Chapter 3: Ionic Compounds

Chemical Formulas

• Atoms are represented by their periodic table abbreviation.

• A subscript indicates the number of atoms of that element in the compound.

  • Example: Water ($H_2O$) consists of two hydrogen atoms and one oxygen atom.

  • Example: Carbon dioxide ($CO_2$) consists of one carbon atom and two oxygen atoms.

  • Example: Sugar (sucrose) ($C<em>{12}H</em>{22}O_{11}$).

Molecules

• Molecules are formed when two or more atoms are bound together.

Elements vs. Compounds

• Elements: Composed of atoms of the same element (e.g., $O_2$).

• Compounds: Composed of atoms of different elements (e.g., $H_2O$).

Monoatomic Molecules

• Consist of single atoms.

• Noble gases exist as monoatomic molecules:

  • Helium (He)

  • Neon (Ne)

  • Argon (Ar)

  • Krypton (Kr)

  • Xenon (Xe)

  • Radon (Rn)

Diatomic Molecules

• Consist of two atoms of the same element.

• Seven elements naturally exist as diatomic molecules:

  • Hydrogen ($H_2$)

  • Nitrogen ($N_2$)

  • Fluorine ($F_2$)

  • Oxygen ($O_2$)

  • Iodine ($I_2$)

  • Chlorine ($Cl_2$)

  • Bromine ($Br_2$)

Introduction to Bonding

• Bonding: The joining of two atoms in a stable arrangement.

• Atoms gain, lose, or share electrons to achieve the electron configuration of the nearest noble gas in the periodic table.

Types of Bonding

• Ionic Bonds: Result from the transfer of electrons from one element to another.

• Covalent Bonds: Result from the sharing of electrons between two atoms.

Octet Rule

• Atoms form octets to become more stable.

• Achieved by losing, gaining, or sharing valence electrons.

• Results in the formation of ionic or covalent bonds.

Covalent Bonds

• Covalent bonds occur when two atoms share electrons (typically two).

• Usually formed between two nonmetals.

Representations of Molecules with Covalent Bonds

• Molecular Formula: Shows the types and numbers of atoms in a molecule ($CH_4$).

• Structural Formula: Shows the arrangement of atoms and bonds in a molecule (H-C-H).

• Perspective Drawing: Represents the three-dimensional arrangement of atoms and bonds using lines, wedges, and dashed wedges.

  • Solid line: bond in the plane of the page.

  • Dashed wedge: bond behind the page.

  • Wedge: bond out of the page.

• Ball-and-Stick Model: Spheres (atoms) joined by sticks (covalent bonds).

• Space-Filling Model: Illustrates the overall molecular shape without explicitly showing covalent bonds.

Examples

• Hydrogen chloride (HCl)

• Water ($H_2O$)

• Ammonia ($NH_3$)

• Methane ($CH_4$)

Covalent Compounds

• Formed when nonmetals share valence electrons, creating covalent bonds.

• Consist of molecules, which are discrete groups of atoms.

  • Water ($H_2O$): 2 H atoms + 1 O atom

  • Carbon dioxide ($CO_2$): 1 C atom + 2 O atoms

  • Glucose ($C<em>6H</em>{12}O_6$): 6 C atoms + 12 H atoms + 6 O atoms

Ionic Compounds

• Formed when metal atoms transfer electrons to nonmetal atoms, resulting in ionic bonds.

• Examples:

  • Table salt (NaCl)

  • Baking soda ($NaHCO_3$)

  • Milk of magnesia ($Mg(OH)_2$)

Ionic Bonding

• Occurs between metals (lose electrons, positive charge - cation) and nonmetals (gain electrons, negative charge - anion).

  • Example: $2Na + Cl_2 \rightarrow 2NaCl$ (sodium chloride crystals)

Ionic vs. Covalent Compounds

• Ionic: Contain both a metal and a nonmetal.

• Covalent: Consist of two or more nonmetals or a metalloid bonded to a nonmetal.

Ions and Ionic Bonds

• Ionic Bond: Transfer of one or more electrons from one atom to another, resulting in a strong electrical attraction between charged particles.

• Typically occurs between a metal and a nonmetal.

• Ion: A charged particle.

Physical Properties of Ionic Compounds

• Crystalline solids.

• High melting points (e.g., NaCl = $801^\circ C$).

• Extremely high boiling points (e.g., NaCl = $1413^\circ C$).

• Water-soluble.

• When dissolved in water, they separate into cations and anions.

• Aqueous solutions conduct an electric current (electrolytes).

Ions: Cations and Anions

• Ions: Charged species with an unequal number of protons and electrons.

• Ionic compounds consist of oppositely charged ions with strong electrostatic attraction.

Types of Ions

1. Cations: Positively charged ions.

2. Anions: Negatively charged ions.

Octet Rule and Noble Gas Stability

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve a full outer electron shell (8 valence electrons), resembling the electron configuration of noble gases.

• Helium (He) is an exception, stable with 2 valence electrons (duet).

• Examples of noble gas electron configurations:

  • He: $1s^2$ (2 valence electrons)

  • Ne: $1s^22s^22p^6$ (8 valence electrons)

  • Ar: $1s^22s^22p^63s^23p^6$ (8 valence electrons)

  • Kr: $1s^22s^22p^63s^23p^64s^23d^{10}4p^6$ (8 valence electrons)

Determining Ionic Charges

• Atoms gain or lose electrons to achieve a noble gas configuration, resulting in a specific charge.

Cations

• Formed when an atom loses one or more electrons.

• Metals typically form cations.

• Monatomic cations are formed by metals.

Anions

• Formed when an atom gains one or more electrons.

• Nonmetals (except noble gases) form monatomic anions.

Common Cations and Anions

• Main-group elements typically form predictable ions based on their group number.

Cations

• Group 1A: +1 charge

• Group 2A: +2 charge

• Group 3A: +3 charge

Anions

• Group 6A: -2 charge

• Group 7A: -1 charge

Transition Metal Cations

• Some transition metals can form more than one type of cation (variable charge).

Naming Cations

• Main group cations are named after the element (e.g., $Na^+$ = sodium ion).

Naming Transition Metal Cations

• Systematic Name: Use a Roman numeral in parentheses to indicate the charge (e.g., $Fe^{2+}$ = iron(II)).

• Common Name: Use the suffix "-ous" for the lower charge and "-ic" for the higher charge (e.g., $Fe^{2+}$ = ferrous, $Fe^{3+}$ = ferric).

Cations with Variable Valence

• Metals that can form cations with different charges (e.g., iron, copper, gold, tin, lead). The systematic name uses Roman numerals to indicate the charge (e.g., Iron(II), Copper(I)), while the older name uses -ous for the lower charge and -ic for the higher charge (e.g., Ferrous, Cupric).

Naming Anions

• Anions are named by replacing the ending of the element name with the suffix "-ide" (e.g., Chlorine → Chloride).

Naming Compounds of Main Group Metals

• Name the cation first, then the anion.

• Do not specify the charge on the cation or the number of ions.

  • Example: NaCl = sodium fluoride

Naming Compounds of Metals with Variable Charge

1. Determine the charge on the cation.

2. Name the cation using either the systematic (Roman numeral) or common (–ous/–ic) name.

3. Name the anion.

Binary Ionic Compounds

Examples of naming binary ionic compounds:

• $Al<em>2S</em>3$: Aluminum sulfide (Al3+, S2−)

• NaCl: Sodium chloride (Na+, Cl−)

• MgO: Magnesium oxide (Mg2+, O2−)

Determining Chemical Formulas from Names

• Identify the cation and anion and their charges for transition metal compounds.

1. tin(IV) oxide: Sn4+, O2-

2. Balance charges: Two −2 anions are needed for each +4 cation.

3. Write the formula with the cation first, and use subscripts to show how many of each ion is needed to have zero overall charge.

• Final answer = $SnO_2$

Polyatomic Ions

• Groups of atoms that gain or lose electrons, resulting in a charge.

Examples

• Polyatomic Cation:

  • Ammonium ($NH_4^+$)

• Polyatomic Anion:

  • Sulfate ($SO_4^{2-}$

Common Polyatomic Ions

Examples of common polyatomic ions and their formulas:

• Acetate ($CH<em>3CO</em>2^-$)

• Cyanide ($CN^-$)

• Hydroxide ($OH^-$)

• Nitrate ($NO_3^-$)

• Carbonate ($CO_3^{2-}$)

• Phosphate ($PO_4^{3-}$)

• Ammonium ($NH_4^+$)

/

Oxyanion Nomenclature Patterns

• When two oxyanions involve the same element:

  • The one with fewer oxygens ends in "-ite" (e.g., nitrite, sulfite).

  • The one with more oxygens ends in "-ate" (e.g., nitrate, sulfate).

Writing Formulas for Ionic Compounds with Polyatomic Ions

• When a cation and anion of equal charge combine, only one of each ion is needed (e.g., $Na^+ + NO<em>2^- \rightarrow NaNO</em>2$).

• When a cation and anion of unequal charge combine, use the ionic charges to determine the relative number of each ion that is needed (e.g., ).

Naming Ionic Compounds with Polyatomic Ions

• Name the cation and then the anion.

• Do not specify the charge on the cation.

• Do not specify how many ions of each type are needed to balance the charge.

  • Example: $NaHCO_3$ = Sodium bicarbonate (or sodium hydrogen carbonate)

  • Example: $Al<em>2(SO</em>4)_3$ = Aluminum sulfate

Inorganic Nomenclature Summary

1. Name the cation first.

   • If the cation can have more than one possible charge, indicate the charge with a Roman numeral in parentheses.

   • If it is a polyatomic cation, it will end in "-ium".

2. Name the anion.

   • If the anion is an element, change its ending to "-ide".

   • If the anion is a polyatomic ion, write the name of the polyatomic ion.