Chapter 6: Ionic and Molecular Compounds – Study Notes

Chapter 6: Ionic and Molecular Compounds

Learning Goals

This chapter introduces the fundamental concepts of ionic and molecular compounds, focusing on the formation, naming, and properties of ions and molecules. Students will learn to write formulas, determine charges, predict molecular shapes, and understand intermolecular forces.

• Describe ions formed by representative elements

• Write formulas for ionic compounds, including those with variable charges and polyatomic ions

• Name ionic and molecular compounds

• Determine bond polarity and molecular shape

• Classify compounds as polar or nonpolar

• Explain intermolecular forces between ions and molecules

Introduction to Compounds

Formation of Compounds

Elements combine to form compounds in nature, except for noble gases, which rarely form compounds due to their filled valence shells. Compounds are formed by either transferring or sharing electrons to achieve a stable electron configuration, typically an octet (eight valence electrons).

• Octet Rule: Atoms tend to gain, lose, or share electrons to obtain eight valence electrons.

• Ionic Bond: Formed by the transfer of electrons from a metal to a nonmetal.

• Covalent Bond: Formed by the sharing of electrons between nonmetals.

Example: Sodium (Na) transfers an electron to chlorine (Cl) to form sodium chloride (NaCl).

Formation of Ions

How Atoms Become Ions

An ion is formed when an atom loses or gains valence electrons. The loss of electrons results in a positively charged ion (cation), while the gain of electrons results in a negatively charged ion (anion).

• Cation: Atom loses electrons; typically metals in Groups 1A, 2A, and 3A.

• Anion: Atom gains electrons; typically nonmetals in Groups 5A, 6A, and 7A.

• Ionization Energy: Determines whether an atom will lose or gain electrons.

Positively Charged Ions (Cations)

Cations are formed when atoms lose one or more valence electrons, resulting in fewer electrons than protons and a net positive charge.

• Valence electrons are decreased to achieve a stable octet.

• Charge is indicated as a superscript (e.g., Na+).

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

Negatively Charged Ions (Anions)

Anions are formed when atoms gain one or more valence electrons, resulting in more electrons than protons and a net negative charge.

• Electrons are gained until the highest energy level is filled with eight valence electrons.

• Charge is indicated as a superscript (e.g., F-).

• Example: Fluorine atom (F) gains one electron to become F-.

Atoms versus Ions

Size and Electron Configuration

Cations are smaller than their parent atoms due to the loss of an energy level, while anions are larger because they have more valence electrons.

• Cations: Fewer electrons, smaller radius

• Anions: More electrons, larger radius

Example: Sodium ion (Na+) is smaller than sodium atom; chloride ion (Cl-) is larger than chlorine atom.

Electron Configuration of Ions

Achieving Noble Gas Configuration

Atoms gain or lose electrons to reach a full octet, resulting in an electron configuration similar to that of a noble gas.

• The number of electrons in an ion matches that of a noble gas.

• Electron configuration notation shows the arrangement of electrons in shells and subshells.

Species	Electron Configuration
Sodium atom (Na)	$1s^2 2s^2 2p^6 3s^1$
Sodium ion (Na+)	$1s^2 2s^2 2p^6$
Chlorine atom (Cl)	$1s^2 2s^2 2p^6 3s^2 3p^5$
Chloride ion (Cl-)	$1s^2 2s^2 2p^6 3s^2 3p^6$
Argon atom (Ar)	$1s^2 2s^2 2p^6 3s^2 3p^6$

Example: Both Na+ and Cl- have the same electron configuration as argon.

*Additional info: The notes above are based on lecture slides and cover the first sections of Chapter 6, focusing on ionic compounds, ion formation, and electron configuration. Later sections (not shown in images but present in the text) would include molecular compounds, naming conventions, Lewis structures, polarity, and intermolecular forces, which are also relevant for GOB Chemistry students.*