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Chapter Overview

• Chapter 6: Electronic Structure and Periodic Properties of Elements.

Outline of Chapter 6

• 6.1 Electromagnetic Energy

• 6.2 The Bohr Model

• 6.3 Development of Quantum Theory

• 6.4 Electronic Structure of Atoms (Electron Configurations)

• 6.5 Periodic Variations in Element Properties

Learning Objectives

6.1 Electromagnetic Energy

• Explain basic wave behaviors: traveling waves, standing waves.

• Describe the wave nature of light.

• Calculate light-wave properties (frequency, wavelength, energy).

• Distinguish between line and continuous emission spectra.

• Describe the particle nature of light.

• Important calculation notes: Memorize conversions between nm and m, Hz to MHz, etc.

Electromagnetic Energy

• Light nature studied since antiquity.

• Isaac Newton (17th Century)

  • Proposed white light consists of rainbow colors.

  • Suggested light composed of tiny particles moving fast.

Light Behavior

• Thomas Young (19th Century)

  • Light through narrow slits creates interference patterns.

  • Found to align more with wave theory than particle theory.

• James Clerk Maxwell

  • Defined light as a visible spectrum part of electromagnetic waves.

  • Established wave-particle duality concept.

Wave Characteristics

• Wave: Oscillatory movement transporting energy.

• Examples: Shaking rope, pebble in pond, air expansion during lightning.

• Energy transported while matter remains stationary.

Speed of Light

• Electromagnetic waves: Electric and magnetic fields oscillate perpendicular to travel direction.

• Speed in a vacuum: c = 2.998 × 10^8 m/s (often rounded to 3.00 × 10^8 m/s).

Wave Properties

• Waves characterized by:

  • Wavelength (λ): Distance between peaks/troughs.

  • Frequency (ν): Number of wavelengths passing a point in a unit time.

Relationship Between Wavelength, Frequency, and Speed

• Wave speed equation: c = λν.

• Inverse relationship: As wavelength increases, frequency decreases.

Frequency Units

• Measured in hertz (Hz): cycles per second.

• Common multiples:

  • 1 MHz = 1 × 10^6 Hz

  • 1 GHz = 1 × 10^9 Hz

Electromagnetic Spectrum

• Range of all electromagnetic radiation types.

• Visible light is a small portion within the spectrum.

Blackbody Radiation

• Blackbody: Ideal emitter, approximates many materials under heat.

• Sunlight has continuous spectrum due to broad wavelength range.

• Max Planck (1900): Developed expression for blackbody radiation fitting observations.

Planck’s Constant

• Energy quantization assumption: E = nhν (where h = 6.626 × 10^-34 J·s).

  • Atoms vibrate at frequencies increasing with temperature.

Photoelectric Effect

• Photon energy fluctuates with frequency: E = hv or E = hc/λ.

Line Spectra Observations

• Emission from solids/gases under sufficient heating leads to continuous spectra.

• Contrast: Line spectra exist with distinct lines for each element, formed by low-pressure gas excitation.

The Bohr Model (Niels Bohr, 1913)

• Explains hydrogen's spectra and propels quantum mechanics development.

• Assumed quantization in orbits and introduced energy levels.

Key Assumptions of Bohr’s Model

• Electrons occupy fixed orbits without emitting radiation unless transitioning between orbits.

• Energy transitions correspond to differences in orbital energies.

• Photons emitted/absorbed match these specific energy changes.

Limitations of the Bohr Model

• Failed for multi-electron atoms due to electron interactions.

• Continued reliance on classical mechanics' precise orbit model.

Quantum Theory Development

• Electrons exist in discrete energy levels with quantized values, defined by quantum numbers.

Atomic Orbitals

• Named based on Principal Quantum Number (n), correlating to energy levels.

• Shapes and orientations defined by Angular Momentum Quantum Number (ℓ) and Magnetic Quantum Number (mℓ).

Electron Spin

• Electrons possess a characteristic rotation described by Spin Quantum Number (mₛ), with values ±1/2.

• The Pauli Exclusion Principle states two electrons can occupy the same orbital only if they have opposite spins.

Electron Configuration Principles

• Electrons fill orbitals based on the Aufbau Principle (building up from lowest energy levels).

• Core and valence electrons defined based on their shells.

Valence and Core Electron Definitions

• Valence electrons: Outermost electrons affecting chemical reactivity.

• Core electrons: Inner shell electrons following noble gas configurations.

Periodic Trends in Element Properties

• Atomic size, Ionization energy, and Electron affinity trends are periodic and are influenced by electron configuration and effective nuclear charge.

• Covalent radius and ionic size vary across groups and periods, influenced by effective nuclear charge and electron distribution.