1. The Atom
Lesson 3 : The Bohr Model

2. Bohr Model of an Atom
Electrons orbit the nucleus in fixed energy ranges called orbits (energy levels)
An electron can move from one energy level to another by gaining or losing discrete amounts of energy
Electrons cannot be found between energy levels (think of energy levels like rungs on a ladder )
The lowest energy level is
closest to the nucleus, the
highest is farthest away
The electron energy levels are
quantized.
Think of it like potential energy, the higher you are the more you have

3. Absorption Vs Emission
When an electron (e-) absorbs (gains) energy (in whole photons or “quanta”) it “jumps” to a higher energy level
This is called the EXCITED STATE
When an e- emits (loses) energy it falls to a lower energy level and the energy emission is given off as photons (light)
This is called the GROUND STATE
The return to ground state is what we see as color in the flame test

4. So how was the “color” made in the flame test?
Scientists use the Bohr model to explain this phenomenon
Hydrogen Spectrum
Flame Test

5. There is NO net change in energy
Energy absorbed = energy released = energy of light produced
Sometimes (like the flame test) this light is in the small section of wavelengths called the visible spectrum and we can see it
Most of the time the human eye cannot

6. Bohr’s Hydrogen Model Turn to page 8 in your ref. packet
When an electron falls from n=6 to n=3 what wavelength of light will be emitted?
1094 nanometers
What region of the spectrum does that wavelength correspond to?
Infrared
Would we see it?
Not with our naked eye

7. Hydrogen’s Line Spectrum
Hydrogen emits four visible wavelengths of light
Visible light is emitted when an excited electron “falls” from n= 3, 4, 5, or 6 back to n=2
434 nm; blue
656 nm; red
Faint lines are UV
410 nm; violet
486 nm; cyan

8. Practice What color of light will be emitted if an e- goes from:
n=6 to n=2?
Violet (410nm)
n=5 to n=2?
Blue (434nm)
n= 3 to n=2?
Red/orange (656 nm)

9. Evidence for Energy Levels
Bohr realized that this was the evidence he needed to prove his theory.
The electric charge temporarily excites an electron to a higher orbit. When the electron drops back down, a photon is given off.
The red line is the least energetic and corresponds to an electron dropping from energy level to energy level 2.

10. Electromagnetic Spectrum (EM)
EM is the complete range of electromagnetic radiation
The spectrum includes most types of radiation, most of which are invisible to the human eye.
The visible spectrum is the range of wavelengths between 400 and 700 nm.
Chapter 5

11. Wavelength increases
Frequency increases
Energy increases

12. Anatomy of a Wave
Wavelength ( λ ) – the shortest distance between equivalent points on a continuous wave. Wavelength is measured is units of length - m, mm, µm, nm
Amplitude – the height from the origin to the crest (or trough)
Frequency ( ν ) – the number of waves that pass a given point in one second

13. Wave Nature of Light
Light travels through space as a wave, similar to an ocean wave.
As frequency increases, energy increases (direct relationship)
Wavelength (λ) and frequency (ν) have an inverse relationship…. As λ increases, ν decreases
Short Wavelength = High Frequency = High Energy
Long wavelength = Low Frequency = Low Energy

14. The longer the wavelength of light, the lower the frequency
The longer the wavelength of light, the lower the frequency. The shorter the wavelength of light, the higher the frequency.

15. The Wave/Particle Nature of Light
In 1900, Max Planck proposed that radiant energy is not continuous, but is emitted in small bundles. This is the quantum concept.
Radiant/Light energy has both a wave nature and a particle nature.
An individual unit of light energy is a photon.