Module 7 Day 2

Page 1: Announcements

• Exam scores will be posted to Blackboard on Tuesday.

• The exam key is posted on Blackboard.

• Akindi bubble sheets have been mailed to you.

Page 2: Module 7 Overview

Module Title

• Periodic Properties Day 1 2/19

Topics Covered

• Introduction to Periodic Properties

• Periodic Variations in Element Properties

• Variation in Covalent Radius

• Variation in Ionic Radii

Upcoming Topics

• Day 2 2/24: Variation in Ionization Energies

• Variation in Electron Affinities

Page 3: Ionization Energy

• Definition: The energy required to remove the most loosely bound electron from a gaseous atom in its ground state.

• Equation: X(g) ⟶ X+(g) + e−

• First Ionization Energy (IE1) values are always positive, indicating energy is always required to remove electrons from atoms or ions.

Page 4: Trends in Ionization Energies

General Trends

• Down a Group: First ionization energy decreases, easier to remove an electron due to increasing atomic size and distance from the nucleus.

• Across a Period: First ionization energy generally increases from left to right as effective nuclear charge increases, making it harder to remove an electron.

Page 5: Valence Orbital Diagrams

• Example Question: Determine which valence orbital diagram has the greatest first ionization energy among Ca, Ge, Ga, Br.

• Key Concept: The closer an atom is to having a full valence shell, the harder it is to remove an electron.

Page 6: First Ionization Energies of Elements

• Graph plotting first ionization energy for the first 20 elements against atomic number.

• Notable elements:

  • Group 2 elements (Be, Mg, Ca) show slightly larger IE values than expected.

  • Group 16 elements (O, S) also show slightly larger IE values than expected.

Page 7: Deviations in Ionization Energy Trend

Key Deviations

• Beryllium (Be) vs Boron (B):

  • Be has a higher IE than B, despite boron's expected trend.

  • Configurations: Be [He]2s², B [He]2s²2p¹.

• Explanation: Removing an s electron from Be is harder than removing a p electron from B, reflecting expected patterns in ionization energies.

Page 8: More Deviations in Trends

Electron Pairing Effects

• Removing one electron from O decreases electron-electron repulsion in the 2p orbital, making it more favorable compared to N.

• Configurations: N [He]2s²2p³, O [He]2s²2p⁴.

• Results in a lower IE than predicted by the trend.

Page 9: Successive Ionization Energies

• Definition: Energy required to remove the most loosely bound electron from a gaseous atom (IE1), followed by removal of successive electrons (IE2, IE3, etc.).

• General Trend: Ionization energies increase successively.

• Example Equations:

  • IE1: X(g) ⟶ X+(g) + e−

  • IE2: X+(g) ⟶ X2+(g) + e−

Page 10: Ionization Energies and Charge

• Trend: IE1 < IE2 < IE3. Removal of electrons from cations requires more energy due to increased electrostatic attraction.

• Note: Removal of electrons becomes more challenging as positive charge increases in cations.

Page 11: Tables of Successive Ionization Energies

Example Elements and Their Ionization Energies

• Various elements with their successive IEs outlined in kJ/mol.

• Note that large increases in energy correspond to the removal of core electrons indicating they are harder to remove than valence electrons.

Page 12: Core Electrons and Large Increases

• For Ga, the transition between IE3 → IE4 shows a significant increase indicative of core electron removal.

Page 13: Prediction of Highest Ionization Energy

• Question: Which of the following will have the highest ionization energy? F, Sc2+, Ca2+, Ge3+

Page 14: Ionization Energy Configurations

• Breakdown of configurations and predictions for ionization energies based on electron configurations for F, Sc2+, Ca2+, Ge3+.

Page 15: Electron Affinity

• Definition: Energy change during the addition of an electron to a gaseous atom.

• Reaction Equation: X(g) + e− ⟶ X−(g)

• Characteristics: Can be exothermic (releasing energy) or endothermic (absorbing energy).

Page 16: Trends in Electron Affinity

• Easier to add electrons as effective nuclear charge increases.

• Electron affinities become more negative across a period, generally favoring gaining an electron.

• Exceptions: Noble gases exhibit low electron affinity due to complete valence shells.

Page 17: Deviation Trends in Electron Affinity

• Group 2 and Group 15 see additional challenges in adding electrons due to electronic configuration and pairing issues.

Page 18: Prediction of Greatest Electron Affinity

• Question: Which element would likely have the greatest electron affinity: Li, Ca, S, Br?

Page 19: Comparison of Ionization Energies

• Question: Which element is likely to have the largest first ionization energy among Na, H, Li, and Be?

Page 20: Electron Configuration Changes

• Configuration: [Og] 8s²7d¹⁰6f¹⁴5g¹⁸p⁵ transforming to [Og] 8s²7d¹⁰6f¹⁴5g¹⁸p⁶ (gain of an electron).

Page 21: Order of Increasing Ionization Energy

• Determine the order of increasing ionization energy among Mg, I, Cl based on periodic trends.

Page 22: Electron Configuration Excitation

• Configuration change from [Og] 8s²7d¹⁰6f¹⁴5g¹⁸p⁵ to [Og] 8s²7d¹⁰6f¹⁴5g¹⁸p⁴9s¹ indicates excitation of an electron to a higher energy level.

Page 23: Cation Formation

• Configuration change from [Og] 8s²7d¹⁰6f¹⁴5g¹⁸p⁵ to [Og] 8s²7d¹⁰6f¹⁴5g¹⁸p⁴ signifies loss of one electron, forming a cation.

Page 24: Quantum Numbers Prediction

• Determine possible quantum numbers for the electron removed in the first ionization of Na, based on its electron configuration.