1. Chapter 13 “Electrons in Atoms”
Mr. Daniel Olympic High School
Credits: Stephen L. Cotton
Charles Page High School

2. Section 13.3 Physics and the Quantum Mechanical Model
OBJECTIVES:
Describe the relationship between the wavelength and frequency
Distinguish between quantum mechanics and classical mechanics. of light.
Identify the source of atomic emission spectra.
Explain how the frequencies of emitted light are related to changes in electron energies.

3. Light Visible light is a type of electromagnetic radiation.
Electromagnetic radiation is a form of energy and includes many types: gamma rays, x-rays, radio waves, visible light…
Speed of light (c) = 3.00 x 108 m/s
All electromagnetic radiation travels at this same rate when measured in a vacuum

4. Parts of a Wave
Crest
Wavelength
Amplitude
Node
Trough

5. - Page 373
“R O Y G B I V”
Frequency Increases
Wavelength Shorter

6. Long Wavelength Low Frequency Low ENERGY Short Wavelength
Wavelength Table
Short
Wavelength
High Frequency
High ENERGY

7. Wavelength and Frequency:
Are inversely related
As one increases the other decreases.
Different frequencies of visible light are different colors.
There is a wide variety of frequencies
Spectrum: A whole range of electromagnetic wavelengths. (e.g. the visible light spectrum)

8. The Electromagnetic Spectrum
Low Energy
High Energy
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Long Wavelength
Short Wavelength
Low Frequency
High Frequency
Visible Light

9. So what is Energy? Not all quanta (plural) are the same size.
All energy is quantized
A quantum is a “packet” of energy.
Not all quanta (plural) are the same size.
(eggs are not all the same size either, but all are eggs)

10. So what is Light Energy?
Light is a form of energy.
Therefore, light must be quantized
A quantum of light energy is called a photon.
Einstein determined that light is not only a wave, but is also a particle!
He demonstrated it in an experiment that showed the photoelectric effect

11. Photoelectric Effect

12. Experiment demonstrates the particle nature of light.
Photoelectric Effect
Experiment demonstrates the particle nature of light.

13. So what is Light Energy? (con’t)
Therefore, light has what is called wave-particle duality. It has characteristics of both waves and particles.

14. Wave-Particle Duality (again)
J.J. Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!

15. Confused? You’ve Got Company!
“No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934

16. The Physics of the Very Small
Quantum mechanics explains how very small particles behave
Quantum mechanics is an explanation for subatomic particles and atoms as waves
Classical mechanics describes the motions of bodies much larger than atoms

17. Section 13.1 Models of the Atom
OBJECTIVES:
Identify the inadequacies in the Rutherford atomic model.
Identify the new proposal in the Bohr model of the atom.
Describe the energies and positions of electrons according to the quantum mechanical model.
Describe how the shapes of orbitals related to different sublevels differ.

18. Ernest Rutherford’s Model
Discovered dense positive piece at the center of the atom- “nucleus”
Electrons would surround and move around the nucleus
Atom is mostly empty space
It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

19. Niels Bohr’s Model
Why don’t the electrons fall into the nucleus?
He agreed with Rutherford that electrons move around the nucleus, But:
In specific circular paths, or orbits (like planets around the sun), at specific energy levels.
An amount of fixed energy separates one electron energy level from another.

20. The Bohr Model of the Atom
I pictured the electrons orbiting the nucleus much like planets orbiting the sun.
However, electrons are found in specific energy levels around the nucleus, and can jump from one level to another.
Niels Bohr

21. Bohr’s Model Electrons occupy specific energy levels
analogous to the rungs of a ladder
The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder
A quantum of energy is the amount of energy required to move an electron from one energy level to another (plural: quanta)
Since the energy of an atom is never “in between” there must be a quantum leap in energy.

22. Changing the energy
Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

23. Changing the energy
Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be in an “excited state”

24. Changing the energy
The electron is unstable at the higher energy level and as it falls back to the ground state, it gives the energy back in the form of light

25. Changing the energy They fall down in specific steps
Each step has a different energy (quantum) and results in a different color of light.

26. Origin of Line Spectra
Balmer series

27. Ultraviolet Visible Infrared
The further electrons fall, the more energy is released and the higher the frequency of light emitted.
This is a simplified explanation!
Remember, the orbitals also have different sublevels within the principle energy levels

28. Atomic Spectra
White light is made up of all the colors of the visible spectrum.
Passing it through a prism separates it.

29. But not all light is white..
By heating a gas with electricity we can get it to give off colors.
Passing this light through a prism does something different.

30. Atomic Spectrum
Each element gives off its own characteristic colors of light.
The colors can be used to identify the atom.
This is how we know what elements stars are made of.

31. This is called a bright line spectrum
Unique to each element, like fingerprints!
Very useful for identifying elements

32. Explanation of Atomic Spectra
ground state - the lowest energy level of the electron.
In summary. When an electron at ground state receives a quantum of energy it jumps directly to a higher energy level. The electron is unstable and immediately drops to a lower energy level. As it drops it gives off the same amount of

33. The Quantum Mechanical Model
Problems with Bohr’s theory :
It was only successful for H- no other elements followed his predictions.
It introduced the quantum idea artificially.

34. Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.”
You can find out where the electron is, but not its energy
OR…
You can know how much energy it has, but not where it is!
Werner Heisenberg

35. Heisenberg Uncertainty Principle
It is impossible to know exactly the location and velocity of a particle simultaneously.
The better we know one, the less we know the other.
Measuring one property, changes the other.

36. After Before Photon wavelength changes Photon Moving Electron
Electron velocity changes

37. In 1926, Erwin Schrodinger derived an equation that described the energy and probable position of the electrons in an atom

38. Schrodinger’s Wave Equation
His equation determined the probability of finding a single electron along a single axis (x-axis)
Erwin Schrodinger
Erwin Schrodinger

39. The Quantum Mechanical Model
Particles which are very small and travel very quickly (like electrons) behave very differently from objects big enough to observe.
The quantum mechanical model is a mathematical solution describing how those particles act.

40. The Quantum Mechanical Model
Describes energy levels for electrons.
Electrons move in an unpredictable manner
We can only determine the probability of finding an electron a certain distance from the nucleus.

41. The Quantum Mechanical Model
The electrons are probably located inside a blurry “electron cloud”
The area where there is the greatest chance of finding an electron.

42. Atomic Orbitals
Principal Quantum Number (n) = the energy level of the electron: 1, 2, 3, etc.
Within each energy level, there are sub-levels (like theater seats arranged in sections): letters s, p, d, and f
The complex math of Schrodinger’s equation describes several shapes
These are the atomic orbitals - regions where there is a 90% probability of finding an electron.

43. Principal Quantum Number
The Principle Quantum Number (n) denotes the shell (energy level) in which the electron is located.
The maximum number of electrons that fit into an energy level can be calculated:
2n2

44. Summary # of orbitals Max. electrons Starts at energy level s 1 2 1 p3 6 2 d 5 10 3 7 14 4 f

45. Types of Orbitals
s orbital
p orbital
d orbital

46. Types of Atomic Orbitals

47. By Energy Level First Energy Level Has only an s sublevel
only 2 electrons
1s2
Second Energy Level
Has s and p sublevels
2 e- in s, 6 e- in p
2s22p6
8 total electrons

48. By Energy Level Third energy level Has s, p, and d sublevels
2 e- in s, 6 e- in p, and 10 e- in d
3s23p63d10
18 total electrons
Fourth energy level
Has s, p, d, and f sublevels
2 e- in s, 6 e- in p, 10 e- in d, and 14 e- in f
4s24p64d104f14
32 total electrons

49. By Energy Level
Beyond the fourth energy level, not all sublevels fill up.
You simply run out of electrons
So only the s, p, d and f sublevels are used
Because the energy levels overlap the orbitals do not fill up in a consistent pattern
However, the lowest energy orbitals fill first.

50. Section 13.2 Electron Arrangement in Atoms
OBJECTIVES:
Describe how to write the electron configuration for an atom.
Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

51. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
aufbau diagram

52. Electron Configurations…
…are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how:
Aufbau principle - electrons enter the lowest energy sublevels first.
This becomes complex because of the overlap of orbitals of different energies – follow the diagram!
Pauli Exclusion Principle – there are at most 2 electrons per orbital - with opposite spins

53. Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
To show the different direction of spin, a pair in the same orbital is written as:
Wolfgang Pauli

54. Electron Configurations
Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until each orbital has one electron.
Let’s write the electron configuration for Phosphorus
We need to account for all 15 electrons in phosphorus

55. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The first two electrons go into the 1s orbital
Notice the opposite direction of the spins
only 13 more to go...

56. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The next electrons go into the 2s orbital
only 11 more...

57. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The next electrons go into the 2p orbital
only 5 more...

58. Increasing energy The next electrons go into the 3s orbital2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
The next electrons go into the 3s orbital
only 3 more...

59. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The last three electrons go into the 3p sublevel.
They each go into separate orbitals (Hund’s)
3 unpaired electrons
= 1s22s22p63s23p3

60. Orbitals fill in an order:
Lowest energy to higher energy.
Adding electrons can change the energy of the orbital. Full sublevels are the most stable arrangement.
Half filled sublevels have a lower energy than partially filled sublevels, and are next most stable.

61. Write the electron configurations for these elements:
Zirconium - 40 electrons
[Kr] 5s2 4d2
Tantalum - 73 electrons
[Xe] 6s2 4f14 5d3
Chromium - 24 electrons
[Ar] 4s2 3d4 (expected)
But this is not what happens with Chromium

62. Chromium is actually: [Ar]4s13d5 Why?
This gives us two half filled orbitals (the others are all still full)
Half full is slightly lower in energy.
The same principal applies to copper…

63. Copper’s electron configuration
Copper has 29 electrons so we expect: [Ar] 4s2 3d9
But the actual configuration is: [Ar]4s13d10
This change gives one more filled orbital and one that is half filled.
Remember these exceptions: Groups ending in d4 and d9

64. Irregular configurations of Cr and Cu
Chromium steals a 4s electron to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel

65. End of Chapter 13