Chemistry 1110: Ionic and Molecular Compounds (Chapters 3–4) Study Notes

Chapter 3: Ionic Compounds

Anions and Cations

An anion is a negatively charged ion formed when an atom gains electrons, while a cation is a positively charged ion formed when an atom loses electrons. The formulas for ions are written by indicating the element symbol and its charge as a superscript (e.g., Na+, Cl-), but in compound formulas, charges are not shown.

• Anions: Typically formed by nonmetals (e.g., Cl-, O2-).

• Cations: Typically formed by metals (e.g., Na+, Ca2+).

• Writing Formulas: The total positive and negative charges must balance to make the compound neutral.

• Example: Sodium chloride: Na+ + Cl- → NaCl

The Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration with eight valence electrons, similar to noble gases. This rule mainly applies to main-group elements (s- and p-block elements).

• Atoms with fewer than 8 valence electrons will react to achieve an octet.

• Transition metals and some heavier elements may not strictly follow the octet rule.

• Effect on Electron Configuration: Atoms gain or lose electrons to fill or empty their outermost shell.

• Example: Oxygen (6 valence electrons) gains 2 electrons to form O2-.

Periodic Trends: Ionization Energy and Electron Affinity

Ionization energy is the energy required to remove an electron from an atom. Electron affinity is the energy change when an atom gains an electron.

• Ionization energy increases across a period (left to right) and decreases down a group.

• Electron affinity becomes more negative (stronger attraction for electrons) across a period.

• Metals have low ionization energies; nonmetals have high electron affinities.

Ionic Bonds and Properties of Ionic Compounds

An ionic bond is the electrostatic attraction between oppositely charged ions (cations and anions). Ionic compounds have distinct properties:

• High melting and boiling points

• Conduct electricity when molten or dissolved in water

• Usually solid and crystalline at room temperature

• Example: NaCl, MgO

Common Ionic Charges (Figure 3.1)

Elements form ions with characteristic charges based on their group:

Group	Common Charge	Example
1A (Alkali metals)	+1	Na+
2A (Alkaline earth metals)	+2	Ca2+
3A (Aluminum)	+3	Al3+
5A (Nitrogen group)	-3	N3-
6A (Oxygen group)	-2	O2-
7A (Halogens)	-1	Cl-

Naming Cations and Anions

• Cations: Named by the element name (e.g., sodium ion). If a metal can form more than one charge, use Roman numerals (e.g., iron(II) ion for Fe2+).

• Anions: Named by replacing the ending with -ide (e.g., chloride for Cl-).

• Polyatomic ions: Use specific names (e.g., sulfate for SO42-).

Polyatomic Ions (Table 3.3)

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a charge. Memorize their formulas and charges.

Name	Formula	Charge
Ammonium	NH4+	+1
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Carbonate	CO32-	-2
Phosphate	PO43-	-3
Hydroxide	OH-	-1
Acetate	C2H3O2-	-1

Writing Names and Formulas of Ionic Compounds

• Formulas: Write the cation first, then the anion. Use subscripts to balance charges; do not show charges in the formula.

• Names: Name the cation, then the anion. Use Roman numerals for metals with variable charges. For polyatomic ions, use their specific names.

• Example: FeCl3 is iron(III) chloride; Ca(NO3)2 is calcium nitrate.

Acids and Bases (Section 3.11)

• Acids: Substances that produce H+ ions in water (e.g., HCl, H2SO4).

• Bases: Substances that produce OH- ions in water (e.g., NaOH, KOH).

• Recognize acids by the presence of H at the beginning of the formula; bases often contain OH.

Key Terms (Page 95)

• Ionic bond: Electrostatic attraction between ions

• Electrolyte: Substance that conducts electricity when dissolved in water

• Monatomic ion: Ion formed from a single atom

• Polyatomic ion: Ion formed from multiple atoms

• Formula unit: Simplest ratio of ions in an ionic compound

Chapter 4: Molecular Compounds

Covalent Bonds

Covalent bonds are formed when two nonmetal atoms share electrons. The shared electrons allow each atom to achieve a stable electron configuration.

• Nonmetals form covalent bonds with each other.

• Single, double, and triple bonds involve sharing 2, 4, or 6 electrons, respectively.

• Nonpolar covalent bond: Electrons are shared equally.

• Polar covalent bond: Electrons are shared unequally, creating dipoles (partial charges).

• Delta notation: $\delta^+$ (partial positive), $\delta^-$ (partial negative).

Lewis Structures

Lewis structures are diagrams that show the arrangement of valence electrons in a molecule. They help predict bonding and molecular shape.

• Count total valence electrons for all atoms.

• Arrange atoms, connect with single bonds, then complete octets with lone pairs.

• Some elements (e.g., P, S, Cl, Br, I, Xe) can have more than an octet (expanded octet) due to available d orbitals.

• Common bonding preferences (Figure 4.3):

  • H: 1 bond

  • O: 2 bonds

  • N: 3 bonds

  • C: 4 bonds

Coordinate Covalent Bonds

A coordinate covalent bond is a covalent bond in which both electrons come from the same atom. This often occurs in polyatomic ions (e.g., NH4+).

Properties of Molecular Compounds (Table 4.1)

Property	Molecular Compounds
State at Room Temp	Often gases, liquids, or low-melting solids
Melting/Boiling Point	Generally low
Electrical Conductivity	Poor (do not conduct electricity)
Solubility	Varies; many dissolve in nonpolar solvents

VSEPR Theory: Molecular Shapes and Bond Angles

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on the repulsion between electron pairs around a central atom.

• Count the number of charge clouds (bonding and lone pairs) around the central atom.

• Common geometries:

  • Linear: 2 charge clouds, 180°

  • Trigonal planar: 3 charge clouds, 120°

  • Tetrahedral: 4 charge clouds, 109.5°

  • Trigonal pyramidal: 3 bonds + 1 lone pair, ~107°

  • Bent: 2 bonds + 1 or 2 lone pairs, ~104.5°

Electronegativity (Figure 4.7)

Electronegativity is a measure of an atom's ability to attract shared electrons in a bond. The difference in electronegativity determines bond polarity.

• Increases across a period, decreases down a group.

• Fluorine is the most electronegative element.

• Bond type by electronegativity difference:

  • 0–0.4: Nonpolar covalent

  • 0.5–1.9: Polar covalent

  • ≥2.0: Ionic

Polarity of Bonds and Molecules

• Bond polarity depends on electronegativity difference.

• Molecular polarity depends on both bond polarity and molecular geometry.

• Symmetrical molecules (e.g., CO2) can be nonpolar even with polar bonds.

Naming Binary Molecular Compounds (Table 4.4)

• Use prefixes to indicate the number of each atom:

  • 1: mono-

  • 2: di-

  • 3: tri-

  • 4: tetra-

  • 5: penta-

  • 6: hexa-

  • 7: hepta-

  • 8: octa-

  • 9: nona-

  • 10: deca-

• The first element keeps its name; the second gets the -ide suffix.

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

Key Terms (Page 130)

• Covalent bond: Shared pair of electrons between atoms

• Lone pair: Nonbonding pair of electrons

• Polar molecule: Molecule with an uneven distribution of charge

• Dipole: Separation of charge in a bond or molecule

• Resonance: Multiple valid Lewis structures for a molecule

Additional info:

• Practice drawing Lewis structures and predicting molecular shapes for exam preparation.

• Work through end-of-chapter problems to reinforce concepts.