Arrangement of Electrons in Atoms

1. Arrangement of Electrons in Atoms Development of a New Atomic Model

2. 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)

3. Electromagnetic Spectrum

4. 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

5. 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

6. 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)

7. Equation Practice • What is the frequency of yellow light with a wavelength of 548 nm?

8. Equation Practice • What is the wavelength of blue light with a frequency of 4.60 x 1023 Hz?

9. Equation Practice • What is the energy of magenta light with a wavelength of 691 nm?

10. 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.

11. 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"

12. 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?

13. 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.

14. 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

15. 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), dB proved that electrons can have specific energies and that Bohr’s quantized orbits were actually correct

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

17. Schrodinger Wave Equation • He gave more support to Bohr’s quantized energy levels • Quantum theory – describes the wave properties of electrons using mathematical equations • Disproved Bohr’s “train tracks” within those energy levels