1. Defining Atoms & Electrons in Atoms

2. Originated idea of the atom
Democritus ( BC)
Originated idea of the atom

3. 1803 Dalton’s Atomic Theory
John Dalton ( )
1803 Dalton’s Atomic Theory

4. Modern Atomic Theory Dalton said:
Several changes have been made to Dalton’s theory.
Dalton said:
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties
Modern theory states:
Atoms of an element have a characteristic average mass which is unique to that element.

5. Modern Atomic Theory Dalton said:
Atoms cannot be subdivided, created, or destroyed
Modern theory states:
Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

6. J.J. Thomson (1856-1940) 1897 Discovered the electron
(“plum pudding” model)

7. Mass of the Electron
1909 – Robert Millikan determines the mass of the electron.
The oil drop apparatus
Mass of the electron is
9.109 x kg

8. Conclusions from the Study of Electrons
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons.
Electrons have so little mass that atoms must contain other particles that account for most of the mass

9. Rutherford (1871-1937) 1911 Discovered the nucleus
(gold foil experiment)

10. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

11. Since opposite charges attract each other, why don’t the electrons fall into the nucleus?

12. 1913 proposed Planetary model
Niels Bohr ( )
1913 proposed Planetary model

13. The Bohr Model of the Atom
I pictured electrons orbiting the nucleus much like planets orbiting the sun.
But I was wrong! They’re more like bees around a hive.
WRONG!!!
Neils Bohr

14. The Quantum-Mechanical Model
1926
The Quantum-Mechanical Model
Based upon the work of several men, a new mathematical model was developed to describe the structure of the atom.
Louis de Broglie ( )
Werner Heisenberg ( )
Erwin Schrodinger ( )

17. Atomic Number 6 6 15 15 79 79 In a neutral atom:
Atomic number of an element is the number of protons in the nucleus of each atom of that element.
Element
Atomic #
# of protons
# of electrons
Carbon 6
Phosphorus 15
Gold 79 6 6 15 15 79 79
In a neutral atom:
# electrons = # protons

18. The Atomic Scale
Most of the mass of the atom is in the nucleus (protons and neutrons)
Electrons are found outside of the nucleus (the electron cloud)
Most of the volume of the atom is the electron cloud.

19. Atomic Particles Particle Charge Mass (g) Location Electron e- -1
9.109 x 10-28
(1/1840 amu)
Electron cloud
Proton p+ +1
1.673 x 10-24
(1 amu)
Nucleus
Neutron n
1.675 x 10-24

20. Reading the Periodic Table3 Li
6.941
Lithium
Atomic Number
# p+
# e- (in a neutral atom)
Element Symbol
Element Name
Atomic Mass
# p+ + # n0
# n0 = Atomic Mass - #p+

21. The Quark…
Oops…
wrong Quark!

22. The Atomic Scale
Most of the mass of the atom is in the nucleus (protons and neutrons)
Electrons are found outside of the nucleus (the electron cloud)
Most of the volume of the atom is empty space
“q” is a particle called a “quark”

23. About Quarks… Protons and neutrons are NOT fundamental particles.
Protons are made of two “up” quarks and one “down” quark.
Neutrons are made of one “up” quark and two “down” quarks.
Quarks are held together
by “gluons”

26. Isotopes Elements occur in nature as mixtures of isotopes.
Isotopes are atoms of the same element that differ in the number of neutrons

27. Mass Number Mass # = p+ + n0
Mass number is the number of protons and neutrons in the nucleus of an isotope.
Mass # = p+ + n0

28. C C Cl Symbols of Isotopes Atomic Mass 12 Carbon-12 Atomic Number 614 6
Atomic Number
Chlorine-35
Atomic Mass Cl 35 17

29. Mass Number Mass # = protons + + neutons0 8 8 16 Arsenic -75 33 75
Element p+ n0 e-
Mass #
Oxygen-16 8 33 42
-31 15 29 64
Neon-20 10 8 8 16
Arsenic
-75 33 75
Phosphorus 16 15 31 35
Copper-64 29 10 10 20

30. Mass Number Mass # = p+ + n0 O As P Isotope p+ n0 e- Mass # Oxygen -10 - 33 42
-31 15 18 O 18 8 8 8 18 As 75 33
Arsenic 75 33 75 P 31 15
Phosphorus 16 15 31

31. Isotopes…Again (must be on the test)
Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

32. Composition of the nucleus
Atomic Masses
Atomic mass is the weighted average of all the naturally occuring isotopes of that element.
Multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products.
Isotope
Symbol
Composition of the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
7 neutrons
1.11%
Carbon-14
14C
8 neutrons
<0.01%
Carbon =

33. Silver has two naturally occurring isotopes:
Isotope name
Isotope mass (amu)
percentage
Silver-107
51.86
Silver-109
remainder
Find the missing percentage.
Find the average atomic mass of an atom of silver.

34. Silicon has three naturally occurring isotopes:
Isotope name
Isotope mass (amu)
Relative Abundance
Silicon-28
27.98
92.21
Silicon-29
28.98
4.70
Silicon-30
29.97
3.09
Look over the data before you begin the problem. Estimate the value of the answer before you begin the calculation. Will the weighted average be closer to 28,  29, or 30?
Find the average atomic mass of silicon.

35. Iron has four naturally occurring isotopes:
Isotope name
Isotope abundance
Isotope mass (amu)
Iron-54
5.90%
53.94
Iron-56
91.72%
55.93
Iron-57
2.10%
56.94
Iron-58
0.280%
57.93
Estimate the average mass.
Find the average atomic mass of iron.

36. The Periodic Table Period: horizontal rows of the periodic table.
Group or Family: vertical columns of the periodic table. Elements within a group have similar chemical and physical properties.

37. Period
(rows)
Group or Family
(columns)

38. The discovery of the STM's ability to image variations in the density distribution of surface state electrons created in the artists a compulsion to have complete control of not only the atomic landscape, but the electronic landscape also. Here they have positioned 48 iron atoms into a circular ring in order to "corral" some surface state electrons and force them into "quantum" states of the circular structure. The ripples in the ring of atoms are the density distribution of a particular set of quantum states of the corral. The artists were delighted to discover that they could predict what goes on in the corral by solving the classic eigenvalue problem in quantum mechanics -- a particle in a hard-wall box.

40. Atomic Structure Where are the electrons? What are electron shells?
How many electrons per shell?

41. Electrons are in shells that circles the nucleus at light speed.

42. Electron Shells 2n2 The letter n, represents the electron shell. 2e-
Number of electrons that can fit in a shell:
2n2
2e-
Electron Shell 1 can hold:
Electron Shell 2 can hold:
8e-
18e-
Electron Shell 3 can hold:
Electron Shell 4 can hold:
32e-

43. 2e- 8e- 18e- 32e- Energy Level 1 can hold: Energy Level 2 can hold:
Things to remember:
The element’s period # = the # of electron shells
There can only be 2 e- in the first energy level
All atoms want to have 8 electrons on their outer shells
Note: There is a lot more to this story. If you want to know more you should take IB Chemistry.

45. Valence Electrons
Valence electrons – Electrons on highest energy level / highest electron shell.

46. COOL FACT about atoms!
Most of the volume (space) of an atom is made up of electrons
Electrons have very little mass and take up very little space
SO, atoms are mainly empty space
We are made of atoms SO we are empty mainly made up of Empty Space

47. Sulfur
# Neutrons:
# Protons:
# Electrons:
Nucleus

48. An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.

49. s sublevels have 1 orbital
The s orbital has a spherical shape centered around the origin of the three axes in space.
s sublevels have 1 orbital
s orbital shape

50. P orbital shape
There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to
its own axis (x, y and z) in space.

52. d orbital shapes
Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells”
…and a “dumbell
with a donut”!

53. Shape of f orbitals

56. Orbital Notation
Tell where electrons are arranged in s, p, d, and f sublevel orbitals in each level around the nuclei of atoms.
Use boxes to represent orbitals
Use  or  to represent e-

57. Aufbau Principle
Electrons occupy lowest energy orbitals first

58. Pauli Exclusion Principle
Only two electrons can occupy one orbital… and they must have opposite spin.
Like This  
Wolfgang
Pauli

59. Hund’s Rule
One electron enters each orbital until all the orbitals contain one electron with the same spin direction…
…then they pair up.
Like This
p orbitals      

60. Writing Electron Configurations
# of e-
Describes e- location.
3p4
Principal Energy Level
Sublevel

62. Orbital filling table

63. Electron configuration of the elements of the first three series

64. Irregular configurations of Cr and Cu
Chromium steals a 4s electron to half
fill its 3d sublevel
Copper steals a 4s electron to FILL
its 3d sublevel

65. Wave-Particle Duality
The electron is a particle!
JJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is an energy wave!

66. The Wave-like Electron
The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.
Louis deBroglie

67. The Electromagnetic Spectrum and Light

68. Wavelength - The distance between two consecutive peaks of a wave.

69. Frequency - The number of cycles in a certain period of time… measured in cycles per second, or Hertz (Hz).
1Hz = 1/sec = 1 sec -1

70. Electromagnetic radiation propagates through space as a wave moving at the speed of light.
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters

71. Types of electromagnetic radiation:

72. Ultraviolet Rays

73. The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.
E = h
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x J·s)
 = frequency, in units of hertz (hz, sec-1)

74. Wavelength Table Long Wavelength = Low Frequency Low ENERGY Short
High Frequency
High ENERGY

75. You broke your big toe! 
The x ray they take of your toe uses waves that have a length 2.19 x 10-10m.     
What is the speed of the wave in m/s?
What is the frequency of the x ray?
What is the Energy of the photons?

76. Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum

77. Refraction of White Light

78. Spectroscopic analysis of the hydrogen spectrum…
…produces a “bright line” spectrum

79. Atomic Spectrum of Hydrogen

80. No two elements have the same emission spectrum
Atomic Spectra
When atoms absorb energy electrons move into higher energy levels…
…these electrons lose energy by emitting light when they return to lower energy levels.
No two elements have the same emission spectrum

82. Hydrogen
Iron

84. Ground State: the lowest possible energy of the electron
Excitation of the electron by absorbing energy raises it from the ground state to an excited state
A quantum of energy (in the form of light) is emitted when the electron drops back to a lower energy level.
The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.

85. Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.”
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
Werner
Heisenberg

87. electron elevates to another level.
As electron falls back to ground state, light is emitted
electron elevates to another level.
Quantum of energy hits electron

88. Electron transitions involve jumps of definite amounts of energy.
This produces bands
of light with definite
wavelengths.

89. Electron excitation & emission

92. Flame Tests

94. Pickle Light