1. Chapter 5 Electrons In Atoms 5.1 Revising the Atomic Model
5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectra and
the Quantum Mechanical Model

2. CHEMISTRY & YOU
Why do scientists use mathematical models to describe the position of electrons in atoms?
Shown here is a life-sized model of a skier, but not all models are physical. In fact, the current model of the atom is a mathematical model.

3. Understand how the atomic model was revised.
Energy Levels in Atoms
Objective
Energy Levels in Atoms
Understand how the atomic model was revised.

4. Limitations of Rutherford’s Atomic Model
Energy Levels in Atoms
Limitations of Rutherford’s Atomic Model
It explained only a few simple properties of atoms.

5. Limitations of Rutherford’s Atomic Model
Energy Levels in Atoms
Limitations of Rutherford’s Atomic Model
It explained only a few simple properties of atoms.
It could not explain the chemical properties of elements.

6. Limitations of Rutherford’s Atomic Model
Energy Levels in Atoms
Limitations of Rutherford’s Atomic Model
It explained only a few simple properties of atoms.
It could not explain the chemical properties of elements.
For example, Rutherford’s model could not explain why an object such as the iron scroll shown here first glows dull red, then yellow, and then white when heated to higher and higher temperatures.

7. 1913, Niels Bohr develops a new atomic model
 Bohr stated that the electrons orbit the nucleus like the planets orbit the sun.

8. Each possible electron orbit in Bohr’s model has a fixed energy.
The fixed energies an electron can have are called energy levels.
Each energy level further from the nucleus is of greater energy

9. There are 7 different energy levels
Each energy level can contain a different amount of electrons

10. Niels Bohr’s Model (1913)
Electrons orbit the nucleus in circular paths of fixed energy (energy levels).
n=1  first energy level
n=2 second energy lev
n=3  third energy level

11. Highest energy level for carbon is n = 2 (2 rings).
Valence electrons – electrons in the outermost energy level

12. Bohr's model:
-electrons orbit the nucleus like planets orbit the sun
-each orbit can hold a specific maximum number of electrons

13. The Rutherford model could not explain why elements that have been heated to higher temperatures give off different colors of light. The Bohr model explains how the energy levels of electrons in an atom change when the atom emits light.

14. Energy and Atoms
Ground State: the lowest energy state of an atom.
- An electron absorbs energy (photon) and moves from the ground state to an excited state.
Excited State: when an atom contains excess energy (has higher potential energy).
When an excited atom returns to ground state it gives off light (the energy it has gained as electromagnetic radiation).
Example: Neon signs

15. Absorption E4 E3 E2 E1
An electron absorbs energy (photon) and moves from the ground state to an excited state.

16. What goes up…must come down! Emission
When an electron in the excited state returns to the ground state it emits a photon. E4 E3 E2 E1
Ephoton= h =E3-E1

17. Absorption and Emission
Emitting photons creates light or electromagnetic radiation
Electromagnetic radiation in the visible light spectrum has color!
These photons have wavelengths that correspond to their color.

18. Unfortunately, Bohr’s model only applied to hydrogen atoms and did not apply to other atoms. That led scientists to question his model

19. Wave Mechanical Model
Today, the modern description of electrons in atoms is called the Quantum Mechanical Model.
The wave model tells you the probability of finding an electron in an atom (the exact path of an electron is not known)

20. Do Now What does the Bohr model say about electrons?
Why is this model incorrect?

21. There are 7 different energy levels
Each energy level can contain a different amount of electrons
There are 4 different types of sublevels

22. Energy levels (n=1, n=2 ….) Sublevels (s,p,d,f) orbitals

23. The sublevels each can contain a different amount of electrons
s – 2 electrons
p – 6 electrons
d – 10 electrons
f – 14 electrons

24. These orbitals have different shapes
An atomic orbital is a region of space in which there is a high probability of finding an electron.
These orbitals have different shapes

25. Energy level 1 has only an s sublevel – total of 2 e- Energy level 2 has the s and p sublevels – total of 8 e- Energy level 3 has the s, p, and d sublevels – total of 18 e- Energy level 4 has the s, p, d, and f sublevels – total of 32 e-

26. Do Now
1. How many principal energy levels are there? ____ 2. What are the four sublevels? _____ 3. How many orbitals does the f sublevel hold? _____ 4. How many electrons can each orbital hold? _____

27. Do Now Answers
1. How many principal energy levels are there? 7 (n=7) 2. What are the four sublevels? s, p, d , f 3. How many orbitals does the f sublevel hold? 7 4. How many electrons can each orbital hold? 2

28. Chapter 5 Electrons In Atoms 5.2 Electron Arrangement in Atoms
5.1 Revising the Atomic Model
5.2 Electron Arrangement
in Atoms
5.3 Atomic Emission Spectra and
the Quantum Mechanical Model

29. Electron Configurations
5.2 Electron arrangement
Electron Configurations
Aufbau Principle
According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram, each box represents an atomic orbital.
Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p
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30. Electron Configurations
5.2 Electron arrangement
Electron Configurations
Hund’s Rule
According to Hund’s rule, electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.
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31. Electron Configurations
5.2 Electron arrangement
Electron Configurations
Hund’s Rule
Three electrons would occupy three orbitals of equal energy as follows.
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32. Electron Configurations
5.2 Electron arrangement
Electron Configurations
Hund’s Rule
Three electrons would occupy three orbitals of equal energy as follows.
Electrons then occupy each orbital so that their spins are paired with the first electron in the orbital.
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33. 5.2 Electron arrangement
Electron configuration describes the placement of the electrons
Example: Hydrogen: 1s1

34. Electron Configurations
5.2 Electron arrangement
Look at the orbital filling diagram of the oxygen atom.
An oxygen atom contains eight electrons.
Electron Configurations of Selected Elements
Element 1s 2s
2px 2py 2pz 3s
Electron configuration H
1s1 He
1s2 Li
1s22s1 C
1s22s22p2 N
1s22s22p3 O
1s22s22p4 F
1s22s22p5 Ne
1s22s22p6 Na
1s22s22p63s1
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35. Chapter 5 Electrons In Atoms 5.3 Atomic Emission Spectra
5.1 Revising the Atomic Model
5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectra
and the Quantum
Mechanical Model
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36. 5.3 Atomic Emission Spectra
What gives gas-filled lights their colors?
An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color.
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37. 5.3 Electromagnetic Radiation and Energy
By the year 1900, there was enough experimental evidence to convince scientists that light consisted of waves. The wavelength, represented by  (the Greek letter lambda), is the distance between the crests.

38. 5.3 Electromagnetic Radiation and Energy
The frequency, represented by  (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time. The SI unit of waves per second is called the hertz (Hz).

39. The frequency () and wavelength () of light are inversely proportional to each other. As the wavelength increases, the frequency decreases.

41. 5.3 Electromagnetic Radiation and Energy
According to the wave model, light consists of electromagnetic waves.
Electromagnetic radiation - a form of energy that exhibits wavelike behavior as it travels through space.
All electromagnetic radiation travels at the speed of light:
c = 3.0 X108 m/s
Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.

42. Rutherfo

43. Wave Description of Light
Equation relating frequency and wavelength:
c = 
c = speed of light (m/s)
 = wavelength (m)
 = frequency (Hz or s-1)
 = c  = c
 
c is constant, so is , so as frequency increases, wavelength decreases (inversely proportional).

44. Light as a Wave: Problems
c = 
If c = 3.00 x 108 m/s and  = 1 x 1019s-1 ,
what does  equal?
2) What is the frequency of light () if its wavelength () is 4.34 X 10-7 m?

45. Visible light of different wavelengths can be separated into a spectrum of colors. In the visible spectrum, red light has the longest wavelength and the lowest frequency. Violet light has the shortest wavelength and the highest frequency.

46. Atomic Emission Spectrum
Atomic Emission Spectrum- a beam of light separated into a series of specific frequencies (and therefore specific wavelengths) of visible light.
produced when electrons fall back to ground state
the energy emitted in the fall give off specific patterns (colors) of light

47. The Hydrogen-Atom Line Emission Spectrum
The Hydrogen-Atomic Emission Spectrum

48. The Hydrogen-Atomic Emission Spectrum

49. Atomic Emission Spectrum of Na, He, Ne, and Mercury

50. A fluorescent lamp or a fluorescent tube is a low pressure mercury-vapor gas-discharge lamp that uses fluorescence to produce visible light. An electric current in the gas excites mercury vapor which produces short-wave ultraviolet light that then causes a phosphor coating on the inside of the bulb to glow.

51. E = hn The Quantization of Energy Energy
Planck’s constant (h)
= 6.63 x Js
E = hn
Energy
Frequency (n)
German physicist Max Planck (1858–1947) showed mathematically that the amount of radiant energy (E) of a single quantum absorbed or emitted by a body is proportional to the frequency of radiation (n).
h is Planck’s constant = 6.63 x Js
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52. Energy Problems
E = hn
h = 6.63 x Js
If the frequency (n) = 1.15 x 1012 s-1 , what is the energy of the radiation?
What is the energy of a photon of microwave radiation with a frequency of 3.20 × 1011s-1?

53. Photons
Einstein proposed that light could be described as quanta of energy that behave as if they were particles. These light quanta are called photons.

54. Key Concepts
Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.
The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus of an atom.
Each energy sublevel corresponds to one or more orbitals of different shapes, which describe where the electron is likely to be found.

55. Glossary Terms
energy level: the specific energies an electron in an atom or other system can have
quantum: the amount of energy needed to move an electron from one energy level to another

56. Glossary Terms
quantum mechanical model: the modern description, primarily mathematical, of the behavior of electrons in atoms
atomic orbital: a mathematical expression describing the probability of finding an electron at various locations; usually represented by the region of space around the nucleus where there is a high probability of finding an electron

57. Bohr's Model of the Atom
e.g. fluorine:
#P =
#e- =
#N =

58. Bohr's Model of the Atom
e.g. fluorine:
#P = atomic #
= 9
#e- =
#N =

59. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = # P
= 9
#N =

60. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = atomic mass - # P
= 10

61. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
draw the nucleus with
protons & neutrons 9P
10N

62. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons can fit in the first orbit? 9P
10N

63. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons can fit in the first orbit? 2 9P
10N

64. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons are left? 9P
10N

65. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons are left? 7 9P
10N

66. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons are left? 7
how many electrons fit in the second orbit? 9P
10N

67. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons are left? 7
how many electrons fit in the second orbit? 8 9P
10N

68. Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10 9P
10N

69. Bohr's Model of the Atom e.g. fluorine: #P = 9 #e- = 9 #N = 10 9P 10N
How many valence electrons? 7 9P
10N