Periodic Trends and Electron Configurations
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Lecture 9 Overview
Title: Periodic Trends / Trends in Chemical Reactivity
Date: September 25, 2025
Sections Covered:
Section 8.2: The Quantum Mechanical Model and the Periodic Table
Section 8.3: Trends in Three Atomic Properties
Announcements
Office Hours:
Wednesdays, 11:00 AM – 12:00 PM in WTHR 261 or by appointment.
Email: mille201@purdue.edu with 3 suggested days and times.
Feasting with Faculty:
TBD for next week
Exam Dates:
Exam 1: Thursday, September 25, 8:00 PM. Arrive by 7:40 PM.
Final Exam: Thursday, December 18, 10:30 AM – 12:30 PM, Elliott Hall of Music for WL students; location TBA for Indianapolis students.
Learning Objectives for Lecture 9
Chapter 8 Concepts
Periodic Table Arrangement
Explain how the arrangement is based on the order of sublevel energies (Section 8.2).
Sublevel Filling
Describe how sublevels are filled in main-group and transition elements.
Hund’s Rule
Describe the importance of Hund’s Rule (Section 8.2).
Electron Types
Understand the distinction among inner, outer, and valence electrons (Section 8.2).
Atomic Radius Definition
Define the meaning of atomic radius (Section 8.3).
Periodic Trends
Explain how the n value and effective nuclear charge contribute to the trend of atomic size (Section 8.3).
Skills
Write full and condensed electron configurations for elements (Skill Practice 8.1).
Use periodic trends to rank elements by atomic size (Skill Practice 8.2).
Significance of Electron Configuration
The recurring patterns in electron configuration provide a basis for understanding the behavior and properties of elements in the Periodic Table.
Concepts such as atomic radius, ionization energy, and electron affinity arise from these patterns.
It is important to note that there may be anomalies in the transition metals and inner transition metals. However, utilizing expected patterns still allows predicting configurations to explain chemical behaviors.
Electron Configurations
Composition of Electron Configurations:
Consist of:
Principal energy level (n)
Letter designation of the sublevel (s/p/d/f)
Number of electrons (#) in the sublevel, written as a superscript.
Important Rules for Writing Electron Configurations
Aufbau Principle
Electrons are added to the atom in the order of increasing energy of the orbitals, starting with the lowest energy orbital.
Pauli Exclusion Principle
Any orbital can hold a maximum of two electrons, but only if they have different spins.
Hund’s Rule
For a set of p, d, or f orbitals, each orbital is singly occupied before any gets a second electron, ensuring maximum unpaired electrons with parallel spins.
Orbital Filling and Periodic Table Relation
The periodic table arrangement reflects the order in which orbitals fill based on sublevel energies.
Examples of Electron Configurations
Hydrogen (H), $Z = 1$:
Configuration: $1s^1$
Helium (He), $Z = 2$:
Configuration: $1s^2$
Lithium (Li), $Z = 3$:
Configuration: $1s^2 2s^1$
Boron (B), $Z = 5$:
Configuration: $1s^2 2s^2 2p^1$
Nitrogen (N), $Z = 7$:
Configuration: $1s^2 2s^2 2p^3$
Oxygen (O), $Z = 8$:
Configuration: $1s^2 2s^2 2p^4$
Hund’s Rule in Practice
Hund’s Rule:
Specifies when orbitals of equal energy are available, the configuration with the maximum number of unpaired electrons with parallel spins is the lowest energy state.
Nitrogen (N) example ($Z = 7$) shows that it follows this rule: $1s^2 2s^2 2p^3$ has 3 unpaired electrons in 2p.
Orbital Diagrams
Construction of Orbital Diagrams:
Each box (or line) represents an orbital in a given energy level grouped by sublevel, with arrows indicating electrons and their spins.
Example for Lithium: $1s^2 2s^1$
Electron Configuration Examples for Period 3 Elements
Partial Orbital Diagrams and Configurations:
Atomic Number
Element
Full Electron Configuration
Condensed Electron Configuration
11
Na
$[1s^2 2s^2 2p^6] 3s^1$
$[Ne] 3s^1$
12
Mg
$[1s^2 2s^2 2p^6] 3s^2$
$[Ne] 3s^2$
13
Al
$[1s^2 2s^2 2p^6] 3s^2 3p^1$
$[Ne] 3s^2 3p^1$
14
Si
$[1s^2 2s^2 2p^6] 3s^2 3p^2$
$[Ne] 3s^2 3p^2$
18
Ar
$[1s^2 2s^2 2p^6] 3s^2 3p^6$
$[Ne] 3s^2 3p^6$
Practice Questions
Determine the atom represented by a given electronic energy-level diagram.
Options: Beryllium, Carbon, Oxygen, Sulfur
What is the electron configuration of Potassium (K)?
Options:
a. $1s^2 2s^2 2p^6 3s^2 3p^6 3d^1$
b. $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1$
c. $1s^2 2s^2 2p^3 3s^2 3p^3 3d^7$
d. $1s^2 2s^2 2p^3 3s^2 3p^3 4s^2 3d^5$
Provide the orbital box diagram for atomic number 19.
Identify the element corresponding to ground-state electron configuration $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2$.
Expected element: Calcium (Ca)
Condensed Configurations
Definition: A condensed electron configuration features the element symbol of the previous noble gas in square brackets.
Example for Aluminum (Al):
Full: $1s^2 2s^2 2p^6 3s^2 3p^1$
Condensed: $[Ne] 3s^2 3p^1$
Stability of Electron Configurations
Half-Filled and Filled Sublevels:
Stability is higher for filled and half-filled outer energy levels due to overlapping sublevel energies allowing easier electron movement.
Examples:
Chromium (Cr):
Expected configuration: $[Ar] 4s^2 3d^4$
Actual: $[Ar] 4s^1 3d^5$ (half-filled d subshell)
Copper (Cu):
Expected: $[Ar] 4s^2 3d^9$
Actual: $[Ar] 4s^1 3d^{10}$ (filled d subshell)
Types of Electrons
Core (inner) Electrons:
Electrons shared with the previous noble gas.
Outer Electrons:
Highest energy level electrons (highest n).
Valence Electrons:
Electrons involved in bond formation.
Main group elements: Valence = outer electrons.
Transition metals: Valence = outer + number of d electrons.
Determining Electrons and Valences
Identify core, outer, and valence electrons
Example for Phosphorus (P):
$1s^2 2s^2 2p^6 3s^2 3p^3$
5 valence electrons.
Example for Manganese (Mn):
$[Ar] 4s^2 3d^5$
7 valence electrons.
Periodic Table Insights:
Elements in a group possess the same valence electrons, leading to similar chemical behaviors.
Group 1 = 1 valence electron, Group 2 = 2, Group 8 = 8, except for Helium which has 2.
Conclusion on Electron Configuration and Reactivity
Elements in the same group exhibit similar outer electron configurations, resulting in comparable chemical behaviors.
Further Practice and Problems
Follow-Up Problems in Chapter: 8.1A, 8.1B and end-of-chapter problems numbered: 8.21, 8.23, 8.25, 8.27, 8.29, 8.31, 8.33, 8.35, 8.37, 8.39.
Atomic Radii Discussion
Atomic radii are difficult to measure for isolated atoms.
Metallic and Covalent Radii:
Determined by measurements between atoms.
Example Measurements:
Metallic radius of Aluminum (Al) = 286 pm
Covalent radius of Chlorine (Cl) = 99 pm
Effective Nuclear Charge (Zeff)
Definition: The effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom.
The pull an electron feels from the nucleus increases as its proximity to the nucleus decreases.
**Trends Across Periods:
Zeff increases as one moves across a period due to an increase in the atomic number while the core electrons remain constant. Core electrons shield the valence electrons more effectively than other valence electrons do among themselves.
Atomic Radius Trends
As Zeff increases, atomic radius typically decreases due to a stronger attractive force pulling electrons closer to the nucleus.
Specific patterns in atomic radius can be observed throughout the main-group elements.
Example Trend Chart of Atomic Radii:
Radii Increase through Periods, especially noticeable across groups of the periodic table.
Practice Problem - Neglecting Zeff
Sample Problem:
Using the periodic table, rank elements in order of decreasing atomic size, for example, with Ca, Mg, Sr, K, Ga, Cl.
For Future Reference
Atomic Radius Problems:
Review in Chapter for follow-up problems: 8.2A, 8.2B, and end-of-chapter problems numbered: 8.50.