Periodic Trends and Element Properties: Study Guide

Unit 4: Element and Compound Properties

Overview

This unit explores the predictable patterns of properties that emerge when elements are arranged according to the periodic table. It covers atomic structure, periodic trends, chemical bonding, and the relationship between atomic properties and chemical reactivity.

Periodic Table and Atomic Structure

The Periodic Table

• Periodic Table: A systematic arrangement of elements based on increasing atomic number and recurring chemical properties.

• Groups/Families: Vertical columns; elements in the same group have similar valence electron configurations and chemical properties.

• Periods: Horizontal rows; elements show gradual changes in properties across a period.

• Example: Alkali metals (Group 1) are highly reactive and have one valence electron.

Periodic Trends

Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electron shell. It varies predictably across periods and groups.

• Trend Across a Period: Atomic radius decreases from left to right due to increasing effective nuclear charge ($Z_{eff}$).

• Trend Down a Group: Atomic radius increases due to the addition of electron shells.

• Effective nuclear charge:

• Equation: $Z_{eff} = Z - S$ (where $Z$ is atomic number, $S$ is shielding constant/number of core electrons)

• Example: Lithium has a larger atomic radius than fluorine in the same period.

Shielding Effect

The shielding effect describes how inner electrons block the attraction between the nucleus and outer electrons, affecting atomic size and ionization energy.

• Shielding increases down a group as more electron shells are added.

• Shielding remains relatively constant across a period.

• Example: Sodium has greater shielding than lithium.

Ionic Radius

Ionic radius refers to the size of an ion after gaining or losing electrons.

• Cations (positive ions) are smaller than their parent atoms due to loss of electrons.

• Anions (negative ions) are larger than their parent atoms due to gain of electrons.

• Example: $Na^+$ is smaller than $Na$, $Cl^-$ is larger than $Cl$.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom.

• Trend Across a Period: Increases due to higher $Z_{eff}$.

• Trend Down a Group: Decreases due to increased distance from nucleus and greater shielding.

• Equation: $X(g) + energy \rightarrow X^+(g) + e^-$

• Example: Fluorine has a higher ionization energy than sodium.

• Exceptions: Grps 2A and 3A, Grps 5A and 6A

Electronegativity

Electronegativity is the tendency of an atom to attract electrons in a chemical bond.

• Trend Across a Period: Increases from left to right.

• Trend Down a Group: Decreases down a group.

• Example: Fluorine is the most electronegative element.

Reactivity

Reactivity depends on whether an element is a metal or nonmetal and its position in the periodic table.

• Metals: Reactivity increases down a group (e.g., alkali metals) and decreases across a period.

• Nonmetals: Reactivity increases up a group (e.g., halogens) and increases across a period.

• Example: Cesium is more reactive than sodium; fluorine is more reactive than iodine.

Melting and Boiling Points

Melting and boiling points depend on the strength of forces between atoms or molecules.

• Generally, melting and boiling points increase across a period for metals and decrease for nonmetals.

• Example: Sodium has a lower melting point than magnesium.

Density

Density tends to increase down a group due to increased atomic mass and volume.

• Example: Potassium is less dense than cesium.

Summary Tables

Periodic Trend Summary Table

Trend	Across a Period	Down a Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases
Shielding Effect	Constant	Increases
Reactivity (Metals)	Decreases	Increases
Reactivity (Nonmetals)	Increases	Decreases

Key Concepts and Applications

Effective Nuclear Charge ($Z_{eff}$)

• Defines the net positive charge experienced by valence electrons.

• Higher $Z_{eff}$ leads to smaller atomic radius and higher ionization energy.

Electron Affinity

• Energy change when an atom gains an electron.

• Generally becomes more negative across a period.

Applications

• Periodic trends help predict element reactivity, compound formation, and physical properties.

• Understanding trends is essential for explaining chemical behavior and designing new materials.

Practice and Writing Prompts

• Explain why atomic radius increases down a group and decreases across a period.

• Describe the difference in ionization energy between Group 1 and Group 17 elements.

• Compare the reactivity of alkali metals and halogens using periodic trends.

Final Thoughts

• Always justify trends using concepts such as energy levels, shielding, and effective nuclear charge.

• Use the periodic table as a predictive tool for chemical properties and reactivity.