1. Chapter 4 Arrangement of Electrons in Atoms
Section 4-1
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.
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 .
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 J∙s)

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

9. 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.

10. 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"

11. 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?

12. 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.

13. 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

14. 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

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

16. 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