1. Atomic Theory

2. Democritus – 460 – 370 BC Matter consisted of tiny particles “atomos”
Ideas were wrong

3. John Dalton (1803)

4. Dalton’s Theory – Explains the Law of Conservation of Matter
Atoms are separated, combined or rearranged in a reaction, they are not created, destroyed or divided.

5. What is an atom?
The smallest particle of an element that retains its original properties

6. Discovering the Existence of the Electron – JJ Thompson
Cathode Ray Tube – Led to the TV

7. Rays consisted of particles.
All types of gases and cathodes produced a beam of particles.
Particles had a negative charge.

8. Which part of Dalton’s theory did he prove wrong?
JJ Thomson – 1890s
Found that the mass of the particle in the cathode ray tube was smaller than the mass of the Hydrogen atom.
Which part of Dalton’s theory did he prove wrong? X

9. Robert Millikan – 1909 – Oil Drop Experiment
Determined the mass and charge of an electron.

10. Oil droplets of different masses
Dropped between charged plates
Became negatively charged
Drops fell due to gravity
Negatively charged lower plate repelled them and they became stationary
The magnitude of the charge could be calculated from the known voltage and mass.

11. More Questions
If electrons are part of all matter and they possess a negative charge, why is matter neutral?
If the mass of an electron is so small, what accounts for the rest of the mass in an atom?

12. JJ Thomson’s Answer – Plum Pudding Model

13. Ernest Rutherford’s Experiment - 1911

14. Evidence Contradicts the Plum Pudding Model
Rutherford Concluded:
Atoms consist mainly of empty space in which electrons move freely.
A tiny dense space in the center of the atom contains the majority of the mass and the positive charges.
The positive charge of the nucleus holds the negative electrons within the atom.

15. 1st Question Answered
If electrons are part of all matter and they possess a negative charge, why is matter neutral?

16. Protons and Neutrons Rutherford (1920) James Chadwick (1932)
Protons are positively charged and found in the nucleus
James Chadwick (1932)
Neutrons are found in the nucleus, has no charge and has a mass equal to the proton
Space Between Atoms - Video

17. Second Question Answered: If the mass of an electron is so small, what accounts for the rest of the mass in an atom?

18. What do atoms look like? Silicon Atoms
How can we take pictures of atoms?- Video

21. Atomic Number Henry Moseley Each element has a unique positive charge
Equal to # of Protons and Electrons

22. Mass Number and Atomic Number
Mass number = protons + neutrons
1H 1H H
Protium Deuterium Tritium
Mass number
Atomic number
1 proton
1 neutron
1 proton
2 neutrons
1 proton
Isotopes – atoms with the same number of protons, but a different number of neutrons

23. Atomic Mass Unit 1 AMU = 1/12 the mass of the Carbon Atom
1 AMU ≠ mass of a proton or a neutron

24. Average Atomic Mass Chlorine-35 (34.969 amu x 75.770%)

27. Chlorine’s average atomic mass is 35. 452 amu
Chlorine’s average atomic mass is amu. Chlorine-35 has an atomic mass of and Chlorine-37 has an atomic mass of What are the percentages of each isotope?

28. Electrons in Atoms

29. Problems with Rutherford’s Model
Chlorine # 17
Reactive
Argon # 18
Not reactive
Potassium # 19
Very reactive

30. The Quest for a Better Model
Electromagnetic radiation behaves like a wave. (video)

31. Characteristics of a Wave
Wavelength = λ
Frequency = v (number of waves that pass a point per second)
1 Hertz (Hz) = 1 wave per second (SI Unit for frequency)

32. Speed and Frequency of Light
c = λv
c = speed of light (3.0 x 108 m/s)
↑ wavelength ↓ frequency
↓ wave length ↑frequency

33. What is the relationship between energy and frequency?

34. Microwave Oven - Simulation

35. Problems

36. Light: Particle or Wave?
Wave model doesn’t address:
Why heated objects emit only certain frequencies of light at a given temperature?
Why some metals emit electrons when a colored light of a specific frequency shines on them?

37. Iron Dark gray = room temp Red = hot temp Blue = extremely hotter temp
↑ temp, ↑ kinetic energy, emit different colors of light
Wave model could not explain this

38. Max Planck
Matter gains or loses energy only in small, specific amounts called quanta
quantum is the minimum amount of energy that can be gained or lost by an atom
Equantum = hv
h – Planck’s constant – x J·s
J = joule, SI Unit for energy

39. Photoelectric Effect – The problem with wave theory. - Simulation
Only certain frequencies of light could emit an electron from a plate of Ag.
Accumulation of low frequencies couldn’t

40. Einstein and the Dual Nature of EMR (1900)
EMR acts as a wave of individual particles (photon)
Ephoton = hv

41. Calculating the energy in a Photon
Ephoton = hv
E = (6.626 x J·s) x (7.23 x 1014 s-1)
E = 4.79 x J

42. Atomic Emission Spectra
The frequencies of the EMR emitted by atoms of the element.
Unique to each element

43. Flame Test Demo

44. Niels Bohr - 1913 Worked in Rutherford’s lab
Proposed a quantum model of the atom
Explain why emission spectra were discontinuous
Predicted frequencies of light in Hydrogen’s atomic emission spectra

45. Bohr’s Explanation Ground state – lowest energy state of an atom
Excited state – when an atom gains energy
Electrons move in circular orbits
Smaller orbit – lower energy state, “energy level”
Larger orbit – higher energy state, “energy level”

46. An explanation for the Emission Spectra
Atoms absorb energy and are excited. As the electron returns to the ground state they give off energy “photon” equal to the difference in energy levels.
Excited States
Ground state

47. photon
Electron absorbs energy e-
Electron in ground state

48. Energy released
“photon”
Electron in Excited State e-
Returns to
Ground state

49. Atomic Spectra - Simulation

50. Problem: Bohr’s Model Only explains Hydrogen
Louis de Broglie (1924) – proposed that the energy levels are based on the wave like nature of electrons

51. Heisenberg Uncertainty Principle
It is impossible to know the velocity and position of a particle at any given time
(Video)
Photon and electron are about the same mass.

52. Erwin Schrodinger
Developed the quantum mechanical model of the atom
Assigns electrons to energy levels like Bohr
Does not predict the path of the electron
It predicts the probability of finding an electron
An electron’s “atomic orbital”
(Video – Quantum Atom)

53. Each dot is a picture of an electron during a given amount of time.
Where does the electron spend most of the time?
Boundary represents the location of an electron 90% of the time.

54. Principle Energy Levels
Lowest energy is 1 – greatest energy 7
Each level consists of sublevels
The second energy level is larger and the electrons are farther from the nucleus.

55. Types of Sublevels
Same energy

56. Putting it together: Principal quantum number “energy level”
Sublevels (Types of orbitals and total number) 1
S – 1 2
S – 1, P – 3 3
S – 1, P – 3, D – 5 4
S – 1, P – 3, D – 5, F - 7

57. Ground State Electron Configurations
Most stable – lowest energy
3 principals to follow

58. Aufbau Principle – each electron must occupy the lowest energy state
Orbitals in an energy sublevel have equal energy
The energy sublevels in a principle energy level have different energies.
The sublevels increase in energy from s,p,d,f
Principal energy levels can overlap

59. Sublevels have different energy levels
Aufbau Diagram
Energy levels overlap
Equal energies – 2 p
Sublevels have different energy levels

60. Pauli Exclusion Principle
Each electron spins
Electrons must spin in opposite directions
2 electrons per orbital
Written as

61. Hund’s Rule
Electrons must occupy each orbital before additional electrons can be added.
Pauli Exclusion Principle

62. Representing Electron Configurations

63. Electron Configurations
Sub level diagram – indicates the order that orbitals are filled
What are the orbital diagrams and electron configuration notation for Al and Cl?

64. Electron Configuration Shorthand
Substitute noble gases from preceding energy levels in the notation
Li – [He] 2s1
C – [He] 2s2 2p2

65. Valence Electrons Electrons in the outer most energy levels
S [Ne] 3s2 3p4
Sulfur has 6 valence electrons
How many valence electrons
do Al, Ne, and Cl have?

66. Electron Dot Structures
Valence electrons are used in reactions and are represented by an electron dot structure.

67. Writing Electron Dot Structures
Fill the valence electrons 1 at a time in any particular order.
Ca C O * * * * * * * * * * * *
What are the electron dot diagrams for K, Ar and F?

68. Modern Atomic Theory
Any electron in an atom can be described by 4 quantum numbers
Principal Quantum Number
Azimuthal Quantum Number
Magnetic Quantum Number
Spin Quantum Number

69. Principal Quantum Number (n)
Related to the size and energy of principal energy level.
The farther away from the nucleus the more energy the electron has
1 < 2 < 3 < 4 < 5 < 6 etc….

70. Azimuthal Quantum Number (Angular Momentum) = l
Refers to the subshells in each principal energy level (n)
S = 0
P = 1
D = 2
F = 3 n l 1 2 3 4

71. Magnetic Quantum Number (ml)
Orbital designation ml 1 1s 2 2s 2p
-1,0,+1 3 3s 3p 3d
-2,-1,0-,1,2 4 4s 4p 4d 4f
-3,-2,-1,0,1,2,3
Specifies the orbital within a energy level where an electron is likely to be found

72. Spin Quantum Number (ms)
+ ½ or – ½
Electrons in the same orbitals must have opposite spins (Pauli Exclusion Principle)

73. n l ml ms Orbital designation 1 1s + ½, - ½ 2 2s 2p 3 3s 3p 3d 4 4s 4p1s
+ ½, - ½ 2 2s 2p
-1,0,+1 3 3s 3p 3d
-2,-1,0-,1,2 4 4s 4p 4d 4f
-3,-2,-1,0,1,2,3

74. A B n l ml ms 2 1 -1
+ ½ or
What are the quantum numbers for A? B?