Chapter 3 JM st

Chapter Overview

• Focus: Electronic Structure and Periodic Properties of Elements

• Key Concepts: Electromagnetic Energy, Quantum Theory, Electronic Structure of Atoms, Properties of Elements, The Periodic Table, Molecular and Ionic Compounds.

Electromagnetic Energy

• Visible Light:

  • Isaac Newton’s experiments: White light consists of rainbow colors.

  • Thought of light as tiny particles.

  • Christiaan Huygens suggested light has wave-like properties.

• Electromagnetic Radiation:

  • Shows visible light is part of a vast spectrum.

  • Developed by James Clerk Maxwell, who discredited the particle model of light.

  • Used to infer electron energies in atoms or molecules.

  • Foundation of modern technologies.

Wave Properties of Light

• Wave Definition:

  • Oscillation or periodic movement that transports energy.

  • Characterized by:

    • Wavelength (λ): Distance between two peaks (measured in nm).

    • Frequency (ν): Number of cycles per second (measured in Hz).

    • Amplitude: Height of the wave's peaks; relates to light intensity.

• Wavelength and Frequency:

  • Wavelength (λ) and frequency (ν) are inversely related: c = λν where c is the speed of light.

The Speed of Light

• Constant Speed of Light:

  • c = 2.9979 x 10^8 m/s in a vacuum.

  • Light speed product with wavelength and frequency remains constant.

Practical Applications

• Example Problem:

  • Calculate frequency from wavelength using c = λν.

    • Example: Arsenic emits light at 193.7 nm, find frequency ν.

Photoelectric Effect

• Electrons emitted from a metal surface when exposed to light (only above threshold frequency).

• Albert Einstein: Proposed light behaves like particles (photons).

• Photon energy formula: E = hν or E = hc/λ.

Wave-Particle Duality

• Light and particles exhibit both wave-like and particle-like behaviors.

• Explains phenomena such as interference patterns.

Line Spectra and Quantum Transitions

• Line Spectra:

  • Produced by gases at low pressures under electrical current.

  • Each element shows distinct lines representing quantized energies.

• Hydrogen Line Spectrum:

  • Only specific energy levels for electrons; transitions produce photons.

• Rydberg-Balmer Equation for calculating spectral lines.

Bohr Model

• Niels Bohr:

  • Electrons in quantized orbits around a nucleus.

  • Photons emitted or absorbed during energy transitions between orbits.

• Key Concepts:

  • Ground state: Lowest energy level (n=1).

  • Excited states: Higher energy levels.

  • Energy conservation principle applies to photon emission/absorption.

Quantum Behavior of Electrons

• Electrons described by quantum numbers:

  • n (principal), l (azimuthal), ml (magnetic), ms (spin).

• Electrons occupy orbitals defined by quantized energy and shape.

Effective Nuclear Charge and Shielding

• Effective nuclear charge (Zeff) affects electron ability to remain in orbitals.

• Shielding influences orbital energy.

Electron Configurations and Orbital Filling

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

• Hund’s Rule: Maximizes unpaired electrons in degenerate orbitals.

• Elements exhibit configurations based on periodic table trends.

Ionization Energy and Electron Affinity

• Ionization Energy: Energy required to remove an electron.

  • Increases across a period and decreases down a group.

• Electron Affinity: Change in energy when gaining an electron.

  • Generally becomes more negative across a period.

Ionic Compounds

• Formed between cations and anions (e.g. NaCl).

• Charge balance dictates the formula structure.

• Specific ion formation based on group trends in the periodic table.