1. The Development of Atomic Theory
Larry Scheffler
Lincoln High School
Portland OR 1

2. The Atom
The term atom is derived from the Greek word atomos (atomos) meaning invisible
Democritius ( BC ) suggested that all matter was made up of invisible particles called atoms 2

3. Law of Constant Composition
A compound always contains atoms of two or
More elements combined in definite proportions
by mass
Example:
Water H2O always contains 8 grams of oxygen to 1 gram of hydrogen 3

4. Law of Multiple Proportions
Atoms of two or more elements may
combine in different ratios to produce
more than one compound.
Examples:
NO NO N2O N2O5 4

5. Dalton’s Atomic Theory
All elements are composed of indivisible and indestructible particles called atoms.
Atoms of the same element are exactly alike, They have the same masses.
Atoms of different elements have different masses.
Atoms combine to form compounds in small whole number ratios.. 5

6. Objections to Dalton’s Atomic Theory
Atoms are not indivisible. They are
composed of subatomic particles.
Not all atoms of a particular element
have exactly the same mass.
Some nuclear transformations
alter (destroy) atoms 6

7. Crookes Experiment
Crookes found that passing an electrical current through a gas at very low pressure caused the gas to glow. Putting a magnet next to the beam caused it to be deflected. 7

8. The Electron
The electron was the first subatomic particle to be identified.
In 1897 J.J Thomson used a cathode ray tube to establish the presence of a charged particle known as the electron
Thomson established the charge to mass ratio
E/m = 1.76 x 108 coulombs/gram 8

9. A Cathode Ray Tube
Thomson found that an electrical field would also deflect an electron beam. He surmised that the ratio of charge to mass is constant.

10. Thomson’s Charge to Mass Ratio
E/m = 1.76 x 108 coulombs/gram

11. Thomsen’s Plum Pudding Model
Thompson proposed that an atom was made up of electrons scattered unevenly through out an elastic sphere. These charges were surrounded by a sea of positive charge to balance the electron's charge like plums surrounded by pudding.
This early model of the atom was called The Plum Pudding Model. A more contemporary American label might be the “chocolate chip cookie” model 11

12. Millikan’s Experiment
By varying the charge on the plates, Millikan found that he could suspend the oil drops or make them levitate. 12

13. Millikan’s Experiment
Millikan used his data to measure the charge of an electron and then to calculate the mass of the electron from Thomson’s charge to mass ratio.
Given the charge =
1.60 x coulomb and the ratio of E/m = 1.76 x 108 coulombs/gram it is possible to calculate the mass
Mass
= x gram 13

14. Protons First observed by E. Goldstein in 1896
J.J. Thomson established the presence of positive charges.
The mass of the proton is
1.673 x grams 14

15. Rutherford’s Experiment
Rutherford oversaw Geiger and Marsden carrying out his famous experiment.
They fired high speed alpha particles (Helium nuclei) at a piece of gold foil which was only a few atoms thick.
They found that although most of them passed through. About 1 in 10,000 hit and were deflected
1910
Ernest Rutherford 15

16. Rutherford’s Experiment16

17. Rutherford’s Experiment17

18. Rutherford’s Experiment
By studying this pattern, Rutherford concluded that atoms have a very dense nucleus, but there are mostly empty space. 18

19. Subatomic Particles The diameter of a single atom ranges
From 0.1 to 0.5 nm. (1 nm = 10-9 m).
Within the atom are smaller particles:
Electrons
Protons
Neutrons 19

20. Neutrons Discovered by James Chadwick in 1932
Slightly heavier than a proton
Mass of a neutron = x grams 20

21. The Bohr Model
Niels Bohr proposed the Planetary Model in Electrons move in definite orbits around the nucleus like planets moving around the nucleus. Bohr proposed that each electron moves in a specific energy level. 21

22. Aspects of the Bohr Model
Bohr put together Balmer’s and Plank’s discoveries to form a new atomic model
In Bohr’s model:
Electrons can orbit only at certain allowed distances from the nucleus.
Electrons that are further away from the nucleus have higher energy levels (explaining the faults with Rutherford’s model). 22

23. The Electromagnetic Spectrum

24. Wave Characteristics Energy of a wave E = hn
Frequency = n = number of peaks per unit of time
Speed of light
c = nl

25. Emission Spectra 25

26. Flame Tests

27. According to Bohr
Atoms radiate energy whenever an electron jumps from a higher-energy orbit to a lower-energy orbit. Also, an atom absorbs energy when an electron gets boosted from a low-energy orbit to a high-energy orbit. 27

28. Problems with the Bohr Model
The Bohr model provided a model that gave precise results for simple atoms like hydrogen.
Using the Bohr model precise energies could be calculated for energy level transitions in hydrogen.
Unfortunately these calculations did not work for atoms with more than 1 electron. 28

29. Weakness of the Bohr Model
According to the Bohr model electrons could be found in orbitals with distinct energies.
When the data for energies measured using spectral methods where compared to the values predicted by the Rydberg equation, they were accurate only for hydrogen.
By the 1920s, further experiments showed that Bohr's model of the atom had some difficulties. Bohr's atom seemed too simple to describe the heavier elements. 29

30. Modern View of the Atom
The wave mechanical model for the atom was developed to answer some of the objections that were raised about the Bohr model. It is based on the work of a number of scientists and evolved over a period of time
The quantum theorists such as Maxwell Planck suggested that energy consists of small particles known as photons. These photons can have only discreet energies
Maxwell Planck 30

31. Modern View of the Atom
Albert Einstein demonstrated the equivalence of matter and energy. Hence matter and energy in Einstein’s theory were not different entities but different expressions of the same thing
Einstein then proposed the equivalence of Matter and Energy given by his famous equation
E = mc2 31

32. Modern View of the Atom
Louis de Broglie suggested that if energy could be thought of as having particle properties, perhaps matter could be thought of as having wave like characteristics
Louis de Broglie 32

33. Modern View of the Atom
Louis de Broglie proposed that an electron is not just a particle but it also has wave characteristics.
E = mc2 = hn 33

34. --Heisenberg, Uncertainty paper, 1927
Modern View of the Atom
The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa.
--Heisenberg, Uncertainty paper, 1927
Heisenberg proposed that it was impossible to know the location and the momentum of a high speed particle such as an electron. 34

35. Modern View of the Atom
The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa Werner Heisenberg,
Uncertainty paper, 1927
The atom cannot be defined as a solar system with discreet orbits for the electrons. The best that we could do was define the probability of finding an electron in a particular location. 35

36. Modern View of the Atom
Edwin Schroedinger proposed that the electron is really a wave. It only exists when we identify its location. Therefore the electrons are best thought of probability distributions rather than discreet particles. 36

37. Modern View of the Atom
The modern view of the atom suggests that the atom is more like a cloud. Atomic orbitals around the nucleus define the places where electrons are most likely to be found. 37

38. Wave Mechanical Model
The location of the electron in a hydrogen atom is a probability distribution. 38

39. Progression of Atomic Models39
Our view of the atom has changed over time

40. ATOMIC STRUCTURE Particle Charge Mass proton + charge 1 neutron
No charge 1
electron
- charge 40

41. ATOMIC NUMBER AND MASS NUMBER Number of electrons = Number of protonsHe
Mass Number 4
the number of protons and neutrons in an atom 2
Atomic Number
the number of protons in an atom
Number of electrons = Number of protons
in a neutral atom 41

42. Atomic Mass
The atomic mass of an atom is a relative number that is used to compare the mass of atoms.
An atomic mass unit is defined as 1/12 of the mass of an atom of carbon The atomic masses of all other atoms are a ratio to carbon 12 42

43. Isotopes Many elements have atoms that have multiple forms
Different forms of the same element having different numbers of neutrons are called isotopes.
For example: Carbon exists as both Carbon 12 and Carbon 14
Carbon 12 Carbon 14
6 electrons 6 electrons
6 protons 6 protons
6 neutrons 8 neutrons 43

44. Isotopes and Atomic Mass
Many elements have atoms that have multiple isotopes.
Isotopes vary in abundance. Some are quite common while others are very rare.
The atomic mass that appears in the periodic table is a weighted average taking into account the relative abundance of each isotope. 44

45. or Na-23
or Na-24
Isotope: one of two or more atoms having the same number of protons but different numbers of neutrons

46. Measuring Atomic Mass --the Mass Spectrometer
The mass spectrometer can be used to determine the atomic mass of isotopes.

47. Mass Spectrum of Neon
The mass spectrum neon shows three isotopes with the isotope at atomic mass = 20 accounting for more than 90% of neon.

48. Mass Spectrum of Germanium
The mass spectrum of germanium shows 5 peaks at relative atomic masses of 70, 72,73,74, and 75

49. Calculating the average relative atomic mass
The average atomic mass that is shown in the periodic table is really the weighted average of the atomic masses of each of the elements isotopes. Germanium has 5 isotopes whose relative atomic masses are shown in the table
Mass Number % Abundance

50. Calculating the Average Relative Atomic Mass
To calculate the average atomic mass multiply the atomic mass of each isotope by its abundance (expressed as a decimal fraction)
Mass Number % Abundance
Average atomic mass
= (0.2055)(70) + (0.2737)(72) + (0.0767)(73) + (0.3674)(74)+ (0.0767)(75) =
Note: atomic masses are ratios so they do not have real units although they are sometimes called atomic mass units or amu

51. Problem
The mass spectrum of an element, A, contained 4 lines at mass/charge ratios of 54, 56, 57 and 58 with relative intensities of 5.84, 91.68, 2.17 and 0.31 respectively. Calculate the relative atomic mass of element A.
Average atomic mass
= (0.0584)(54) + (0.9168)(56) + (0.0217)(57) + (0.0031)(58) =

52. The Nucleus
The nucleus is very small — of the order of meter whereas the atom is of the order of 10-9 meters. By analogy, the nucleus occupies as much of the total volume of the atom as a fly in a cathedral

53. Protons and Neutrons
Protons and neutrons have nearly equal masses, and their combined number, the mass number, is approximately equal to the atomic mass of an atom.
The combined mass of the electrons is very small in comparison to the mass of the nucleus, since protons and neutrons weigh roughly 2000 times more than electrons.
Charge
Mass
Relative Mass
Electron -1
x g
1/1837
Proton +1
x g 1
Neutron
x g

54. Atomic Mass Units
An atomic mass unit (amu) is equal to exactly 1/12 of the mass of an atom of Carbon 12.
One atomic mass unit is equal to x grams. Note that this is slightly less than the mass of a proton or a neutron.
An atomic mass unit is sometimes called a Dalton (D).
1.00 g = x 1023 amu. This number is also known as Avogadro’s Number and it defines the size of a quantity we call a mole.

55. Radioactive Nuclei
The presence of neutrons in the nucleus tends to buffer the repulsions of multiple protons in the nucleus.
There appears to be an optimal number of neutrons for the number of protons in a given atom in a stable atom.
In a radioactive element the nucleus may disintegrate releasing either an alpha particle or a beta particle as well as some high energy gamma radiation.

56. Alpha Particles 238 U --> Th + 4 He 92 2
An alpha particle consists of two protons and two neutrons. This makes it equivalent to a helium nucleus.
When a radioactive element undergoes alpha decay its nucleus is decreased in mass by 2 protons and 2 neutrons.
Since the number of protons changes, it has a new atomic number and hence it is a different element. The mass number decreases by 4.
238
U >
234
Th + 4 He 92 90 2
146 neutrons
92 protons
Ratio = 1.52
144 neutrons
90 protons
Ratio = 1.60
Alpha decay raises the neutron to proton ratio. It occurs in radioactive nuclei where the ratio is too low

57. Beta Particles
A beta particle consists of a high-speed electron released from the nucleus.
When a radioactive element undergoes beta decay, the number of protons increases by one and the number of neutrons decreases by one. The mass number remains the same. 14
C >
N + e- 6 7 -1
8 neutrons
6 protons
Ratio = 1.333
7 neutrons
7 protons
Ratio = 1.00
Beta decay lowers the neutron to proton ratio. It occurs in radioactive nuclei where the ratio is too high

58. Beta Particles
A beta particle consists of a high-speed electron released from the nucleus.
When a radioactive element undergoes beta decay, the number of protons increases by one and the number of neutrons decreases by one. The mass number remains the same. 14
C >
N + e- 6 7 -1
8 neutrons
6 protons
Ratio = 1.333
7 neutrons
7 protons
Ratio = 1.00
Beta decay lowers the neutron to proton ratio. It occurs in radioactive nuclei where the ratio is too high

59. The Half-Life
The rates at which various radioactive elements undergo decay vary considerably. The half-life of a radioactive element if the time required for half of the nuclei in a sample of radioactive nuclei to disintegrate.
Isotope
Type
Half-life
Uranium 238
Alpha
4.51 x 109
Carbon 14
Beta
5730 years
Iodine 131
8 days
Radon-222
3.825 days
Cesium -137
30 years
Polonium 210
138 days