1. Let’s look at our goal
Objectives 1-4, possibly 5

2. Unit 8
 Outline
 Pg 365 1-6 & pg 370 1-6 (assigned at the end of day 1)
 Models thus far
 What was wrong with Rutherford’s model
 First came Bohr… then ??
 Electrons – are they waves or particles? Yes!
 e- how do they lets us see light of different colors? What’s up with
that?
 e- and their location in atoms? Ions? Etc
 NEXT  What is e- configuration?

3. •1. EVOLUTION OF MODEL
 Dalton, Thomson, Rutherford, Bohr

4. The Development of Atomic
Models
 The timeline shoes the development of atomic models from
1803 to 1911.
5.1

5. •Dalton
 Solid indivisible (no parts inside)
 Combine in simple ratios
 Same element are identical…can’t change chemically
into other element
 Can mix chemically, but elements are not changed into
other atoms.
 DID EXPERIMENTS TO PROVE HIS THEORIES AS
OPPOSED TO PREDICESSORS.

6. •THOMSON
 Used CRT to discover electrons because it bent toward
a positive plate.
 Later discovered proton too.
 Plum pudding model

7. •Rutherford
 Gold foil exp. Discovered nucleus large positive mass
and that atom was mostly empty space
 Rutherford’s model had problems though. Could not
explain why metals, like iron, changed colors when
they were heated.

8. The Development of Atomic
Models
Rutherford’s atomic model could
not explain the chemical
properties of elements.
 Rutherford’s atomic model could not
explain why objects change color when
heated.
5.1

9. •Bohr
 Planetary model with electrons in fixed energy
locations (but as we will soon see, his model could
only describe some things)
 It is like the 2, 8, 8 that we described earlier this year.
See video if time

10. The Bohr Model
 Each possible electron orbit in Bohr’s model has a fixed
energy.
 The fixed energies an electron can have are called
energy levels.
 A quantum of energy is the amount of energy required to
move an electron from one energy level to another
energy level.
5.1

11. The Bohr Model
 Like the rungs of the strange
ladder, the energy levels in an
atom are not equally spaced.
 The higher the energy level
occupied by an electron, the
less energy it takes to move
from that energy level to the
next higher energy level.
5.1

12. Wave Lecture
 The current model could not describe emission
spectra which were not continuous. (like the gas
spectrum tubes.)pg 364-369
 SEE
 DEMOS of colors from salts.
 spectroscopy video
 Is light a Wave or a particle? SEE BOOK PG 363
 Light is like a particle (cathode ray)
 Light is like a wave too.
 Plank and Einstein described this.

13. Wave Terms
 Wavelength
 Frequency
 speed

15. Spectroscopy demo
 Use gas tubes & flames
 Mr. K will insert a little about the electromagnetic
spectrum here.
 Basically:
 Shorter wavelength = more energy = violet or UV or
gamma
 Longer wavelength = less energy – red, IR microwave,
radio.

16. Electromagnetic spectrum

17. Pictures of spectroscopy
 When atoms and molecules are exposed to light or
some form of energy they are said to be excited.
 They absorb that energy and can produce a unique
energy change that can identify that atom or
molecule
 This study of substances that are exposed to some
sort of continuous exciting energy is called
spectroscopy.

18. Na (see pg 374)
 Bright line spectra (made from a element (this gives
us an atomic emission spectrum)
 In lab, we’ll see an image of a flame and determine
the wavelength. l

19. The process pg 366-367
 1. The atom gets excited.
 2. an e- jumps to a higher energy level
 3. it comes back to its ground state
 4. and a photon of light is emitted.
(maybe in the visible spectrum or not)

20. +
Ground state is where the
electron is normally located.
1
2
3
Emitted photon..

21. Stair step model

23. What
happens
in
spectroscopy?
 Longer jumps
mean higher
energy and
shorter
wavelengths.

24. Light & spectra
 C= (speed of light = wavelength *
frequency)
 Red light is longer  & lower freq.
 Blue/violet light is shorter  & higher freq. &
HAS MORE ENERGY!
 Diff. color = diff. freq.
 e-s in atoms are what give off diff. colors

25. Photo electric effect & quantum
mechanics
 Planck’s constant discussion. Diff. colors meant diff. e-
s jumps in atoms (Fe when heated goes from black to red,
yellow, white, blue)
 Einstein proved it with photoelectric effect. Blue light
higher energy.
 Simulation:
http://phet.colorado.edu/en/simulation/photoelectric
 DEMO w/ LED and “glow in the dark material”

26. Again, Bohr
could explain
only the
mathematical
model for H,
so the model
changed.
SEE pg 367

27. NOTE:
When more energy is absorbed that means there are
larger quantum leaps and more energy that can be
given off.
When atoms become ions, they must gain energy to
lose an electron! That is called the ionization energy.
More on this later.

28. See Phet
 http://phet.colorado.edu/en/simulation/discharge-lamps

29. Assignment
 Begin vocab and page 3 in packet as well as…
 IN OUR BOOK
 Read pg 360-365 do 1-6
 Read pg 366-370 do 1-6
 (these will be collected at our next class meeting after a
brief discussion)
 Next… the quantum model & e- configuration

30. Collect and discuss questions
 365 & 370 # 1-6 & 1-6 (collect/ walk by)
 Packet pages 3 discuss
 Lecture configuration.
 Key things thus far: NEXT SLIDE

31. KEY IDEAS
 Light observations led to the model change
 Rutherford couldn’t explain light & color changes. Bohr
could only explain hydrogen. Pg 367
 Bohr used Quantized energy levels 368
 colors are determined by wavelength, Shorter
wavelength has more energy (wavelength and
frequency are inverses of each other)

32. NOTE:
When more energy is absorbed that means there are
larger quantum leaps and more energy that can be
given off.
When atoms become ions, they must gain energy to
lose an electron! That is called the ionization energy.
More on this later.

33. Move on to the model now

34. The Quantum Mechanical
Model
Determines the allowed energies an
electron can have and how likely it is
to find the electron in various
locations around the nucleus.
5.1

35. The Quantum Mechanical
Model
 Austrian physicist Erwin Schrödinger (1887–1961)
used new theoretical calculations and results to devise
and solve a mathematical equation describing the
behavior of the electron in a hydrogen atom.
 The modern description of the electrons in atoms, the
quantum mechanical model, comes from the
mathematical solutions to the Schrödinger equation.
5.1

36. The Quantum Mechanical Model
 The propeller blade has the same probability of being
anywhere in the blurry region, but you cannot tell its location
at any instant. The electron cloud of an atom can be
compared to a spinning airplane propeller.
5.1

37. The Quantum Mechanical Model
 It’s the probability of finding an electron
within a certain volume of space surrounding
the nucleus See pg 370 of our book about the
lightning bug.

38. Atomic Orbitals
 An atomic orbital is often thought of as a
region of space in which there is a high
probability of finding an electron.
 Each energy sublevel corresponds to an orbital of a
different shape, which describes where the electron
is likely to be found.
 Each energy level we go out away from the
nucleus, we add an energy level and also add an
additional sublevel.
5.1

39. Atomic Orbitals
 Different atomic orbitals are denoted by letters. The s orbitals
are spherical, and p orbitals are dumbbell-shaped.
5.1

40. Where are electrons
(a new approach S2012)
Bohr originally thought e-
s were in shells.
Because of the light experiments
(spectroscopy) we actually have
sublevels.

41. See text pg 371-374
 Kicker draw the overlapping orbitals here.
 Student should know these two shapes
 S shaped like a sphere
 P like a dumbell

42. Each sublevel adds a new
orbital shape. Each orbital
holds 2 electrons. (see book for filling order)

44. Atomic Orbitals
 Four of the five d orbitals have the same shape but different
orientations in space.
5.1

45. Orbital diagram template

46. Filling rules.
1.Fill a sublevel
before going to
another (except S
before D & F)
2.Single up before
doubling up.
3.Opposite spin
Aufbau & poly. Rules.

47. You try Oxygen
orbital diagram, then list configuration.

48.  Here are the
energy levels with
the sub levels,
labeled s, p, d & f

49. BOOK IMAGE
 See the book for the filling order
 PG 382

50. See pg 378 & look at sample
11-2 page 379
 Try a couple
 C
 N
 O
 Mg
 Ti (4s fills before the 3d!)
Hey Kicker,
there has to be
an easier way!

51. Let’s make this easy by using
our PT to write configurations
 See moodle
 Where are e-s located.

52. s pd
f
s
Areas or Blocks
 Shade in these areas or notate them on the PT I will
give you!

53. s pd
f
s
Filling Order
1 2
3 4 5 6 7 8 9 10
11 12 13 14 15 16 17 18
19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36

54. Another picture

55. Keep in mind that this is not as
difficult as it may seem
 Basically we go in order from left to right across a period.
 Let’s try a couple.
 N
 Si
 K
 Ti
 Cl
 The only problem with this method: when we get to d filling we
subtract one from the energy level, at f, we subtract 2. Because
they are actually further inside. S & P are the valence electrons
 Try Ti

56. YOU TRY
 Zn
 Fe
 Mg
 Br
 K
 Hg
 Rb

57. Orbital diagrams
 How about an orbital diagram for the ones we just
did? N Si K Ti Cl
 See textbook pg 378

58. Valence electrons (see pg 384)
 are the outer electron shell, which we learned earlier
can only hold up to 8 total electrons.
 Valence is the outer shell – only s & p
 How about the valence configuration for these?

59. Most atoms
 Want to have a full valence of 8 electrons (called octet)
this would be s2 & p 6
 Metals lose electrons when there is sufficient energy to
“pull one off” this is called the Ionization energy.

60. Configuration for ions?
 Write these configurations.
 Ca
 Ca +2
 F
 F-1
 O-2

61. Spring 2014
 Discuss notes pg 3-5
 Discuss answers pg 6-7 -Try pg 11 (valence pg 384 image is wrong)
 Highlight spectroscopy steps
 Discuss vocab
 Discuss exceptions & ions (See page 8 for notes and probs)
 Assign pg 12-13 & colored PT areas
 IF TIME –wavelength lab (pg 9-10) else –read and do
tomorrow

64. Ready for OVERKILL?

65. B) Quantum numbers
1 s 2
Energy level Refered to as n
Sublevel (orbital shape)
# electrons
Discuss
2p4

69. VOCABULARY NOTES
 electromagnetic radiation Wavelength
 Frequency Photons
 Quantized wave mechanical model
 Orbital principal energy levels
 Sublevels Pauli exclusion principle
 electron configuration orbital diagram
 valence electrons core (shielding) electrons
 Groups representative elements
 atomic size ionization energy
 Electronegativity stability

70. Continuous spectrum:
What is seen when white light goes
through a prism. (not discrete lines)
ROYGBIV

71. Bright line spectra
 Also called emission spectra. This
is the fingerprint of the atom.
Electrons are excited and jump to a
higher energy level, when the return
to ground state, they give off a
discrete line of a particular
wavelength. Each atoms spectrum
is different.

72. What is the explanation for the discrete
lines in atomic emission spectra?
 Electrons absorb energy and leap from one
orbital in an atom to an orbital of higher
energy. When these excited electrons fall
back down to lower energy levels, they emit
light. The lines result from the fact that the
electrons can move only between discrete
energy levels; they cannot have intermediate
energies. Electron energies are quantized,
not continuous. (LIKE STAIRSTEPS & A
BALL)

73. The process pg 366-367
 1. The atom gets excited.
 2. an e- jumps to a higher energy level
 3. it comes back to its ground state
 4. and a photon of light is emitted.
(maybe in the visible spectrum or not)

74. +
Ground state is where the
electron is normally located.
1
2
3
Emitted photon..

75. Wavelength
Distance from crest to crest
INVERSE OF FREQUENCY

76. Photon
discrete bundle of electromagnetic
energy

77. Photoelectric effect
 This is the situation where light hits a photo cell
and creates a direct current. (electron flow)
What happens In the photoelectric effect?
Electrons are ejected by metals when light
shines on them. The effect is only observed if
the frequency of the incident light is above a
certain threshold frequency.

78. Energy level
Locations where electrons can be.
Each has a specific orbital

79. Quantum
energy needed to move an electron
from one energy level to another

80. Quantum mechanical
model
 It is a model that describes the motions of
electrons in atoms as probabilistic motions
within a certain region. It is depicted as
electron clouds, the density of which represents
the probability of finding the electron in that
region. The electron cloud of the quantum
mechanical model is centered on the atomic
orbital as proposed by Bohr, but Bohr could not
describe the discrete spectral lines on an
emission spectrum with his planetary model.

81. Atomic orbitals
 The shapes of the clouds that electrons tend to make.
There are specific shapes and numbers of orbitals for
each atom. SHAPES: An s orbital has the shape of a
sphere and is the orbital having the lowest energy. A p
orbital is dumbbell-shaped and has the next highest
energy.

82. Electron configuration
The filling order of electrons into the
orbitals
Two methods.
 Orbital filling
 Electron config. (1s2 2s3 sp6…

83. Ground state
The lowest energy state of the atom.
(The original location of electrons in
their orbitals.)

84. Valence electrons
This is the highest energy levels S &
P electrons.

85. Octet rule
All atoms want a full valence of
electrons. This is the highest energy
levels S & P electrons.

86. Still to discuss this unit.
 Valence of ions?
 Exceptions to electron configuration
 Do the Na wavelength lab & practice worksheets.

87. Packet Pg 8 Stability levels. (ordered from
most stable to least stable)
 1. full octet s2p6
 2. full sub level s2 d10 F14
 3. ½ filled sub level d5
 Which is more stable?
 2S
2 or 2s
1
 3s
2 3p
3 or 3s
2 3p
4
 2s
1 2p
6 or 2s
1 2p
2
 Fe or Mn? Si or P Ca or Sc

88. Work thru pg 8
 Then assign: pg 12-13 & colored packet
 Lab tomorrow

89. Exception for e- configuration
 Some atoms rearrange their electrons from what we
would predict the e- configuration to be to become
more stable.
 Ex. Cu

90. Stable ions- explain them
 Elements will lose e-s to become more stable. It takes
energy to do this though. Some elements take on a
couple different oxidation numbers because of the
options in stability.
 Be able to predict stable ions OR What are possible
oxidation numbers that these atoms can have)
 Ca Ti Sc Ta

92. What information does the PT give us?
 Demo trend of Alkali metals
 discuss (assign) label a table and periodically yours unit 8 quiz ws
 Discuss why they react and why the trend.
 Lecture lots discuss periodic table ws 1 pg 18 packet
 Ionization energy & atomic size ws
 Review using worksheets
 Layout various metals, Si, Non metals, Alkali metals, Ca. Show locations & variety.

93. Goals
 Describe the history of the design of PT
 How, location of certain groups/regions
 Size of atoms (AR), Cations, Anions
 How do they become ions (cations)
 IE
 Trend of IE. (exceptions)
 EN

94. Periodic table
•Designed by Mendeleev and put in
increasing atomic mass. He put elements
that had similar properties in vertical
columns.
•Later it was put in order of increasing
atomic number. (moseley)
•Packet pg 15 (some notes can be filled in as we go.

95. Terms & concepts TO KNOW
 Mendeleev & his process
 Period
 Group
 Periodic law
 Metal, nonmetals, metalloids and some of their
properties
 Alkali metals
 Alkaline earth metals
 Halogens
 Noble gases
 Transition metals
 Inner transition metals
 Representative elements

96. Terms & concepts TO KNOW
 Mendeleev & his process –mass order, periodicity.
 Period - row
 Group (family – vertical column)
 Periodic law when elements are arranged in atomic number a
repeating pattern of chemical AND physical properties exist.
 Metal, nonmetals, metalloids and some of their properties
 Metals: conduct heat & electricity, maleable, ductile, luster, left side.
 Non-metals: don’t conduct, brittle, dull, right side
 Metalloids (semi-metals) some properties of both.
 Alkali metals (group 1A)
 Alkaline earth metals (group II A)
 Halogens (group VII A)
 Noble gases (group VIII A)
 Transition metals (d filling)
 Inner transition metals (f filling)
 Representative elements ( s & p filling)

97. categories

98. See the mini PT pages
Metalloids
Halogens
Alkali metals
Alkaline Earth Metals
Transition Metals
Representative elements
Inner transition metals
noble gases

99. More from LABEL THE TABLE
 atom size trend
 ion size trend.
 How does the size of a
cation compare to it’s
atom? (explain why)
 How does the size of a
anion compare to it’s
atom? (explain why)

100. Reactivity of Alkali Metals

101. Alkali metals
See video

102. KEY IDEA
The periodic table position
and the chemical properties
of the elements arise from
their electron configuration.

103. Predict PAIR SHARE
 What do you think happens to atom size as you
 A) go down a family
 B) go across a period
 WHY???

104. Atomic Size (AR)
 General trends
 Atoms get larger as you
add more energy levels
 Atoms tend to get
smaller as you go
across a period from
left to right. ( Be able
to explain WHY?)

105. ATOM Size Trend Rational
 As you go down, we add an energy level.
 As we go across, from L to R, we add
protons which make the nuclear charge
stronger and effectively sucking in the
electron cloud making the atom smaller

106. So which is larger?
 Ca or Mg
 F or Br
 Si or C
 Sr or Ne

108. What about the size of ions?
 When atoms gain or lose electrons, the
atom becomes an ion.
 What would you predict would be true
about the size of cations compared to their
atom? What about the anions?
 Trend of ions going across? Going down?

111. Ion size from atom size. You need to think about if it
is a cation or an anion?

112. Can you explain which
is larger of each of
these pairs?
Li to Li+1
Li+1 to Be +2
F to F-1
O-2 to F-1

113. ION Size Trend Rational
 As you go down, we add an energy level.
 As we go across, from L to R, we add protons
which make the nuclear charge stronger and
effectively sucking in the atom to make it smaller
 The pattern is the same as for atoms, however
when you get to where atoms become anions,
there is a big jump in size.

114. You predict
 Which is larger?
 Use only your periodic
table
 Na, Na +1
 Cl, Cl-1
 Na+1, Cl-1
 H+, H, H-1
 F -1, Na+1, Ne

115. Put these in order from largest
to smallest and be able to justify
your answer.
Mg+2 Na+1 Ne F-1 O-2
When an atom, or ion has an equal
number of electrons or has the same
number of electrons when compared
to the atoms of another element, the
two species are called
isoelectronic.

116. Ionization Energy packet pg 23-24 See chapter
 IE is the energy required to lose one electron.
 TWO THINGS EFFECT IE
 Size of atom (shielding) and nuclear charge.
More shielding
less shielding

117. So let’s think about this
Those valence electrons that
are closer to the nucleus
are harder to pull away.
So an e- from Na is
harder to pull away than
it is to steal an e- from K.
Therefore K is more
reactive!
What would you predict is
true as you go across a
period? Which has lower
IE: Mg or Na?

118. IE
1312 2371
520 900 800 1086 1402 1314 1681 2080
495.8 737.6 577.4 786.2 1012.0 999.6 1255.0 1520.0
418.0 589.5

120.  We have various IE’s
First IE, Second IE, Third IE as you try to take away a
second electron from an atom the IE increases even more.

121.  Generally, IE increases as we go across a period from
left to right. What happened between N & O? Why?
 Also IE1 and IE 2, what seems to be true? Why is there
a huge jump from the first to the second IE for lithium?

122. Are there other areas where the trend
(increasing IE to the right) does not fit?
Remember the exceptions
to configuration.
 Octet
 Filled sub level
 Half filled sublevel.

123. Show my Straw Example

124. A closer look at ionization
 Boron is further to the
right on the PT, but it
has a lower IE. Why?
 Because it would
rather have a more
stable config. so it lets
the e- go with a little
more ease than Be
does.

125. Also
 Look at the activity
chart (single replacement rxns.)
on your Periodic table.
With the exception of
Lithium what trend
exists for the first 5 or
6 metals?
• Ionization tends to
decrease as we go
down the periodic
table. This is
because the atoms
become so big, that
the attraction of the
outer e- is so small.

126. Use your PT to predict
 Which has the higher IE Ca or Mg
 B or N
 Why would IE for a nonmetal, such as Cl, be
higher than that of a metal, such as K?
 REMEMBER WHAT IE IS!!!

127. Electronegativity
 E.N. Is the affinity for electrons.
 What would be characteristics of an
element with a high E.N.?

128. Metallic Trends
 An element is more metallic if it tend to lose
electrons easier

129. Test will have a different PT

130. Predict the most reactive
 Metal? (low IE and lots of shielding)
 Non metal? (high IE and not much shielding)

131. summary

132. Time in class to try
Unit 8 WS 1 pg 15 (15 mins)
 You will have about 10 minutes or so to work
through this page and we’ll discuss it. So get
to it! Pg 69 & 386 might help

133. Unit 8 WS 1 pg 15 ANSWERS
 Who did the pioneer work on the periodic table we used
today?
 DEMETRI MENDELEEV later it was Moseley
 In what order are the elements listed in our present
periodic table
 BY ATOMIC # (originally by atomic mass)

134.  State the periodic law
 When the elements are arrange in order of increasing
atomic #, there is a periodic repetition of their physical and
chemical properties.
 What is the name given to the elements in a vertical column
of the periodic table?
 GROUP (or Family)
 What name is given to the elements in a horizontal row on the
periodic table?
 PERIOD
Unit 8 WS 1 pg 15 ANSWERS

135. What part of the periodic table represents the filling of the d
sublevel
 Transition metals
 Which groups, by number, are the representative elements?
 A GROUPS
What part of the periodic table represents the filling of the f
sublevel
 INNER TRANSITION METALS
Unit 8 WS 1 pg 15 ANSWERS

136.  What is true of the valence electron configuration of the
elements with similar properties?
 THEY ARE THE SAME

 State the octet rule (or when atoms become un-reactive)
 ALLATOMS THAT HAVE A FULL VALENCE
(mainly 8 e-)
Unit 8 WS 1 pg 15 ANSWERS

137.  Atoms of the alkali group have 1 valence electron and
tend to LOSE 1 e-s

 Alkaline earth metals have 2 valence electrons and tend to
LOSE 2 e-s

 Halogens have 7 valence electrons and tend to GAIN 1 e-s
Unit 8 WS 1 pg 15 ANSWERS

138.  Isoelectronic refers to ions that have the same electron
configuration of each other (usually that of a noble gas.
For example. Mg+ 2, Na +1 Ne, F-1, O-2)
 Can you tell me some that are isoelectronic to Kr?
Unit 8 WS 1 pg 15 ANSWERS

139. Unit 8 WS 1 pg 15 ANSWERS
 What is the basic trend as you go across a period?
 Increase an electron in the valence and increase nuclear
charge. The atom size becomes smaller as a result
 What can be said about elements that are in the same family or
group?
 They have similar chemical and physical properties.

140. LAB (s)
 Na vapor wavelength (collect data, complete and turn
in)
 Alkaline Earth LAB
 Homework: plantet quack pages
 Pg 25
 Next day in class (26-27 and discussion) begin the
next lab.

141. Discussion Protocol
 See the handout
 Expectations: everyone
participates!
 Each one write some.
 Groups 1 & 2 (oral protocol)
 Discuss part 1#1&2
 Discuss part 2