can accommodate more than eight electrons when acting as the central atom by utilizing empty valence d-orbitals.

• Bonding Electron Pairs: Pairs of valence electrons shared between atoms; can exist as single bonds ($2$ electrons), double bonds ($4$ electrons), or triple bonds ($6$ electrons).

• Lone Pairs: Nonbonding valence electrons localized on a specific atom.

Ionic Compounds and Lattice Energy

• Formula Unit: The basic repeating unit of an ionic compound representing the smallest electrically neutral collection of ions, not an isolated molecule.

• Lattice Energy: The energy released when separated gaseous ions coalesce to form an ionic solid crystal lattice; lattice energy values are always exothermic ($\Delta H^{\circ} < 0$).

• Coulomb's Law Dependence: Lattice energy is directly proportional to ion charge magnitude and inversely proportional to internuclear distance between ions (e.g., $MgO$ has a higher melting point than $NaCl$ because of $2+$ and $2-$ charges versus $1+$ and $1-$ charges).

• Physical Properties of Ionic Solids:

  • High hardness and brittleness (displacing ions aligns like charges, causing repulsive fracture).

  • Electrical insulators in solid state ($NaCl(s)$) because ions are immobilized in the crystal lattice.

  • Electrical conductors when molten or dissolved in water ($NaCl(aq)$) due to free-flowing mobile ions.

Chemical Nomenclature

• Type I Binary Ionic Compounds (Invariant Metal Charge):

  • Formula: [Name of metal cation] + [Base name of nonmetal anion] + -ide

  • Common invariant cations: Group 1A ($1+$), Group 2A ($2+$), $Al^{3+}$, $Zn^{2+}$, $Cd^{2+}$, $Ag^{+}$.

  • Example: $CaO$ is calcium oxide; $KCl$ is potassium chloride.

• Type II Binary Ionic Compounds (Variable Metal Charge):

  • Formula: [Name of metal cation] + [Roman numeral of charge in parentheses] + [Base name of nonmetal anion] + -ide

  • Used for transition metals and select main-group metals like $Pb$ and $Sn$.

  • Example: $FeCl_2$ is iron(II) chloride; $FeCl_3$ is iron(III) chloride; $CrBr_3$ is chromium(III) bromide.

• Polyatomic Oxyanions:

  • Two-member series: Suffix -ate denotes the ion with more oxygens; -ite denotes the ion with fewer oxygens (e.g., $NO_3^{-}$ nitrate, $NO_2^{-}$ nitrite).

  • Four-member halogen series: per---ate (most $O$), --ate, --ite, hypo---ite (least $O$):

  • $ClO_4^{-}$ (perchlorate)

  • $ClO_3^{-}$ (chlorate)

  • $ClO_2^{-}$ (chlorite)

  • $ClO^{-}$ (hypochlorite)

• Hydrated Ionic Compounds: Named using the anhydrous compound name followed by a Greek numerical prefix (hemi-, mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-) and the word hydrate (e.g., $MgSO_4 \cdot 7H_2O$ is magnesium sulfate heptahydrate).

• Binary Molecular Compounds (Nonmetal + Nonmetal):

  • Name the element with the lower group number first; if in the same group, name the element with the higher period number first.

  • Prefix the second element with a Greek numerical indicator and append -ide. Prefix mono- is omitted on the first element (e.g., $N_2O_5$ is dinitrogen pentoxide; $PCl_5$ is phosphorus pentachloride).

Quantitative Composition and Chemical Formulas

• Formula Mass (Molecular Mass): The sum of atomic masses of all atoms present in a chemical formula:   $\text{Formula Mass} = \sum (\text{number of atoms of element} \times \text{atomic mass of element})$

• Molar Mass Conversions:   $\text{Mass } (g) \xrightarrow{\div \text{Molar Mass}} \text{Moles } (mol) \xrightarrow{\times 6.022 \times 10^{23}} \text{Number of Molecules}$

• Mass Percent Composition:   $\text{Mass \% of element } X = \frac{\text{mass of element } X \text{ in } 1\,mol \text{ of compound}}{\text{molar mass of } 1\,mol \text{ of compound}} \times 100\%$

• Determining Empirical Formulas:

  • Convert elemental masses or percentages (assuming a $100\,g$ sample) to moles.

  • Divide each mole value by the smallest calculated mole value to obtain pseudo-subscripts.

  • Multiply fractional ratios by small integers ($0.5 \times 2$, $0.33 \times 3$, $0.25 \times 4$, $0.20 \times 5$) to achieve whole-number subscripts.

• Determining Molecular Formulas:   $n = \frac{\text{molar mass}}{\text{empirical formula molar mass}}$   $\text{Molecular Formula} = (\text{Empirical Formula})_n$

• Combustion Analysis:

• Used for organic compounds containing $C$, $H$, and often $O$.

• All original Carbon converts to $CO_2$; all original Hydrogen converts to $H_2O$.

• The mass of Oxygen is determined by difference:     $\text{mass } O = \text{mass of sample} - (\text{mass } C + \text{mass } H)$

Organic Compounds and Hydrocarbons

• Organic Compounds: Primarily composed of carbon and hydrogen, frequently bonded with oxygen, nitrogen, phosphorus, and sulfur.

• Carbon Bonding Versatility: Carbon forms four covalent bonds (single, double, or triple) and readily concatenates into straight chains, branched structures, and rings.

• Common Hydrocarbons:

  • Methane ($CH_4$): Single carbon alkane; primary component of natural gas.

  • Propane ($C_3H_8$): Three-carbon straight-chain alkane; LP fuel.

  • n-Butane ($C_4H_{10}$): Four-carbon straight-chain alkane; lighter fluid.

  • n-Pentane ($C_5H_{12}$): Five-carbon straight-chain alkane; gasoline component.

  • Ethene ($C_2H_4$): Two-carbon alkene containing a double bond ($C=C$); plant ripening agent.

  • Ethyne ($C_2H_2$): Two-carbon alkyne containing a triple bond ($C \equiv C$); welding torch fuel.