Chapter 9 Lecture Notes - CHEM 1113 Broering
Page 1: Introduction to Periodicity and Ionic Bonding
The slide mentions various chemistry textbooks and copyright information.
The focus is on Chapter 9: Periodicity and Ionic Bonding.
Page 2: Valence and Core Electrons
Valence Electrons:
Outermost electrons that display periodicity in their configurations.
Involved in bonding, elements with similar configurations exhibit similar chemical properties.
Core Electrons:
Electrons not involved in bonding, remaining in lower energy levels.
Example of Electron Configurations:
Oxygen (O): 1s² 2s² 2p⁴
Sulfur (S): 1s² 2s² 2p⁶ 3s² 3p⁴
Selenium (Se): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴
Tellurium (Te): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁴
All four elements share the ns²np⁴ configuration in their highest energy levels.
Page 3: Electrostatic Principles
Principles:
Opposite charges attract, like charges repel.
Attraction/repulsion increases with charge magnitude.
Closer charged entities experience stronger forces.
Interactions between electrons and nucleus follow these principles.
Page 4: Energy Concepts of Electrons
Coulomb’s Law:
Potential energy of charged particles depends on charge magnitude and distance.
Opposite charges have negative potential energy that decreases as they approach.
Shielding Effect:
Electrons in outer shells are partially shielded from nuclear charge by inner shell electrons.
Page 5: Penetration Effect
Penetration:
Describes how outer electrons experience greater nuclear attraction when they come closer to the nucleus.
Electrons that penetrate inner electron shells feel more nuclear charge.
Page 6: Energy Levels and Penetration Effects
Sublevel Energies:
Due to penetration, subshell energies are not degenerate. Example: 4s is lower than 3d.
As more shells are added, energy levels of subshells become closer and may overlap.
Page 7: Effective Nuclear Charge (Z_eff)
Definition: The net positive charge experienced by an electron in an atom.
Calculated as: Z_eff = Z - S, where Z is nuclear charge and S is shielding.
Core electrons are more efficient at shielding than valence electrons, affecting Z_eff.
Page 8: Atomic and Ionic Sizes
Size Trends:
Atomic size increases down a group due to higher n values.
Size decreases across a period due to increased nuclear charge pulling electrons closer.
Page 9: Trends in Ionic Size: Cations
Cations:
Formed by losing electrons, resulting in a smaller ionic radius than the neutral atom due to increased nuclear attraction.
Example: Na → Na⁺ (loss of 3s electron).
Page 10: Trends in Ionic Size: Anions
Anions:
Formed by gaining electrons, resulting in a larger ionic radius than the neutral atom due to increased electron repulsion.
Example: F → F⁻ (gaining an electron).
Page 11: Ionization Energy
Defined as the energy needed to remove electrons from gaseous atoms or ions.
First Ionization Energy (IE1): Energy required to remove the first electron.
Second Ionization Energy (IE2): Energy required to remove an electron from a cation.
Page 12: Trends in Ionization Energy
General Trend:
IE increases up a group and across a period; larger atoms with smaller Z_eff have lower IE.
Provided examples show relationships between elements and their ionization energy.
Page 13: Exceptions in Ionization Energy Trends
Notable decreases in ionization energy occur when removing paired electrons in p orbitals compared to unpaired electrons.
Page 14: Ionization Energies Reference
Discusses the pattern in ionization energies as more electrons are removed, highlighting significant jumps after valence electrons.
Page 15: Electron Affinity (EA)
Defined as the energy change associated with adding an electron to a neutral atom.
Negative EA denotes an exothermic process (energy release), while positive denotes endothermic (energy requirement).
Page 16: EA Trends and Unique Cases
EA generally becomes more negative across a period, with exceptions for noble gases.
Halogens have high negative EA due to their strong tendency to gain an electron to fill their valence band.
Page 17: Formation of Ionic Compounds
Discusses electron transfer from metals to nonmetals to form ions, subsequently creating ionic bonds through electrostatic attractions.
Page 18: Lattice Energy Overview
Lattice energy describes the energy released when gaseous ions form a solid ionic compound.
Influenced by ionic charge and distance between ions.
Page 19: Lattice Energy Rank and Comparison
Examines the comparison of lattice energies based on the ratio of ionic charges and sizes, emphasizing how ionic characteristics affect lattice formation.
Page 20: Future Concepts in Lattice Energy
Concept tests address predicting melting points and lattice energy relationships.
Page 21: Born-Haber Cycle
Describes using the Born-Haber cycle to calculate lattice energy through a series of enthalpy changes.
Lattice energy cannot be directly measured and is calculated using Hess's law.
Page 22: Example Calculation of Lattice Energy
Step-by-step example of calculating lattice energy of KCl using various enthalpy values, illustrating the necessary calculations.
Introduction to Periodicity and Ionic Bonding
Focus on Chapter 9: Periodicity and Ionic Bonding.
Valence and Core Electrons
Valence Electrons: Outermost electrons involved in bonding, show periodicity in configurations.
Core Electrons: Electrons not involved in bonding, present in lower energy levels. Examples:
O: 1s² 2s² 2p⁴
S: 1s² 2s² 2p⁶ 3s² 3p⁴
Se: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴
Te: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁴
Electrostatic Principles
Opposite charges attract; like charges repel. Attraction increases with charge magnitude, closer entities experience stronger forces.
Energy Concepts of Electrons
Coulomb’s Law: Potential energy depends on charge magnitude and distance. Opposite charges have negative potential energy decreasing as they approach.
Shielding Effect: Outer shell electrons are partially shielded from nuclear charge by inner shell electrons.
Penetration Effect
Describes how outer electrons feel greater nuclear attraction when closer to the nucleus.
Energy Levels and Penetration Effects
Sublevel Energies: Due to penetration, subshell energies are not degenerate (e.g., 4s < 3d).
Effective Nuclear Charge (Z_eff)
Net positive charge experienced by an electron, calculated as Z_eff = Z - S.
Atomic and Ionic Sizes
Atomic size increases down a group, decreases across a period due to nuclear charge.
Trends in Ionic Size
Cations: Formed by losing electrons, resulting in a smaller ionic radius.
Anions: Formed by gaining electrons, resulting in a larger ionic radius.
Ionization Energy
Energy required to remove electrons from gaseous atoms or ions.
Trends: IE increases up a group and across a period; larger atoms with smaller Z_eff have lower IE.
Electron Affinity (EA)
Energy change associated with adding an electron. Negative EA indicates an exothermic process.
Formation of Ionic Compounds
Involves electron transfer from metals to nonmetals, creating ions and ionic bonds through electrostatic attractions.
Lattice Energy
Energy released when gaseous ions form a solid ionic compound; influenced by ionic charge and distance. Calculated using Born-Haber Cycle for precision.