Electrons in Atoms and the Periodic Table: Chapter 9 Study Notes

Electrons in Atoms and the Periodic Table

Introduction

This chapter explores the structure of atoms, the behavior of electrons, and how these concepts explain the properties and trends observed in the periodic table. Understanding atomic models and electron configurations is fundamental to predicting chemical reactivity and the organization of elements.

Models of the Atom

Historical Context: The Hindenburg Disaster

• Hydrogen Gas: Used for buoyancy in airships due to its low density, but highly reactive and flammable.

• Helium Gas: Modern blimps use helium, which is inert and non-flammable, due to its stable electron configuration.

• Reactivity of Hydrogen: Hydrogen atoms are so reactive that they combine to form diatomic hydrogen molecules (H2).

• Inertness of Helium: Helium's electron configuration (two electrons) makes it chemically stable and nonreactive.

Additional info: The difference in reactivity between hydrogen and helium is explained by their electron arrangements.

Periodic Law and Element Groups

• Periodic Law: When elements are arranged by increasing atomic number, recurring sets of properties appear periodically.

• Group 1A Elements: (e.g., lithium, sodium) share reactivity similar to hydrogen.

• Noble Gases: (e.g., neon, argon) share inertness similar to helium.

Atomic Models: Bohr and Quantum-Mechanical

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

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

• Key Contributors: Niels Bohr, Erwin Schrödinger, and Albert Einstein contributed to the development of quantum mechanics.

Electromagnetic Radiation and Light

Nature of Light

• Electromagnetic Radiation: Light is energy that travels at a constant speed ($c = 3.0 \times 10^8$ m/s).

• Wave-Particle Duality: Light exhibits both wave-like and particle-like (photon) properties.

Properties of Light

• Wavelength ($\lambda$): Distance between adjacent wave crests.

• Frequency ($\nu$): Number of cycles passing a point per second.

• Relationship: Wavelength and frequency are inversely related: $c = \lambda \nu$

• Energy of a Photon: $E = h\nu$ (where $h$ is Planck's constant)

Visible Light and Color

• Visible Spectrum: Ranges from red (750 nm) to violet (400 nm).

• Color Perception: Objects appear colored due to selective reflection and absorption of wavelengths.

Electromagnetic Spectrum

• Gamma Rays: Shortest wavelength, highest energy; can damage biological molecules.

• X-Rays: Used in medical imaging; also ionizing and potentially harmful.

• Ultraviolet (UV): Causes sunburn; energetic enough to damage molecules.

• Visible Light: Enables vision; does not damage biological molecules.

• Infrared: Felt as heat; emitted by warm objects.

• Microwaves: Used in ovens and communication; absorbed by water.

• Radio Waves: Longest wavelength; used for communication.

Atomic Spectra and Models

Emission Spectra

• White Light: Produces a continuous spectrum.

• Elemental Emission: Each element emits light at specific wavelengths, producing a line spectrum.

• Hydrogen Spectrum: Lines correspond to electron transitions between energy levels.

Bohr Model of the Atom

• Quantized Orbits: Electrons occupy fixed energy levels specified by quantum number $n$.

• Energy Transitions: Electrons absorb energy to move to higher orbits (excitation) and emit photons when relaxing to lower orbits.

• Photon Energy: $E = h\nu = \frac{hc}{\lambda}$

• Limitations: Accurately predicts hydrogen spectrum but fails for multi-electron atoms.

Quantum-Mechanical Model

• Orbitals: Probability maps showing where electrons are likely to be found.

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

• Subshells: Indicated by letters (s, p, d, f) specifying orbital shape.

• Ground State: Lowest energy configuration; electrons occupy lowest available orbitals.

• Excited State: Electrons occupy higher energy orbitals after absorbing energy.

Orbital Types and Shapes

• s Orbitals: Spherical shape.

• p Orbitals: Dumbbell shape.

• d Orbitals: More complex shapes; five per shell starting at $n=3$.

• f Orbitals: Even more complex; seven per shell starting at $n=4$.

Electron Configurations

Writing Electron Configurations

• Notation: Indicates the number of electrons in each orbital (e.g., H: 1s1).

• Orbital Diagrams: Use arrows to represent electrons and their spins in boxes for each orbital.

• Pauli Exclusion Principle: No more than two electrons per orbital, with opposite spins.

• Hund's Rule: Electrons fill orbitals of equal energy singly before pairing.

• Energy Ordering: Lower-energy orbitals fill before higher-energy ones (e.g., 4s before 3d).

Core and Valence Electrons

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

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

• Noble Gas Notation: Use brackets to abbreviate core electron configuration (e.g., Na: [Ne]3s1).

Periodic Table and Electron Configuration

• Groups: Elements in the same group have the same number of valence electrons.

• Blocks: s block (left), p block (right), d block (transition metals), f block (lanthanides/actinides).

• Trends: Main-group elements: group number equals number of valence electrons (except helium).

• Transition Metals: d orbital principal quantum number is row number minus one.

• Exceptions: Chromium (Cr) and Copper (Cu) have unusual electron configurations due to stability of half-filled or fully-filled subshells.

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 a gaseous atom.

• Across a Period: Ionization energy increases as atoms approach noble gas configuration.

• Down a Group: Ionization energy decreases as atomic size increases.

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.

Applications: Chemistry and Health

Ion Pumps in Biology

• Sodium (Na+) and Potassium (K+) Ions: Essential for nerve signal transmission and cellular function.

• Ion Gradients: Created by pumps in cell membranes; crucial for biological processes.

Radiation and Health

• Ionizing Radiation: X-rays and gamma rays can ionize atoms and molecules, damaging biological tissue.

• Medical Use: Radiation therapy targets cancer cells but can also affect healthy cells.

Summary Table: Electromagnetic Spectrum

Type	Wavelength	Energy	Applications/Effects
Gamma Rays	Shortest	Highest	Produced by stars, nuclear decay; damages biological molecules
X-Rays	Short	High	Medical imaging; ionizing
Ultraviolet	Shorter than visible	Moderate-High	Sunburn, skin damage
Visible Light	400-750 nm	Moderate	Vision
Infrared	Longer than visible	Lower	Heat, night vision
Microwaves	Long	Low	Cooking, communication
Radio Waves	Longest	Lowest	Radio, TV, cell phones

Key Equations

• Speed of Light: $c = \lambda \nu$

• Photon Energy: $E = h\nu$

• Relationship of Energy and Wavelength: $E = \frac{hc}{\lambda}$

Learning Objectives

• Understand electromagnetic radiation and its properties.

• Predict relative wavelength, energy, and frequency of different types of light.

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

• Write electron configurations and orbital diagrams for atoms.

• Identify valence and core electrons.

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