CHEM10 Chapter 4 Part 1

Chapter Four: Electronic Structure and Organization of Electrons

Introduction to Light and Electrons

• Understanding electrons requires knowledge of light and its behavior.

• Electrons produce light and colors, connecting to atomic structure introduced in Chapter Three.

• Light helps to understand electronic structures and their configurations around the nucleus.

Light and Electromagnetic Radiation

• Light: A type of electromagnetic radiation traveling in waves produced by vibrating charged particles.

• Photons: Light exists in small packets of energy called photons.

• The electromagnetic spectrum includes a range of waves from low energy (radio waves) to high energy (gamma rays).

• Visible spectrum: A small range of electromagnetic waves visible to the human eye, with colors from red (700 nm) to violet (350 nm).

Key Definitions

• Wavelength (λ): Distance between successive crests of a wave, measured in meters or nanometers (1 nm = 10⁻⁹ m).

• Frequency (ν): Number of waves passing a point per second, expressed in hertz (Hz), where 1 Hz = 1 s⁻¹.

Relationship Between Wavelength and Frequency

• Inverse relationship: as wavelength decreases, frequency increases and vice versa.

• Speed of Light (C): Approximately 2.998 x 10⁸ m/s.

• Equation: C = λν, where λ is in meters.

Example Calculation

• For green light with a wavelength of 5.00 x 10² nm:

  • Convert to meters: 5.00 x 10² nm = 5.00 x 10⁻⁷ m.

  • Calculate frequency: ν = C / λ = (3.00 x 10⁸ m/s) / (5.00 x 10⁻⁷ m) = 6.00 x 10¹⁴ s⁻¹.

Energy of Light

• Energy (E) of light relates directly to its frequency and can be calculated with:

  • E = Hν (H = Planck's constant = 6.63 x 10⁻³⁴ J·s)

  • E = HC / λ (derived from E = Hν and C = λν).

Example Calculation for Energy

• For a photon frequency of 7.50 x 10¹⁴ Hz:

  • Find wavelength using λ = C / ν = (3.00 x 10⁸ m/s) / (7.50 x 10¹⁴ s⁻¹) = 4.00 x 10⁻⁷ m = 400 nm (violet).

  • Calculate energy: E = Hν = (6.63 x 10⁻³⁴ J·s) * (7.50 x 10¹⁴ s⁻¹) = 4.97 x 10⁻¹⁹ J.

Relationship Between Light and Electronic Structure

• Fireworks: Colors from burning metal ions (e.g., strontium = red, barium = green).

• Flame tests: Used to observe colors emitted by metal ions heated in flame.

• Gas lamps: Produce characteristic colors when an electric current passes through gases (e.g., neon = red, argon = blue).

• Line spectra: Unique patterns of colors associated with each element (like fingerprints).

Historical Perspective on Atomic Structure

• High-energy light can knock off electrons from atoms (photoelectric effect).

• Niels Bohr's Model (1913): Electrons orbit the nucleus; energy levels correspond to orbital closeness.

• Electrons can absorb energy to move to higher levels and emit photons when relaxing back, which corresponds to visible light.

• Bohr's model was effective for hydrogen but not more complex atoms.

Quantum Mechanics and Electron Behavior

• Quantum Model: Electrons as both particles and waves; uncertainty principle states exact velocity and position cannot be known.

Rules for Electron Arrangement in Atoms

1. Electrons occupy energy levels labeled by principal quantum number (n). Higher n = greater distance from nucleus.

2. Energy levels have sub-levels (s, p, d, f) with increasing energy.

3. Orbitals hold a maximum of two electrons (with opposite spins).

4. Maximum electrons per level = 2n²:

   • Level 1: 2 (s)

   • Level 2: 8 (s, p)

   • Level 3: 18 (s, p, d)

   • Level 4: 32 (s, p, d, f)

Energy Level Diagram

• Diagram shows the energy levels (n) and sub-levels (s, p, d, f), with shape representation:

  • e.g., yellow for 2s and 2p, green for 3s, 3p, 3d, etc.

Conclusion

• Understanding these principles will help explore electron configurations and behavior of elements in future discussions.