4. Ch. 4 Electron Arrangement
Part 1: “What does light have to do with electrons???”

5. Ch. 4 Electrons in atoms – a new atomic model
Rutherford’s model of the atom featured a highly dense, positively charged nucleus surrounded by extremely small, fast- moving, negatively charged electrons that occupied most of the atom’s volume.

6. His model did not provide:
1. Detail about how electrons were arranged around the nucleus
2. an explanation for why the negatively-charged electrons were not sucked into the positively- charged nucleus.
3. insight into the chemical properties of matter.

7. Breakthrough into chemical behavior
The puzzle of the distinct chemical behavior of different elements began to unravel when scientists noticed that certain elements emitted visible light when heated in a flame.
Analysis of this light offered insight into the arrangement of the electrons in these atoms, which was then linked to their chemical behavior.

8. Review of the properties of light.
Visible light is a form of electromagnetic radiation, just like microwaves, gamma rays, x-rays and radio waves.
All electromagnetic radiation, including visible light, travels at a speed of 3.00 x 108 m/s in a vacuum. Nothing can travel faster.
We now know that light has a dual nature as it travels. As it moves, it has properties of both a wave and a particle.

9. Wave properties

10. Review of Wave Properties (copy all)
Wavelength (l)-distance between equivalent points on a wave.(units-length) meter
Frequency(n)- number of waves that pass a set point per second.(units-waves/sec or Hertz- Hz)
Amplitude – the height of the wave from its origin to its crest.
Velocity- (c = velocity of light)
c = ln = 3.00 x 108 m/s in a vacuum.
The speed of light is constant because as frequency decreases / wavelength increases, and vice versa.

11. Particle Nature of Light
Section 4-1 continued

12. Duality of Light
Many of the interactions between light and matter observed by physicists and chemists could not be explained if light had only wave-like properties.
In order to explain these occurrences, scientists proposed a revolutionary idea:
That light could have properties of waves and particles.

13. The Quantum Concept
In 1900, a German Physicist, Max Planck was searching for an explanation for why heated objects produced light of only specific frequencies (colors) rather than of all frequencies (white light).
He proposed that matter can only gain or lose energy in small specific amounts he called quanta.

14. A quantum is the minimum amount of energy that can be gained or lost by an atom.
Prior belief was that matter could absorb or emit energy in continuous and variable quantities, and would have to, if light traveled only as a wave.

15. Relationship between frequency and energy
Planck proposed that the frequency of the light produced by a heated object corresponded to the specific amount of energy in a quantum.
Equantum = hn
Planck’s constant (h) = x J x s

16. Equation showed that the energy of a quantum of light increases with increased frequency.
Thus violet light has greater energy than red.
Planck believed that the emission of light of only specific frequencies was proof that matter released and absorbed energy in packets.(quanta)

17. The Photoelectric Effect
Another death blow to the wave-only view of light was the explanation of the photoelectric effect.
The photoelectric effect is the emission of electrons from a metal when light shines on the metal.

18. The mystery was that light of only a certain minimum frequency was capable of knocking these electrons off.
If light traveled in waves, then even low frequency light (shined on the metal for a long enough time) should have been capable of building up enough energy to knock off an electron.

19. Einstein makes the leap
In 1905, Albert Einstein proposed an explanation for Planck’s quantum concept and the photoelectric effect.
He proposed that light had a dual, wave- particle nature.
He proposed that light travels in waves of individual particles. Each particle of light carries a quantum of energy.

20. He called these particles photons
He called these particles photons. The energy of a particular photon depends on the frequency of the electromagnetic radiation. Ephoton = hn

21. Photons were the key
Einstein proposed that in order for an electron to be knocked from a metal, the electron must be struck by a single photon of light with enough energy to knock it loose.
If struck by a photon with a frequency too low, and thus having too little energy, the electron would stay bound to the metal.

22. Line-Emission Spectrum and the Bohr Model of the H atom
4-1 continued

23. Excited vs Ground States
Atoms normally exist in a low energy state called the ground state.
However, if energy is passed into an atom, it is capable of absorbing the energy and entering an excited state.
When the atom releases the energy and returns to ground state, the energy is emitted as electromagnetic radiation, some of which is visible.
Ex. Neon and sodium lights

24. Line Emission Spectrum
If this light is separated with a prism, it separates into a series of specific frequencies of light.
This series of frequencies is highly specific for each type of atom and is called a line- emission spectrum.

25. These line-emission spectrums are often used to identify unknown elements.
Prior belief held that excited atoms should return to their ground states smoothly and emit light in a continuous spectrum rather than in bands of specific frequencies.

26. Why bands rather than a continuous spectrum?
Characteristic wavelengths of light are produced by atoms because there are specific amounts of energy between the various excited states and the ground state.
Each jump produces a photon of light equal in energy to the difference between the energy levels.

27. Each photon has a specific frequency and therefore a specific wavelength. Ephoton = hn
Each type of atom has its own series of specific excited energy states. The pattern is fixed and specific for each element.

28. Bohr Model of the H atom
Niels Bohr explained the line-emission spectrum of hydrogen in 1913.
He linked the various energy states of hydrogen to the position of its electron.
When the atom was in its ground state, the electron was in an orbit close to the nucleus.

29. As energy was provided, the electron jumped to an orbit farther from the nucleus.
This new orbit was a specific distance from the nucleus. The electron could not exist between the orbits.

30. Bohr’s model continued
Even more energy could knock the electron still further from the nucleus to other orbits.
As the electron falls back toward the nucleus, a photon of light (equal to the amount of energy required to knock the electron out of orbit) is emitted from the atom.

31. This is why the light produced by hydrogen was in such precise wavelengths.
However, Bohr’s model did not explain the line emission spectra for atoms with more than one electron. This was left to the Quantum theory.

32. The Quantum Model of the Atom

33. Louis de Broglie
The behavior of light as both a particle and a wave led one scientist, Louis de Broglie, in to hypothesize that electrons behaved in a similar fashion.
He said that the definite energy states of electrons corresponded to how waves behave when confined to an enclosed space.

34. In the case of the atom, the enclosed space was the region around the nucleus.
The wave/particle behavior of electrons was later confirmed by experiments.

35. Werner Heisenberg
de Broglie’s led to a new interest in where electron actually were located around the nucleus.
A new hypothesis by Werner Heisenberg in offered a paradox for scientists trying to locate electrons.

36. Heisenberg uncertainty principle- states that it is impossible to absolutely simultaneously determine both the position and velocity of an electron or any other particle.
This hypothesis turned out to be absolutely correct.

37. What did the uncertainty principle mean:
In other words, the act of looking for an electron distorted its position.
Scientists realized that the best they could do was establish areas around the nucleus that had a high probability of containing an electron.
These areas of “probable” electron location are called orbitals.

38. Quantum Theory
The charting and prediction of the probable paths of electrons around the nucleus is called the Quantum Theory.
It uses mathematical formulas to describe the wave- like movement of electrons around the nucleus and determines the shape and location of their orbitals.
The first scientist to mathematically describe the probable positions of electrons was Erwin Schrodinger, in 1926.
Orbitals, however, are not the simple planet-like orbits their name suggests.

39. Assignment:
Continue working on timeline given earlier. (Names on board) – due Thursday.
Homework: Back of 2nd page of note outline:
Review:
Do #’s 3-5, 9-19.

40. Assignment:
#7. Make a timeline showing all names, dates & major contribution.
May work alone or w/ three partner.
Use white or goldenrod paper; may use construction paper, markers, etc.

41. Atomic Orbitals and Quantum Numbers
Section 4-2 continued

42. What are orbitals and quantum numbers?
An orbital is a three-dimensional region around the nucleus that indicates the probable location of an electron.
Orbitals are described and identified by their quantum numbers.
Quantum numbers specify the properties of atomic orbitals and the properties of the electrons in the orbitals.

43. Types of quantum numbers
Principle quantum number- indicates main energy level
Angular momentum quantum number- indicates the shape of the orbital
Magnetic quantum number- indicates the orientation of the orbital.

44. Principle Quantum number
Identifies the energy level of an electron
Values are 1,2,3, etc.
As the number increases, the average distance of the energy level from the nucleus increases.
To calculate the total # of orbitals that can exist in an energy level, square the principle quantum number. Example: How many orbitals in energy level 3? Answer: 32 = 9 (1s, 3p, 5d)
Energy levels are also called shells.

45. Angular momentum quantum number
Indicates the shape of an orbital
There are as many different orbital shapes in an energy level as the number of the energy level itself. Example: Energy level 4 has how many different orbital shapes?
Answer: 4 (s, p, d, f)
A group of orbitals of the same shape are called a sublevel. (p sublevel; d sublevel)

46. Different orbital shapes
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s orbital is spherical in shape
p orbital is dumbbell-shaped
d orbital is varied in shape
f orbital is so complex in shape that we won’t worry about it.
There are of course other orbital shapes, as many as there are energy levels, we just won’t go beyond these 4 in this course

47. How to use these quantum numbers.
An atomic orbital is labeled first with the principal quantum number, followed by the letter of the sublevel.
3p orbital would designate one of the three p orbitals in the 3rd energy level.
4d sublevel would designate all five of the d orbitals in the 4th energy level.

48. Magnetic Quantum Numbers
indicates the 3-D orientation of the orbital
s orbital has only one possible orientation because it is spherical, so no magnetic quantum number is necessary. (3s, 5s)
p orbital has 3 possible orientations, along the x, y, and z axis. They are labeled px, py, and pz. (2px, 2py, 2pz)
d orbital has 5 possible orientations.
f orbital has 7 possible orientations.

49. Electron Configurations
Section 4-3

50. What are they for?
An electron configuration describes the arrangement of electrons in an atom.
Because each different element has a different number of electrons, each has its own distinct electron configuration.
Electrons always want to exist in their lowest energy state. This state is called the ground-state electron configuration.

51. Rules for writing electron configurations
Aufbau principle- an electron will occupy the lowest-energy orbital that can receive it.
Lowest energy Highest energy
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s

52. -Orbital Notations 2/12/13 2 Rules first!
Pauli exclusion Principle = no 2 electrons in the same atom can have the same set of 4 quantum numbers (electron config.)
Hund’s rule = first fill orbitals w/ one electron before pairing w/ another. All unpaired electrons must have the same spin.

53. Steps for orbital configurations (diagrams) (copy)
1. Write the electron configuration (long way)
2. Use lines to represent orbitals; put # & letter underneath.
3. Draw arrows over lines to show electrons. (follow Hund’s rule)

54. Example: Boron 1s22s22p1 ___ ____ _____ ______ ______
___ ____ _____ ______ ______
1s s 2px 2py 2pz
Assignment: Draw the orbital config. for #1 – a, f, g, h, l & m on your electron config. sheet.

55. Assignment:
Finish noble gas configurations & 6 orbital configurations’

56. Rules continued…
Hund’s rule- orbitals with the same amount of energy (example, all the 2p orbitals, or all of the 3d orbitals) will be occupied by one electron before any of these orbitals has two.
In addition, all electrons in single orbitals will have the same spin.

57. Using the rules Write the quantum numbers for each orbital in carbon.
1s2 2s2 2px1 2py1
Write the quantum numbers for each electron of oxygen.
1s2 2s2 2px2 2py1 2pz1