3.5 Periodic Variations in Element Properties | Chemistry: Atoms First 2e

Learning Objectives

Describe and explain the observed trends in atomic size, ionization energy, and electron affinity of the elements

Core Concepts & Principles

The periodic table is arranged so that elements in the same vertical column (group) share similar chemical behaviors because they possess identical valence electron configurations. However, moving across periods or down groups reveals systematic variations in physical and chemical properties, driven by changes in atomic structure.

Atomic and Ionic Radii

Atomic Radius (Covalent Radius): Defined as one-half the distance between the nuclei of two identical bonded atoms.
- Down a group: Atomic radius increases. Each row adds a principal quantum number ( $n$ ), placing outer electrons in larger orbitals farther from the nucleus.
- Across a period: Atomic radius decreases. Although electrons are added, the effective nuclear charge ( $Z_{\text{eff}}$ ) increases, pulling valence electrons tighter toward the nucleus.
Ionic Radius:
- Cations are always smaller than their parent neutral atoms because losing electrons decreases electron-electron repulsion while increasing $Z_{\text{eff}}$ per remaining electron.
- Anions are always larger than their parent neutral atoms because adding electrons increases electron-repulsion and expands the valence shell.
- Isoelectronic species (atoms/ions with identical electron configurations) decrease in size as nuclear charge ( $Z$ , proton count) increases.

Effective Nuclear Charge ( $Z_{\text{eff}}$ )

Core inner-shell electrons shield outer valence electrons from the full attractive pull of the nucleus. The net pull experienced by an electron is given by the effective nuclear charge: $Z_{\text{eff}} = Z - \text{shielding}$ Because core electrons shield outer electrons inefficiently, $Z_{\text{eff}}$ increases steadily from left to right across a period, drawing the electron cloud inward.

Ionization Energy (IE)

Ionization energy is the endothermic energy required to remove the most loosely bound electron from a gaseous atom or ion: $\text{X}(g) \rightarrow \text{X}^+(g) + e^- \quad (\text{IE}_1 > 0)$

Down a group: $\text{IE}_1$ decreases because larger atoms hold their outer electrons farther away and more loosely.
Across a period: $\text{IE}_1$ increases due to increasing $Z_{\text{eff}}$ .
Exceptions: Minor dips occur when entering a new subshell (e.g., B has a lower IE than Be because a $p$ electron is easier to remove than an $s$ electron) or when half-filled subshell stability is broken (e.g., O has a lower IE than N due to paired-electron repulsion).
Successive Ionization Energies: Removing electrons becomes progressively harder ( $IE_1 < IE_2 < IE_3 \dots$ ). Massive energy jumps occur once core electrons are breached.

Electron Affinity (EA)

Electron affinity is the energy change when a gaseous atom accepts an electron to form an anion: $\text{X}(g) + e^- \rightarrow \text{X}^-(g)$

Across a period: EAs tend to become more negative (more exothermic/favorable) as $Z_{\text{eff}}$ increases.
Group Anomalies: Noble gases, Group 2, and Group 15 show disruptions due to stable filled/half-filled subshells. Additionally, fluorine has a less negative EA than chlorine because adding an electron to fluorine's tiny $n=2$ shell creates intense electron-electron repulsion, whereas chlorine's larger $n=3$ shell accommodates the extra electron more comfortably.

Core Periodic Trends Law

Size (Radius): Increases $\downarrow$ , decreases $\rightarrow$ .
Ionization Energy & $Z_{\text{eff}}$ : Decreases $\downarrow$ , increases $\rightarrow$ .
Electron Affinity: Generally becomes more negative $\rightarrow$ .

Problem-Solving Routines & Methods

Comparing Atomic and Ionic Radii

1
Locate the elements or ions on the periodic table.
2
Check if they are isoelectronic. If yes, the species with the highest proton count ( $Z$ ) has the smallest radius.
3
If not isoelectronic, check group placement: radii increase down a group due to a higher principal quantum number ( $n$ ).
4
Check period placement: radii decrease from left to right due to a higher effective nuclear charge ( $Z_{\text{eff}}$ ).

Pro-Tip: Always remember that cations are smaller and anions are larger than their neutral parent atoms.

Effective Nuclear Charge

Z_{\text{eff}} = Z - \text{shielding}

Calculates the net positive pull experienced by valence electrons.

Variables & Constants

Z

=atomic number (protons);

\text{shielding}

=screening effect provided by core inner-shell electrons

Practice & Concept Checks

Concept Check

Why is the first ionization energy of oxygen slightly lower than that of nitrogen, even though oxygen has a higher atomic number?

Concept Check

Arrange the following isoelectronic species in order of increasing radius: S²⁻, Cl⁻, Ar, K⁺, Ca²⁺.

Key Terms & Vocabulary

covalent radiusAtomic Structure

One-half the distance between the nuclei of two identical atoms joined by a covalent bond.

Example: Used to gauge relative atomic sizes.

effective nuclear charge, ZeffPeriodic Trends

The net positive charge experienced by an outer electron after accounting for core electron shielding.

Example: Zeff = Z - shielding

isoelectronicAtomic Structure

Describing atoms or ions that possess the exact same electron configuration.

Example: N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, and Al³⁺ all share 1s²2s²2p⁶.

ionization energyPeriodic Trends

The energy required to remove the most loosely bound electron from a gaseous atom or ion in its ground state.

Example: X(g) → X⁺(g) + e⁻

electron affinityPeriodic Trends

The energy change associated with the addition of an electron to a gaseous atom to form an anion.

Example: X(g) + e⁻ → X⁻(g)

Periodic Variations in Element Properties