Periodic Trends and Properties: Electron Configurations, Shielding, and Periodic Properties

Periodic Trends and Properties

Introduction

This chapter explores how the arrangement of electrons in atoms leads to periodic trends in the properties of elements. Understanding these trends is essential for predicting chemical behavior and rationalizing the structure of the periodic table.

Dealing with More Than One Electron in the Atom

Penetration of Orbitals to the Nucleus – Effect on Energy

• Penetration refers to how close an electron in a given orbital can get to the nucleus.

• Electrons in orbitals with higher penetration (e.g., s orbitals) experience a greater attraction to the nucleus and are lower in energy compared to those in orbitals with less penetration (e.g., p, d, f).

• For a given principal quantum number (n), the order of penetration is: s > p > d > f.

Probability Distributions for the Wavefunctions: The probability of finding an electron at a certain distance from the nucleus varies by orbital type, influencing energy levels and chemical properties.

Multi-Electron Energy Level Diagrams

Hydrogen-Like vs. Non-Hydrogen-Like Atoms

• In hydrogen-like atoms (one electron), all orbitals with the same n are degenerate (same energy).

• In multi-electron atoms, electron-electron repulsions and shielding cause splitting of energy levels: E(s) < E(p) < E(d) < E(f) for a given n.

• Energy diagrams show that as you move to higher n and l, orbitals become closer in energy and more complex in their arrangement.

Review of Orbitals and Energy Placement

Types of Orbitals

Type of Orbital	Shape	Placement in Energy Diagram
s	Spherical	Lowest energy for each n
p	Dumbbell	Higher than s, same n
d	Cloverleaf	Higher than p, same n
f	Complex	Highest for given n

On energy diagrams, more negative values indicate greater stability.

Filling the Electrons into the NRG (Energy) Level Diagram

Rules for Electron Configuration

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers; each orbital holds a maximum of two electrons with opposite spins.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing up, maximizing unpaired electrons and minimizing electron repulsion.

• Paramagnetic: Atoms or ions with unpaired electrons; attracted to magnetic fields.

• Diamagnetic: Atoms or ions with all electrons paired; weakly repelled by magnetic fields.

Periodic Table Blocks and Energy Diagrams

Defining the Blocks

• s-block: Groups 1A and 2A (alkali and alkaline earth metals)

• p-block: Groups 3A to 8A (main group elements)

• d-block: Transition metals (groups 3B to 2B)

• f-block: Lanthanides and actinides

Energy level diagrams can be constructed using the periodic table as a guide, following the order of orbital filling.

Writing Long and Short Electron Configurations

Electron Configuration Notation

• Long form: Lists all occupied orbitals in order of filling (e.g., 1s2 2s2 2p6 ...).

• Short form (Noble Gas Core): Uses the previous noble gas in brackets to represent filled inner shells (e.g., [Ne] 3s2 3p4).

Some transition metals have exceptions due to extra stability of half-filled or fully-filled d subshells (e.g., Cr, Cu).

Electron Configurations of Ions

Formation of Cations and Anions

• When forming cations, electrons are removed first from the highest n value orbital (usually s before d for transition metals).

• Anions gain electrons to fill the next available orbital.

• Isoelectronic species: Atoms and ions with the same electron configuration.

Reviewing Coulomb's Law

Forces Between Charged Particles

• Coulomb's Law: $F = k \frac{q_1 q_2}{r^2}$

• The force of attraction or repulsion between two charged particles increases with greater charge and decreases with greater distance.

Defining Shielding and Effective Nuclear Charge (Zeff)

Shielding Effect

• Core electrons shield valence electrons from the full charge of the nucleus, reducing the effective nuclear charge felt by valence electrons.

• Effective nuclear charge (Zeff): $Z_{eff} = Z - S$, where Z is the atomic number and S is the number of core electrons.

• Zeff increases across a period and decreases down a group.

Periodic Trends

Trend 1: Effective Nuclear Charge (Zeff)

• Down a group: Zeff decreases due to increased shielding.

• Across a period: Zeff increases as protons are added but shielding remains relatively constant.

Trend 2: Atomic Radius

• Down a group: Atomic radius increases due to addition of electron shells.

• Across a period: Atomic radius decreases as Zeff increases, pulling electrons closer to the nucleus.

Trend 3: Ionization Energy

• Ionization energy (IE): The energy required to remove an electron from a gaseous atom or ion.

• Down a group: IE decreases as atomic radius increases and electrons are farther from the nucleus.

• Across a period: IE increases as Zeff increases and atomic radius decreases.

Successive ionization energies increase for each electron removed, with large jumps when removing core electrons.

Trend 4: Ionic Radii

• Anions are larger than their parent atoms due to increased electron-electron repulsion.

• Cations are smaller than their parent atoms due to loss of electrons and decreased repulsion.

Trend 5: Electron Affinity (EA)

• Electron affinity: The energy change when an electron is added to a gaseous atom.

• EA is generally more negative (more favorable) across a period, especially for halogens.

• Noble gases and group 2 elements have positive or less negative EA due to stable configurations.

Tables and Data

Successive Ionization Energies (kJ/mol) for Na to Ar

Element	1st IE	2nd IE	3rd IE	4th IE	5th IE	6th IE	7th IE	8th IE
Na	496	4560	6910	9540	13300	16600	20800	25400
Mg	738	1450	7730	10500	13600	18000	21700	25600
Al	578	1817	2745	11600	14800	18300	23300	27400
Si	786	1577	3231	4360	16000	19800	23700	29200
P	1012	1907	2910	4960	6270	21200	25400	29800
S	1000	2250	3350	4550	7000	8490	27100	31700
Cl	1251	2297	3822	5150	6540	9360	11000	34800
Ar	1521	2665	3931	5770	7238	8781	11995	13842

Additional info: Large jumps in ionization energy indicate removal of core electrons.

Electron Affinities (kJ/mol)

Group	1A	2A	3A	4A	5A	6A	7A	8A
H	–73
Li	–60
Na	–53
K	–48
Be		+240
Mg		+230
B			–27
C				–122
N					+7
O						–141
F							–328
Ne								+29

Additional info: Negative values indicate energy is released; positive values indicate energy is required.

Summary Table: Periodic Trends

Trend	Down a Group	Across a Period (Left to Right)
Atomic Radius	Increases	Decreases
Ionization Energy	Decreases	Increases
Electron Affinity	Generally less negative	Generally more negative
Zeff	Decreases	Increases

Key Terms and Concepts

• Valence Electrons: Electrons in the outermost shell, involved in chemical bonding.

• Core Electrons: Inner electrons, not involved in bonding.

• Isoelectronic: Species with the same electron configuration.

• Shielding: Reduction in effective nuclear charge on valence electrons due to presence of core electrons.

Examples and Applications

• Example 1: The electron configuration of Fe2+ is [Ar] 3d6.

• Example 2: Cl– and Ar are isoelectronic, both with the configuration [Ne] 3s2 3p6.

• Example 3: The atomic radius of Na is larger than that of Mg due to lower Zeff.

Additional info: These periodic trends are foundational for understanding chemical reactivity, bonding, and the organization of the periodic table.