Atomic Structure, Electron Configuration, and Periodic Trends

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Atomic Structure and the Periodic Table

Scientific Contributions to the Periodic Table

The development of the periodic table was a major milestone in chemistry, allowing scientists to organize elements based on their properties and atomic structure.

Dmitri Mendeleev (1834–1907) is credited with arranging elements in order of increasing atomic weight and grouping elements with similar properties in columns.
Mendeleev left empty spaces for elements not yet discovered, predicting their properties based on their position in the table.
Other contributors include Julius Lothar Meyer, Alexandre Béguyer de Chancourtois, and John Newlands.

Periodic Law: When elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.

Sometimes, elements are reordered by other properties (e.g., tellurium and iodine) to better fit observed chemical behavior.
The periodic law allows prediction of element properties based on table position, but quantum mechanics explains the underlying reasons for these patterns.

Quantum Mechanics and Electron Configuration

Quantum Mechanical Model of the Atom

Quantum mechanics describes the behavior of electrons in atoms, which exist in specific orbitals defined by quantum numbers.

Electron Configuration: Describes the arrangement of electrons in an atom's orbitals.
Electrons in multi-electron atoms experience electron-electron interactions, making their energy levels more complex than hydrogen-like atoms.
Orbitals are labeled as s, p, d, and f.

Subshell Energies and Penetration

In multi-electron atoms, subshell energies depend on both the principal quantum number (n) and the angular momentum quantum number (l).

Electrons in orbitals with greater penetration (closer to the nucleus) feel a stronger effective nuclear charge.
Order of subshell energies:
Within a shell, subshell energies increase with increasing l.
When two subshells have the same value, the subshell with lower n is filled first.

Order of Subshell Energies for Multielectron Atoms

Subshell	n	l	n + l
1s	1	0	1
2s	2	0	2
2p	2	1	3
3s	3	0	3
3p	3	1	4
3d	3	2	5

Rules for Electron Configuration

Pauli Exclusion Principle

No two electrons in an atom can have the same four quantum numbers. Each orbital can hold a maximum of two electrons with opposite spins.

Aufbau Principle

Electrons fill the lowest-energy orbitals first before occupying higher-energy orbitals.

Hund's Rule

Electrons occupy orbitals singly as far as possible before pairing up.
Electrons in half-filled orbitals have the same spin.

Orbital Box Diagrams

Orbital box diagrams visually represent electron configurations, showing electrons as arrows in boxes for each orbital.

Example Table: Electron Configurations for First 10 Elements

Element	Number of Electrons	Simplified Orbital Diagram	Electron Configuration
Li	3	1s: ↑↓ 2s: ↑	1s22s1
Be	4	1s: ↑↓ 2s: ↑↓	1s22s2
B	5	1s: ↑↓ 2s: ↑↓ 2p: ↑	1s22s22p1
C	6	1s: ↑↓ 2s: ↑↓ 2p: ↑ ↑	1s22s22p2
N	7	1s: ↑↓ 2s: ↑↓ 2p: ↑ ↑ ↑	1s22s22p3
O	8	1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑ ↑	1s22s22p4
F	9	1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑↓ ↑	1s22s22p5
Ne	10	1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑↓ ↑↓	1s22s22p6
Na	11	1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑↓ ↑↓ 3s: ↑	1s22s22p63s1

Electron Configurations: Notation

spdf notation:
Noble gas notation:
Orbital box notation: [Ne] ↑↓ ↑ ↑ ↑ (for 3s and 3p)

Transition Metals and Exceptions

Electron Configurations for Transition Metals

For transition metals, the ns subshell is filled before the (n-1)d subshell. However, some elements (e.g., chromium and copper) have exceptions to this pattern.

Example: filled before
Chromium:
Copper:

Valence Electron Configurations for First Transition Series

Element	No. of electrons	Core	Simplified Orbital Diagram	Electron Configuration
Sc	21	[Ar]	4s: ↑↓ 3d: ↑	[Ar] 3d14s2
Ti	22	[Ar]	4s: ↑↓ 3d: ↑ ↑	[Ar] 3d24s2
Cr	24	[Ar]	4s: ↑ 3d: ↑ ↑ ↑ ↑ ↑	[Ar] 3d54s1
Ni	28	[Ar]	4s: ↑↓ 3d: ↑↓ ↑↓ ↑↓ ↑↓	[Ar] 3d84s2
Cu	29	[Ar]	4s: ↑ 3d: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓	[Ar] 3d104s1
Zn	30	[Ar]	4s: ↑↓ 3d: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓	[Ar] 3d104s2

Periodic Trends in Atomic Properties

Atomic Radius

Atomic radius decreases across a period (left to right) due to increasing effective nuclear charge.
Atomic radius increases down a group due to addition of electron shells.

Ion Size (Ionic Radius)

Cations (positive ions) are smaller than their parent atoms due to loss of electrons and increased nuclear attraction.
Anions (negative ions) are larger than their parent atoms due to electron-electron repulsion.
Isoelectronic ions (ions with the same number of electrons) can be compared by their nuclear charge; higher nuclear charge results in a smaller radius.

Ionization Energy

First ionization energy is the energy required to remove the outermost electron from a neutral atom in the gas phase.
Ionization energy increases across a period and decreases down a group.
Successive ionization energies increase as more electrons are removed.

Electron Affinity

Electron affinity is the energy change when an atom gains an electron to form an anion.
Generally becomes more negative across a period (atoms more likely to accept electrons).

Properties of Metals and Nonmetals

Metals

Malleable, ductile, shiny, and good conductors of heat and electricity.
Tend to lose electrons to form cations.
Form basic oxides.

Nonmetals

Brittle, dull, poor conductors.
Tend to gain electrons to form anions.
Form acidic oxides.

Key Terms and Definitions

Pauli exclusion principle: No two electrons in an atom can have the same four quantum numbers.
Orbital box diagram: Visual representation of electron configuration using boxes and arrows.
Aufbau principle: Electrons fill lowest energy orbitals first.
Effective nuclear charge: Net positive charge experienced by valence electrons.
Valence electrons: Electrons in the outermost shell, involved in chemical bonding.
Core electrons: Electrons in inner shells, not involved in bonding.
Hund's rule: Electrons occupy orbitals singly before pairing.
Isoelectronic: Species with the same number of electrons.

Summary of Electron Configuration Notation

spdf notation:
Noble gas notation:
Orbital box notation: [Ne] ↑↓ ↑ ↑ ↑ (for 3s and 3p)

Practice and Application

Write electron configurations for main group and transition elements using all three notations.
Predict periodic trends and explain exceptions using quantum mechanical principles.
Compare properties of metals and nonmetals based on electron configuration and periodic position.

Additional info: These notes are based on a college-level General Chemistry course, focusing on atomic structure, electron configuration, and periodic trends. All equations are provided in LaTeX format for clarity.