Skip to main content
Back

Ionic Compounds: Structure, Formation, and Properties

NotesPracticeVideo lessons

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Chapter 3: Ionic Compounds

3.1 Ions

Ions are charged particles formed when atoms gain or lose electrons. They play essential roles in biological and chemical processes, such as nerve signal transmission.

  • Cation: A positively charged ion formed when a neutral atom loses one or more electrons.

  • Anion: A negatively charged ion formed when a neutral atom gains one or more electrons.

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

  • Properties of Ionic Compounds (e.g., NaCl):

    • High melting points

    • Stable, white, crystalline solids

    • Soluble in water

    • Conduct electricity when dissolved in water (due to free-moving ions)

3.2 Ions and the Octet Rule

The octet rule states that main group elements tend to react to achieve eight valence electrons, resembling the electron configuration of noble gases.

  • Metals lose electrons to form cations with noble gas configurations (preceding noble gas).

  • Nonmetals gain electrons to form anions with noble gas configurations (following noble gas).

  • Example (Magnesium):

    • Electron configuration:

    • Mg loses 2 electrons to form Mg2+ with a neon configuration:

  • Example (Nitrogen):

    • Electron configuration:

    • N gains 3 electrons to form N3– with a neon configuration:

3.3 Ions of Some Common Elements

The periodic table helps predict the charges of ions formed by main group elements.

  • Group 1A: Form +1 cations (e.g., Na+)

  • Group 2A: Form +2 cations (e.g., Ca2+)

  • Group 3A: Form +3 cations (e.g., Al3+)

  • Group 5A, 6A, 7A: Form anions; charge = 8 – group number (e.g., O2–, Cl–)

  • Transition Metals: Can form cations of multiple charges by losing s and d electrons; octet rule not strictly followed.

  • Noble Gases (Group 8A): Generally do not form ions.

3.4 Periodic Properties and Ion Formation

Ion formation is influenced by periodic trends such as ionization energy and electron affinity.

  • Ionization Energy: Energy required to remove an electron from a gaseous atom. Lower values mean easier cation formation (e.g., alkali metals).

  • Electron Affinity: Energy released when an atom gains an electron. Higher values mean easier anion formation (e.g., halogens).

  • Trends:

    • Metals (left side) lose electrons easily (low ionization energy).

    • Nonmetals (right side) gain electrons easily (high electron affinity).

    • Noble gases do not form ions easily.

3.5 Naming Monoatomic Ions

Monoatomic ions are named based on their element and charge.

  • Main Group Metal Cations: Name = element + "ion" (e.g., potassium ion, K+).

  • Transition Metal Cations:

    • Old system: -ous (lower charge), -ic (higher charge) (e.g., ferrous, ferric).

    • New system: Roman numeral in parentheses (e.g., iron(II) ion, iron(III) ion).

  • Anions: Replace the ending with "-ide" + "ion" (e.g., chloride ion, Cl–).

Transition Metal

Old Name

New Name

Fe2+

Ferrous ion

Iron(II) ion

Fe3+

Ferric ion

Iron(III) ion

Cu+

Cuprous ion

Copper(I) ion

Cu2+

Cupric ion

Copper(II) ion

Cr2+

Chromous ion

Chromium(II) ion

Cr3+

Chromic ion

Chromium(III) ion

3.6 Polyatomic Ions

A polyatomic ion is composed of two or more atoms covalently bonded, carrying a net charge. These ions are common in chemistry and must be memorized.

Name

Formula

Charge

Ammonium

NH4+

+1

Nitrate

NO3–

–1

Sulfate

SO42–

–2

Carbonate

CO32–

–2

Phosphate

PO43–

–3

Hydroxide

OH–

–1

3.7 Ionic Bonds

Ionic bonds are the electrostatic attractions between oppositely charged ions in an ionic compound. These bonds hold the ions together in a crystal lattice structure.

  • Formed by electron transfer from a metal to a nonmetal.

  • Result in ionic solids with repeating patterns of cations and anions.

  • Example: NaCl crystal lattice consists of alternating Na+ and Cl– ions.

3.8 Formulas of Ionic Compounds

Ionic compounds are electrically neutral. Their formulas reflect the lowest whole-number ratio of cations to anions that results in zero net charge.

  • Steps to Write Formulas:

    1. Identify the cation and anion, including their charges.

    2. Determine the ratio needed for total charge to be zero.

    3. Write the cation first, then the anion. Do not include charges in the formula.

    4. Use parentheses for polyatomic ions if more than one is needed.

  • Example: Calcium nitrate: Ca2+ and NO3– combine to form Ca(NO3)2.

  • Formula Unit: The simplest ratio of ions in an ionic compound (e.g., NaCl, CaF2).

3.9 Naming Ionic Compounds

The name of an ionic compound lists the cation first, followed by the anion. The naming convention depends on whether the metal forms only one type of ion (Type I) or more than one (Type II).

  • Type I: Main group metals (fixed charge). Do not specify the charge.

    • Example: NaCl is sodium chloride.

    • MgCO3 is magnesium carbonate.

  • Type II: Transition metals (variable charge). Specify the charge with a Roman numeral or old naming system.

    • FeCl2: iron(II) chloride or ferrous chloride.

    • FeCl3: iron(III) chloride or ferric chloride.

  • Do not use prefixes (e.g., di-, tri-) in ionic compound names.

Compound

Name

NaBr

Sodium bromide

CaBr2

Calcium bromide

Na2SO4

Sodium sulfate

CaSO4

Calcium sulfate

Na3PO4

Sodium phosphate

Ca3(PO4)2

Calcium phosphate

3.10 Some Properties of Ionic Compounds

Ionic compounds have characteristic physical properties due to the strong electrostatic forces between ions.

  • High melting and boiling points (e.g., NaCl melts at 801 °C, boils at 1413 °C).

  • Hard, brittle solids that shatter when struck.

  • Conduct electricity when dissolved in water (electrolytes).

  • Solubility in water varies; not all ionic compounds are water-soluble.

3.11 H+ and OH– Ions: An Introduction to Acids and Bases

The hydrogen ion (H+) and hydroxide ion (OH–) are central to acid-base chemistry.

  • Acid: A substance that provides H+ ions in water. The H+ ion (proton) often associates with water to form the hydronium ion (H3O+).

  • Base: A substance that provides OH– ions in water.

  • Examples of Acids:

    • Hydrochloric acid (HCl): provides one H+ per molecule.

    • Sulfuric acid (H2SO4): provides two H+ per molecule.

    • Phosphoric acid (H3PO4): provides three H+ per molecule.

  • Examples of Bases:

    • Sodium hydroxide (NaOH): provides one OH– per formula unit.

    • Barium hydroxide (Ba(OH)2): provides two OH– per formula unit.

Acid

Formula

Anion Derived

Hydrochloric acid

HCl

Cl– (chloride)

Sulfuric acid

H2SO4

SO42– (sulfate)

Phosphoric acid

H3PO4

PO43– (phosphate)

Acetic acid

CH3COOH

CH3COO– (acetate)

Additional info: The above tables and examples are inferred from standard GOB Chemistry content and the context of the provided notes.

Pearson Logo

Study Prep