2. Brief Timeline of Atomic Theory

3. Democritus
400BC
Greek philosopher

4. Hard Particle (Cannonball)Theory
Proposed that they world was made up of tiny, indivisible particles moving through a void of empty space
“atom” comes from the Greek word “atomos”, meaning indivisible (cannot be divided)

5. John Dalton
1808 AD
First modern atomic theory

6. Daltons Atomic Theory
All matter is composed of tiny, indivisible particles called atoms
All atoms of an element are identical
Atoms of different elements are all different
Atoms combine in simple ratios to form compounds

7. J.J. Thomson 1897-1904 “Plum Pudding Model”
Cathode Ray tube experiment
demo

8. demo

9. Cathode Ray Tube
Thompson showed that cathode rays (electrons) were composed of negatively charged particles that separated from the gas atoms inside the tube
Significant because: this meant that atoms are not hard, indivisible particles. Atoms are composed of smaller “subatomic” particles

10. Thomson’s Plum Pudding Model
The atom was a hard sphere that was positively charged with negatively charged electrons that “dotted” the atom like raisins in plum pudding

11. The discovery of radioactivity
Henri Becquerel
1896
Discovered that uranium
ore released rays that could
expose photographic film

12. The discovery of radioactivity
Marie & Pierre Curie
Extracted 2 new elements
from uranium (U)ore:
radium (Ra) and
polonium (Po)
Marie Curie

13. Magnetic Field Experiment Gold Foil Experiment (1911)
Ernest Rutherford
Magnetic Field Experiment
Gold Foil Experiment (1911)
Was able to separate radioactive rays into 2 types: alpha (α) & beta (Β)
Determined that α rays were composed of helium nuclei (He +2 charge)
Lead to discovery of the nucleus, as a positively charged center of atom, containing the mass
Most of the atom is negatively charged empty space, electrons are outside the nucleus

14. Magnetic Field Experiment

16. Gold Foil Experiment

17. Gold Foil Experiment

18. Gold Foil Experiment

19. Gold Foil Experiment

20. Rutherford’s Atomic Model

21. Rutherford’s “Nuclear Model”
Most of the atom is negatively charged empty space, surrounding a small, positively charged nucleus, containing most of the mass of the atom

22. Modern Theory of Atomic Structure
Developed by Niels Bohr, based on the science of nuclear physics
Bohr determined that an element's position on the periodic table was related to its electron configuration.

23. Electron configuration
 Electron configuration – shows how many electrons are in each energy level or “ring”
Ex: Carbon 2-4

24. Bohr’s Planetary Atomic Model
Niels Bohr (1922)
Determined that electrons
rotate around the nucleus
in discrete paths or rings

25. Planetary Model of Atomic Structure

26. Wave-Mechanical Model
Current (modern) theory of atomic structure
Moseley used x-ray analysis to calculate an integer for each element: these integers are the atomic numbers

27. Wave-Mechanical Model
There is a tiny, dense positively charged nucleus at the center of a huge negatively charged electron cloud

28. Wave-Mechanical Model

29. Orbital
Region of probability of finding an electron

30. “The whole point:”
The modern model of the atom is the result of many investigations that have been revised over a long period of time by many scientists
Atomic theory song

31. Place the models of atomic structure in order from earliest to the modern theory:

32. Basic Atomic Structure
The nucleus occupies less than 0.01% of the total volume of an atom but accounts for 99.97% of its mass. Thus most of an atom is EMPTY SPACE where the ELECTRONS are found, this is called an ELECTRON CLOUD.
One atomic mass unit is 1/12TH THE MASS OF A CARBON-12 ATOM. This is the standard by which the masses of all other elements are determined. It is abbreviated “u”.

33. Start here with lecture notes!!

34. Subatomic Particles Proton p+ In Nucleus +1 1 Neutron n0 In nucleus
Symbol
Location
Charge
Mass (amu)
Mass (g)
Proton p+ In
Nucleus +1 1
Neutron n0
In nucleus
Electron e-
Outside
nucleus -1

35. Nucleons: Def: particles found in the nucleus : protons & neutrons
Number of Nucleons = Atomic Mass

36. Nuclear Charge: Def: equal to the number of protons
(b/c only protons have charge, neutrons are neutral)
For this atom the nuclear charge would be +6, because there are 6 protons

37. protons
neutrons
protons
neutral
protons
electrons 6 6 6

38. 2713Al 3517Cl 11H 20782Pb Signature Atomic Number Mass Number
Nuclear charge
# of PROTONS
# of
NEUTRONS
# OF ELECTRONS
2713Al
3517Cl
11H
20782Pb

39. Signature
Atomic Number
Mass
Number
Nuclear charge
# of PROTONS
# of
NEUTRONS
# OF ELECTRONS
2713Al 13 27
+13 14
3517Cl 17 35
+17 18
11H 1 +1
20782Pb 82
207
+87 87
125

40. The only number that never changes for an element is
ATOMIC
NUMBER !!

41. 24?N 64?Cu Atomic number:___________ # protons: ______________
Atomic number:___________
# protons: ______________
# electrons: _____________
Mass number: ___________
# neutrons: _____________
Name-mass:_______________________________
64?Cu

42. 147N 3115P 2713Al 10847Ag Atomic number:___________
Atomic number:___________
# protons: ______________
# electrons: _____________
Mass number: ___________
# neutrons: _____________
Name-mass:______________________________ _
3115P
Name-mass:_______________________________
2713Al
10847Ag

43. Atomic Structure Practice: Determining Subatomic Particles
1.) Complete the following definitions:
a. ATOM
b. ATOMIC NUMBER – equals the number of ________ in an atom
c. MASS NUMBER – equals the number of ________ + _________
d. Charge of a PROTON - ________ Mass of a PROTON- ________
e. Charge of a NEUTRON - ______ Mass of a NEUTRON -_______
f. Charge of an ELECTRON- _____ Mass of an ELECTRON- ______

44. WORD BANK: proton, neutron, electron, nucleus, electron cloud

45. For the atom shown above:
f. Number of Protons= _______________
g. Atomic Number =________________
h. Number of Neutrons= _____________
i. Mass Number = _______________
Using the ATOMIC NUMBER, identify the element name___________ and SYMBOL _____

46. Atomic Structure 1 Name Symbol Atomic Number Mass Number Charge # of
Protons
# of Neutrons
# of electrons
199F
Helium-4 11 23
Nitrogen-14 14 32 16
6429Cu
 Manganese-60
 6025Mn 25
60 
25  35
Barium-137
 13756Ba
56 
137  81 56
 Iodine-131
 13153I 53
131 78
53 
 Iodine-127
 12753I
127  74

47. Atomic Structure 1 Name Symbol Atomic Number Mass Number Charge # of
Protons
# of Neutrons
# of electrons
199F
Helium-4 11 23
Nitrogen-14 14 32 16
6429Cu 25 35 81 56 53
131 74

48. **Shade the columns representing the nucleons light blue
Phosphorus-32
146C
Potassium-39 16 8
5626Fe 18 40 29 35 79
197
2412Mg
**Shade the columns representing the nucleons light blue

49. Changes in number of subatomic particles
Isotopes
Ions
Change in number of neutrons
Same atomic number, different mass
Same number protons, different number neutrons
Change in number of electrons
A cation is positive ion, results from loss of electrons, reducing radius
An anion is negative ion, results from gain of electrons, increasing radius

50. Name-mass:_______________________________
Atomic number:___________
# protons: ______________
# electrons: _____________
Mass number: ___________
# neutrons: _____________
Name-mass:_______________________________
2311?
??Li
??I

51. ISOTOPE
Forms of the same element having different mass due to different number of neutrons.
Indicated by “element name-mass”

52. 158O 168O Name: _______________ Mass: ________________
Protons: ______________
Neutrons: _____________
Name: _______________
Mass: ________________
Protons: ______________
Neutrons: _____________

53. Practice: Name Symbol Atomic # Mass # # Protons # Neutrons # Electrons
235U
238U
Carbon-12
Carbon-13

54. The mass on the periodic table is the MOST ABUNDANT mass!
** You can estimate which isotope is found in the highest abundance as the one with a mass closest to the mass listed on the periodic table
Example:
Chlorine-35 mass g
Chlorine-37 mass g
Look on the periodic table for the mass of chlorine____________________________
The more abundant isotope has a mass closer to the mass given on the periodic table_____________

55. Practice:
Which isotope of silicon would be found in the highest percentage?
2814Si, mass Si, mass Si, mass
Why? _______________________________________ _______________________________________

56. Atomic Structure 2 Isotopic Notation Number of protons
Number of protons
Number of neutrons
Number of electrons
Mass number
1.Oxygen-16
O-16
16O
2.Oxygen-18 3.
Ar-40 4. 18 5. 16 32 6.
34S 7. 19 20 8. 41
9. Iron-
10.
57Fe
11. 26
12.
Ne-20
13. 10 22
14.Hydrogen- 1
15.
H-2
16. 3H

57. 2.) Calculate the weighted average of the following naturally occurring isotopes. SHOW ALL WORK!
a.) 95.50%7Li & 7.50% 6Li
d.) %14N & 0.37%15N
b.)80.20%11B & 19.80%10B
e.) 78.9%24Mg, 10.00%25Mg, & 11.01%26Mg
c.)95.02%32S, 0.75%33S, & 4.21%34S
f.) 92.23%28Si, 4.67%29Si, & 3.10%30Si

58. IONS
A charged part of an atom, resulting from the loss or gain of electrons
VALENCE electrons: outermost electrons, the last number in an electron configuration
KERNEL electrons: all electrons except valance electrons

59. Electron configuration
 Electron configuration – shows how many electrons are in each energy level or “ring”
Ex: Carbon 2-4

60. Electron configuration of sodium:

61. 2 diagrams of atomic structure:
Bohr diagrams
Lewis electron dot diagrams
Bohr realized that the rows on the periodic table corresponded to the number of shells of electrons
Lewis realized that the groups/families on the periodic table correspond to the number of valence electrons
This model shows the nucleus, indicating the number of protons and neutrons, surrounded by rings, representing each energy level
This model shows the element symbol surrounded by dots, representing the valence electrons. You must place one dot at each (3, 6,9,12 o’clock) location before “doubling up” (exception: Helium)
189F electron configuration 2-7
F electron configuration 2-7

62. Bohr Atomic Structures1 18
Bohr Atomic Structures
tables to fill in the electron 4
configurations, as shown, then
draw the Bohr Atomic Structure
for each element 1-20. 2 13 14 15 16 17 7 9 11 12 19 20 3 5 6 8 10
2-1
2-2
2-3
2-4 23 24 27 28 31 32 35 40
2-8-1
2-8-2
2-8-3
2-8-4 39
Rules:
1.) Show placement
2.) The nucleus is
3.) Indicate the number
of ALL electrons
represented by a center
of electrons in each
circle showing the
energy level, by writing
*use atomic #
# of protons & the
the number on each ring.
OR the entire
# of neutrons
electron configuration
** closest to nucleus is 1st

64. 2 Main Types of Ions: anion A negative ion Ex: Cl-, O-2 cation
A positive ion
Ex: Na+, Al+3

65. The octet rule
Atoms will gain or lose electrons in order to have a full valence shell of 8 electrons.
Exception: Helium can have a maximum of 2 valance electrons

66. When an atom gains 1 or more electrons
It becomes a negative ion and it’s radius increases. A negative ion is an anion.

67. When an atom loses 1 or more electrons
It becomes a positive ion and it’s radius decreases. A positive ion is a cation.

69. Change ending of element to “-ide”
CATION
ANION
Definition
positive ion
negative ion
Results from
Loss of electron(s)
Gain of electron(s)
Indicated by
(+) charge
(-) charge
What happens to radius???
Gets smaller
Gets bigger
Na Na+
Naming
“Element name-ion”
Change ending of element to “-ide”
Lewis Dot Structure
[Na] + ..
[:.F.:]-

70. How to predict if an element will form an anion or cation:
The “electron clock”:
8/0
7 1
6 2
5 3 4
# valance electrons

71. Atomic Structure 3: Predicting Ions
Element
Electron configuration
Lewis dot structure of atom
Lose or gain electrons?
How many electrons lost or gained?
Ionic Charge **
Lewis dot structure of ion
Radius increase or decrease? F
2-7
gain 1 -1
increase Mg
2-8-2
lose 2 +2
decrease O Al N Fr C

72. Electron configuration Lewis dot structure of atom
Element
Electron configuration
Lewis dot structure of atom
Lose or gain electrons?
How many electrons lost or gained?
Ionic Charge **
Lewis dot structure of ion
Radius increase or decrease?
2-8-7
2-8-6
2-8-5
2-3
**In the “ionic charge” column only: shade the cation charges red and the anion charges blue

73. Atomic Structure 4 3517Cl 2311Na 94Be 6530Zn 147N 3216S 2010Ne 12753I
# of Protons
# of Neutrons
# of Electrons
Nuclear Charge
Bohr Diagram of Atom
Lewis Dot of Atom
Predict
Ionic Charge
Lewis Dot of Ion
Name of Ion ex
3517Cl 17 18
+17 Cl -1
Chloride 1
2311Na 2
94Be 3
6530Zn 4
147N 5
3216S 6
2010Ne 7
12753I 8
10847Ag 9
7031Ga 10
126C

74. Atomic Spectra

75. Radiant Energy
Energy that travels through space as electromagnetic waves at the speed of light

76. Electromagnetic Spectrum
Includes all types of radiant energy from gamma rays (hi E) to radiowaves (lo E)
Visible light is only a small portion of the spectrum

77. 1 photon = 1 quantum
Quanta: tiny packets of energy released or absorbed by objects
*Einstein and Plank determined that energy is released or absorbed in a continuous flow of small packets or quantum/photons

78. Release or Absorption of Energy:

80. Higher energy levels
(excited state)
Electrons absorb energy when jumping to
Electrons release energy when falling to
Lower energy levels
(ground state)

82. Bohr used the emission spectrum as proof of planetary model
But his model only works for hydrogen because he didn’t account for electrons moving between energy levels

83. Spectral Lines
Characteristic wavelengths (λ) of photons of energy released as electrons fall from hi to lo energy

84. Spectral lines demo: Salt of Element Color of Flame Strontium Chloride
Strontium Chloride
Barium Chloride
Copper (II) Chloride
Lithium Chloride
Potassium Chloride
Identity
Unknown Element
Unknown Mixture

85. Emission Spectrum:

86. Each element has it’s own characteristic spectrum:

87. Compare H & He:
hydrogen
helium

88. Because electrons do move between energy levels, emitting “spectral lines”, we had to change our view of atomic structure:

89. Excited State Electron Configurations
Occurs when elements absorb energy and jump to a higher energy level.
** it will not look like it is written on periodic table, be sure they add to the correct number!
Ground state: 2-8-1
Excited state : 2-7-2

94. “Crib Sheet” #p+ = atomic number *#n0 = mass-atomic number
#e- = #p+ - charge (use the sign of the charge)
Isotope: same #p+, different #no OR same atomic number, different mass
To calculate weighted average: (%/100 x atomic mass) + (%/100 X atomic mass) + …..
*Ion: same # p+, different #e-
Charge= #p+ - #e-

95. Atomic Structure Review p. 1711 9 43 92
118 13 4
9. Br
10. C
11. Sn
12. Zn
13. Cl
14. 40
15. 16

96. Atomic Structure Review p. 17
17)
=(.789x24)+(.10x25)+(.1101x26) = =
= 24.30g
16.)
=(.925x7) + (.0750x6) =
=6.925
=6.93g

97. Atomic Structure Review p. 18
18.) 2-8-1
19.) Na
20.) 2-7-2
21.) 19
22.) 1
23.) Y
24.) Ar
25.) Not possible
27.) as electrons fall from excited state to ground state energy is released as radiant energy (spectral lines).
28.) you can ID the gas element using spectral line analysis.
29.) electrons are negatively charged particles. B has 5 e-, its e- config. is 2-3, with 2 e- in the 1st energy level and 3 e- in the 2nd (valence) level

98. Atomic Structure Review MC?s
1.) ) 1
2.) ) 4
3.) ) 1
4.) ) 3
5.) ) 4
6.) ) 3
7.) ) 2
8.) ) 2
9.) ) 3
10.) 3
11.) 4
12.) 3 pg 19-20
1.) ) 2
2.) ) 4
3.) ) 3
4.) ) 3
5.) ) 2
6.) ) 1
7.) ) 1
8.) ) 4
9.) ) 4
10.) ) 2
11.) ) 3
12.) 1 pg 21-22

99. Atomic Structure Review p. 23
1.) 19p, 20n, 18e 2.) 9p,10n,10e
3.) 5p,6n,2e 4.) 15p,16n,18e
5.) 16p,16n,18e 6.) 14p,14n,10e
7.) 7p,7n,10e 8.) 20p,20n,20e
9.) 37p,48n,36e 10.) 53p,75n,54e
11.) 30p,35n,28e 12.) 6p,6n,10e