Periodic Trends and Atomic Properties

Event Details

• Event Name: Purdue Pulse Fall Tailgate

• Date and Time:

  • Saturday, October 4, 2025

  • Start: 10:00 AM EDT

  • End: 11:30 AM EDT

• Location: Stadium Mall

• Description:

  • Join us on Stadium Mall before the Purdue vs. Illinois Football game for food, games, gameday gear, and fun!

• Perks:

  • Free Food

  • Free Stuff

• Host Organization:

  • NPSUB - Purdue Student Union Board

Lecture Overview

• Lecture 10: Periodic Trends / Trends in Chemical Reactivity

• Date: September 30, 2025

• Sections Covered:

  • 8.3 Trends in Three Atomic Properties

  • 8.4 Atomic Properties and Chemical Reactivity

Announcements

• Office Hours:

  • Wednesdays, 11:00 AM – 12:00 PM, WTHR 261 or by appointment (email: mille201@purdue.edu with 3 suggested days and times).

• Feasting with Faculty:

  • Friday, October 3, about 12:50 – 1:50 PM, near the desserts.

• Exam 1 Results: Expected to take a long time.

• Final Exam: Scheduled for Thursday, December 18, 2025, from 10:30 AM – 12:30 PM, Elliott Hall of Music for WL students, location TBA for Indianapolis students.

Learning Objectives (Chapter 8 Concepts)

• Define the meanings of:

  • Atomic Radius (8.3)

  • Ionization Energy (8.3)

  • Electron Affinity (8.3)

• Explain how the n value and effective nuclear charge (Z_eff) relate to periodic trends in:

  • Atomic size

  • Ionization energy

• Understand core electrons' importance in the pattern of successive ionization energies (8.3)

• Explain the relationship between atomic properties and the tendency to form ions. (8.3)

• List general properties of metals and nonmetals (8.4)

• Describe why main-group ions are either isoelectronic with the nearest noble gas or have a pseudo-noble gas electron configuration (8.4)

• Describe the relation between ionic and atomic size and the trends in ionic size (8.4)

Skills

• Use periodic trends to

  • Rank elements by atomic size and first ionization energy (SPs 8.2, 8.3)

  • Identify an element from its successive ionization energies (SP 8.4)

  • Write electron configurations of main-group and transition metal ions (SPs 8.5, 8.6)

  • Rank ions by size using periodic trends (SP 8.7)

Periodic Trends

• Definition: Properties that exhibit consistent changes within a group or period.

  • Main-group Elements: Includes

  • Alkali Metals (Group 1)

  • Alkaline Earth Metals (Group 2)

  • Halogens

  • Noble Gases (Group 18)

Trends in Properties

1. Atomic Radius

2. Ionization Energy

3. Electron Affinity

For each property, discussion includes:

• Definition

• Trend

• Reason for the trend (n, Z_eff)

Influences on Periodic Trends

• Principal Quantum Number (n):

  • Identifies the principal energy level or shell.

  • Positive integers starting with 1.

  • Associated with the Period.

• Effective Nuclear Charge (Z_eff):

  • Each electron is attracted to the positively charged nucleus and repelled by other negatively charged electrons.

  • Z_eff takes shielding effects into account to represent the net electric field experienced by an electron.

  • Significance of attraction (protons) and repulsion (electrons).

Atomic Radii

• Challenge: Identifying atomic radius for isolated atoms is complex.

Defining Atomic Size

1. Metallic Radius:

   • Defined as half the shortest distance between nuclei of adjacent, individual atoms in a crystal of a metal.

2. Covalent Radius:

   • Defined as half the shortest distance between nuclei of identical, covalently bonded atoms. Applicable to nonmetals.

3. Variations: Atomic size varies slightly from substance to substance.

Effective Nuclear Charge

• Trend: Increases across a period. Reason:

  • Atomic number (Z) increases by one for each element, while core electrons remain constant.

  • Core electrons significantly influence shielding of valence electrons.

Examples of Effective Nuclear Charge

• Sodium (Na):

  • Z = 11, 10 core electrons, 1 valence electron.

  • Z_eff = +2.5

• Magnesium (Mg):

  • Z = 12, 10 core electrons, 2 valence electrons.

  • Z_eff = +3.15

  • Comparison: A valence electron in Mg experiences more attraction from the nucleus than in Na.

Group-Based Trends

• Down a Group:

  • n value increases, while Z_eff is approximately constant.

  • Outer electrons are farther from the nucleus, leading to an increase in radius due to shielding from additional inner electrons.

• Across a Period:

  • n constant while Z_eff increases.

  • Outer electrons are pulled closer to the nucleus, resulting in a decrease in radius.

Ionization Energy (IE)

• Definition: The energy required for the complete removal of an electron from a gaseous atom.

  • Represented as: X(g) \rightarrow X^+(g) + e^-

• Trends:

  • First Ionization Energy (IE₁):

  • There are as many IEs as there are electrons. Subsequent IEs are higher due to the removal of negative charge from a more positive parent atom.

  • Large increase occurs in IE after all valence electrons are removed (core electron IEs are much higher).

Summary of Trends in Ionization Energy

• Down a Group:

  • n increases, Z_eff is approximately constant → IE decreases.

• Across a Period:

  • n is constant while Z_eff increases → IE increases.

Electron Affinity (EA)

• Definition: The energy change accompanying the addition of one mole of electrons to one mole of gaseous atoms or ions.

  • Represented as:

  • X(g) + e^- \rightarrow X^-(g)

  • Often, \Delta E{rxn} = EA1 < 0

  • Usually exothermic when first electron is added.

• Trends: Not as regular as other properties.

Summary of Trends in Electron Affinity

• Down a Group:

  • n increases, while Z_eff is approximately constant → EA decreases.

• Across a Period:

  • n constant while Z_eff increases → EA increases.

Ion Formation

• Key Point: Atoms with low IE tend to form cations while those with high IE tend to form anions, with exceptions for noble gases.

Isoelectronic Species

• Definition: Species having the same number and configuration of electrons as another species.

  • Example: Na extsuperscript{+}, F extsuperscript{-}, and Ne are isoelectronic:

  • Na: 1s²2s²2p⁶

  • Na extsuperscript{+}: 1s²2s²2p⁶

  • F: 1s²2s²2p⁵

  • F extsuperscript{-}: 1s²2s²2p⁶

  • Ne: 1s²2s²2p⁶.

Transition Metal Ions

• Transition metals usually do not reach noble gas electron configuration; they typically lose ns electrons along with (n-1)d electrons to form cations.

• Example: Cobalt, Co forms Co²⁺ and Co³⁺ ions.

Ionic Radii and Comparison

• Cations: Smaller than their neutral atoms.

• Anions: Larger than their neutral atoms.

• As positive charge increases in isoelectronic series, radius decreases:

  • Order: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺ within similar configurations.

Summary of Properties

• Metals:

  • Shiny, malleable, ductile, good conductors of heat and electricity, mostly solids at room temperature.

• Nonmetals:

  • Dull, brittle, nonconductors, presence of gases, and one liquid (Br₂).

• Metalloids:

  • Characteristics of both metals and nonmetals (e.g., shiny but brittle, semiconductors).

Periodic Table Overview

• Organized based on main-group elements, transition elements, and properties navigating from metals to nonmetals.

Additional Practice

• Problems from the text for deeper understanding and application of periodic trends in atomic properties.