1. The Atom & Modern Atomic Theory
Unit 3
The Atom & Modern Atomic Theory

2. Theories of the Atom
Early Models & Thoughts: Democritus Matter is made up of tiny particles called atoms. Smallest unit that retains the identity of the element

3. Theories of the Atom John Dalton: Hard Sphere model
The main points of Dalton's atomic theory were:
Elements are made of extremely small particles called atoms.
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
Atoms cannot be subdivided, created, or destroyed.
Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.

5. Theories of the Atom
J.J. Thomson: Cathode Ray Tube experiment. Found the atom to be a positively charged sphere, negatively charge electrons embedded within the sphere. Discovered the electron Plum pudding model of the atom.

6. Cathode Ray Tube Experiment

8. Gold Foil Experiment

9. Theories of the Atom
Ernest Rutherford Gold Foil experiment : Shot alpha particles at gold foil. Some went straight through others bounced back & deflected at odd angles. His 3 conclusions: 1. Atoms are mostly empty space, 2. Dense positive center. 3. Most of the mass of the atom is in the center (nucleus) Model of the atom has a nucleus with orbiting electrons.

10. Theories of the Atom
Bohr Model: Planetary Model of the atom The atom contained a nucleus w/ protons and Neutrons Each electron in an atom is allowed to be in certain orbits that correspond to different amounts of energy. Electrons can jump orbits (energy levels)

11. Theories of the Atom
Electron Cloud/ Wave Mechanical Model: Current model Protons & neutrons in the nucleus Electrons are in areas called orbitals and the model is based on the probability of finding an electron. Greater cloud density = increased probability of finding an electron. There is a greater chance of finding an electron in certain orbitals.

12. Theories of the Atom
The modern model of the atom is based on the work of many scientists over a long period of time.

13. Venus Explains the Atom

14. Atomic Structure

16. Atomic Structure
Moseley: Devised the concept of atomic number Said atomic number of an atom is equal to the number of protons in the atom's nucleus.

17. The Periodic Table can tell you about an atom.

18. Atomic Structure The atomic number is the ID of the element unique
to each element). 1 2 3
3 Important Ideas:

19. Atomic Structure The atomic number is the ID of the element unique
to each element).
In a neutral atom: # of Protons = # of electrons 1 2 3
3 Important Ideas:

20. Atomic Structure The atomic number is the ID of the element unique
to each element).
In a neutral atom: # of Protons = # of electrons
When an atom gains or loses an electron, it is now
an ion 1 2 3
3 Important Ideas:

21. Atomic Structure
Charge = # of protons – # of electrons

22. Atomic Structure
Example 1. What is the name of the element whose atoms have 11 protons?
Example 2: How many protons and electrons in an oxygen atom?
Example 3: How many electrons and protons in Na+?

23. Atomic Structure
Example 4: Write the symbol for the ion with 9 protons and 10 electrons.
Example 5: Write the symbol for the ion with 13 protons and 10 electrons.
Example 6: How many protons & electrons in S-2 ion?

24. Atomic Structure, Part II
Reflect:
Why use the a.m.u. scale to measure the mass of an atom?

25. Atomic Structure, Part II
a.m.u. : Atomic Mass Unit
 To determine the number of neutrons:
# of neutrons = mass number – atomic number

26. Atomic Structure, Part II
Example 1: How many neutrons, protons, and electrons are in an ion?
Example 2: Determine the number of protons, electrons, and neutrons this ion:

27. Atomic Structure, Part II
Example 3: Write the complete chemical symbol of an element with 21 protons, 24 neutrons, and 18 electrons.
Example 4: Write the complete chemical symbol of an element with 53 protons, 74 neutrons, and 54 electrons.

28. Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron 5 Na 11 +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar
Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron 5 Na 11 +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar

29. Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron 5 Na 11 +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar
Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron B 11 5 6 Na +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar

30. Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron 5 Na 11 +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar
Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron B 11 5 6
Sodium Na 23 12 10 +1 1 17 18 Al 13 +3 9
Oxygen O -2
Calcium 20
Sulfur 32 Ar

31. Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron 5 Na 11 +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar
Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron B 11 5 6
Sodium Na 23 12 10 +1
Hydrogen H 1 17 18 Al 13 +3 9
Oxygen O -2
Calcium 20
Sulfur 32 Ar

32. Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron 5 Na 11 +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar
Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron B 11 5 6
Sodium Na 23 12 10 +1
Hydrogen H 1
Chlorine Cl 35 17 18 -1 Al 13 +3 9
Oxygen O -2
Calcium 20
Sulfur 32 Ar

33. Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron 5 Na 11 +1 1 17 18 Al 13 +3 9 10
Oxygen O -2
Calcium 20
Sulfur 32 Ar
Element
Symbol
Mass
Number
# of
Protons
Neutrons
Electrons
Ion
Charge
Boron B 11 5 6
Sodium Na 23 12 10 +1
Hydrogen H 1
Chlorine Cl 35 17 18 -1
Aluminum Al 27 13 14 +3
Fluorine F 19 9
Oxygen O 16 8 -2
Calcium Ca 40 20 +2
Sulfur S 32
Argon Ar 22

34. Atomic Structure, Part II
Isotope: Same element (same # of protons), but different number of neutrons. (So the mass of each isotope is different).
Remember: atomic mass is the weighted average of masses of these isotopes.

35. Atomic Structure, Part II
Example 1: Chlorine has two isotopes, chlorine –35 and chlorine
37. Approximately 75% is Cl-35 and 25% is Cl-37.
Which isotope should the weighted average be closer to?__________
(isotope mass)(% abundance) + (isotope #2 mass)(% abundance) ……..
100

36. Atomic Structure & Isotopes
Example 2: Boron-10 has a relative abundance of 19.78%. Boron-11 has a relative abundance of 80.22%. Calculate the average atomic mass.
(isotope mass)(% abundance) + (isotope #2 mass)(% abundance) ……..
100

37. Atomic Structure & Isotopes
Example 3: The relative masses and abundances of Iodine are given below. Calculate the average atomic mass.
(isotope mass)(% abundance) + (isotope #2 mass)(% abundance) ……..
100
Isotope
% Abundance
I-127
80.0%
I-126
17.0%
I-128
3.0 %

38. The Wave-Mechanical Model of the Atom
Reviewing the Atom
The Wave-Mechanical Model of the Atom

39. Reviewing the Atom Protons Neutrons Electrons Charge Mass Where found
If you change the number of these particles, you get
How to figure out how many an atom has

40. Bohr Model & Electron States

41. Bohr Model & Electron States
Electron Configurations:
Ordered sequence showing how many
electrons are in each Principle Energy Level for a particular atom.
Given in your reference table, starting from energy level 1
Note the * for elements 72 and up

42. Bohr Model & Electron States
Valence Electrons: Outer most electrons found in the last Principle Energy Level that are used for bonding.
Hint: Look at your periodic table. In each box is the Bohr e- configuration. The last number are the valence electrons.
Ex: or 2-4 or 2-7

43. Bohr Model & Electron States

44. Bohr Model & Electron States
Practice: Write the electron configurations of the following atom and then state how many principle energy levels each element has.
Element
Electron Configuration
# of Principle Energy Levels
Neon
Strontium
Bromine
Magnesium
Copper

45. Bohr Model & Electron States
Directions: Draw a Bohr model of the following neutral atoms using the electron configuration on the Periodic Table. Include the number of protons and neutrons in the nucleus and write the electron configurations in the box below.

46. Bohr Model & Electron States
Directions: Draw a Bohr model of the following neutral atoms using the electron configuration on the Periodic Table. Include the number of protons and neutrons in the nucleus and write the electron configurations in the box below.

47. Bohr Model & Excited States

48. Bohr Model & Electron States
Excited State: An electron has jumped to a higher principle energy level. Ex: Magnesium
Ground State: All electrons are in their lowest possible state. Ex: Magnesium

49. Bohr Model & Electron States

50. Electromagnetic Spectrum

51. In different scenarios, light behaves as a
continuous wave AND as individual
particles.

52. Light & Quantum Theory Much of the understanding of how
electrons behave in
atoms comes from
studies of how light
interacts with matter.

53. One photon interacts with 1 atom of matter at a time.

54. Depending on what orbital the electrons jumped to and fell from, there will be photons of different energy emitted. We can see the color of those photons.

55. Electron Transitions
Key Idea #1
Longer wavelengths (red end of spectrum) have less energy.
Shorter wavelengths (violet end of spectrum) have more energy

56. Electron Transitions
Key Idea #2
When electrons move up energy levels, this requires energy (remember – protons and electrons attract!). (Ground  Excited)
When electrons move down energy levels (Excited  Ground), they emit energy as light waves
Different transitions result in different energy emissions (i.e. different colors)

57. Electron Transitions
Key Idea #3
Each atom emits a specific number and pattern of bands based on its electron configuration when you pass its light through a prisim.

59. Bohr Model & Electron States
Ground State: When electrons are at their lowest possible energy state.
Example electron configuration:
2 – 8 – 1
Excited State: When an electron gains energy (heat or electricity) and jumps to a higher PEL.
Example electron configuration:
2 – 7 – 3

60. Bohr Model & Electron States
How Light is Produced:
Electron(s) absorb energy in the form of heat or electricity
Electron(s) then jump up to a higher PEL & are in the excited state.
The electron(s) then fall back down to the ground state.
When the electron(s) falls back to the ground state it emits a photon of light (energy absorbed = energy emitted).

61. Bohr Model & Electron States
Flame Test:
Metallic salts are put in a flame
A specific color of light is given off

62. Bohr Model & Electron States
Bright Line Spectrum:
Electricity is run through a gas tube
A specific color of light is given off
Light is passed through a prism
A unique patter of colored lines are produced
Ca be used to ID an element!

63. Red light has the least energy.
Violet light has the greatest energy.

64. Review of Light & Quantum Theory
Much of the
understanding of how
electrons behave in
atoms comes from
studies of how light
interacts with matter.

65. In different scenarios, light behaves as a
continuous wave AND as individual
particles.

66. One photon interacts with 1 atom of matter at a time.
Energy, wavelength,
frequency are
associated with each
photon.

67. Bohr took his model of the atom and stated that electrons can absorb certain amounts of energy and then release that energy in the form of a photon.

68. Review of Light & Quantum Theory
When all electrons are at their lowest possible energy, it is called the Ground State. When the electron absorbs a specific amount of energy, it moves to a higher orbital and is in the excited state.

69. Depending on what orbital the electrons jumped to and fell from, there will be photons of different energy emitted. We can see the color of those photons.

70. Every atom has a range of colors emitted
when their electrons get excited and fall
back down.
Sodium: Orange
Copper: Green
Strontium: Drk. Orange
Mercury: Blue

71. The colorful lights of fireworks are emitted by "excited" atoms; that is, by atoms that have absorbed extra energy.
Different compounds within the fireworks give them their signature colors.

72. Light can be further analyzed by passing it through a prism
Light can be further analyzed by passing it through a prism. White light when passed through a prism breaks up into all the colors of the rainbow (ROYGBIV)

73. When the color of light emitted by an atom is passed through a prism, only certain colors are observed. These specific colors are called a Bright Line Spectrum. Each line spectra is different for every element and can serve to identify a particular element.
Hydrogen:

74. Electron Configurations
Reflect:
Think of electrons, all negatively charged, in a cloud around the nucleus. How can they be organized to minimize repulsion?

75. Electron Configuration3
Wave Mechanical Model: Electrons are organized in ‘orbitals’, or areas of probable electron location

76. Electron Configurations
Heisenberg Uncertainty Principle: We cannot know both the direction and location of electrons – measuring one changes the other

77. Organized by: which is like…
Principle Energy levels  Floors in an apartment
(7 total, numbered 1, 2, etc)
Sublevels  The actual apartment
(Named s, p, d, f)
Orbitals  Rooms in the apartment
(Can hold only 2 electrons)

78. 3 Orbitals
Holds 6 e- P S
Orbital
Holds 2 e- d
5 Orbitals
Holds 10 e-

79. Each sublevel holds a different number of orbitals:
s = 1 orbital = 2 electrons
p = 3 orbitals = 6 electrons
d = 5 orbitals = 10 electrons
f = 7 orbitals = 14 electrons

80. Electron Configurations
Principal Energy Level
Type of Sublevel
(shape of orbitals)
Number of Orbitals
in Sublevel
Electron Capacity of
Sublevel

81. Electron Configurations

82. Electron Configurations
Sublevel: Sublevels are contained within a PEL. There are 4 “shapes”/types that are labeled with letters (s, p,d, f).
*sublevels are regions of space where the electrons w/ specific energies “live”.
s sublevel

83. Electron Configurations
p sublevel

84. Electron Configurations
d sublevel

85. Electron Configurations
f sublevel

86. Electron Configurations
Orbital: Orbitals make up the sublevels. Each orbital can hold only 2 electrons max.
‘s’ sublevel has 1 orbital. 2 e- max
The ‘p’ sublevel contains 3 orbitals. Each holds 2 electrons for a total of 6 electrons in the ‘p’ sublevel
‘d’ sublevel has 5 orbitals. 10 e- max
‘f’ sublevel has 7 orbital2. 14 e- max

87. Electron Configurations
When writing an electron configuration, the general format is to list the information in the following order: PEL then sublevel then # of Electrons 1,2,3… s,p,d,f write as an exponent You start with the innermost electrons and continue describing the location of all electrons as you move away from the nucleus. General Trend: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

88. Analogy
Electron Apartment Complex: Each floor represents a PEL. On the first floor (n =1)there is only 1 apartment. Apartment ‘s’. It is the cheapest apartment (lowest in energy). It is a studio (one room) and can only hold 2 occupants (electrons). s
n = 1

89. Analogy
Electron Apartment Complex: On the second floor (n = 2)there are 2 apartments. One studio apartment (s) that can hold 2 people (which costs more since it is on the 2nd floor) and a ‘p’ apartment that has 3 bedrooms (2 people to a room) for a total of 6 occupants, which costs more than the studio because it is bigger. (*cost = energy) s p p p
n = 2 s

90. Analogy
Electron Apartment Complex: On the 3rd floor (n = 3)there are 3 apartments. One studio apartment (s) that can hold 2 people (which costs more since it is on the 3nd floor) and a ‘p’ apartment that has 3 bedrooms (2 people to a room) for a total of 6 occupants, which costs more than the studio because it is bigger. Then there is the ‘d’ apartment that has 5 bedrooms, 2 people to a room for a total of 10 occupants. s p p p d d d d d
n = 3 s p p p s

91. Analogy
Electron Apartment Complex: On the 4th floor (n = 4) there are a total of 4 apartments. 1 studio, 1 three bedroom, a 5 bedroom and lastly the Penthouse with a total of 7 bedrooms (2 occupants to a room) that can house up to 14 people. This apartment costs the most to occupy. (The s, p and d apartments are also more expensive than the ones on lower floors).
n = 4 s p p p d d d d d f f f f f f f s p p p d d d d d s p p p s

92. Analogy

93. Practice:
Be: C: F: Ne: * Exponents should add up to the total # of electrons in the atom!

95. Using the Periodic Table to Write Sublevel Electron Configurations

96. Using the Periodic Table to Write Sublevel Electron Configurations

98. Practice Examples
H: _______________________________________________ He: _______________________________________________ Li: _______________________________________________ Be: _______________________________________________ B: _______________________________________________ C: _______________________________________________ N: ______________________________________________

99. Orbital Notation
Electron Configurations: 1s2 Orbital Diagrams:

100. Orbital Notation
Orbital Diagram: A pictorial representation (using arrows to represent electrons) of sublevel electron configurations. Examples of Orbital Diagrams: Orbital diagram for Oxygen

101. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Aufbau: Electrons are added 1 at a time to the lowest energy orbitals available until all the electrons are used.

102. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Pauli Exclusion Principle: An orbital can only hold a maximum of two electrons.

103. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Hund’s Rule: Electrons of the same energy spread out before pairing up.

104. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

105. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

106. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

107. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

108. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

109. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

110. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

111. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

112. Orbital Notation
Rules for Electron Configurations & Orbital Diagrams: Example: Orbital Diagram for Fluorine 9 electrons 1s2 2s2 2p5

113. Practice
Sodium:

114. Practice
Sodium +1 ion:

115. Practice
Chlorine:

116. Atomic Structure Protons Neutrons Electrons Found in the nucleus
(+) charge
Mass of 1 amu
# of protons = Atomic #
# of protons = # of
electrons in a NEUTRAL
atom
# of protons = the
nuclear charge.

117. Atomic Structure Protons Neutrons Electrons Found in the nucleus
(+) charge
Mass of 1 amu
# of protons = Atomic #
# of protons = # of
electrons in a NEUTRAL
atom
# of protons = the
nuclear charge.
(0) charge – neutral
mass of 1 amu

118. Atomic Structure Protons Neutrons Electrons Found in the nucleus
(+) charge
Mass of 1 amu
# of protons = Atomic #
# of protons = # of
electrons in a NEUTRAL
atom
# of protons = the
nuclear charge.
(0) charge – neutral
mass of 1 amu
Found in orbitals ( e- clouds) surrounding nucleus
( - ) charge
Mass 1/1836th of a proton
# of e- = # of protons in
a NEUTRAL atom