Chapter 3: Ionic Compounds

Ions and Ion Formation

Ionic compounds are formed when metals and nonmetals combine, resulting in the transfer of electrons. This process creates ions: metals lose electrons to become cations (positively charged), while nonmetals gain electrons to become anions (negatively charged).

• Cation: An ion with a positive charge, formed by the loss of electrons (e.g., Na+, Al3+).

• Anion: An ion with a negative charge, formed by the gain of electrons (e.g., Cl-, O2-).

• Example: Sodium (Na) loses one electron to become Na+; chlorine (Cl) gains one electron to become Cl-.

Ionic Bonds and Properties

Ionic bonds occur between metals and nonmetals, where electrons are transferred from the metal to the nonmetal.

• Bonding Types:

  • Metallic: Metal + metal

  • Covalent: Nonmetal + nonmetal

  • Ionic: Metal + nonmetal

• Properties of Ionic Solids:

  • High melting points

  • Brittle

  • Good electrolytes when dissolved in water

Writing Formulas for Ionic Compounds

The formula for an ionic compound reflects the simplest whole-number ratio of cations to anions that results in a neutral compound.

• Charge Balance: The total positive charge must equal the total negative charge.

• Formula Examples:

  • Al3+ + F- → AlF3

  • Al3+ + O2- → Al2O3

  • Al3+ + N3- → AlN

• Swap & Drop Method: Swap the charges and use them as subscripts for the opposite ion.

Naming Ionic Compounds

• The cation is named first and retains the element name.

• If the cation is a transition metal, its oxidation number is indicated in parentheses with Roman numerals (e.g., FeCl3: iron(III) chloride).

• The anion name ends in -ide (e.g., chloride, oxide).

• Examples:

  • NaCl: sodium chloride

  • AuCl3: gold(III) chloride

  • PbO2: lead(IV) oxide

Polyatomic Ions

Polyatomic ions are groups of atoms that carry a charge and act as a single unit in compounds.

• Examples:

  • Phosphate: PO43-

  • Acetate: C2H3O2-

  • Carbonate: CO32-

  • Ammonium: NH4+

• When writing formulas, polyatomic ions are treated as a unit and their charge is used for balancing.

Naming Ionic Compounds with Polyatomic Ions

• The cation (element or ammonium) is named first.

• The polyatomic anion retains its name (e.g., calcium sulfate, ammonium chloride).

Chapter 4: Covalent Bonding

Covalent Bonds and Properties

Covalent bonds form when two nonmetals share electrons.

• Bond Formation: Electrons are shared between atoms.

• Properties:

  • Low boiling points

  • Soft solids or liquids

  • Electrons are localized between nuclei

• Example: H2O (water) is formed by sharing electrons between hydrogen and oxygen.

Lewis Dot Structures

Lewis dot structures represent the arrangement of valence electrons in molecules.

• Central atom is usually the least numerous in the formula.

• Pair up unpaired electrons to form bonds.

• Check electron count to satisfy the octet rule (exceptions: H needs 2e-, B needs 6e-).

• Example:

  • H2O: 2(1) + 6 = 8 valence electrons

  • BF3: 3 + 3(7) = 24 valence electrons

Octet Rule and Exceptions

Atoms gain, lose, or share electrons to achieve a full set of eight valence electrons.

• Hydrogen only needs 2 electrons.

• Boron only needs 6 electrons.

Multiple Bonds and Unpaired Electrons

• Double and triple bonds are formed when atoms share more than one pair of electrons.

• Example: O2 has a double bond; N2 has a triple bond.

The Seven Diatomic Elements

• H2, N2, O2, F2, Cl2, Br2, I2

Naming Covalent Molecules

• The first element retains its name; the second element ends in -ide.

• Prefixes indicate the number of each element (mono-, di-, tri-, etc.).

• Example: CO2 is carbon dioxide; N2O4 is dinitrogen tetroxide.

Molecular Geometry and VSEPR Theory

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) Theory predicts the shape of molecules based on minimizing repulsion between electron pairs.

• Molecules adopt shapes that minimize electron-electron repulsions.

• AXE notation:

  • A = central atom

  • X = number of atoms bonded to central atom

  • E = number of lone pairs on central atom

Main Molecular Shapes

• Linear: 180°

• Trigonal planar: 120°

• Tetrahedral: 109.5°

• Trigonal bipyramidal: 120°, 180°

• Octahedral: 90°, 180°

Determining Molecular Geometry

• Draw the Lewis structure.

• Apply AXE notation to determine geometry.

• Examples:

  • PH3: AX3E (trigonal pyramidal)

  • H2S: AX2E2 (bent)

Hybrid Orbitals

Hybridization

Hybrid orbitals are formed by blending atomic orbitals to accommodate the number of electron pairs around a central atom.

• AXE notation helps determine the number and type of hybrid orbitals.

• Examples:

  • AX2: sp hybridization

  • AX3: sp2 hybridization

  • AX4: sp3 hybridization

Bond Types: Sigma and Pi Bonds

Sigma Bonds

• Formed by end-on overlap of atomic orbitals.

• Single bonds are always sigma bonds.

Pi Bonds

• Formed by side-to-side overlap of atomic orbitals.

• Double and triple bonds contain one sigma bond and one or more pi bonds.

Polarity of Bonds and Molecules

Polar and Nonpolar Covalent Bonds

• Electrons are not always shared equally; differences in electronegativity cause polarity.

• Polar bonds have permanent positive and negative charges (e.g., HF).

• Nonpolar bonds share electrons equally (e.g., F2).

Polarity of Molecules

• Asymmetric molecules (e.g., H2O) are polar.

• Symmetric molecules (e.g., CO2) are nonpolar.

Acids and Bases

Acid and Base Characteristics

• Acids: sour taste, corrosive, react with indicators, increase [H3O+]

• Bases: bitter taste, slippery feel, corrosive, react with indicators, increase [OH-]

pH Scale

• pH = 7: neutral

• pH < 7: acidic

• pH > 7: basic

Acid Examples

• HC2H3O2: acetic acid (vinegar)

• HNO3: nitric acid

• H2CO3: carbonic acid

• HCl: hydrochloric acid

• H3PO4: phosphoric acid

• H2SO4: sulfuric acid

Base Examples

• Al(OH)3: aluminum hydroxide

• Ca(OH)2: calcium hydroxide

• Mg(OH)2: magnesium hydroxide

• NaOH: sodium hydroxide

• NH3: ammonia

Acid and Base Strength

• Strong acids: dissociate completely in water.

• Weak acids: only partially dissociate.

• Strong bases: ionize completely in water.

• Weak bases: only partially ionize.

Naming Acids

• Binary acids: hydro- + root + -ic (e.g., hydrochloric acid)

• Oxyacids: root of anion + suffix + acid (e.g., nitric acid, nitrous acid)

Additional info:

• Some images (e.g., image_1, image_2, image_11) were not included as they are not directly relevant to the chemistry concepts explained.

• Tables and periodic tables (image_5, image_6) are referenced where appropriate for element and ion information.