Chapter 6 Student Notes

Chapter 6: Ionic Compounds – Periodic Trends and Bonding Theory

Periodic Trends

Periodic trends describe how certain atomic properties change in a predictable way across the periodic table. These trends are essential for understanding chemical reactivity and bonding.

• Ionization Energy: The energy required to remove an electron from an atom in the gas phase.

• Electron Affinity: The energy change when an electron is added to a neutral atom.

• Atomic Radius: The size of an atom, typically measured as half the distance between nuclei in a diatomic molecule.

General Trends:

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

• Electron affinity becomes more negative across a period and less negative down a group.

• Atomic radius decreases across a period and increases down a group.

Exceptions: There are exceptions to these trends, often due to electron configurations and subshell filling.

Ionization Energy

Ionization energy is a measure of how strongly an atom holds onto its electrons. It is influenced by nuclear charge and electron shielding.

• Definition: $IE = \text{energy required to remove an electron}$

• First ionization energy: removal of the first electron.

• Second ionization energy: removal of the second electron, and so on.

• Successive ionization energies increase due to increased effective nuclear charge.

Example: The first ionization energy of Na is lower than that of Mg, but the second ionization energy of Na is much higher than Mg due to removal from a stable noble gas configuration.

Electron Affinity

Electron affinity reflects the tendency of an atom to accept an electron. It is most negative for elements that are one electron away from a full shell (e.g., halogens).

• Definition: $EA = \text{energy change when an electron is added to a neutral atom}$

• High electron affinity = more negative value.

• Halogens have the most negative electron affinities.

• Noble gases have positive electron affinities (do not want to gain electrons).

Example: Cl has a more negative electron affinity than O.

Atomic Radius

Atomic radius is the distance from the nucleus to the outermost electron shell. It is affected by the number of electron shells and effective nuclear charge.

• Decreases across a period due to increased nuclear charge pulling electrons closer.

• Increases down a group due to addition of electron shells.

• Cations are always smaller than their parent atoms; anions are always larger.

Example: Na+ is smaller than Na; Cl- is larger than Cl.

Metallic Character

Metallic character refers to the tendency of an element to lose electrons and form positive ions. It increases down a group and decreases across a period.

• Metals have low ionization energies and are good conductors.

• Nonmetals have high ionization energies and tend to gain electrons.

Ionic and Covalent Bonding

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Ionic Bonding: Transfer of electrons from metals to nonmetals, resulting in oppositely charged ions.

• Covalent Bonding: Sharing of electrons between nonmetals.

Electronegativity: The tendency of an atom to attract electrons in a bond. Higher electronegativity means stronger attraction for electrons.

• Electronegativity increases across a period and decreases down a group.

• Fluorine is the most electronegative element.

Lattice Energy

Lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions. It is a measure of the strength of the ionic bonds in a crystal.

• Definition: $\text{Lattice Energy} = k \frac{Q_1 Q_2}{r}$

• Depends on the charges of the ions ($Q_1$, $Q_2$) and the distance between them ($r$).

• Higher charges and smaller ionic radii result in higher lattice energies.

Example: MgO has a higher lattice energy than NaCl due to higher ionic charges.

Coulomb's Law

Coulomb's Law describes the force of attraction between charged particles. It is fundamental to understanding lattice energy.

• Equation: $F = k \frac{q_1 q_2}{r^2}$

• Force increases with higher charges and decreases with greater distance.

Bond Strength and Bond Energy

Bond strength is related to the energy required to break a bond. Stronger bonds have higher bond energies.

• Triple bonds are stronger and shorter than double bonds; double bonds are stronger and shorter than single bonds.

• Bond breaking is endothermic; bond formation is exothermic.

• Bond Energy: $\text{Bond Energy} = \text{energy required to break one mole of bonds in gaseous molecules}$

Summary Table: Periodic Trends

Property	Across a Period (→)	Down a Group (↓)
Ionization Energy	Increases	Decreases
Electron Affinity	Becomes more negative	Becomes less negative
Atomic Radius	Decreases	Increases
Metallic Character	Decreases	Increases
Electronegativity	Increases	Decreases

Practice Problems and Applications

• Arrange elements in order of increasing ionization energy.

• Predict which compound has the highest lattice energy.

• Identify the most electronegative element in a group.

• Write balanced equations for electron affinity processes.

Additional info: These notes include both conceptual explanations and worked examples, as well as practice questions to reinforce understanding of periodic trends and bonding theory. The content is suitable for exam preparation in a General Chemistry college course.