Ionic and Covalent Compounds: Structure, Bonding, and Properties

Ionic and Covalent Compounds

Overview

This chapter introduces the fundamental types of chemical compounds—ionic and covalent—focusing on their formation, properties, and significance in chemistry and biology. The main topics include the nature of chemical bonds, the structure of ionic and covalent compounds, and the importance of ions in biological systems.

3.1 Ionic Compounds Containing Monatomic Ions

Definition and Formation

• Ionic compounds are substances composed of two or more different types of atoms held together by ionic bonds.

• Ionic bonds form when valence electrons are transferred from one atom (typically a metal) to another (typically a nonmetal).

• The resulting oppositely charged ions (cations and anions) are attracted to each other, forming a stable compound.

Key Terms

• Cation: A positively charged ion (e.g., Na+).

• Anion: A negatively charged ion (e.g., Cl-).

• Monatomic ion: An ion consisting of a single atom with a positive or negative charge.

Example: Formation of Sodium Chloride (NaCl)

• Sodium (Na) loses one electron to become Na+.

• Chlorine (Cl) gains one electron to become Cl-.

• The electrostatic attraction between Na+ and Cl- forms the ionic compound NaCl.

Properties of Ionic Compounds

• High melting and boiling points.

• Conduct electricity when dissolved in water (as electrolytes).

• Often form crystalline solids (salts).

• Properties are very different from those of their constituent elements.

Electrolytes in the Body

• Electrolytes are ions dissolved in water that conduct electricity.

• Important biological electrolytes include sodium (Na+), potassium (K+), calcium (Ca2+), and chloride (Cl-).

3.2 Ionic Compounds Containing Polyatomic Ions

Definition

• Polyatomic ions are charged species composed of two or more covalently bonded atoms acting as a single ion.

• Examples include ammonium (NH4+), carbonate (CO32-), and sulfate (SO42-).

Properties

• Polyatomic ions participate in ionic bonding similarly to monatomic ions.

• Each polyatomic ion has a unique name, formula, and charge.

Examples of Common Polyatomic Ions

Formula	Name
NH4+	Ammonium
CO32-	Carbonate
NO3-	Nitrate
SO42-	Sulfate
OH-	Hydroxide

3.3 Covalent Compounds

Definition and Formation

• Covalent compounds are composed of two or more nonmetal atoms joined by covalent bonds.

• Covalent bonds form when valence electrons are shared between atoms.

• These compounds include a wide variety of molecules, such as water, carbon dioxide, and biological macromolecules (proteins, DNA).

Diatomic Elements

• Certain elements exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.

Properties of Covalent Compounds

• Lower melting and boiling points compared to ionic compounds.

• Do not conduct electricity in water (generally non-electrolytes).

• Can exist as gases, liquids, or solids at room temperature.

Naming Binary Covalent Compounds

• Name the first element, then the second element with the ending changed to -ide.

• Use prefixes to indicate the number of each atom (mono-, di-, tri-, tetra-, etc.).

• The prefix mono- is only used for the second element.

Number of Atoms	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-

Example

• CO2: Carbon dioxide

• SF6: Sulfur hexafluoride

• N2O: Dinitrogen monoxide

3.4 Writing Lewis Dot Structures of Covalent Compounds

Lewis Dot Structures

• Lewis dot structures represent the arrangement of valence electrons in a molecule.

• The octet rule states that atoms tend to form bonds to achieve eight electrons in their valence shell (except hydrogen, which achieves two).

• Bonding pairs (shared electrons) and nonbonding pairs (lone pairs) are shown as dots or lines.

Steps for Drawing Lewis Structures

1. Count total valence electrons for all atoms in the molecule.

2. Determine the central atom (usually the least electronegative).

3. Connect atoms with single bonds and distribute remaining electrons as lone pairs to complete octets.

4. Form double or triple bonds if necessary to satisfy the octet rule.

Bonding Patterns

Element	Typical Number of Bonds	Typical Number of Lone Pairs
Carbon (C)	4	0
Nitrogen (N)	3	1
Oxygen (O)	2	2
Halogens (F, Cl, Br, I)	1	3
Hydrogen (H)	1	0

Example: Lewis Structure for Water (H2O)

• Oxygen forms two single bonds with hydrogen atoms and has two lone pairs.

Example: Lewis Structure for Ammonia (NH3)

• Nitrogen forms three single bonds with hydrogen atoms and has one lone pair.

Multiple Bonds

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

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

Potential Energy and Compounds

Stability of Compounds

• Compounds have lower potential energy than the separate atoms that form them.

• Atoms form compounds to attain filled valence energy levels (noble gas configuration).

• Noble gases have filled valence shells and are generally unreactive.

Ionic vs. Covalent Bonding: Summary Table

Property	Ionic Compounds	Covalent Compounds
Bond Formation	Transfer of electrons	Sharing of electrons
Constituent Elements	Metal + Nonmetal	Nonmetals
Physical State	Solid (crystalline)	Solid, liquid, or gas
Melting/Boiling Point	High	Low to moderate
Electrical Conductivity	Conducts when dissolved	Does not conduct

Additional info: The notes also reference the importance of ions in biological systems, the use of Roman numerals for transition metal cations, and the naming conventions for both ionic and covalent compounds. For a complete understanding, students should practice writing formulas and names for a variety of compounds, and draw Lewis structures for molecules of increasing complexity.