1. Quantum Model

2. Modern View The atom is mostly empty space Two regions Nucleus
protons and neutrons
Electron cloud
region where you might find an electron

3. (positive and negative charges) + + Rutherford (1911) (the nucleus) +
Dalton (1803)
Thomson (1904)
(positive and negative charges) + +
Rutherford (1911)
(the nucleus) + + + + .
Bohr (1913)
(energy levels - orbits) .
Schrödinger (1926)
(electron cloud model – orbitals)
From the time of Dalton to Schrödinger, our model
of the atom has undergone many modifications.

4. Quantum Mechanical Model
Niels Bohr &
Albert Einstein
Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals)

5. Bohr Model
Planetary
model
Niels Bohr
After Rutherford’s discovery, Bohr proposed that electrons travel in definite orbits around the nucleus

6. Bohr’s Work: Quantized Energy
Bohr realized that the atom has various levels of energy (states of energy)
The first level is the ground state
Although below this is the nucleus
The second and following states are “excited”
How does it go from ground state to an excited state?
BY GAINING ENERGY FROM THE UNIVERSE

7. Bohr Model Nucleus Ground State First Excited State
Second Excited State

8. Absorption of Photon by Atom
When a photon of light is absorbed by an atom, it causes an increase in the energy of the atom.
The photon disappears.
The energy of the atom increases by exactly the amount of energy contained in the photon.
The photon can be absorbed ONLY if it can produce an “allowed” energy increase in the atom.

9. Absorption Spectrum
When an atom absorbs photons, it removes the photons from the white light striking the atom, resulting in dark bands in the spectrum
Therefore, a spectrum with dark bands in it is called an absorption spectrum

10. Atomic Spectra
When an atom absorbs energy from a photon of light, it is absorbed in an electron and thus moves further from the nucleus.
Since Bohr explained this principle, the opposite must be also be true and the electron will return to its prior potential energy state when it releases the energy (“loses it”) as light energy.

11. Emission of Photon by Atom
When a photon of light is emitted by an atom, it causes a decrease in the energy of the atom.
A photon of light is created.
The energy of the atom decreases by exactly the amount of energy contained in the photon that is emitted.
The photon can be emitted ONLY if it can produce an “allowed” energy decrease in an excited atom.

12. Emission Spectrum
When an atom emits photons, it glows! The photons cause bright lines of light in a spectrum.
Therefore, a spectrum with bright bands in it is called an emission spectrum.

13. Bohr Model of Atom
Increasing energy
of orbits
n = 3
n = 2
The Bohr model of the atom, like many ideas in the history of science, was at first prompted by and later partially disproved by experimentation.
n = 1 e-
A photon is emitted
with energy E = hf

14. Bohr’s Model
Although his model only worked correctly for the hydrogen atom, it further provided an explanation of some of the chemical properties of the elements
The idea that atoms have electron arrangements unique to each element is the foundation of much of our chemical reactions/bonding knowledge
Electron configurations!!!

15. Atoms and Their Lights
When an atom absorbs energy they are temporarily energized
As it loses this energy, it emits light
Since each element has its own unique atomic structure, each element has its unique COLOR!

16. Bohr’s Contributions:
Electrons can occupy only certain regions of space, called orbits
Orbits closer to the nucleus are more stable — they are at lower energy levels
Electrons can move from one orbit to another by absorbing or emitting energy, giving rise to characteristic spectra
Bohr’s model only worked for the hydrogen atom, so it needed work
His concept of electrons moving in fixed orbits was later abandoned.

17. Bohr Model of the Atom
The planetary model was suggested but this is not plausible because if the electrons orbited the nucleus, then it would collide after ~10-9 s
Niels Bohr joined Rutherford’s team in and took the planetary model from Rutherford and combined it with Einstein’s Theory of Light and Planck’s quantized energy levels

18. Advances in Atomic Theory
Major advances didn’t begin until scientists applied the properties of waves and light to discovering the atom structure
Visible light is a type of electromagnetic radiation

19. Electromagnetism
All of these forms of electromagnetism will travel as waves and move at the speed of light
3.0 x 108 m/s
Represented by the variable c
The significant feature of these, besides their similar speed, is their repetitive nature
They will travel in waves continuously repeating

20. Anatomy of a Wave

21. Wavelength Measurements
Wavelength is measured as the distance between corresponding points of adjacent waves
e.g. crest to crest, trough to trough, etc.
Wavelength is equal to speed divided by frequency
l = v/f
v is velocity/speed thus it is actually c (speed of light)
𝑐=𝜆𝑓
where c always equals 3.0∙ 𝑚/𝑠

22. Photon Energy
If all light moves a wave at the speed of light, then only its wavelength and frequency can vary.
Therefore, the differences in these values are what separate the 7 EM waves on the spectrum
The frequency (albeit the wavelength factors in too) will determine the amount of energy found in the wave
Energy of an EM wave can be calculated by taking the frequency f (in Hertz) and multiplying by a universal constant known as Planck’s Constant h
𝐸=ℎ∙𝑓 or sometimes written as 𝐸=ℎ∙𝑣

23. I. Waves and Particles de Broglie’s Hypothesis
Particles have wave characteristics
Waves have particle characteristics
λ = h/mv
Wave-Particle Duality of Nature

24. Electrons as Waves QUANTIZED WAVELENGTHS Louis de Broglie (1924)
~1924
Louis de Broglie (1924)
Applied wave-particle theory to electrons
electrons exhibit wave properties
QUANTIZED WAVELENGTHS
Fundamental mode
Second Harmonic or First Overtone
Standing Wave
200
150
100 50
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-100
-150
-200
200
150
100 50
- 50
-100
-150
-200
200
150
100 50
- 50
-100
-150
-200

25. Dual Nature of Light Three ways to tell a wave from a particle…
Waves can bend
around small obstacles…
…and fan out
from pinholes.
Particles effuse from pinholes
Three ways to tell a wave from a particle…
wave behavior particle behavior
waves interfere particle collide
waves diffract particles effuse
waves are delocalized particles are localized

26. Werner Heisenberg – The Uncertainty Principle
Think of this…
a free electron moves into the focus of a hypothetical microscope and is struck by a photon of light; the photon transfers momentum to the electron
The reflected photon is seen in the microscope, but the electron has moved out of focus
The electron is not where it appears to be
Think of this like seeing stars in space that may not be there anymore because the light is still being seen due to great distances
A wave is a disturbance that travels in space and has no fixed position

27. Quantum Mechanics Heisenberg Uncertainty Principle g
Werner Heisenberg
~1926
Heisenberg Uncertainty Principle
Impossible to know both velocity and position of an electron at the same time g
Microscope
Electron

28. New Atomic Model
The Heisenberg Uncertainty Principle states in simple terms that it is impossible to describe precisely both the location and the speed of particles that exhibit wavelike behavior
The upshot of all this is that we cannot state with absolute certainty the velocity and position of electrons and so we must replace the Bohr model with another which considers the probability of an electron being at a certain point
In effect, all we can say is that we are pretty certain that the electron is within a particular region for some (most) of the time
Of course outside that time, it could be anywhere

29. II. The Electron as a Wave
Schrödinger’s Wave Equation
Used to determine the probability of finding the single electron in a hydrogen atom at any given distance from the nucleus
Electron best described as a cloud
Effectively covers all points at the same time
(fan blades)

30. Quantum Mechanics Schrödinger Wave Equation (1926)
Erwin Schrödinger
~1926
Schrödinger Wave Equation (1926)
finite # of solutions  quantized energy levels
defines probability of finding an electron

32. Quantum Mechanics Orbital (“electron cloud”)
Region in space where there is 90% probability of finding an electron
90% probability of
finding the electron
Orbital
Electron Probability vs. Distance 40 30
Electron Probability (%) 20 10 50
100
150
200
250
Distance from the Nucleus (pm)

33. Quantum Numbers Four Quantum Numbers:
Specify the “address” of each electron in an atom
UPPER LEVEL

34. Writing & Reading Quantum Numbers
Since this is the specific “address” for each an every electron found in an atom,
This will be done shortly after learning about electron configurations

35. Quantum Model of the Atom
Electron Configuration & Orbitals
1s22s22p63s23p64s23d104p65s24d104p65s24d105p66s24f145d106p6…

36. Ground-State Electron Configurations
The arrangement of electrons in an atom is called the electron configuration
Because atoms of different elements have different number of electrons, each atom has its own unique configuration
Since low-energy systems are more stable than high-energy systems, electrons in atoms tend to assume arrangements that give the lowest possible energy
This is known as Aufbau’s Principle
Helium with 2 valence e- in the s orbital

37. Order of Electron Subshell Filling: It does not go “in order” of reading a book
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
7s2
7p6
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
5p6
6s2
4f14
5d10
6p6
7s2
5f14
6d10
7p6

38. s p d f 1 2 3 4 5 6 7 1A 2A 3B 4B 5B 6B 7B 8B 1B 2B 3A 4A 5A 6A 7A 8A
Group Number = # valence e- s 1 2 3 4 5 6 7 p
Row =
# shells d f

39. Electron Configurations
Orbital Energy Diagram
Reading the Diagram
The energy levels are approximate
Meaning that sometimes the higher sub-orbital will fill despite the lower orbital not completely full
Beginning in the third energy level (third shell), the sublevels begin to overlap
This is because it’s easier to put two electrons together in 4s than 10 in 3d

40. Electron Configurations
There are many different ways to show an atom’s electron configuration:
Orbital Diagrams (Arrows)
Electron Configuration Notation (1s22s2…)
A shortcut has been used to make notations faster too
Noble Gas Configurations
Lewis Dot
There are some exceptions, but we will ignore those rules for now
n = # of orbit (or layer of electron cloud)
n2 = # of orbitals per orbit (boxes that hold 2 electrons each)
2n2 = # of electrons per orbit (layer)

41. Pauli and Hund Pauli Exclusion Principle Hund’s Rule
Electrons in an atom have a specific spin, just like a top spins on its axis
The electron can only spin two ways, these are represented by an up arrow (↑) and a down arrow (↓)
Only two electrons can fill a single atomic orbital, but only if they are spinning in opposite directions
Hund’s Rule
***Electrons are negative and therefore, must repel one another
Hund’s Rule states that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can be paired

42. Electron Configuration
1s1
# valence e-
possibilities are:
s: 1 or 2
p: 1-6
d: 1-10
f: 1-14
Total e- should equal
Atomic # (not for ions)
row #
shell #
possibilities are 1-7
7 rows
subshell
possibilities are
s, p, d, or f
4 subshells
What element has an electron configuration of 1s1?

43. Practice
Write the standard electron configuration for a neutral phosphorus-32 atom
1s22s22p63s23p3

44. 1s22s22p63s23p64s23d104p65s24d7 [Kr]5s24d7 Practice
Write the standard electron configuration for a neutral rhodium-102 atom
Well that takes too long, what about the shortcut…
Look at the row or period that it belongs to and count down from the start
This atom is in the 5th period, so start the electron configuration on the s-orbital of that period with the most recent Noble Gas
1s22s22p63s23p64s23d104p65s24d7
[Kr]5s24d7

45. 1s22s22p63s23p64s23d1 [Ar]4s23d1 Practice
Write the standard electron configuration and the Noble Gas electron configuration for Scandium-45
Noble Gas configurations will always be [Noble Gas] followed by row # “s”
1s22s22p63s23p64s23d1
[Ar]4s23d1

46. Practice
Write the standard electron configuration for the anion O 2−
1s22s22p6

47. Valence Shells
The electrons of the outermost occupied shell in any atom are directly exposed to the external environment and are the first to interact with other atoms.
The electrons in the outermost shell, therefore, are quite important
They are called valence electrons (valent – combining power of an atom)

48. Valence Electrons Again
Since protons determine the type of atom/matter and electrons determine the reactivity of the atoms…
Valence electrons determine the chemical properties of an element
(e- in outermost valence shell)
As we will cover in the not too distant future, these valence electrons determine the chemical bonds that will form between atoms

49. Valence e- – Electron-Dot Structures
The element’s symbol represents the nucleus and inner valence shell electrons
The dots are the valence electrons and must be placed one at a time on the four sides (any sequence) and then paired
You only represent the s & p orbital electrons!
This method was designed by G.N. Lewis ( ) and thus is also called Lewis Dot Structures

50. Apparent Valence Electrons
What if the atom is from the d or f orbital?
As noted earlier, it is sometimes easier to fill shells partially before completely for equal stability
You do not need to memorize this an any way Sc Ti V Cr Mn Fe Co Ni Cu Zn
Outer Configuration
4s23d1
4s23d2
4s23d3
4s13d5
4s23d5
4s23d6
4s23d7
4s23d8
3d104s1
3d104s2
Apparent Valence Electrons 3
2-4
2-5
2-6
2-7
2 or 3
1 or 2 2

51. Feeling Overwhelmed? Read the Book!
Chemistry
"Teacher, may I be excused? My brain is full."