1. The Electron
6.0 Chemistry

3. Development of the Periodic Table
History of the Periodic Table – By the end of the 1700’s, scientists had identified only 30 elements (ex. Cu, Ag, Au, H2, N2, O2, C).
By the mid 1800’s, about 60 elements had been identified.
Sept 1860 – chemists assembled at the First International Congress of Chemists in Germany to settle the controversial issues such as atomic mass. Standard values set for atomic mass and improved communication for research.

4. Johann Dobereiner: 1817
Organized the elements into sets of three with similar properties.
He called these groups triads. The middle element is often the average of the other two.
Ex) Cl – 35.5
Br – 79.9
I – 126.9 Ca
Avg Sr Ba

5. Triads on the Periodic Table

6. B. John Newlands: 1866
Arranged elements in order of increasing atomic mass.
Noticed repeating patterns in the elements’ properties every 8th element.
Law of Octaves - properties of elements repeated every 8th element.
There were 62 known elements at the time.

7. C. Dmitri Mendeleev: 1869
Arranged elements in order of increasing atomic mass.
Similar properties occurred after periods (horizontal rows) of varying lengths.
Organized the 1st periodic table according to increasing atomic mass and put elements with similar properties in the same column.
Periodic – repeating properties or patterns
Noticed inconsistencies in the arrangement.

8. Mendeleev’s 1st Periodic Table

9. He arranged some elements out of atomic mass order to keep them together with other elements with similar properties. (Notice Te and I)
He also left several blanks in his table.
In 1871, he correctly predicted the existence and properties of 3 unidentified elements – Sc, Ga and Ge
These elements were later identified and matched his predictions.

10. Mendeleev is known as the
1st Periodic Law
Properties of the elements repeat periodically when the elements are arranged in increasing order by atomic mass
Mendeleev is known as the
Father of Chemistry
#101 honors Mendeleev

11. D. Henry Moseley: 1911
Studied X-ray spectral lines of 38 metals. Each element had a certain amount of positive charge in the nucleus which are called protons.
Analyzed data and found that the elements in the PT fit into patterns better when arranged in increasing nuclear charge, which is the Atomic Number.
The Modern Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.

12. Glenn Seaborg “Seaborgium” Sg #106
Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII.
He suggested pulling the “f-block” elements out to the bottom of the table.
He was the principle or co-discoverer of 10 transuranium elements.
He was awarded the Noble prize in 1951 and
died in 1999.

13. Seaborgium is the exception…
After some argument between the USA and the rest of the world, element 106 was named Seaborgium shortly before he died. This was a matter of some controversy because the International Union of Pure and Applied Chemistry, IUPAC, the body that deals with naming in chemistry, had previously ruled that elements should not be named after living people.
Atomic #
104
105
106
107
IUPAC
Unnilquadeum
Unq
Unnilpentium
Unp
Unnilhexium
Unh
Unnilseptium
Uns
Agreed in 1995
Dubnium
(Dubna, Russia)
Joliotium
(Frederic Joliot)
Rutherfordium
(Earnest Rutherford)
Bohrium
(Neils Bohr)
Agreed in 1996
Seaborgium

14. Parts of the Periodic Table
Horizontal Rows – PERIODS
There are 7 periods in the periodic table
Elements in a period do NOT have similar properties.
Vertical Columns – GROUPS or FAMILIES
Labeled 1-18
IA-VIIIA are the Main-group or representative elements.
Elements in a group have similar properties.
Why?

15. Family Names write these on your P.T.
Transition elements or metals (3-12): d-block
Inner transition elements or metals (f-block)
Lanthanides or lanthanide series
Actinides or actinide series
Transuranium elements
Hydrogen (1)
Alkali metals (1) – most reactive metals; reactivity increases down the group
Alkaline earth metals (2)
Boron family (13)
Carbon family (14)
Nitrogen family (15)
Oxygen or Chalcogen family (16)
Halogens (17)
Noble gases (18) - inert

16. Parts of the Periodic Table

17. E. Metal, Nonmetals and Metalloids (Semimetals):
Found on the
LEFT side of
the PT
2. Nonmetals
Located on the
RIGHT side of
3. Metalloids - Properties of both metals & nonmetals
- Good conductors of heat & electricity
- High melting points most solids at room temperature
- High luster (shiny)
- Ductile (can be drawn into thin wire)
- Malleable (bends without breaking)
- High densities
- Reacts with acids
- Brittle (easy to break)
- No luster (dull)
- Insulators nonconductors
- Neither ductile nor malleable
- Nonreactive with acids
(Semimetals)

18. Review of Early Atomic Theories
Dalton 
Thomson 
(plum pudding)

19. Atomic Theories

20. Quantum Mechanical (Schrödinger) Model
Bohr – electrons in a particular path have a fixed energy called energy levels
Rungs of a ladder
Quantum Mechanical (Schrödinger) Model
Electrons better understood as WAVES
Does not tell where the electrons are located
Electrons have a certain amount of energy - QUANTIZED

21. Parts of a Wave

22. ex) Radio FM 93.3 megahertz (MHz) is 93.3 x 106 Hz (cycles/sec)
Light as a Wave
Characteristics of a Wave
Amplitude: Height of the wave from the baseline. The higher the wave the greater the intensity.
Wavelength: (λ , “lambda”) in nanometers (1 x 10-9 m). Distance between similar points on 2 consecutive waves.
Frequency: (ν , “nu”) The number of waves that pass a fixed point per unit of time. Measured in cycles/second (1/s) 1 cycle/second = Hertz (Hz)
ex) Radio FM 93.3 megahertz (MHz) is x 106 Hz (cycles/sec)

23. Wavelength
Trough
Baseline
Crest
Amplitude 3
Wavelength
Amplitude

24. D. Electromagnetic Radiation - a form of energy that exhibits wavelike behavior as it travels through space - all forms of EM radiation move at the speed of light

25. This is an indirect relationship.
Speed of Light (c)
E x 108 m/s or 186,000 miles/sec.
The relationship between wavelength and
frequency can be shown with the following
equation:
This is an indirect relationship.
If λ  then ν .
c = λ ν

27. High Low Low High Long Short Energy Energy Visible Light Radio/TV
Microwaves
Ultraviolet
Gamma Rays
Radar
Infrared
X-Rays
Low
High
Long
Short
High
Low
Energy
Energy
Red
Orange
Yellow
Green
Blue
Violet

29. X Quantum Theory Planck’s Hypothesis: (Max Planck 1900)
Studied emission of light from hot objects
Observed color of light varied with temperature
Suggested the objects do not continuously emit E, but emit E in small specific amounts
Light is absorbed or emitted in a little packet or bundle called a quantum (quanta –plural).
Quantum = minimum amount of E that can be lost or gained by an atom
Energies are quantized.
(Think steps not a ramp) e- e- X e- e-

30. Incandescent light bulbs give off most of their energy in the form of heat-carrying infrared light photons -- only about 10 percent of the light produced is in the visible spectrum. This wastes a lot of electricity. Cool light sources, such as fluorescent lamps and LEDs, don't waste a lot of energy generating heat -- they give off mostly visible light. For this reason, they are slowly edging out the old reliable light bulb.

31. Max Planck’s Energy Equation
Proposed that energy is directly proportional to frequency.
E = h
Planck’s equation for each quantum
h = Plank’s constant = x J.s
This is a direct relationship.
As energy increases, frequency increases.

32. Albert Einstein
While well-known for the equation E=mc2 , Einstein’s work on the photoelectric effect resulted in being awarded the 1921 Nobel Prize in Physics.
(1879 – 1955)
German Physicist

33. Albert Einstein and the Photoelectric Effect
Refers to the emission of electrons from a metal when light shines on the metal
Observations:
Electrons are ejected by light of sufficient energy. Energy minimum is different for different metals.
The current (# of electrons emitted/s) increases with brightness of the light.
my.hrw.com

34. Albert Einstein and the Photoelectric Effect
Conclusions:
Proposed that light consists of quanta of energy that behaves like particles.
Quantum of light = photon = massless particle that carries a quantum of energy.
Proposed the Dual Nature of Light: its wave and particle nature.
Light travels through space as waves
Light acts as a stream of particles when it interacts with matter.
my.hrw.com

35. Light (Electromagnetic Radiation)

36. Spectroscopy
Definition: a method of studying substances that are exposed to some sort of continuous exciting energy.
A. Emission Line Spectra: contains only certain colors or wavelengths (  ) of light.
1. Every element has its own line spectrum (fingerprint).

37. Continuous Spectrum – White Light
Line Spectrum – Excited Elements

38. B. White light gives off a Continuous Spectrum a blending of every possible wavelength

39. Gas Discharge Tubes
Electricity is added to the gas which causes the electrons to jump to a higher or excited state. They immediately fall back to the ground state and give off particular wavelengths of light. We see a blending of wavelengths without the spectroscopes.

40. Flame Tests
used to test qualitatively for the presence of certain metals in chemical compounds. 
the heat of the Bunsen flame excites electrons that emit visible light.
Copper(II) sulfate
Lithium chloride
Potassium chloride
Barium nitrate

42. Spectroscope
Uses a diffraction grating to diffract the light into particular wavelengths of light.

43. A Line Spectra result from excited elements - as electrons of an element gain energy and rise to an excited state they then fall back to their ground state in the same pattern producing the same energy drop each time which we see as individual wavelengths of light.

44. III. Atomic Spectra and the Bohr Model of Hydrogen (1913)
Neils Bohr - Danish Scientist
Explained the bright-line spectrum of hydrogen
Study:
Added E as electricity to H gas at low pressure in a tube.
Emitted E as visible light, was observed through a prism
Result: Hydrogen emitted 4 distinct bright lines of color, aka bright line spectrum

45. Electrons release energy as they fall back to a lower energy level

46. Electrons absorb energy to rise to a higher or excited state and emit energy in the form of a photon of light as they fall back to their ground states.

47. Path of an excited electron as it “falls” back to the Ground State
When electrons gain energy, they jump to a higher energy level (excited state).
Electrons are not stable at the excited state and will immediately fall back to a lower level or ground state.
As they fall, they emit electromagnetic radiation.
Depending on how far they fall determines the type of radiation (light) released.

48. Bohr Model of Hydrogen
Conclusion:
*Unique line spectrum is due to quantized electron energies.
*Electrons are in specific orbits related to certain amounts of energy known as stationary states.
*Orbits are related to energy levels.
*Energy levels are identified as E1, E2, E3, … (n = 1, 2, 3, …)
*Lowest energy level = ground state
*Electrons absorb certain amounts of energy to move to a higher energy level farther away from the nucleus = excited state
*Electrons return to the more stable ground state and release a photon that has energy equal to the difference in energy between the energy levels.
from E2 to E1: Ephoton = E2 – E1 (difference in energy)

49. The Bohr Atom for Hydrogen - a Model
Successful in calculating the wavelength, frequency, & energy of hydrogen’s line spectrum.
Successful in calculating the energy needed to remove hydrogen’s electron
H(g) + energy  H+1(g) e-
Calculated ionization E = observed ionization E
= kJ/mol

50. Lyman, Balmer and Paschen series of the Hydrogen Atom
Lyman series: electrons fall to n = 1 and give off UV light.
Balmer series: electrons fall to n = 2 and give off visible light.
Paschen series: electrons fall to n = 3 and give off infrared light.

51. When electrons absorb energy they jump to a higher (excited) state.
n=2 n=3 n=4 n=5 n=6 n=7
Electrons are not stable. Radiation (light) is emitted when an electron falls back from a higher level to a lower level.

53. Infrared Light
Visible Light
Ultraviolet Light

54. Atomic Spectra
Hydrogen
Although Bohr’s atomic model explained the line spectra of hydrogen, it failed for heavier elements.
Helium
Lithium
Mercury

55. Limitations of the Bohr Model
Model could not calculate the wavelengths of observed spectra of multi-electron atoms.
Model could not explain the chemical behavior of atoms.
Bohr used classical mechanics to understand the behaviors of small particles.
The Bohr model is also known as the planetary, solar system, or satellite model.

56. Quantum Mechanical Model of the Atom
Louie De Broglie (1924-5)
Took Einstein’s idea that light can exhibit both wave and particle properties
Very small particles (like electrons) display properties of waves.
Behavior of electrons in Bohr’s quantized orbits was similar to behavior of waves
French scientist
Known: any wave confined to a space can only have specific frequencies
De Broglie suggested electrons are waves confined to the space around the atomic nucleus.
Electrons could exist only at specific frequencies which correspond to specific energies (E = h quantized E of Bohr)

57. Quantum Mechanical Model of the Atom
Experimentally proven in 1927 by diffraction of electrons by Davisson & Germer (showed diffraction of electrons by a crystal of Ni)
Wave-Particle Duality of Nature
Light and electrons (very small particles like electrons, atoms, molecules) have properties of waves and particles  QUANTUM MECHANICS (based on WAVE properties)
**Large objects obey the laws of classical mechanics**

58. C. Werner Heisenberg: (1927)
Heisenberg’s Uncertainty Principle: states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.
You cannot predict future locations of particles.
He found a problem with the Bohr Atom - no way to observe or measure the orbit of an electron.

59. D. Erwin Schrödinger Wave Equation (1926)
Wave nature of an electron is described by a mathematical equation.
Four quantum numbers in the equation are used to describe an electron’s behavior – location and energy.
Electron is treated as a wave with quantized energy.
Describes the probability of the electrons found in certain locations around the nucleus.
(1887 – 1961) Austrian Physicist

60. Electron Density An orbital is a region in which an electron with a
particular energy is likely
to be found.
Where the density of an
electron cloud is high there
is a high probability that is
where the electron is
located. If the electron
density is low then there is
a low probability.

64. E. Atomic Orbitals - region around the nucleus where an electron with a particular energy is likely to be found (not the same as Bohr’s orbits!)
Orbitals have characteristic shapes, sizes, &
energies.
2. Orbitals do not describe how the electron moves.
3. The drawing of an orbital represents the dimentional surface within which the electron is found 90% of the time.

65. 4. Sublevels can have 4 different shapes
s – orbital spherical

66. 1s, 2s & 3s orbitals Superimposed on one another

67. Electron-Cloud Models

68. p-orbital – dumbbell shaped

69. p-orbital - dumbbell shaped

70. d-orbital - double dumbbell or fan blades

71. s,p and d orbitals s orbital p orbitals d orbitals
For a more complete representation and presentation of atomic orbitals go to

72. Models of d-orbitals

73. f-orbital – more complex!

74. f orbitals

75. f – orbitals (3D)

76. Quantum Numbers
Each quantum number provides more specific information on the probable location of an electron.
Each electron within an atom can be described by a unique set of 4 quantum numbers.

77. Quantum Numbers - Finding an address for each electron:
“state” Principle Quantum Number (n) or the energy level;
Describes the relative size of the electron cloud.
Positive integer values (n = 1 to n = 7)
“city” Sublevel (l)
Describes the shape of the electron cloud.
The maximum number of sublevels within a level = n
Shapes are s, p, d,or f.
Lowest energy = s Highest energy = f

78. Quantum numbers cont. “street” Orbital (ml) odd # of orbitals
Describes the orientation or direction in space
s – 1 orbital
p – 3 orbitals (x, y, z)
d – 5 orbitals (xy, yz, xz, x2 – y2, z2)
f – 7 orbitals (y3 – 3yx2, 5yz2-yr2, x3-3xy2,
zx2-zy2, xyz, 5xz2-3xr2, 5z3-3zr2)
Orbitals within the same sublevel have the same energy are called degenerate orbitals
An orbital can hold a maximum of 2 electrons

79. Quantum numbers cont. “house” Spin (ms)
Describes the direction of electron spin in an orbital.
The clockwise or counterclockwise motion of electrons.
Only electrons with opposite spins can occupy the same orbital.
The opposite spin is written as
+1/2 or -1/2 or  or

80. E. Electron Configurations:
Shorthand notation for indicating the number of electrons in each level, sublevel, and orbital.
1s2
Shows the distribution of electrons among the orbitals. Describes where the electrons are found & what energy they possess.

81. Electron Configuration Rules
The Aufbau Principle: electrons are added one at a time to the lowest energy orbital available.

82. Pauli Exclusion Principle:
1. Each orbital can only hold 2 electrons.
2. The electrons must have opposite spins.
s-sublevel = max 2 electrons
p-sublevel = max 6 electrons
d-sublevel = max 10 electrons
f-sublevel = max 14 electrons
incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓

83. Hund’s Rule:
Electrons will remain unpaired in degenerate orbitals before they pair up.
incorrect ↑↓ ↑ __
correct ↑ ↑ ↑

85. Electron Blocks on the Periodic Table

87. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
Pauli Exclusion Principle: No more than 2 e- are put in each orbital and they must have opposite spin.
Hund’s Rule: electrons spread out among equal energy orbitals in a sublevel (like charges repel)
Aufbau Principle: Electrons fill lowest energy levels first (n=1)

88. Electron Configuration Examples:
Ex) electron configuration for Na:
1s2 2s2 2p6 3s1
Ex) orbital filling box diagram for Na:

89. Electron Dot Diagrams:
Write the symbol for the element.
Place dots around the symbol to represent the
valence s & p electrons only.
Do NOT include d & f orbitals in diagram.
p orbital electrons s orbital electrons

90. Electron Configuration Orbital Box Diagram Electron-dot Diagram
1s22s22p4
1s22s22p63s23p5
1s22s22p63s23p64s23d104p65s24d105p4
What does the Tellurium electron-dot resemble???

91. Mark your Periodic Tables1 2 13 14 15 16 17 18

92. Unpaired vs. Paired Electrons Filled and Half-filled orbitals
Atoms with unpaired electrons are said to be paramagnetic. These are weakly attracted to a magnetic field.
Atoms with all paired electrons are said to be diamagnetic. These are weakly repelled from a magnetic field.
½ filled and filled orbitals have special stability

93. Noble Gas or Shorthand Electron ConfigurationsRb Se At
Draw the Dot Diagrams for these elements

94. Exceptions to the Rules
Max stability - ½ filled and filled orbitals Cr Mo Cu Ag Au

95. Exceptions to the Rules
Max stability - ½ filled and filled orbitals Cr Mo Cu Ag Au

96. Electron Configuration for IonsK P Al Se
K+1
P-3
Al+3
Se-2

97. Excited vs. Ground State
If an electron absorbs energy, it is in an EXCITED state
Ne: 1s22s22p53s1
How is this different from the ground state configuration?
Ne: 1s22s22p6

98. C. Electron Configuration & Families
Valence electrons – outermost electrons (s and p); responsible for bonding and chemical behavior.
Elements in the same group have the same number of valence electrons.
Carbon has 4 valence electrons

99. Electron Configurations s, p, d, f blocks

100. Electron Configurations
Stable Octet: 8 electrons in the outer level is very stable (includes He)
Ions – gain/lose electrons to achieve a stable octet
Isoelectronic – same electron configuration
Examples: N, O, F, Na, Mg, Al are isoelectronic with Ne – this is called an isoelectronic series
Pseudoisoelectronic – same electron configuration but includes the d orbitals
Fe+2 is pseudoisoelectronic with Ar