1 st - Light & Electrons

1. 1st- Light & Electrons Chapter 4Arrangement of Electrons in Atoms

2. The Development of A New Atomic Model • Rutherford’s model was an improvement over previous models, but still incomplete. • Where exactly are electrons located? • What prevented the electrons from being drawn into the nucleus?

3. Wave Description Of Light • Electromagnetic Radiation: • form of energy that exhibits wavelike behavior as it travels through space. • EX: visible light, X-ray, Ultraviolet and inferred light, microwaves, and radio waves. • Travels at a constant speed of 3.0 x 108 m/s • Electromagnetic Spectrum: All the electromagnetic radiation form the ES. (fig 4-1, p. 92)

4. Electromagnetic Spectrum

5. Wave Calculations • Wavelength (λ) - distance between two peaks . Measured in meters • Frequency (v) - number of peaks that pass a point each second. • Hz = Hertz = s-1 • c = λ v  • where c = 3.0 x 108 m/s

6. Is light really a wave? • Max Planck – did experiments with light-matter interactions where light did not act like a wave • Photoelectric Effect - emission of electrons from a metal when light shines on the metal. • Only emitted at certain energies; wave theory said any energy should do it. • Led to the particle theory of light

7. Planck suggested that objects emit energy in specific amounts called QUANTA • Quantum - minimum quantity of energy that can be lost or gained by an atom. • led Planck to relate the energy of an electron with the frequency of EMR             • E = hv  • E= Energy (J, of a quantum of radiation) • v= frequency of radiation emitted • h= Planck’s constant (6.626 x 10-34 J∙s)

8. leads to Einstein’s dual nature of light (EMR behaves as both a wave and a particle) • Photon - particle of EMR having zero mass and carrying a quantum of energy.

9. Hydrogen Emission Spectrum • Ground State - Lowest energy state of electron. • Excited State - higher energy than ground state. • Bright-line Spectrum (emission spectrum) • Series of specific light frequencies emitted by elements "spectra are the fingerprints of the elements"

10. Bohr Model Of H Atom • Bohr explained how the electrons stay in the cloud instead of slamming into the nucleus • Definite orbits; paths • The greater the distance from the nucleus, the greater the energy of an electron in that shell.

11. Electrons start in lowest possible level - ground state. • Absorb energy - become excited and shift upward. • Dropping back down - emits photons (packets of energies equal to the previously absorbed energy). • Hydrogen Emission Spectrum

12. Quantum Model of the Atom • Bohr’s model was great, but it didn’t answer the question “why?” • Why did electrons have to stay in specific orbits? • Why couldn’t the electrons exist anywhere within the electron cloud? • Louis de Broglie pointed out that electrons act like waves • Using Planck’s equation (E=hv), Louis proved that electrons can have specific energies and that Bohr’s quantized orbits were actually correct

13. Heisenberg Uncertainty Principle • Impossible to determine both the exact location and velocity of an electron

14. Schrodinger Wave Equation • He gave more support to Bohr’s quantized energy levels • Quantum theory – describes the wave properties of electrons using mathematical equations

15. Equation Practice • What is the energy of yellow light with a wavelength of 548 nm? • Convert nm  m • Use wave equation to calculate frequency • Use Planck’s equation to calculate energy

16. Equation Practice • What is the energy of blue light with a wavelength of 460 nm? • Convert nm  m • Use wave equation to calculate frequency • Use Planck’s equation to calculate energy

17. Equation Practice • What is the energy of magenta light with a wavelength of 691 nm? • Convert nm  m • Use wave equation to calculate frequency • Use Planck’s equation to calculate energy

18. 2nd- Flame Tests How colors of Fireworks are made

19. Background Information • electromagnetic radiation: • form of energy that acts as a wave as it travels • includes: visible light, X rays, ultraviolet and infrared light, microwaves, and radio waves • All forms are combined to form electromagnetic spectrum

20. Electromagnetic Spectrum

21. Background Information • all forms of radiation travel at a speed of 3.0 x 108 m/s (c: speed of light) • wavelength: () • distance between points on adjacent waves • in nm (109 nm = 1 m) • frequency: () • number of waves that passes a point in a second • in waves/second Inversely proportional!

23. Background Information • we also have an equation to relate Energy (E) of the radiation and frequency () • where h is Planck’s constant: 6.626 x 10-34 J*s

24. Flame Tests • used to identify metal ions in unknown compounds • usually for Group 1 and 2 metals • when the metal ions in compounds are heated, they release certain wavelengths of visible radiation • colors are dependent on energy of electrons in ion

25. Step C2: Convert  in nm to m • begin with wavelength in nm • use the statement of equality: 1m = 109 nm • Example: 360 nm

26. Step C3: Find  • Example:

27. Step C4: Find E • Example: