1. Properties of Light

2. Light is an electromagnetic wave
EM wave- a form of energy that exhibits wavelike behavior as it travels through space
All the forms of EM radiation form the electromagnetic spectrum

3. Properties of EM waves
Speed- all forms of EM radiation travel at 3 x 108 m/s in a vacuum
Wavelength-the distance between 2 consecutive waves
Frequency- the # of waves that pass a stationary pt. in one second
Amplitude- the height of a wave measured from the origin to its crest

4. Frequency and Wavelength
Are mathematically related
c=λv
Where :
c=speed of light
λ=wavelength
v= frequency

5. Photoelectric Effect
Refers to the emission of electrons from a metal when light of a specific frequency (or energy) shines on the metal
Wave theory predicted that light of any
frequency should be capable of knocking
off electrons. However, it was observed
that the light had to be of a minimum
frequency for the electrons to be
emitted.
In early 1900s, wave theory of light was questioned because it could not explain the interaction of light and matter. Wave theory predicted that light of any energy could cause the emission of electrons-wave theory offered no explanation of why light had to be of a minimum frequency to emit electrons

6. Particle Description of Light
Quantum- the minimum quantity of energy that can be lost or gained by an atom
This relationship is expressed as:
E = hv
Where E = energy (J)
h= Planck’s constant
(6.626 x J.s)
v= frequency of light
Max Planck- proposed the idea that a hot object does not emit energy continuously as it would if energy traveled as a wave-Planck suggested that objects emit energy in small, specific amounts called quanta

7. Example Problems: c =λv E=hv
Determine the energy in joules of a photon whose frequency is 3.55 x 1017 Hz.
The energy for a quantum of light is
2.84 x J. What is the wavelength for this light?
A certain blue light has a frequency of
6.91 x 1014 Hz. What is the wavelength of this light?

8. The Dual Nature of Light
Light exhibits many wavelike properties but can also be thought of as a stream of particles called photons.
Photon- a particle of EM radiation having zero mass and carrying a quantum of energy
Einstein elaborated on Planck’s theory suggesting that EM radiation has a dual wave-particle nature-Einstein’s theory explained the photoelectric effect by proposing that EM radiation is absorbed only in whole #’s-for an electron to be emitted it must be struck by a single photon of minimum energy-if below this minimum the electron remains on the surface

9. Line Emission Spectra
As the excited electron falls back to its ground state, it releases EM radiation of an energy that corresponds to the amount of energy gained to reach the excited state.
The lowest energy state of an electron is the ground state.
A state in which that atom has a higher energy potential is called an excited state.

10. energy of emitted photon = (atom energy before) - (atom energy after)

11. This is referred to as a line emission spectra
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When the EM radiation released is passed through a prism or diffraction grating- it is separated into a series of specific frequencies of visible light.
This is referred to as a line emission spectra

12. Bohr’s Model
The electron can circle the nucleus only in allowed paths or orbits
In an orbit, an electron has a fixed energy
The lowest energy state is closest to the nucleus
An electron can move to a higher orbital if it gains the amount of energy equal to the difference in energy between the initial orbit and the higher energy orbit
Bohr used his planetary model of the atom to explain the line spectrum for the hydrogen atom

14. DeBroglie’s Idea
Suggested that electrons could also have a wave nature much like light.
This was based on the fact that:
Electrons can be diffracted (bending)
Electrons exhibit interference (overlapping that results in a reduction of energy)

15. Heisenberg Uncertainty Principle
It is impossible to determine simultaneously both the position and velocity of an electron.
Electrons are detected by the interaction with photons-because photons have about the same energy as electrons-any attempt to locate an electron knocks it off course.

16. Schrodinger’s Wave Equation
Laid the foundation for modern quantum theory.
Quantum Theory describes mathematically the wave properties of electrons.
The solutions to this equation give only the probability of finding an electron at a given location.

17. Conclusion Electrons do not travel in neat orbits
They exist in three-dimensional regions called orbits that indicate the probable location of an electron.

18. Quantum Numbers
The location of electrons within the atom can be described using quantum numbers:
Principle
Orbital (Angular Momentum)
Magnetic
Spin

19. Principle Quantum Number
Gives the principle energy level
n= 1, 2, 3, etc.
Maximum # of Electrons for Orbitals:
1st- 2
2nd – 8
3rd – 18
4th - 32

20. Orbital Quantum Number
Tells the shape or type of orbital
s orbital is doughnut shaped
p orbital is dumbbell shaped

21. Suborbital
# of electrons
# of orbitals s 2 1 p 6 3 d 10 5 f 14 7

22. Magnetic Quantum Number
Designates specific regions of space w/in each energy sublevel
(s, p, d, f)
Ex.
p sublevel has 3 orbitals
(px, py, pz)

23. Spin Quantum Number Designates direction of electron spin
Electrons w/in an orbital spin in opposite directions

24. Governing Rules and Principles
Pauli Exclusion Principle- only 2 electrons in each orbital
Aufbau Principle- electrons must occupy lower energy orbitals first
Hunds Rule- a second electron can not be added to an orbital until each orbital in a sublevel contains an electron

25. Principle Energy Level
Summary:
Principle Energy Level
Orbitals
Max. Electrons 1 s 2
s p 8 3
s p d 18 4
s p d f 32