Electrons Unit Review

Electrons Unit Review

Relationships between Energy, Wavelength, and Frequency

  • Wavelength and Frequency are Inversely Related

    • As the wavelength increases, the frequency decreases.

  • Wavelength and Energy are Inversely Related

    • As the wavelength increases, the energy decreases.

  • Energy and Frequency are Directly Related

    • As the frequency increases, the energy also increases.

Characteristics of Electromagnetic Waves

  • High Energy

    • Associated with Short Wavelength and High Frequency.

  • Low Energy

    • Associated with Long Wavelength and Low Frequency.

The Electromagnetic Spectrum

  • Range of Wavelengths and Frequencies:

    • Gamma Rays: 10−16extm10^{-16} ext{ m} to 10−14extm10^{-14} ext{ m}

    • X-Rays: 10−14extm10^{-14} ext{ m} to 10−10extm10^{-10} ext{ m}

    • Ultraviolet: 10−10extm10^{-10} ext{ m} to 10−8extm10^{-8} ext{ m}

    • Visible Light: ranges from 400 nm (blue) to 700 nm (red)

    • Infrared: 10−8extm10^{-8} ext{ m} to 10−5extm10^{-5} ext{ m}

    • Microwaves: 10−5extm10^{-5} ext{ m} to 1extm1 ext{ m}

    • TV and Radio Waves: 1extm1 ext{ m} to 100extm100 ext{ m}

Wavelength Conversion Factor

  • Conversion Factor:

    • 1extm=109extnm1 ext{ m} = 10^{9} ext{ nm}

Discharge Tube and Flame Phenomena

  • Key Terms:

    • Photon: Packet of light.

    • Excited: Electron moves up in energy level.

    • Ground State: All electrons are closest to the nucleus.

    • Emission: Release of energy as an electron drops to a lower energy level.

    • Absorption: Process in which electrons take in energy to move to a higher energy level.

Electron Locations in Atoms

  • Electrons exist in energy levels around the nucleus, designated as n = 1 to n = 7.

  • In each energy level, there are sublevels designated as s, p, d, f.

  • Orbitals within Sublevels:

    • Each orbital can hold up to 2 electrons:

    • s = 1 orbital

    • p = 3 orbitals

    • d = 5 orbitals

    • f = 7 orbitals

Energy Levels and Sublevels Configuration

  • The correlation between energy levels and sublevels:

    • n = 1: 1 sublevel (s)

    • n = 2: 2 sublevels (s and p)

    • n = 3: 3 sublevels (s, p, and d)

    • n = 4-7: Each has 4 sublevels (s, p, d, and f)

Electron Configurations and Their Representations

  • Electron configurations illustrate the arrangement of electrons in various levels and orbitals.

  • Different types of electron configurations include:

    • Aufbau diagrams: Show the energy filling order.

    • Full spdf configurations: Provide a complete electron arrangement.

    • Noble gas configurations: Use noble gases to simplify electron configurations.

Orbital Filling Principles

  • Three Principles Governing Electron Filling:

    • Aufbau Principle: Electrons fill from lowest to highest energy orbitals first.

    • Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons, represented as one up-arrow and one down-arrow.

    • Hund's Rule: When filling orbitals of the same energy, one electron enters each orbital until all orbitals have one electron, followed by pairing.

Orbital Filling Diagrams & Aufbau Diagrams

  • Aufbau Diagrams: Visually depict the increasing energy of sublevels.

    • The configuration increases from 1s to higher orbitals (6d, 7p, etc.).

Noble Gas Configuration Examples

  • Representations:

    • Helium: 1s² → [He]

    • Lithium: 1s²2s¹ → [He] 2s¹

    • Beryllium: 1s²2s² → [He] 2s²

    • Boron: 1s²2s²2p¹ → [He] 2s²2p¹

    • Fluorine: 1s²2s²2p⁵ → [He] 2s²2p⁵

    • Neon: 1s²2s²2p⁶ → [Ne]

    • Sodium: 1s²2s²2p⁶3s¹ → [Ne] 3s¹

    • Magnesium: 1s²2s²2p⁶3s² → [Ne] 3s²

    • Aluminum: 1s²2s²2p⁶3s²3p¹ → [Ne] 3s²3p¹

    • Chlorine: 1s²2s²2p⁶3s²3p⁵ → [Ne] 3s²3p⁵

Understanding Excited State vs Ground State

  • Difference between states:

    • Ground State: Electrons in their lowest energy state.

    • Excited State: Electrons have absorbed energy and moved to higher energy levels.

Ionic Considerations

  • Examples of Ions:

    • O²⁻: Ion with a charge of -2, gaining two electrons.

    • Mn³⁺: Ion with a charge of +3, losing three electrons.