CHAPTER OVERVIEW: Ionic Compounds

  • Summary of ionic compounds and chemical bonding
    • Chemical bonds are a force of attraction between atoms or ions.
    • Bonds form when atoms share or transfer valence electrons.
    • Valence electrons: electrons in the outer energy level of atoms involved in chemical interactions.

3.1: Ions

3.1.1: Definition of Ions

  • Ions are formed when atoms gain or lose electrons.
  • Differences between cations and anions:
    • Cation: Positively charged ion, formed by the loss of electrons (e.g., Na+ from Na).
    • Anion: Negatively charged ion, formed by the gain of electrons (e.g., Cl- from Cl).

3.1.2: Formation of Ions

  • Atoms have equal numbers of electrons and protons; when they become ions, they deviate from this balance.
    • Sodium atom (Na) loses an electron → Sodium cation (Na+).
    • Chlorine atom (Cl) gains an electron → Chloride anion (Cl-).
  • Pronunciation of ions:
    • Cations: pronounced as CAT-eye-ons
    • Anions: pronounced as ANN-eye-ons

3.1.3: Examples

  • Sodium (Na) loses one electron: Na → Na+ (11 protons and 10 electrons, +1 charge).
  • Chlorine (Cl) gains one electron: Cl → Cl- (17 protons and 18 electrons, -1 charge).

3.2: Ions and the Octet Rule

3.2.1: Role of the Octet Rule in Ion Formation

  • The octet rule states that atoms gain or lose electrons to achieve a stable electron configuration (noble gas configuration).
  • Predicting whether an atom will gain or lose electrons:
    • Main group elements on the left (metals) tend to lose electrons; the ones on the right (nonmetals) tend to gain electrons.

3.2.2: Electron Configurations

  • Example: Sodium (Na) has 1 valence electron (3s1), loses it to become Na+.
  • Example: Chlorine (Cl) has 7 valence electrons (3s2 3p5), gains 1 electron to become Cl-.
  • Main Group trends with respect to electron configurations.

3.2.3: Exceptions to the Octet Rule

  • Some atoms can remain stable with fewer or more than eight electrons, such as in the case of sodium.

3.3: Ions of Some Common Elements

3.3.1: Group Trends on the Periodic Table

  • Elements in the same group typically form ions with the same charge because of their similar valence electron configurations.
    • Alkali metals: +1 charge.
    • Alkaline earth metals: +2 charge.
    • Halogens: -1 charge.

3.3.2: Variable Charges in Transition Metals

  • Transition metals can have varying charges, which complicates ion formation predictions. (e.g., Fe can be Fe2+ or Fe3+).

3.4: Periodic Properties and Ion Formation

3.4.1: Ionization Energy

  • Ionization Energy (IE): Energy required to remove an electron from an atom in the gas phase.
    • Generally increases across a period and decreases down a group.
    • Affected by atomic size and attraction between electrons and protons.

3.4.2: Electron Affinity

  • Electron Affinity (EA): Energy change when an electron is added to a neutral atom.
    • May be negative (more common), positive, or zero, indicating whether energy is released or required.

3.5: Naming Monoatomic Ions

3.5.1: Naming Cations and Anions

  • Cations: Named after the element + “ion” (e.g., Na+ = sodium ion).
  • Anions: Stem of the element name + “-ide” + “ion” (e.g., Cl- = chloride ion).

3.5.2: Iron's Variable Charges

  • Iron can form more than one ion: Fe2+ (iron(II) or ferrous) and Fe3+ (iron(III) or ferric).

3.6: Polyatomic Ions

3.6.1: Definition and Examples

  • Polyatomic ions consist of two or more covalently bonded atoms with an overall charge.
  • Examples: Ammonium ion (NH4+), Hydroxide ion (OH-).

3.6.2: Naming Conventions

  • Polyatomic ions with oxygen may use suffixes -ate and -ite to indicate the number of oxygen atoms.

3.7: Ionic Bonds

3.7.1: Definition of Ionic Bonds

  • An ionic bond is the electrostatic attraction between cations and anions.
  • Strength of ionic bonds depends on charge and distance between ions.

3.7.2: Formation Example: Sodium Chloride

  • NaCl forms when Na+ and Cl- attract due to opposite charges.

3.8: Formulas of Ionic Compounds

3.8.1: Writing Ionic Formulas

  • The total positive charge must equal the total negative charge in ionic compounds (e.g., Na+ + Cl- = NaCl).

3.8.2: Ratio of Ions

  • The charge of the ions determines the ratio in ionic compounds (e.g., Mg2+ requires two Cl-: MgCl2).

3.9: Naming Ionic Compounds

3.9.1: General Naming Rules

  • Name cation first, then anion.
  • Use Stock system for variable charge cations (e.g., Iron(II) chloride for FeCl2).

3.9.2: Examples of Common Ionic Compounds

  • NaCl: sodium chloride (table salt).
  • KI: potassium iodide; NaHCO3: sodium bicarbonate (baking soda).

3.10: Properties of Ionic Compounds

3.10.1: Physical Properties

  • High melting and boiling points due to strong ionic bonds.
  • Ionic compounds are generally hard but brittle.

3.10.2: Conductivity

  • Ionic compounds conduct electricity when dissolved in water or melted, but not as solids.
    • Dissociation example: NaCl(s) → Na+(aq) + Cl-(aq).

3.11: H⁺ and OH⁻ Ions - Introduction to Acids and Bases

3.11.1: Properties of Acids and Bases

  • Acids produce H+ ions in solution, while bases produce OH- ions.
  • Acids are generally sour and bases are bitter; examples: citric acid, hydrochloric acid, sodium hydroxide.

3.11.2: Common Acids and Their Anions

  • List and identify ratios of hydrogen ions to their corresponding anions in various acids.