1. II. Bohr Model of the Atom (p. 94 - 97)
Ch. 4 - Electrons in Atoms
II. Bohr Model of the Atom (p )
C. Johannesson

2. Ernest Rutherford’s Model
Discovered dense positive piece at the center of the atom- “nucleus”
Electrons would surround and move around it, like planets around the sun
Atom is mostly empty space
It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

3. Niels Bohr’s Model Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In specific circular paths, or orbits, at different levels.
An amount of fixed energy separates one level from another.

4. The Bohr Model of the Atom
I pictured the electrons orbiting the nucleus much like planets orbiting the sun.
However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another.
Niels Bohr

5. Bohr Model
C. Johannesson

6. Bohr’s model Energy level of an electron
analogous to the rungs of a ladder
The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder
A quantum of energy is the amount of energy required to move an electron from one energy level to another

7. B. Bohr Model
e- exist only in orbits with specific amounts of energy called energy levels
Therefore…
e- can only gain or lose certain amounts of energy
only certain photons are produced
C. Johannesson

8. Before After Photon wavelength changes Photon Moving Electron
Electron velocity changes
Fig. 5.16, p. 145

9. A. Line-Emission Spectrum
Hydrogen Gas
excited state
ENERGY IN
PHOTON OUT
ground state
Atoms are “excited” with an electric current
C. Johannesson

10. These are called the atomic emission spectrum
Unique to each element, like fingerprints!
Very useful for identifying elements

11. Light is a Particle? Energy is quantized. Light is a form of energy.
Therefore, light must be quantized
These smallest pieces of light are called photons.
Photoelectric effect? Albert Einstein
Energy & frequency: directly related.

12. Explanation of atomic spectra
When we next write electron configurations, we are writing the lowest energy.
The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level.

13. Changing the energy
Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

14. Changing the energy
Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be “excited”

15. Changing the energy
As the electron falls back to the ground state, it gives the energy back as light

16. Changing the energy They may fall down in specific steps
Each step has a different energy

17. This is a simplified explanation!
Ultraviolet
Visible
Infrared
The further they fall, more energy is released and the higher the frequency.
This is a simplified explanation!
The orbitals also have different energies inside energy levels
All the electrons can move around.

18.  = h/mv (from Louis de Broglie)
What is light?
Light is a particle - it comes in chunks.
Light is a wave - we can measure its wavelength and it behaves as a wave
If we combine E=mc2 , c=, E = 1/2 mv2 and E = h, then we can get:
 = h/mv (from Louis de Broglie)
called de Broglie’s equation
Calculates the wavelength of a particle.

19. A. Wave-properties Wavelength () - length of one complete wave
Frequency () - # of waves that pass a point during a certain time period
hertz (Hz) = 1/s
C. Johannesson

20. A. Waves
C. Johannesson

21. B. EM Spectrum
HIGH
ENERGY
LOW
ENERGY
C. Johannesson

22. B. EM Spectrum HIGH LOW ENERGY ENERGY R O Y G. B I V red orange yellow
green
blue
indigo
violet
C. Johannesson

23. c =  B. EM Spectrum c: speed of light (3.00  108 m/s)
Frequency & wavelength are inversely proportional
c = 
c: speed of light (3.00  108 m/s)
: wavelength (m, nm, etc.)
: frequency (Hz)
C. Johannesson

24. Long Wavelength = Low Frequency Low ENERGY Short Wavelength =
Wavelength Table
Long
Wavelength =
Low Frequency
Low ENERGY
Short
Wavelength =
High Frequency
High ENERGY

25. “wave-particle duality”
C. Quantum Theory
Einstein (1905)
Concluded - light has properties of both waves and particles
“wave-particle duality”
Photon - particle of light that carries a quantum of energy
C. Johannesson

26. C. Quantum Theory
The energy of a photon is proportional to its frequency.
E = h
E: energy (J, joules)
h: Planck’s constant (  J·s)
: frequency (Hz)
C. Johannesson

27. Practice Problems
What is the wavelength of radiation with a frequency of 1.50 x 1013 Hz?
w =
What is the frequency of radiation with a wavelength of 5.00 x 10-8 m? In what region of the electromagnetic spectrum is this radiation?
f =

28. Examples
What is the wavelength of blue light with a frequency of 8.3 x 1015 hz?
What is the frequency of red light with a wavelength of 4.2 x 10-5 m?
What is the energy of a photon of each of the above?