1. Chapter 4 Arrangement of Electrons in Atoms
Development of a new atomic model

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 infrared 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. 98)

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. v is the greek symbol “nu”
Measured in Hz = Hertz = s-1 = 1/s
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. E = hv (Planck’s equation)
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 electromagnetic radiation (EMR)            
E = hv  (Planck’s equation)         
E= Energy (J, of a quantum of radiation)
v= frequency of radiation emitted
h= Planck’s constant (6.626 x 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. Hydrogen Emission Spectrum
Electrons start in lowest possible level - ground state.
Absorb energy - become excited and move outward.
Release energy- 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
Can’t determine position and speed simultaneously

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. How colors of Fireworks are made
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: x 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: