Ionic and Molecular Compounds

Ions and Ionic Bonds

Atoms gain, lose, or share valence electrons to achieve a stable octet ($8$ valence electrons) matching noble gas configurations:

• Metals: Lose valence electrons to form positively charged cations (e.g., $\text{Na} \rightarrow \text{Na}^+ + e^-$).

• Nonmetals: Gain valence electrons to form negatively charged anions (e.g., $\text{Cl} + e^- \rightarrow \text{Cl}^-$).

• Ion Charge Balance: Total positive charge must equal total negative charge so the compound has a net charge of $0$.

### Ionic Compounds and Nomenclature

Ionic compounds consist of cations and anions held together by strong ionic bonds:

• Physical Properties: High melting points and solid state at room temperature due to strong electrostatic attractions.

• Net Charge: Sum of positive ion charges equals sum of negative ion charges (net charge = $0$).

Naming Ionic Compounds

1. Fixed Charge Metals:

   • Name the metal cation first, followed by the nonmetal anion with an "-ide" suffix (e.g., $\text{K}_2\text{O}$ is potassium oxide).

2. Variable Charge Metals:

   • Transition metals (except $\text{Zn}^{2+}$, $\text{Cd}^{2+}$, and $\text{Ag}^+$ which have fixed charges) can form multiple cations.

   • Use Roman numerals in parentheses to denote the ionic charge (e.g., $\text{FeCl}_2$ is iron(II) chloride, $\text{MnO}_2$ is manganese(IV) oxide).

3. Polyatomic Ions:

   • Groups of covalently bonded nonmetal atoms that carry an overall ionic charge (e.g., $\text{NH}_4^+$ ammonium, $\text{NO}_3^-$ nitrate, $\text{SO}_4^{2-}$ sulfate, $\text{PO}_4^{3-}$ phosphate).

   • Oxoanion Naming: Ions with more oxygen end in -ate; ions with one less oxygen end in -ite (e.g., $\text{SO}_4^{2-}$ is sulfate, $\text{SO}_3^{2-}$ is sulfite).

   • Multiple Polyatomic Ions: Enclose the polyatomic ion in parentheses when adding a subscript (e.g., $\text{Mg}(\text{NO}_3)_2$).

### Molecular Compounds and Lewis Structures

Molecular compounds form when nonmetal atoms share valence electrons through covalent bonds to attain noble gas configurations.

Naming Molecular Compounds

• Greek Prefixes: Indicate the number of atoms of each element: mono- ($1$), di- ($2$), tri- ($3$), tetra- ($4$), penta- ($5$), hexa- ($6$), hepta- ($7$), octa- ($8$), nona- ($9$), deca- ($10$).

• Naming Rules:

  • Omit mono- for the first element (e.g., $\text{CO}_2$ is carbon dioxide; exception: $\text{CO}$ is carbon monoxide).

  • Drop the vowel a or o at the end of a prefix when preceding oxide (e.g., $\text{P}_2\text{O}_5$ is diphosphorus pentoxide).

Lewis Structures

• Diatomic Molecules: Seven elements exist naturally as diatomic molecules: $\text{H}_2$, $\text{N}_2$, $\text{O}_2$, $\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$.

• Bonding Representation:

  • Shared electron pairs (bonding pairs) are drawn as lines between atoms.

  • Unshared electron pairs (lone pairs) are drawn as pairs of dots.

  • Double or triple covalent bonds form when atoms share $2$ or $3$ pairs of electrons to fulfill octet requirements.

• Octet Rule Exceptions:

  • Hydrogen requires only $2$ valence electrons (a single duet).

  • Expanded Octets: Central atoms in Period $3$ or higher (e.g., $\text{P}$, $\text{S}$, $\text{Cl}$, $\text{Br}$, $\text{I}$) can hold $10$ or $12$ valence electrons (e.g., $\text{SF}_6$ has $12$ valence electrons around sulfur).

Electronegativity, Bond Polarity, and Molecular Shapes

Electronegativity measures an atom's attraction for shared electrons in a chemical bond. It increases from left to right across a period and decreases down a group, with Fluorine being the most electronegative element.

Bond Types by Electronegativity Difference ($\text{EN}$)

• Nonpolar Covalent: $\text{EN} = 0.0\text{ to }0.4$ (equal sharing of electrons).

• Polar Covalent: $\text{EN} = 0.5\text{ to }1.8$ (unequal sharing, creating bond dipoles with partial charges $\text{delta}^+ \rightarrow \text{delta}^-$).

• Ionic: $\text{EN} = 1.9\text{ to }3.3$ (complete electron transfer from metal to nonmetal).

VSEPR Theory and Molecular Shapes

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts that electron groups surrounding a central atom orient themselves as far apart as possible to minimize repulsion.

- 2 Electron Groups:

• Linear geometry ($180^\text{o}$ bond angle, e.g., $\text{CO}_2$).

  • 3 Electron Groups:

• 3 bonded, 0 lone pairs: Trigonal planar ($120^\text{o}$ bond angle, e.g., $\text{H}_2\text{CO}$).

• 2 bonded, 1 lone pair: Bent ($120^\text{o}$ bond angle, e.g., $\text{SO}_2$).

• 4 Electron Groups:

  • 4 bonded, 0 lone pairs: Tetrahedral ($109^\text{o}$ bond angle, e.g., $\text{CH}_4$).

  • 3 bonded, 1 lone pair: Trigonal pyramidal ($109^\text{o}$ bond angle, e.g., $\text{NH}_3$).

  • 2 bonded, 2 lone pairs: Bent ($109^\text{o}$ bond angle, e.g., $\text{H}_2\text{O}$).

Molecular Polarity and Intermolecular Forces

Molecular Polarity

• Nonpolar Molecules: Contain nonpolar bonds, or contain polar bonds whose dipoles cancel out due to symmetrical shapes (e.g., $\text{CO}_2$, $\text{CCl}_4$).

• Polar Molecules: Contain polar bonds with dipoles that do not cancel due to asymmetrical shapes, resulting in a net dipole moment (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$).

Intermolecular Forces and Melting Points

Intermolecular forces attract molecules to one another and directly influence physical properties such as melting point:

1. Ionic Bonds: Strongest attraction between full positive and negative charges, resulting in high melting points (e.g., $\text{NaCl}$ melts at $801^\text{o}\text{C}$).

2. Hydrogen Bonds: Strongest intermolecular force between covalent molecules; occurs when $\text{H}$ is bonded directly to $\text{F}$, $\text{O}$, or $\text{N}$ (e.g., $\text{H}_2\text{O}$).

3. Dipole-Dipole Attractions: Attractions between opposite partial charges of permanent dipoles in polar molecules (e.g., $\text{HCl}$).

4. Dispersion Forces: Weak, temporary dipole attractions present between all molecules, particularly nonpolar molecules (lowest melting points, e.g., $\text{F}_2$).