1. Atomic Structure and Periodicity Part 1
Chapter 7

2. Warm-up Question
Be prepared to share out your response to the following questions. What is a photon? What is the source of electromagnetic waves? Is the color spectrum simply a small segment of the electromagnetic spectrum? Defend your answer.

3. Electromagnetic Radiation

4. Warm-up Continued What is a photon?
A particle of light.
Particle vs Wave Theory
Video 1
Video 2
What is the source of electromagnetic waves?
Accelerating electric charges
Is the color spectrum simply a small segment of the electromagnetic spectrum? Defend your answer.
Yes; the spectrum is also made up of radio waves, IR, UV, X-rays, and gamma rays.

5. CHARACTERISTICS OF WAVES
Waves are described according to their
Amplitude measures DISPLACEMENT size of the disturbance (from rest to crest)
Wavelength  distance of a “repeating unit” Also called a cycle
Velocity v speed = how fast wave travels

6. Frequency v
How often number of wavelengths that pass any point per second
measured in wavelengths/second or cycles/second Hertz (Hz) = number of wavelengths in 1 second
Frequency is related to velocity: c = v 

7. Electromagnetic Wave
a transverse wave with an electric component and a magnetic component at right angles to each other
How are electromagnetic waves (ex: light) different from mechanical waves (ex: sound and slinky)?
micro.magnet.fsu.edu

8. Electromagnetic Waves
Electromagnetic waves are special in the fact that they do not need a medium to propagate through. But what is creating the disturbance? What is emitting this energy? ELECTRONS
© 2003 Mike Maloney

9. Electromagnetic Waves
Electrons in materials are vibrated and emit energy in the form of photons, which propagate across the universe. Photons have no mass, but are pure energy. Electromagnetic Waves are waves that are made up of these “photons”. When these photons come in contact with boundaries, E-M waves interact like other waves would.
© 2003 Mike Maloney

10. Electromagnetic Spectrum
© 2003 Mike Maloney

11. Speed of E/M Waves
It has been found that the speed of E-M waves and light is ---
3 x 108 or 300,000,000 m/s
671,000,000 mph
186,000 miles per second
We call this value “c”
© 2003 Mike Maloney

12. C is constant throughout the universe, as long as light is in a vacuum
C is constant throughout the universe, as long as light is in a vacuum. When it is in other materials, c can change, but can never be larger than its value in a vacuum. Since “c” is constant, all of E-M waves will have a corresponding frequency to go along with their wavelength.
© 2003 Mike Maloney

13. Energy in E-M Waves
Which waves have more energy, Radio waves or gamma waves? The greater the frequency of an E-M wave, the more crests pass a point in a certain amount of time, therefore the more photons pass that point. This means that more energy moves past that point in a certain amount of time or that the wave contains more energy.
© 2003 Mike Maloney

14. Electromagnetic Spectrum “Check-up”
True or False…
Blue light has a shorter wavelength than red light.
X-rays have lower frequencies than radio waves.
Microwaves have higher frequencies than gamma rays.
Visible radiation composes the major portion of the electromagnetic spectrum.
True; False; False; False

15. Wavelength-frequency relationship example
Photosynthesis uses light with a frequency of 4.54x1014s-1. What wavelength does this correspond to?
A: 660nm

16. Wavelength-frequency relationship Practice
Calculate the frequency of blue light of wavelength 4.5 x 102nm. Calculate the wavelength of green light of frequency 5.7 x 1014Hz. A:6.7x1014Hz ; 5.3 x 10-7m or 530nm

17. The Nature of Matter ΔE = hv = hc/λ
ΔE is the change in energy for a system (in Joules per photon)
h is Planck’s constant (6.626 x 10-34J s)
experimentally determined
v is the frequency of the wave (s-1 or Hz)
**Energy can be gained or lost only in integer multiples of hv. (quanta)

18. Energy, Frequency, wavelength Example
Sodium atoms have a characteristic yellow color when excited in a flame. The color comes from the emission of 589.0nm.
What is the frequency of this radiation?
What is the change in energy associated with this photon? Per mole of photons?

19. Energy, Frequency, wavelength Practice
It takes 382 kJ of energy to remove one mole of electrons from gaseous cesium. What is the wavelength associated with this energy? Would we be able to “directly” observe this energy change? Why or why not.

20. The photoelectric Effect
Emission of electrons from a metal when light shines on the metal Electromagnetic radiation (light) strikes the surface of the metal ejecting electrons from the metal and causing an electric current, if the frequency was below a certain minimum. Analysis of the kinetic energy and numbers of the emitted electrons led Einstein to suggest that electromagnetic radiation can be viewed as a stream of photons. *Note that the apparent mass of a photon depends on its wavelength. The mass of a photon at rest is thought to be zero, although we never observe it at rest.*

21. Big ideas from Einstein and Planck
Energy is quantized. It can occur only in discrete units called quanta.
Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well. (dual nature of light)

22. Wave-like behavior Diffraction
Light is scattered from a regular array of points or lines.
Constructive interference
In-phase (bright)
Destructive interference
Out-of phase (dim/dark)

23. Atomic Spectrum of Hydrogen
Continuous Spectrum
Contains all the wavelengths over which the spectrum is continuous
Line Spectrum
Contains certain specific wavelengths that are characteristic of the substance emitting those wavelengths
*Hydrogen’s line spectrum shoes that only certain energy transfers are allowed in hydrogen.
*Specific energy levels among which the hydrogen electron can shift, thus energy levels are quantized.

25. Individual Practice
15, 31, 33, 35, 39

26. The Bohr Model
1913 Niels Bohr developed the Quantum Model for the hydrogen atom.
The electron in the hydrogen atom moves around the nucleus only in certain allowed circular orbits.
Hydrogen atom energy levels consistent with the hydrogen emission spectrum. (different wavelength/color associated with the different levels of emission)
Ground state
The lowest possible energy state of an atom or molecule
Excited state
Higher potential energy state than ground state of an atom or molecule

27. Although Bohr’s model fits the energy levels for hydrogen, it is a fundamentally incorrect model for the hydrogen atom. Bohr’s model paved the way for later theories on the quantization of energy in atoms. Electrons do NOT move around the nucleus in circular orbits (planetary model).

28. Quantum Mechanics
de Broglie and Schrodinger – wavelike properties of electrons
A specific wave function (function of the coordinates x, y, and z of the electron’s position in 3-D space) is often called an orbital.
The wave function corresponding to the lowest energy for the hydrogen atom is called the 1s orbital (no association to the Bohr “orbit”).
Nature of an orbital takes into consideration the work of Heisenberg.
Heisenberg uncertainty principle: There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time.

29. 1s orbital
The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron probability. (90% of the time the electron I in this sphere)

30. Summary
In the quantum (wave) mechanical model, the electron is viewed as a standing wave. This representation leads to a series of wave functions (orbitals) that describe the possible energies and spatial distributions available to the electron. In agreement with the Heisenberg uncertainty principle, the model cannot specify the detailed electron motions. Instead, the square of the wave function represents the probability distribution of the electron in that orbital. This allows us to picture orbitals in terms of probability distributions, or electron density maps. The size of an orbital is arbitrarily defined as the surface that contains 90% of the total electron probability. The hydrogen atom has many types of orbitals. In the ground state, the single electrons resides in the 1s orbital. The electron can be excited to higher-energy orbitals if energy is put into the atom.

31. Quantum Mechanics http://www. meta-synthesis
Better than any previous model, quantum mechanics does explain how the atom behaves. Quantum mechanics treats electrons not as particles, but more as waves (like light waves) which can gain or lose energy. But they can’t gain or lose just any amount of energy. They gain or lose a “quantum” of energy.
A quantum is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level.
In this case it is losing the energy and dropping a level.

32. Atomic Orbitals http://milesmathis.com/bohr2.jpg
Much like the Bohr model, the energy levels in quantum mechanics describe locations where you are likely to find an electron. Remember that orbitals are “geometric shapes” around the nucleus where electrons are found. Quantum mechanics calculates the probabilities where you are “likely” to find electrons.

33. Atomic Orbitals http://courses. chem. psu. edu/chem210/quantum/quantum
Of course, you could find an electron anywhere if you looked hard enough. So scientists agreed to limit these calculations to locations where there was at least a 90% chance of finding an electron. Think of orbitals as sort of a "border” for spaces around the nucleus inside which electrons are allowed. No more than 2 electrons can ever be in 1 orbital. The orbital just defines an “area” where you can find an electron. What is the chance of finding an electron in the nucleus? Yes, of course, it’s zero. There are not any electrons in the nucleus.

34. Energy Levels http://www.chem4kids.com/files/art/elem_pertable2.gif
Quantum mechanics has a principal quantum number. It is represented by a little n. It represents the “energy level” similar to Bohr’s model.
n=1 describes the first energy level
n=2 describes the second energy level
Etc.
Each energy level represents a period or row on the periodic table. It’s amazing how all this stuff just “fits” together.
Red n = 1
Orange n = 2
Yellow n = 3
Green n = 4
Blue n = 5
Indigo n = 6
Violet n = 7

35. Sub-levels = Specific Atomic Orbitals
Each energy level has 1 or more “sub-levels” which describe the specific “atomic orbitals” for that level.
n = 1 has 1 sub-level (the “s” orbital)
n = 2 has 2 sub-levels (“s” and “p”)
n = 3 has 3 sub-levels (“s”, “p” and “d”)
n = 4 has 4 sub-levels (“s”, “p”, “d” and “f”)
There are 4 types of atomic orbitals:
s, p, d and f
Each of these sub-levels represent the blocks on the periodic table.
Blue = s block
Yellow = p block
Red = d block
Green = f block

36. Orbitals http://media-2. web. britannica
In the s block, electrons are going into s orbitals. In the p block, the s orbitals are full. New electrons are going into the p orbitals. In the d block, the s and p orbitals are full. New electrons are going into the d orbitals. What about the f block?

37. Property of the Orbital
Quantum Numbers
Describe the properties of the orbital.
Name
Symbol
Property of the Orbital
Range of Values
Principal Quantum Number n
Related to size and energy of the orbital
Integers
1 to ∞
Angular Momentum Quantum Number l
Related to the shape of the orbital
“subshell”
0 is s; 1 is p; 2 is d; 3 is f; 4 is g; 5 is h
Integers from
n-1 to 0
Magnetic Quantum Number ml
Related to the position of the orbital in space relative to other orbitals
-l to 0 to +l

38. degenrate
All orbitals having the same value of “n” have the same energy. 3s; 3p; 3d Energy is required to transfer an electron to a higher-energy orbital (excited state). **In polyelectronic atoms we find that the s, p, and d have different levels of potential energy.

39. The 4th quantum number The electron spin quantum number. Electron spin
Two spin states + ½ and – ½
Produce two oppositely directed magnetic moments
Pauli Exclusion Principle
In a given atom no two electrons can have the same set of four quantum numbers (n, l, ml, ms)
Thus, an orbital can hold only TWO electrons, and they must have opposite spins.

40. Practice with Quantum numbers
Which of the following quantum numbers are allowed? For each that is incorrect state why.
Principal, Angular Momentum, Magnetic Quantum Numbers (n, l, ml)
1, 0, 1
2, 2, 1
5, 3, 2
6, -2, 2
6, 2, -2

41. Quantum numbers and levels of orbitals
Table 7.2 on page 294 in text

42. Energy Level
Sub-levels
Total Orbitals
Total Electrons
Total Electrons per Level
n = 1 s
1 (1s orbital) 2
n = 2 p
1 (2s orbital)
3 (2p orbitals) 6 8
n = 3 d
1 (3s orbital)
3 (3p orbitals)
5 (3d orbitals) 10 18
n = 4 f
1 (4s orbital)
3 (4p orbitals)
5 (4d orbitals)
7 (4f orbitals) 14 32
Complete the chart in your notes as we discuss this. The first level (n=1) has an s orbital. It has only 1. There are no other orbitals in the first energy level. We call this orbital the 1s orbital.

43. Where are these Orbitals. http://www. biosulf2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 5d 6p 7s 6d 7p 4f 5f

44. Electron Configurations
What do I mean by “electron configuration?” The electron configuration is the specific way in which the atomic orbitals are filled. Think of it as being similar to your address. The electron configuration tells me where all the electrons “live.”

45. Rules for Electon Configurations https://teach. lanecc
In order to write an electron configuration, we need to know the RULES.
3 rules govern electron configurations.
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Using the orbital filling diagram at the right will help you figure out HOW to write them
Start with the 1s orbital. Fill each orbital completely and then go to the next one, until all of the elements have been acounted for.

46. Fill Lower Energy Orbitals FIRST http://www. meta-synthesis
Each line represents an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
The Aufbau Principle states that electrons enter the lowest energy orbitals first.
The lower the principal quantum number (n) the lower the energy.
Within an energy level, s orbitals are the lowest energy, followed by p, d and then f. F orbitals are the highest energy for that level.
High Energy
Low Energy

47. No more than 2 Electrons in Any Orbital…ever. http://www. fnal
The next rule is the Pauli Exclusion Principal.
The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full.
The spins have to be paired.
We usually represent this with an up arrow and a down arrow.
Since there is only 1 s orbital per energy level, only 2 electrons fill that orbital.
Wolfgang Pauli, yet another German Nobel Prize winner
Quantum numbers describe an electrons position, and no 2 electrons can have the exact same quantum numbers. Because of that, electrons must have opposite spins from each other in order to “share” the same orbital.

48. Hund’s Rule http://intro. chem. okstate
Hunds Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons.
What are degenerate orbitals?
Degenerate means they have the same energy.
So, the 3 p orbitals on each level are degenerate, because they all have the same energy.
Similarly, the d and f orbitals are degenerate too.
Don’t pair up the 2p electrons until all 3 orbitals are half full.

49. Paramagnetic
Diamagnetic
unpaired electrons
all electrons paired 2p 2p

50. Application
NOW that we know the rules, we can try to write some electron configurations. Remember to use your orbital filling guide/PERIODIC TABLE to determine WHICH orbital comes next. Lets write some electron configurations for the first few elements, and let’s start with hydrogen. H; Li; B; N; F; Na; K; Fe

51. Electron Configurations
Element
Configuration
H Z=1
1s1
He Z=2
1s2
Li Z=3
1s22s1
Be Z=4
1s22s2
B Z=5
1s22s22p1
C Z=6
1s22s22p2
N Z=7
1s22s22p3
O Z=8
1s22s22p4
F Z=9
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K Z=19
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
Fe Z=26
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), = 10

52. Conceptual Check
One last thing. Look at the previous slide and look at just hydrogen, lithium, sodium and potassium. Notice their electron configurations. Do you see any similarities? Since H and Li and Na and K are all in Group 1A, they all have a similar ending. (s1)

53. Electron Configurations
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1
This similar configuration causes them to behave the same chemically.
It’s for that reason they are in the same family or group on the periodic table.
Each group will have the same ending configuration, in this case something that ends in s1.

54. Noble gas notation… “short cut”
Be Al Br Mo Ag

55. Orbital notation
Be Al N Br Mo Ag

56. Ion electron configuration
Be2+ Al3+ Br- Ag1+

57. Individual Practice Bond Energies- page 384 (on Unit 3 Test!)
# 53 and 54
Periodicity and Atomic Structure-starting on page 322
# 57, 58, 59, 60, 62, 67, 69, 70, 71, 72, 73, 74, 85, 87, 89, 95, 97, and 123
(Please note that there are multiple questions over the same concept(s). You do not need to do them all but need to KNOW how to do them.)