Electrons in Atoms and the Periodic Table: Models, Orbitals, and Periodic Trends

Models of the Atom

Introduction to Atomic Models

Understanding the structure of atoms is fundamental to chemistry. Early models, such as the Bohr model, and more advanced models, like the quantum-mechanical model, help explain the behavior and properties of elements.

• Bohr Model: Electrons travel in fixed, circular orbits around the nucleus at specific distances.

• Quantum-Mechanical Model: Electrons exist in orbitals, which are probability maps rather than fixed paths.

• Mendeleev's Periodic Law: When elements are arranged by increasing atomic number, sets of properties recur periodically.

Historical Context and Key Figures

• Niels Bohr: Developed the Bohr model to explain emission spectra.

• Erwin Schrödinger: Contributed to the development of quantum mechanics.

• Albert Einstein: Played a role in the early development of quantum theory.

The Bohr Model of the Atom

Key Characteristics

The Bohr model describes electrons as moving in circular orbits around the nucleus, similar to planets around the sun, but only at specific, fixed distances.

• Success: Predicted the lines of the hydrogen emission spectrum.

• Limitation: Failed to predict spectra for elements with more than one electron.

The Quantum-Mechanical Model

Orbitals and Probability Maps

The quantum-mechanical model replaced the Bohr model, introducing orbitals as regions of space where electrons are likely to be found.

• Orbitals: Probability maps, not fixed paths.

• Electron Behavior: Electrons do not act like classical particles; their position is described statistically.

• Principal Quantum Number (n): Specifies the principal shell and energy level.

• Subshells: Indicated by letters (s, p, d, f), which specify orbital shapes.

Ground State and Excited State

• Ground State: Lowest energy state, e.g., hydrogen's electron in the 1s orbital.

• Excited State: Electron absorbs energy and moves to a higher-energy orbital.

Energy and Orbital Shapes

• Energy Increases with n: Higher principal quantum number means higher energy.

• Orbital Shapes:

  • s orbital: Spherical shape.

  • p orbital: Dumbbell shape.

  • d orbital: More complex, cloverleaf shapes.

  • f orbital: Even more complex shapes.

• Dot Density: Represents probability of finding an electron; highest near the nucleus.

Table: Number of Subshells per Principal Shell

Shell (n)	Number of Subshells	Letters Specifying Subshells
n = 1	1	s
n = 2	2	s, p
n = 3	3	s, p, d
n = 4	4	s, p, d, f

Electron Configurations

How Electrons Occupy Orbitals

• Electron Configuration: Shows the occupation of orbitals by electrons for an atom.

• Orbital Diagram: Uses arrows in boxes to represent electrons and their spins.

• Pauli Exclusion Principle: Orbitals hold no more than two electrons with opposing spins.

• Hund's Rule: Electrons fill orbitals singly first, with parallel spins, before pairing.

Periodic Table and Electron Configuration

Blocks and Patterns

• s block: First two groups (left side).

• p block: Six groups (right side).

• d block: Transition metals (center).

• f block: Lanthanides and actinides (bottom).

Valence and Core Electrons

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

• Core Electrons: All other electrons not in the outermost shell.

Periodic Trends

Atomic Size

• Across a Period: Atomic size decreases due to increased nuclear charge pulling electrons closer.

• Down a Group: Atomic size increases as principal quantum number (n) increases, placing electrons farther from the nucleus.

Ionization Energy

• Definition: Energy required to remove an electron from an atom in the gaseous state.

• Across a Period: Ionization energy increases.

• Down a Group: Ionization energy decreases.

Metallic Character

• Definition: Tendency of an element to lose electrons and form cations.

• Across a Period: Metallic character decreases.

• Down a Group: Metallic character increases.

Summary and Learning Objectives

Key Points to Learn

• Understand the Bohr and quantum-mechanical models of the atom.

• Write electron configurations and orbital diagrams.

• Identify valence and core electrons.

• Write electron configurations based on periodic table position.

• Explain how valence electrons determine chemical properties.

• Recognize periodic trends in atomic size, ionization energy, and metallic character.

Example: Electron Configuration of Sodium

• Full configuration: 1s2 2s2 2p6 3s1

• Noble gas core notation: [Ne] 3s1

Example: Atomic Size Trend

• Atomic radius of sodium is larger than that of chlorine in the same period.

• Atomic radius of sodium increases compared to lithium in the same group.

Important Equations

• Principal Quantum Number: $n = 1, 2, 3, ...$

• Electron Configuration Notation: $1s^2 2s^2 2p^6 3s^1$

Additional info: These notes are based on textbook slides for Chapter 9, "Electrons in Atoms and the Periodic Table," and cover atomic models, electron configuration, and periodic trends, which are central to introductory college chemistry.