Chapter 3: Ionic Compounds

Formation of Ions and Ionic Bonds

Ionic compounds are formed when atoms transfer electrons to achieve stable electron configurations. This process results in the formation of ions, which are held together by electrostatic forces known as ionic bonds.

• Ion Formation: Atoms lose or gain electrons to form charged particles called ions. Metals typically lose electrons to form cations (positively charged), while nonmetals gain electrons to form anions (negatively charged).

• Ionic Bond: The attraction between oppositely charged ions forms an ionic bond.

• Characteristics of Ionic Compounds: High melting and boiling points, crystalline structure, and ability to conduct electricity when dissolved in water.

• Example: Sodium (Na) loses one electron to become Na+, and chlorine (Cl) gains one electron to become Cl-. They combine to form NaCl.

Predicting Ion Charge

The charge of an ion can be predicted based on the number of valence electrons or the element's position on the periodic table.

• Group 1: Elements form +1 ions.

• Group 2: Elements form +2 ions.

• Group 17: Elements form -1 ions.

• Group 16: Elements form -2 ions.

• Example: Magnesium (Group 2) forms Mg2+.

Writing Formulas of Ionic Compounds

Formulas are written by combining cations and anions in ratios that result in a neutral compound.

• Steps:

  1. Write the symbol and charge for each ion.

  2. Balance the charges so the total positive and negative charges are equal.

  3. Write the formula without charges, using subscripts as needed.

• Example: Calcium ion (Ca2+) and chloride ion (Cl-):

Polyatomic Ions

Polyatomic ions are groups of atoms covalently bonded that carry a charge. They are common in many ionic compounds.

• Common Polyatomic Ions:

  • Ammonium: NH4+

  • Nitrate: NO3-

  • Sulfate: SO42-

  • Phosphate: PO43-

• Naming Compounds with Polyatomic Ions: Name the cation first, then the polyatomic anion. Example: NaNO3 is sodium nitrate.

Acids and Bases (Preview)

Some common acids and bases are introduced here, but detailed study occurs in later chapters.

• Acids: Compounds that release H+ ions in solution (e.g., HCl, H2SO4).

• Bases: Compounds that release OH- ions in solution (e.g., NaOH, KOH).

Additional info: Detailed properties and reactions of acids and bases are covered in Chapters 5 and 10.

Chapter 4: Molecular Compounds

Covalent Bonding

Covalent bonds form when two atoms share electrons to achieve stable electron configurations. Molecular compounds are composed of molecules held together by covalent bonds.

• General Concept: Atoms share pairs of electrons; each shared pair constitutes a covalent bond.

• Example: In H2O, each hydrogen shares one electron with oxygen.

Comparison of Ionic and Covalent Compounds

Ionic and covalent compounds differ in structure and properties.

• Ionic Compounds: Formed from ions, high melting points, conduct electricity in solution.

• Covalent Compounds: Formed from molecules, lower melting points, do not conduct electricity.

Electronegativity and Periodic Trends

Electronegativity is the ability of an atom to attract shared electrons. It increases across a period and decreases down a group.

• Definition: Electronegativity is a measure of an atom's attraction for electrons in a bond.

• Trend: Highest in the upper right of the periodic table (e.g., fluorine).

• Example: Oxygen is more electronegative than carbon.

Bond Types: Polar, Nonpolar, and Ionic

Bonds are classified based on the difference in electronegativity between atoms.

• Nonpolar Covalent: Electrons are shared equally (e.g., H2).

• Polar Covalent: Electrons are shared unequally (e.g., H2O).

• Ionic: Electrons are transferred (e.g., NaCl).

• Electronegativity Difference:

  • Nonpolar: < 0.5

  • Polar: 0.5 – 1.7

  • Ionic: > 1.7

Predicting Covalent Bonds

The number of covalent bonds an atom forms is related to its group number.

• Group 14: Forms 4 bonds (e.g., carbon).

• Group 15: Forms 3 bonds (e.g., nitrogen).

• Group 16: Forms 2 bonds (e.g., oxygen).

• Group 17: Forms 1 bond (e.g., fluorine).

Lewis Structures and the Octet Rule

Lewis structures represent the arrangement of electrons in molecules. The octet rule states that atoms tend to have eight electrons in their valence shell.

• Drawing Lewis Structures:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to satisfy the octet rule.

• Example: Lewis structure for CO2:

Multiple Covalent Bonds

Double and triple bonds form when atoms share more than one pair of electrons.

• Double Bond: Two pairs of electrons shared (e.g., O2).

• Triple Bond: Three pairs of electrons shared (e.g., N2).

Molecular Geometry and VSEPR Theory

The shape of molecules is predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• VSEPR Theory: Electron pairs repel each other, determining molecular shape.

• Common Geometries:

  • Linear: 180° (e.g., CO2)

  • Bent: <120° (e.g., H2O)

  • Trigonal planar: 120° (e.g., BF3)

  • Tetrahedral: 109.5° (e.g., CH4)

Polarity of Molecules

Molecular polarity depends on both bond polarity and molecular geometry.

• Polar Molecules: Have an uneven distribution of charge (e.g., H2O).

• Nonpolar Molecules: Have an even distribution of charge (e.g., CO2).

• Determining Polarity: Consider both electronegativity and shape.

Naming Binary Molecular Compounds

Binary molecular compounds are named using prefixes to indicate the number of atoms.

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc.

• Example: CO2 is carbon dioxide; N2O is dinitrogen monoxide.

• Writing Formulas: Use the name to determine the number and type of atoms.

Comparison Table: Ionic vs. Covalent Compounds