1. Chapter 5
Electrons in Atoms

2. Wave Nature of Light
Electromagnetic radiation which is a form of energy that exhibits wavelike behavior as it travels through space.
Examples: light, radio waves, x-rays, etc

3. Parts of a Wave crest wavelength amplitude origin amplitude wavelength
trough

4. Wavelength Waves have a repetitive nature. Wavelength- ( lambda)
shortest distance between corresponding points on adjacent waves.
Measured in units like meters, centimeters, or nanometers depending on the size.
1 x 10-9 meters = 1 nanometer

5. c = lf Frequency # of waves that pass a given point per second.
Units are waves/sec, cycles/sec or Hertz (Hz)
Abbreviated n the Greek letter nu or by an f
c = lf

6. Frequency and wavelength
Are inversely related
As one goes up the other goes down.
High frequency, Short Wavelength
Low frequency, Long Wavelength

7. Speed of light = (wavelength) x (frequency)
Wave Formula
All electromagnetic waves, including visible light, travel at the speed of 3.00 x 10 8 m/s in a vacuum.
Speed of light = c = 3.00 x 108 m/s
c=f
Speed of light = (wavelength) x (frequency)

8. Example Problem
What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz?
Formula: c=f
= ?
f = 3.44 x 109 Hz
c = 3.00 x 108 m/s
3.00 x 108 m/s =  (3.44 x 109 s-1)
3.00E8 / 3.44E9 = 8.72 x 10-2 m

9. Practice
What is the frequency of green light, which has a wavelength of 5.90 x 10-7m?
A popular radio station broadcast with a frequency of 94.7MHz, what is the wavelength of the broadcast? ( frequency needs to be is Hz)

10. Different frequencies produce different types of waves.
The entire range of frequencies is called the electromagnetic spectrum
We are only able to see with our eyes a small portion of the spectrum = visible light
ROY G BIV
Different colors mean different frequencies/wavelengths

11. Energy & The Spectrum
The energy of a wave increases with increasing frequency
High Frequency = High Energy
Low Frequency = Low Energy
Blue light has more energy than Red light

12. Low energy
High energy
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Visible Light
Low Frequency
High Frequency
Long Wavelength
Short Wavelength

13. Quanta Max Planck suggested the idea of quanta or packets of energy.
Quanta is the minimum amount of energy that can be lost or gained by an atom.
Energy is quantized = it comes in packets (like stairs or pennies only whole numbers)

14. Energy = (Planck’s constant)(frequency)
h = x J.s (Joule seconds)
Energy = (Planck’s constant)(frequency)
E = h f
Example: What is the energy in Joules of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 Hz?
E = ?
h = x Js
f = 7.23 x 1014 Hz (or s-1)
E = (6.626 x Js)(7.23 x 1014 s-1)
E = 4.79 x J

15. Photoelectric Effect
In the 1900s, scientist studied interactions of light and matter.
One experiment involved the photoelectric effect, which refers to the emission of electrons from a metal when light shines on the metal.
This involved the frequency of the light. It was found that light was a form of energy that could knock an electron loose from a metal.

16. Photon Light waves can also be thought of as streams of particle.
Einstein called these particles photons (He won a Nobel Prize for this)
A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum energy.

17. Bohr’s Model Why don’t electrons fall into nucleus?
Bohr suggested that they move like planets around sun.
Certain amounts of energy separate one level from another.

18. Nucleus is found inside a blurry “electron cloud”

19. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels

20. } Bohr’s Model Further away from nucleus means more energy. Fifth
There is no “in between” energy
Energy Levels
Fifth
Fourth
Third
Increasing energy
Second
First
Nucleus

21. Bohr Model of the Atom
Ground state- the lowest energy state of an atom.
Excited state – state in which an atom has a higher potential energy than its ground state.
Energy is quantized. It comes in chunks.
quanta - amount of energy needed to move from one energy level to another.
Since energy of an atom is never “in between” there must be a quantum leap in energy.

22. Bohr Energy Levels K = 2 electrons – 1st L = 8 electrons – 2nd
M = 18 electrons – 3rd
N = 32 electrons – 4th

23. Heisenberg Uncertainty Principle
This is the theory that states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

24. Quantum Theory
Schrodinger derived an equation that described energy & position of electrons in atom
Schrodinger along with other scientists laid the foundation for the modern quantum theory, which describes mathematically the wave properties of electrons and other very small particles.