1. Atoms and their structure Mr. Bruder
Atomic Structure
Atoms and their structure
Mr. Bruder

2. Democritus (400 B.C.)
proposed that matter was composed of tiny indivisible particles called atomos
not scientifically tested
"Nothing exists but atoms and empty space; everything else is opinion."

3. Democritus (400 B.C.)
rejected by Aristotle and others who believed that matter could be endlessly divided

4. Evidence for Atoms Law of Conservation of Mass
Mass is not gained or lost in a chemical reaction.
Proposed by Antoine Lavoisier in 1787.
What would happen to the mass reading if the reaction was done without the balloon (an open system)?
Figure 2.2

5. Evidence for Atoms Law of Definite Proportions
Proposed by Joseph Proust between 1797 and 1804
A compound always has the same relative amounts of the elements that compose it.
For example, when water is broken down by electrolysis into oxygen and hydrogen, the mass ratio is always 8 to 1.
Figure 1.2

6. Dalton’s Atomic Laws
  Law of Multiple Proportions – if two elements can combine to form more than one compound, then the ratio of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.
Example:
CO vs. CO2
Formula
Ratio of N:O

7. John Dalton (1807) British Schoolteacher
based his theory on others’ experimental data and developed the atomic theory
Billiard Ball Model
atom is a uniform, solid sphere

8. Dalton’s Atomic Theory
John Dalton ( ) had four theories
All elements are composed of submicroscopic indivisible particles called atoms
Atoms of the same element are identical. The atoms of anyone element are different from those of any other element
Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds
Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction

9. Atoms & Subatomic Particles
Atom- smallest particle of an element that retains the properties of that element

10. Proof of Atoms: STM Image of Gold
The scanning tunneling microscope, invented in 1981, allows us to create images of matter at the atomic level.
Figure 2.4

11. Electron J.J Thomson (1856-1940) – discovered the electron in 1897
Electron is the negative charged subatomic particle
An electron carries exactly one unit of negative charge & its mass is 1/1840 the mass of a hydrogen atom

12. Cathode Ray
The Cathode Ray tubes pass electricity through a gas that is contained at a very low pressure

13. Thomson’s Experiment
Voltage source - +

14. Thomson’s Experiment
Voltage source - +

15. Thomson’s Experiment
Voltage source - +

16. - + Thomson’s Experiment Voltage source
Passing an electric current makes a beam appear to move from the negative to the positive end

17. - + Thomson’s Experiment Voltage source
Passing an electric current makes a beam appear to move from the negative to the positive end

18. - + Thomson’s Experiment Voltage source
Passing an electric current makes a beam appear to move from the negative to the positive end

19. - + Thomson’s Experiment Voltage source
Passing an electric current makes a beam appear to move from the negative to the positive end

20. Thomson’s Experiment
Voltage source
By adding an electric field

21. Thomson’s Experiment
Voltage source + -
By adding an electric field

22. Thomson’s Experiment
Voltage source + -
By adding an electric field

23. Thomson’s Experiment
Voltage source + -
By adding an electric field

24. Thomson’s Experiment
Voltage source + -
By adding an electric field

25. Thomson’s Experiment
Voltage source + -
By adding an electric field

26. Thomson’s Experiment Voltage source + -
By adding an electric field he found that the moving pieces were negative

27. Thomson’s Atomic Model
Thomson though electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model

28. Thomsom’s Model Found the electron
Couldn’t find positive (for a while)
Said the atom was like plum pudding
A bunch of positive stuff, with the electrons able to be removed

29. Mass of Electron
In 1909 Robert Millikan determined the mass of an electron with his Oil Drop Experiment
He determined the mass to be x kg
The oil drop apparatus

30. Millikan’s Experiment
Atomizer - +
Oil
Microscope

31. Millikan’s Experiment
Atomizer
Oil droplets - +
Oil
Microscope

32. Millikan’s Experiment
X-rays
X-rays give some drops a charge by knocking off electrons

33. Millikan’s Experiment+

34. Millikan’s Experiment- - + +
They put an electric charge on the plates

35. Millikan’s Experiment- - + +
Some drops would hover

36. Millikan’s Experiment- - - - - - - + + + + + + + +

37. Millikan’s Experiment- - + +
Measure the drop and find volume from 4/3πr3
Find mass from M = D x V

38. Millikan’s Experiment- - + +
From the mass of the drop and the charge on
the plates, he calculated the charge on an electron

39. Proton In 1886 Eugen Goldstein discovered the Proton
Proton is a positively charged subatomic particle found in the nucleus of a atom

40. Radioactivity Discovered by accident Bequerel Three types
alpha- helium nucleus (+2 charge, large mass)
beta- high speed electron
gamma- high energy light

41. Henri Becquerel (1896) Discovered radioactivity
spontaneous emission of radiation from the nucleus
Three types:
alpha () - positive
beta () - negative
gamma () - neutral

42. Ernest Rutherford
Rutherford ( ) proposed that all mass and all positive charges are in a small concentrated region at the center of the atom
He used the Gold-Foil Experiment to prove his theory
In 1911 he discovered the Nucleus
Nucleus- central core of an atom, composed of protons and neutrons
The nucleus is a positively charged region and it is surrounded by electrons which occupy most of the volume of the atom

43. Rutherford’s Experiment
Used uranium to produce alpha particles
Aimed alpha particles at gold foil by drilling hole in lead block
Since the mass is evenly distributed in gold atoms alpha particles should go straight through.
Used gold foil because it could be made atoms thin

44. HISTORY OF THE ATOM helium nuclei helium nuclei gold foil
They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back.

45. Florescent
Screen
Lead block
Uranium
Gold Foil

46. What he expected

47. Because

48. Because, he thought the mass was evenly distributed in the atom

49. What he got

50. How he explained it
Atom is mostly empty
Small dense, positive piece at center
Alpha particles are deflected by it if they get close enough +

51. +

52. Nuclear Atom Viewed in Cross Section
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53. Neutron James Chadwick (1891-1974) – discovered the neutron in 1932
Neutron is a subatomic particle with no charge but their mass is nearly equal to that of a proton

54. The Neutron
Because the protons in the atom could account for only about half the mass of most atoms, scientists knew there was another heavy particle in the nucleus.
Neutrons were proposed by Ernest Rutherford in 1907 (to account for a mass discrepancy in the nucleus) and discovered in 1932 by James Chadwick.
The neutron has about the same mass as a proton but with no charge.

55. James Chadwick (1932) Discovered neutrons
neutral particles in the nucleus of an atom
Joliot-Curie Experiment
based his theory on their experimental evidence

56. James Chadwick (1932) Joliot-Curie Experiment 5 MeV Alpha Particles
Gamma rays?
50 MeV Be
Energy values weren’t consistent.

57. Chadwick’s Explanation
James Chadwick (1932)
Chadwick’s Explanation
5 MeV
Alpha Particles
Neutrons
5 MeV Be
Energy values were consistent.
Also accounted for extra mass in the nucleus.

58. James Chadwick (1932)
Neutron Model

59. Bohr Model
Bohr changed the Rutherford model and explained how the electrons travel.
Bohr explained the following in his model:
Electrons travel in definite orbits with a certain energy around the nucleus. They must gain or lose energy in certain packages called Quanta
He explained that accelerating particles should get pulled into the nucleus but that does not occur because of the stability of the atom
His model was patterned after the motion of the planets around the sun. It is often called the Planetary model.

60. What do these particles consist of?
HELIUM ATOM
Shell
proton N + - + N -
neutron
electron
What do these particles consist of?

61. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels

62. The Quantum Mechanical Model
A totally new approach
Several different people made important contributions to the development of the model.

63. De Broglie De Broglie said matter could be like a wave.
De Broglie said particles like the electron could now how properties of both a waves and particles

64. Heisenberg Uncertainty Principle
It is impossible to know exactly the position and velocity (momentum) of a particle.
The better we know one, the less we know the other.
The act of measuring changes the properties.
More precisely the velocity is measured, less precise is the position (vice versa).

65. The Quantum Mechanical Model Schrodinger
Things that are very small behave differently from things big enough to see.
The quantum mechanical model is a mathematical solution. He applied these solutions to waves of electrons
It is not like anything you can see.

66. Erwin Schrödinger (1926) Quantum mechanics
electrons can only exist in specified energy states
Electron cloud model
orbital: region around the nucleus where e- are likely to be found

67. Electron Cloud Model (orbital)
Erwin Schrödinger (1926)
Electron Cloud Model (orbital)
dots represent probability of finding an e-

68. Modern View The atom is mostly empty space Two regions
Nucleus- protons and neutrons
Electron cloud- region where you have a chance of finding an electron

69. Quark
Protons & Neutrons can still be broken down into a smaller particle called the Quark
The Quark is held together by Gluons

70. Density and the Atom
Since most of the particles went through, it was mostly empty.
Because the pieces turned so much, the positive pieces were heavy.
Small volume, big mass, big density
This small dense positive area is the nucleus

71. Chapter 2: Atoms, Molecules, and Ions
Subatomic Particles
Protons and neutrons are located at the center of an atom called the nucleus.
Electrons are dispersed around the nucleus.
EOS
Chapter 2: Atoms, Molecules, and Ions

72. Atomic Particles Particle Charge Mass (kg) Location Electron -1
9.109 x 10-31
Electron cloud
Proton +1
1.673 x 10-27
Nucleus
Neutron
1.675 x 10-27

73. Subatomic particles Relative mass Actual mass (g) Name Symbol Charge
Electron e- -1
1/1840
9.11 x 10-28
Proton p+ +1 1
1.67 x 10-24
Neutron n0 1
1.67 x 10-24

74. X Symbols Mass number Atomic number
Contain the symbol of the element, the mass number and the atomic number
Mass
number X
Atomic
number

75. Sub-atomic Particles
Z - atomic number = number of protons determines type of atom
A - mass number = number of protons + neutrons
Number of protons = number of electrons if neutral

76. Symbols A X Z 23 Na 11

77. Atomic Structure Symbols
Proton = p+
Electron = e-
Neutron = n0
Atomic # - Subscript
Mass # - Superscript

78. Rules for Atomic Structure
Atomic # = # of Protons
# of Protons = # of Electrons
Mass # = # of Protons + # of Neutrons
# of Neutrons = Mass # - # of Protons
If you know the Mass # & Atomic # you know the composition of the element

79. Br Symbols 80 35 Find the number of protons number of neutrons
number of electrons
Atomic number
Mass Number 80 Br 35

80. Symbols
if an element has an atomic number of 34 and a mass number of 78 what is the
number of protons
number of neutrons
number of electrons
Complete symbol

81. Symbols if an element has 78 electrons and 117 neutrons what is the
Atomic number
Mass number
number of protons
Complete symbol

82. Example Atomic # Mass # Protons Electrons Neutrons K 19 11 5 16 17 46
Element
Atomic #
Mass #
Protons
Electrons
Neutrons K 19 11 5 16 17 46 23 35

83. Isotopes
Isotope- atoms that have the same number of protons but different number of neutrons
Since isotopes have a different number of neutrons the isotope has a different mass number.
Isotopes are still chemically alike because they have the same number of protons and electrons

84. Two Isotopes of Sodium
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85. Examples of Isotopes

86. Heavy Water
One ice cube is made with water that contains only the hydrogen-2 isotope. The other ice cube is composed of water with normal water which contains mostly hydrogen-1.
Which is which?
Figure 2.13

87. Isotopes
Isotopes are atoms of the same element with different masses.
Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C
© 2009, Prentice-Hall, Inc.

88. Naming Isotopes Put the mass number after the name of the element
carbon- 12
carbon -14
uranium-235

89. Electrical Charges
Electrical charges are carried by particles of matter
Atoms have no net electrical charges
Given the number of negative charges combines with the number of positive charges = Electrically Neutral
All elements are Electrically Neutral

90. Atomic Mass vs. Atomic Weight
Atomic Mass is for a single element
Most elements are Isotopes
How do we find their mass?
We use Atomic Weight

91. Chapter 2: Atoms, Molecules, and Ions
Atomic Masses
An atomic mass unit (amu) is defined as exactly one-twelfth the mass of a carbon-12 atom
1 u = × 10–24 g
The atomic mass of an element is the weighted average of the masses of the naturally occurring isotopes of that element
EOS
Chapter 2: Atoms, Molecules, and Ions

92. Measuring Atomic Mass Unit is the Atomic Mass Unit (amu)
One twelfth the mass of a carbon-12 atom
Each isotope has its own atomic mass. We need the average from the percent abundance
Each isotope of an element has fixed mass and a natural % abundance
You need both of these values to find the Atomic Weight

93. Calculating Atomic Weight
Cl amu and 75.77% abundance
Cl amu and 24.23% abundance
To solve for Cl-35
AMU x Abundance
x .7577 =
You solve for Cl-37

94. Atomic Weight Cont. Cl-37 AMU x Abundance 36.966 x .2423 = 8.957
Now you combine your two answers =
35.453
Look at Cl on the table. What is the Atomic Weight?

95. Example
Calculate the atomic weight of copper. Copper has two isotopes. One has 69.1% and has a mass of amu. The other has a mass of amu. What is the atomic weight???

96. Atomic Weight & Decimals
Atomic Weight- of an element is a weighted average mass of the atoms in a naturally occurring sample of an element
Atomic Weights use decimal points because it is an average of an element