1. Properties of Light

2. Light is an electromagnetic wave
Electromagnetic wave- a form of energy that exhibits wavelike behavior as it travels through space

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

4. Relationship between frequency and wavelength?
Shorter wavelength = ________ frequency
Longer wavelength = ________ frequency

5. Frequency and Wavelength…
…are mathematically related
c=λv
Where :
c = speed of light
λ = wavelength
v = frequency

6. Photoelectric Effect
Refers to the emission of electrons from a metal when light of a specific frequency (or energy) shines on the metal
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
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.

7. 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

8. Particle Description of Light
Quantum- “a single energy packet of light” -the minimum quantity of energy that can be lost or gained by an atom (a photon is a single quantum of light)
This relationship is expressed as:
E = hv
Where E = energy
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

9. 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)

10. Line Emission Spectra
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.
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.

11. Ground State vs. Excited State
As an atom is hit by energy, its electrons jump to a higher energy level (excited state). As the electrons fall back down, they release energy in the form of light!

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

13. This is referred to as a line emission spectra
hyperlink
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

14. 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

15. Click below for a useful Youtube video!
Line Emission Spectra

16. 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.

17. 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.

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

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

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

21. Orbital Quantum Number (“ℓ”)
Tells the shape or type of orbital
s orbital is a spherical shape
p orbital is dumbbell shaped

22. Orbital Quantum Numbers “ℓ”
Corresponding Orbital
Shape of those orbitals s
Spherical shaped 1 p
Dumbbell shaped 2 d
Cloverleaf shaped mostly 3 f
Complex

23. D Orbital Orientations
F Orbital Orientations

24. Magnetic Quantum Number “m ℓ”
Designates specific regions of space within each energy level
(s, p, d, f)
S has 1 suborbital
P has 3 suborbitals
D has 5 suborbitals
F has 7 suborbitals
Represent them as: p orbital has 3 suborbitals
(px, py, pz)

25. Magnetic Quantum Number “m”
S has 1 suborbital
P has 3 suborbitals (“orientations”  px, py, pz)
D has 5 suborbitals
F has 7 suborbitals
***Each suborbital can hold 2 electrons
Ex: S orbital: ____
P orbital: ____ _____ ____

26. Magnetic Quantum Number “m”
S has 1 suborbital (“orientation”)
P has 3 suborbitals
D has 5 suborbitals
F has 7 suborbitals
***Each suborbital can hold 2 electrons
D orbital: __ __ __ __ __
F orbital: __ __ __ __ __ __ __

27. Magnetic Quantum Number
Orbital
# of suborbitals
# of total electrons s 1 2 p 3 6 d 5 10 f 7 14

28. Spin Quantum Number Designates direction of electron spin
Electrons within an orbital spin in opposite directions
S sublevel: _____
***opposite arrows symbolize opposite spin

29. 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

30. D sublevel: __ __ __ __ __
Hunds Rule
Hunds Rule- a second electron can not be added to an orbital until each orbital in a sublevel contains an electron
D sublevel: __ __ __ __ __

31. 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

32. Electron Configuration
How do we actually write where the electrons are located for each atom??

33. Electron Configurations
Electron configurations tells us in which orbitals the electrons for an element are located.
Three rules: (Review)
electrons fill orbitals starting with lowest n and moving upwards;
no two electrons can fill one orbital with the same spin (Pauli);
for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule).

34. Filling Diagram for Sublevels
Aufbau Principle
Copy this down in your notebook!!! With the arrows included…

35. Writing Electron Configurations “Superscript Notation”
First, determine how many electrons are in the atom. Iron has 26 electrons.
Arrange the energy sublevels according to increasing energy:
1s 2s 2p 3s 3p 4s 3d …
Fill each sublevel with electrons until you have used all the electrons in the atom:
Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6
The sum of the superscripts equals the atomic number of iron (26)

36. Write the Superscript Notation…
Nitrogen (N)
Silicon (Si)
Chlorine (Cl)
Sodium (Na)
Beryllium
Calcium
Carbon
Neon

37. Orbital Notation
Electron configuration with showing the electrons actually in their occupied orbitals.
Ex: Oxygen protons, 8 electrons
Start by using your filling diagram to know the order that electrons occupy orbitals. (1s, 2s, 2p, 3s, 3p…)
Write your superscript notation: 1s2, 2s2, 2p4
Now draw out the orbitals. Place electrons in.
____ _____ _____ _____ _____
1s s px 2py pz

38. Orbital Notation Ex: Magnesium 12 protons, 12 electrons
Use your filling diagram to know the order that electrons occupy orbitals. (1s, 2s, 2p, 3s, 3p…)
Superscript notation: 1s2, 2s2, 2p6, 3s2
Now draw out the orbitals. Place electrons in.
____ _____ _____ _____ _____ _____
1s s px 2py pz s

39. Orbital Notation Ex: Chlorine 17 protons, 17 electrons
Use your filling diagram to know the order that electrons occupy orbitals. (1s, 2s, 2p, 3s, 3p…)
Superscript notation: 1s2, 2s2, 2p6, 3s2, 3p5
Now draw out the orbitals. Place electrons in.
____ _____ _____ _____ _____ _____
1s s px 2py pz s
_____ _____ _____
3px 3py pz

40. Blocks and Sublevels
We can use the periodic table to predict which sublevel is being filled by a particular element.

42. Noble Gas Core Electron Configurations
Recall, the electron configuration for Na is:
Na: 1s2 2s2 2p6 3s1
We can abbreviate the electron configuration by indicating the innermost electrons with the symbol of the preceding noble gas.
The preceding noble gas with an atomic number less than sodium is neon, Ne. We rewrite the electron configuration:
Na: [Ne] 3s1

43. Let’s Look at it backwards
Let’s Look at it backwards! ID the elements with the following configurations…
1s22s22p63s23p4
1s22s22p63s23p64s23d104p65s1 
[Kr] 5s24d105p3
[Xe] 6s24f145d6
[Rn] 7s25f11
Sulfur
Rubidium
Antimony
Osmium
Einsteinium