AUDIENCE THEORY -CULTIVATION THEORY - GERBNER.pptx
Chapters 4,5,6
1. The Atom and Its
Properties
Chapter 4 – Nucleus
Chapter 5 – Electron
Configuration
Chapter 6 - Periodic Table
2. Chapter 4 Objectives
• Describe an atom’s structure and
differentiate among the particles that make
it up.
• Identify the numbers associated with
elements and explain their meaning .
• Realize that the number of protons in a
nucleus defines an element.
• Calculate the average atomic mass given
isotopes and relative abundance
2
3. Chapter 4.1
• Dalton’s Atomic
Theory
• Atom
Chapter 4.2
• Electron
• Nucleus
• Proton
• Neutron
Chapter 4.3
• Atomic Number
• Isotope
• Mass Number
• AMU (Atomic Mass Unit)
• Average Atomic Mass
3
Chapter 4 Vocabulary
4. Brief History
• Many ancient scholars believed
matter was composed of such things
as earth, water, air, and fire.
• Many believed
matter could be
endlessly divided
into smaller and
smaller pieces.
5. Brief History
• Democritus (ancient
Greek philosopher in
about 460-370 BCE)
believed matter was
made up of tiny
particles, ‘atomos’ in
Greek, which gave us
our modern word atom.
• In other words, matter
could not be infinitely
divided.
5
6. Brief History
• Democritus had
many other ideas
that were left
unexplored until
school teacher
John Dalton (1766
– 1844) revived
them in 1803 to
coincide with the
industrial
revolution.
6
7. Dalton’s Atomic Theory - 1803
1. Each element is composed of atoms.
2. Atoms of a given element are identical.
(Atoms of a specific element are different from
atoms of another element.)
3. Atoms cannot be created or destroyed, only
rearranged, combined or separated.
4. Different atoms combine in simple whole
number ratios to form compounds.
7
8. Definitions
• Electrons (e-)are negatively charged particles.
They were the first subatomic particle discovered.
• Protons (p+) are positively charged particles
found in the nucleus. Discovered by Ernest
Rutherford (1871 – 1937) in 1911. (More on him
in the movie)
• Neutrons (n°) are subatomic particles about the
size of a proton but carry no charge. Discovered
by James Chadwick (1891-1974) in 1932.
8
9. Modern View of the Atom
The nucleus is where the protons and neutrons
are located and contain most of the atom’s mass.
9
10. Protons, Neutrons and Electrons
Particle Symbol Charge Relative
Mass
Electron e- 1- 1/1840
Proton p+ 1+ 1
Neutron n0 0 ~1
10
14. Mass Number
• Mass Number is the number of protons
plus the number of neutrons of a particular
isotope of an element.
(Mass # = #p+ + #n°)
• Thus we have atoms called:
• Potassium -39
• Potassium -40
• Potassium -41
• What is the number of neutrons in each of
these isotopes?
14
16. What’s all this amu business?
• To simplify a system of indicating atomic masses
since protons and neutrons have such extremely
small masses, scientists have assigned the
carbon-12 atom a mass of exactly 12 atomic
mass units. (amu)
• The mass of 1 amu (1/12 the mass of carbon-12)
is very nearly equal to the mass of a single proton
or neutron but not the same.
• 1 amu = 1.66 x 10-24 grams
16
17. Weighed Averages
• Your QPA is a weighted average. The more
credit a course is worth, the more an “A”
will help your grade, and a “D” will hurt it.
• Let’s look at the example in the handbook.
17
18. •Your total number of credits is 41 and total quality points
is 3667.
This makes your QPA average 3667/41 = 89.44
18
19. Atomic Mass
• Why is potassium’s atomic mass 39.098 in
the periodic table?
• How about
• Atomic mass for Cl is 35.453?
• Or atomic mass for Li at 6.941?
• The atomic mass is the weighted average
of mass numbers of the isotopes.
• Based on abundance of each isotope.
19
20. •Isotopes and Mass Number
• Your text (p. 119) shows how to calculate
the mass number for Cl given the %
abundance of the isotopes.
• Let’s do this for another element: Li
• 6Li is 7.59 % abundant; 6.015 amu
• 7Li is 92.41% abundant; 7.015 amu
• Method 1: Use percentages. Think of this
as a sample of 100 atoms.
20
21. •In Tabular Form
Species Mass
(amu)
Abundance
%
Mass x
Abundance
(Weighted
share)
6Li
(isotope)
6.015 7.59 45.65
7Li
(isotope)
7.015 92.41 648.26
Li
(100 atoms)
100.00 694.41
Li (atom) 6.94
21
22. • The average mass of an atom is found by
weighting the natural abundances of its
isotopes.
• Lithium (Method 2): Change % to fraction.
• 6Li 6.015 amu 7.59% = 0.0759
• 7Li 7.015 amu 92.41% = 0.9241
Mass (amu) Frac abund Mass share
Avg mass = 6.015 amu x 0.0759 = 0.46 amu
7.015 amu x 0.9241 = 6.48 amu
6.94 amu/atom
22
Average Atomic Mass
24. Bead Activity
• Note what set you have.
• Count the number of protons (red beads)
and neutrons (green beads) each isotope
(bag) has.
• Use the abundance listed on each bag to
calculate the weighted share.
• Use the weighted share of each isotope to
determine the average atomic mass.
• Also, determine which element you have.
24
27. •Chapter 5 Objectives
• Compare wave and particle matters of light
• See how frequency of light emitted by an atom
is unique to that atom
• Compare and contrast the Bohr and quantum
mechanical models of the atom
• Express the arrangements of electrons in
atoms through orbital notations, electron
configurations, and electron dot structures
27
28. Chapter 5 Vocabulary
Chapter 5.1
• Electromagnetic Radiation
• Wavelength
• Frequency
• Amplitude
• Electromagnetic Spectrum
• Quantum
• Photoelectric Effect
• Photon
• Atomic Emission Spectrum
28
Chapter 5.2
• Ground State
• Quantum Number
• Quantum Mechanical
Model of the Atom
• Atomic Orbital
• Principal Quantum
Number
• Principal Energy
Level
• Energy Sublevel
29. •Wave Nature of Light
• Electromagnetic radiation displays wavelike
behavior as it travels through space
• Waves can be described by several
common characteristics
29
30. •Characteristics of a Wave
• Waves transfer energy
• Properties of waves:
• Frequency (ν –
pronounced ‘nu’) -
Number of vibrations
per unit time – Hz
(cycle/second)
• Wavelength (λ) -
Distance between
points on two
consecutive waves
• Speed of wave is
Frequency x
wavelength
Speed = ν x λ
-1.5
-1
-0.5
0
0.5
1
1.5
0 200 400 600 800 1000 1200
λ
Frequency is the
number of waves that
hit this point in one
second.
Amplitude
30
34. •Electromagnetic Spectrum
• Gamma Rays – Highest frequency, shortest
wavelength. Can pass through most
substances
• X Rays – Lower frequency than Gamma
rays. Can pass through soft body tissue but
can’t pass through bone.
• Ultraviolet (UV) Rays – Part of sunlight that
causes sunburn
35. •Electromagnetic Spectrum
• Visible Light – Sensitive to our eyes.
Allows us to see color
• Infrared – Less energy and longer
wavelength than visible light. Felt as heat
given off a heater or near a fire
• Radio Waves – Lowest frequencies on the
EM spectrum. Used by radio and over-the-
air TV.
36. l x n = c
l = c/n
l = 3.00 x 108 m/s
6.0 x 104 /s
l = 5.0 x 103 m
No, it’s a radio wave
(~103 meters)
An electromagnetic wave has a frequency of 6.0 x 104 Hz.
Convert this frequency into its corresponding wavelength.
Does this frequency fall in the visible region?
7.1
36
37. •Examples
Problems
• What is the frequency
of green light, which
has a wavelength of
520 nm.
• A radio station
broadcasts at 94.7
MHz. What is the
wavelength of the
broadcast?
Answers
C = λ*ν
ν = C/ λ
ν = 3.00 x 108 m/s
520 x 10-9 m
= 5.77 x 1014 /s
λ = C/ν
= 3.00 x 108 m/s
94.7 x 106 /s
= 3.17 m
Hint: nm =
nanometer which is
10-9 meters
M = mega which is
106 Hz
37
38. •Nature of Light
• Max Planck (1858-
1947) studied the
different light
emitted from heated
objects
• Matter can only gain
or lose energy in
small specific
amounts
38
39. •Nature of Light
• A quantum is the minimum amount of
energy that can be gained or lost by an
atom
• The energy of EM radiation is proportional
to its frequency (E α ν)
39
40. •Photons
• Albert Einstein (1879-
1955) proposed that
while a beam of light had
wavelike characteristics,
it also can be thought of
as a stream of tiny
particles (or bundles of
energy) called photons
• Each photon carries a
quantum of energy
40
41. •Particle Nature of Light
The photoelectric effect is when electrons are
emitted from a metal’s surface when light
of a certain frequency shines on it.
42. Each element has only certain specific
frequencies of light that are emitted when
atoms absorb energy and become excited
Where do we see this?
fireworks
neon signs
stars
Ch. 5.2 – Quantum Theory of the Atom
43
44. •Balmer Plot
• In 1885, Johann Balmer observed the lines of
the spectrum fit this surprisingly simple
formula:
• Where n1 =2 and n2 = 3, 4, 5, etc.
2 2
1 2
1 1 1
H
R
n n
l
45
45. •Balmer Plot
y = 1.0972E+07x + 4.0238E+02
R² = 1.0000E+00
1500000
1600000
1700000
1800000
1900000
2000000
2100000
2200000
2300000
2400000
2500000
0.12 0.14 0.16 0.18 0.2 0.22 0.24
1/Labmda,
m-1
1/2^2 - 1/n^2
RH is the slope of this line, 1.0972 x 107 m-1
46
46. •Electronic Energy Transitions
• Neils Bohr (1885-
1962) proposed
the model the
hydrogen atom
(1913) to explain
the discreet nature
of the hydrogen
spectrum.
47
47. •Electronic Energy Transitions
• Neils Bohr’s model the atom (1913)
• Electrons exist only in discrete, “allowable”
energy levels
• Energy is involved in moving electrons from
one energy level to another
• Principal quantum number (n) - specifies
the electron’s major energy level
• The lowest energy is n=1, the next lowest in
n=2, etc.
48
48. •Bohr’s Model of the Atom (cont’d)
50
Bohr suggested that an electron moves around the
nucleus in only certain allowed circular orbits.
n = 1
n = 2
54. Schrodinger applied idea of e-
behaving as a wave to the
problem of electrons in atoms.
He developed the WAVE
EQUATION
Solution gives set of math
expressions called WAVE
FUNCTIONS.
Treated electrons as wavelike
particles that became the
Quantum Mechanical Model of
the Atom.
E. Schrodinger
1887-1961
•Quantum or Wave
Mechanics
56
55. •Wave Function
• The wave function predicts a three-
dimensional region around the nucleus
called the atomic orbital.
58
56. •Wave motion: wave length and nodes
“Quantization” in a standing wave
59
Waves
57. •Hydrogen Atom Solution
Where:
a0 is the Bohr Radius given by
Generalized Laguerre Polynomial
Constant = 2.18 x 10-18 J
m is mass of electron
a0 = 4πεoh2/me2
m here is quantum number
60
60. •Orbitals
• No more than 2 e- assigned to an orbital
• Orbitals grouped in s, p, d (and f) sublevels
s orbitals
d orbitals
p orbitals
63
61. s orbitals
d orbitals
p orbitals
s orbitals p orbitals d orbitals
No.
orbs.
No. e-
1 3 5
2 6 10
64
62. Energy Levels and
Sublevels
• Sublevels are grouped in energy
level.
• Each energy level has a number
called the PRINCIPAL QUANTUM
NUMBER, n which indicates
relative size and energy of the
orbitals
• Row on PT indicates n
65
64. •QUANTUM NUMBERS
The shape, size, and energy of each orbital
is a function of quantum numbers:
n (principal) Energy Level
l (angular) sublevel (s, p, d, or
f) which is its shape
• Note: There are other quantum numbers
that we will NOT discuss in detail. The ‘n’
and ‘l’ are sufficient.
67
65. Symbol Values Description
n (principal) 1, 2, 3, .. Orbital size
and energy level
l (angular) 0, 1, 2, 3, … n-1 Orbital shape
s, p, d, f, …n-1 or type
(energy sublevel)
QUANTUM NUMBERS
More commonly
noted as:
68
68. •s Orbitals— Always Spherical
Dot picture of
electron cloud
in 1s orbital.
Surface density
4πr2y versus
distance
Surface of 90%
probability sphere
See Active Figure 6.13
72
71. •Hydrogen-like Orbitals
(at most two electrons/orbital)
n Sublevel
(l)
Orbitals Max.
Orbital
n2
Max
Elec
2n2
1 s s 1 2
2 s
p
s
px, py, pz
4 8
3 s
p
d
s
px, py, pz
dxy,dxz,dyz,dx
2
-y
2, dz
2
9 18
4 s
p
d
f
s
px, py, pz
dxy,dxz,dyz,dx
2
-y
2, dz
2
And 7 f orbitals
16 32
75
72. Chapter 5.3 Vocabulary
76
• Electron configuration
• Aufbau principle
• Pauli Exclusion Principle
• Hund’s Rule
• Valence Electron
• Electron Dot Structure
73. •Electron Configurations
• An atom’s electron configuration is
the arrangements of electrons in the
atom.
• Electrons are arranged to minimize
energy.
• In other words, electrons fill up the
lowest energies possible first. This is
the Aufbau Principle.
77
74. •Assigning Electrons to Atoms
• Electrons generally assigned to
orbitals of successively higher energy.
• For H atoms, E depends only on n.
• For many-electron atoms, energy
depends on both n and l.
78
76. •Assigning Electrons to Subshells
• In H atom all subshells of
same n have same energy.
In many-electron atom:
a) subshells increase in
energy as value of n + l
increases.
b) for subshells of same n +
l, subshell with lower n is
lower in energy.
PLAY MOVIE
80
78. •Aufbau Diagram -- Filling Electron
Orbitals
1s 2s 3s 4s 5s 6s 7s
8s
2p 3p 4p 5p 6p 7p
3d 4d 5d 6d
4f 5f
82
Haven’t gotten this
far. What orbitals
are being filled with
elements 110-118?
Start here
n + l = 1
n + l = 2
n + l = 3
n + l = 4
n + l = 5
n + l = 6
n + l = 7 n + l = 8
The orbital with the lower ‘n’ is lower in energy if the n+l number is the same.
79. •Writing Atomic Electron Configurations
Two ways of writing
configs. One is
called the spdf
notation.
1
1 s
value of n
value of l
no. of
electrons
spdf notation
for H, atomic number = 1
83
80. •Pauli Exclusion Principle
No two electrons in the
same orbital can have
the same spin. One
electron is spin up, the
other is spin down.
84
81. •Writing Atomic Electron
Configurations
Two ways of
writing configs.
Other is called
the orbital box
notation.
Arrows
depict
electron
spin
ORBITAL BOX NOTATION
for He, atomic number = 2
1s
2
1 s
It would be a violation of the Pauli
exclusion principle to have both of
these electrons as spin up or both
as spin down.
85
85. •Carbon
Group 14 (4A)
Atomic number = 6
1s2 2s2 2p2
6 total electrons
Here we see for the first time
HUND’S RULE. When
placing electrons in a set of
orbitals having the same
energy, we place them singly as
long as possible.
1s
2s
3s
3p
2p
89
86. Electron Configuration for Elements 11-18
90
Noble gas notation uses noble gas symbols in
brackets to shorten inner electron configurations
of other elements.
87. Sodium
Group 1 (1A)
Atomic number = 11
1s2 2s2 2p6 3s1 or
“neon core” + 3s1
[Ne] 3s1 (uses noble gas notation)
Note that we have begun a new period.
All Group 1A elements have [core]ns1
configurations.
91
92. Electrons in Energy Levels
• Electrons fill up levels from lowest energy
to highest energy (Aufbau Principle)
• Outermost electrons are called Valence
Electrons.
• When atoms come close together it is the
Valence Electrons that interact.
96
93. Valence Electrons
•How to determine which electrons
are in outer shell?
• Write down electron configuration of
an element in noble-gas
configuration
• Whatever electrons are displayed in
the highest energy shell (n) only
are valence electrons (Main Group
elements)
97
94. Lewis Dot Diagrams
• How do we represent Valence Electrons?
• By a Lewis Dot Diagram
• Rules for Lewis Dot Diagrams:
• Use the Elemental Symbol
• Use 1 dot to represent each valence electron
• The symbol represents the nucleus and all the
inner (core) electrons.
• Examples: . .. ..
• Li·, Be: , ·C·, ·Cl:, :Ne:
. .. ..
98
95. Valence Electrons
• Examples
• O given by [He]2s22p4 so it has 6 valence
electrons.
O
• Ga given by [Ar]3d104s24p1 has 3 valence
electrons.
•Ga•
•
•
•
••
••
99
98. Draw the periodic table and label
the electron blocks and areas of
non-metals, metals, and metalloids.
Relate the Lewis dot structure to its
place in the periodic table.
Explain periodic trends as one moves
along periods and down groups in
the periodic table
102
99. Chapter 6.1-6.2
Periodic Law
Group
Period
Representative
Element
Transition Element
Metal
Alkali Metal
Alkaline Earth Metal
Transition Metal
Inner Transition
Metal
Lanthanide Series
Actinide Series
Nonmetal
Halogen
Noble Gas
Metalloid
103
100. Antoine Lavoisier (1743-1794)
in the 1790s compiled a list
of 23 known elements.
Many known since ancient
times (copper, gold, silver, etc)
Others had been more recently
discovered (oxygen, hydrogen,
carbon, and others)
104
101. Scientists in 1860
agreed upon a common
method to determine
elemental mass first
published by Stanislao
Cannizzaro (1826-
1910).
105
102. Early attempts organized the elements by
increasing mass
The properties were roughly periodic, but
some elements were out of order
Mendeleev’s Table (ca. 1870) was the most
notable effort at organizing the elemental
properties
106
103. Dmitri Mendeleev
noticed in his table
that there were
repetitions of
physical and
chemical properties
when the elements
were arranged by
atomic mass.
107
105. Property Predicted (1869)
Atomic Mass 72 u
Color Dark gray
Density 5.5 g/mL
Melting Point High
Density of Oxide 4.7 g/mL
Oxide solubility in
HCl
Slightly dissolved
by HCl
Formula of chloride EsCl4
Properties of Germanium (Ge)
109
Actual (1886)
72.6 u
Gray-white
5.32 g/mL
937⁰ C
4.70 g/mL
Not dissolved
by HCl
GeCl4
106. Periodic Law states that chemical
and physical properties repeat in
regular cyclic patterns when they
are arranged by increasing atomic
number.
Starts with metals at left and goes to
non-metal (noble gas) on right
Properties change in orderly progression
across a period.
110
109. What are some of the elemental
properties that make the periodic table,
well, periodic?
Classification by metals, nonmetals and
metalloids
Metals - shiny ductile, malleable solids, good
conductors of heat and electricity
Nonmetals - dull, brittle solids; or gas, poor
conductors of heat and electricity
Metalloids - have chemical and physical
properties of both metals and nonmetals
113
110. Representative Elements (Sometimes
called A Group)
Group # = number of valence electrons
Means similar Lewis dot structure and
similar properties.
s-block elements have 1-2 electrons in
s-orbital
p-block elements have 1-6 electrons in
p-orbitals
Noble gases have filled valence shells
Energy level of valence electrons is at
energy level given by period (row)
number
114
111. Transition Elements (Sometimes called
B Group)
d-block elements have 1-10 electrons in d-
orbitals
Columns 3-12 in periodic table
Energy level of valence electrons at n and
partially filled n-1 d orbitals (example: 4s
and 3d)
f-block (Lanthanides and Actinides) have
1-14 electrons in f-orbitals
115
113. Fill in the missing info for the following elements:
Identify the element fitting the description.
a) Group 2 (2A) element in 4th period: Calcium
b) Noble gas in 5th period: Xenon
c) Group 12 (2B) element in 4th period: Zinc
d) Group 16 (6A) element in 2nd period: Oxygen
117
Configuration Group Period Block
[Ne]3s2 2 (2A) 3 S
[He]2s2 2 (2A) 2 S
[Kr]5s24d105p5 17 (7A) 5 P
[Ar]4s23d5 7 (7B) 4 D
115. Effective Nuclear Charge (Z*) – Not in book!
Shielding
Ion
Ionization Energy
Octet Rule
Electronegativity
119
116. Atomic and ionic size
Ionization energy
Electronegativity
Metallic Character
120
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
117. Z* is the nuclear charge experienced by the
outermost electrons. (Note: not in book!)
Z* increases across a period owing to shielding by
inner electrons.
Shielding is blocking by inner electrons.
For a period (row), the number of shielding electrons
remain the same, but the number of protons in the
nucleus increases.
Example: All elements in the second period have the
same underlying [He] noble gas configuration.
However, the number of protons is greater from left to
right.
121
118. So we can estimate as
Z* = [ Z - (no. inner electrons) ]
or
Z* = Z – S (inner electrons)
Z is total number of electrons
S is the number of electrons blocking the valence
shell electrons, the underlying noble gas electrons.
Charge felt by 2s e- in Li Z* = 3 - 2 = 1
Be Z* = 4 - 2 = 2
B Z* = 5 - 2 = 3 and so on!
122
120. Atomic size is a periodic trend influenced by
electron configuration.
For metals, atomic radius is half the distance
between adjacent nuclei in a crystal of the
element.
124
121. For other elements, the atomic radius is half
the distance between nuclei of identical
atoms that are bonded together.
125
123. Size (radius) goes UP on going
down a group. See previous slide.
Because electrons are further
from the nucleus, there is less
attraction.
Size (radius) goes DOWN on
going across a period.
127
124. Size (radius) decreases across a period owing
to increase in Z*. Each added electron feels a
greater and greater positive charge.
Note: Electrons in the same energy level don’t
shield each other too much.
128
Large Small
Increase in Z*
125. 129
0
50
100
150
200
250
0 5 10 15 20 25 30 35 40
Li
Na
K
Kr
He
Ne
Ar
2nd period
3rd period 1st transition
series
Radius (pm)
Atomic Number
126. The radius of an atom when it has
become an ion.
An ion is an atom or bonded group of
atoms that has a positive or negative
charge.
An atom acquires a positive charge by
losing electrons or negative charge by
gaining electrons!!
130
127. To form positive ions from elements remove 1
or more e- from subshell of highest n [or
highest (n + l)].
Al: [Ne] 3s2 3p1 - 3e- Al3+: [Ne] 3s0 3p0
131
1s
2s
3s
3p
2p
1s
2s
3s
3p
2p
128. 132
Atoms tend to gain, lose, or share
electrons to get
8 valence electrons
(except small atoms up to Boron)
129. 1. Write the electron configuration and orbital
box diagram for Mg when it is an ion. Hints:
What is its noble gas configuration? What will
they do to get an octet?
2. Write the electron configuration and orbital
box diagram for O when it is an ion.
133
130. Positive ions are SMALLER than the atoms
from which they come.
The electron/proton attraction has gone
UP and so size DECREASES.
Electron Configuration as ion is: [He] 2s0
134
Li,152 pm
3e and 3p
Li +, 78 pm
2e and 3 p
+
Forming a
positive
ion.
131. Negative ions are LARGER than the atoms from
which they come.
The electron/proton attraction has gone
DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes.
Electron configuration as ion: 1s22s22p6 (just
like neon.)
135
Forming a
negative
ion.
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
-
133. Why do metals lose
electrons in their
reactions?
Why does Mg form Mg2+
ions and not Mg3+?
Why do nonmetals take
on electrons?
137
134. IE = energy required to remove an electron
from an atom in the gas phase.
138
Mg (g) + 738 kJ Mg+ (g) + e-
PLAY MOVIE
135. Mg (g) + 738 kJ Mg+ (g) + e-
139
Mg+ (g) + 1451 kJ Mg2+ (g) + e-
Mg+ has 12 protons and only 11 electrons.
Therefore, IE for Mg+ > Mg.
IE = energy required to remove an electron
from an atom in the gas phase.
PLAY MOVIE
136. 1st: Mg (g) + 735 kJ Mg+ (g) + e-
2nd: Mg+ (g) + 1451 kJ Mg2+ (g) + e-
140
3rd: Mg2+ (g) + 7733 kJ Mg3+ (g) + e-
Energy cost is very high to dip into a shell
of lower n.
PLAY MOVIE
141. IE increases across a period
because Z* increases.
Metals lose electrons more
easily than nonmetals.
Nonmetals lose electrons with
difficulty.
145
High ionization energy: atoms want
to hold on to electrons; likely to form
negative ion
Low ionization energy: atom gives up
electron easily; likely to form positive
ion
142. IE decreases down a group
Because size increases.
Ability to lose electrons
generally increases down
the periodic table.
See reactions of Li, Na, K
146
143. Which element in each pair has the
larger 1st ionization energy?
A. Na or Al
B. Ar or Xe
C. Ba or Mg
147
145. *Note: ‘metallic character’ not in book.
An element with metallic character is one
that loses electrons easily.
Metallic character:
• is more prevalent in metals on left side of
periodic table
• is less for nonmetals on right side of
periodic table that do not lose electrons
easily
149
147. Relative ability of an element to attract
electrons in a chemical bond.
Ionization energy reflects ability of atom to
attract electrons in an isolated atom
Generally, the higher the ionization energy of an
atom, the more electronegative the atom will be
in a molecule
There are many electro negativity scales –
we’ll use the one by Linus Pauling (values
dimensionless)
Will be used to determine things like
polarity of a chemical bond.
151
149. Decreases down a group
Why? Due to greater atomic radius
Increases across a period
Why? Increased positive charge in
nucleus (Greater Z*)
Same trend as for ionization
energy. Surprised?
153
150. Moving Left Right (periods)
Z* Increases
Atomic & ionic Radius Decrease
Ionization Energy Increases
Electronegativity Increases
Metallic Character Decreases
Moving Top Bottom (groups)
Z* is roughly constant, but val e- distance
increases
Atomic & Ionic Radius Increase
Ionization Energy Decreases
Electronegativity Decreases
Metallic Character Increases
154
151. a) Electronegativity
b) Ionic Radius
c) Atomic Radius
d) Ionization Energy
e) Metallic character
Fluorine
Bromine
Bromine
Fluorine
Bromine
155
Editor's Notes
Last pointwill be Illustrated with Bead Activity later in the week.
In #3, they can be rearranged also includes combined and separated.
I will show “The Atom – Episode 1” in an upcoming class which shows the history of the atom. Students are responsible for answering the questions and knowing it for the test.
What the above relative masses show is that an electron is ~1820 times smaller (by mass) than an proton or neutron. Use the money equivalent to explain. A proton/neutron is worth $20 and an electron is worth $0.01.
After this slide have students do Atomic Number/Mass Number activity for CW/HW.In Chromium row, #p is 1st, #e is 2nd, #n is thirdIn 2nd row (As) row, the order is Atomic # (33), #electrons (33), Mass #(75), As-75 IsotopeIn 3rd row (Rn-222), the order is # protons (86), atomic #(86), Isotope (Rn-222), #N°
QPA means Quality Point Average.
Note: An 85 in chemistry would give 510 Quality points and raise your average to 90.90. Because a difference of 10 in chem means a little bit more than and difference in 10 of English. 60 QP vs 50 QPs.
Going back to our QPA analogy, the mass is like the grade and the abundance is like the credit. The mass x abundance is like the QP points. When you divide the Quality Points (grade*credit) by credit you get a weighed grade. When you divide the sum of the Mass x abundance numbers by total abundance (which should always be 100) you get the weighted atomic mass.
The tabular format in the previous slide shows the method assuming 100 atoms. This is a more direct method, first show changing the percents to a fraction and then multiplying by mass. You are just going to get the mass of 1 atom in this way.Either way is acceptable. It’s really just a matter of where you are dividing by 100.
Other Atomic Mass abundances:Boron-10 : 19.8% Mass: 10.01 Boron-11 : 80.2% 11.01Silicon-28 : 92.23% Mass: 27.98Silicon-29 : 4.67% 28.98Silicon-30 : 3.10% 29.98Iron-54 : 5.8% Mass: 53.94Iron-56 : 91.72% 55.93Iron-57 : 2.2% 56.93Iron-58 : 0.28% 57.93Before they start bead activity (next slide) give for HW The Atom and Its Properties WS 2 – Atomic MassAnd give out The Atom and Its Properties WS 3 – Ch. 5.1 and 5.2 Definitions
Make sure you note which bag has which abundance.Set up a table w/#p+ and #n0 as well as mass number. For simplicity, we’ll consider each proton and each neutron to be 1 amu. Submit as group activity.
Go over how to read this chart. Wavelength in meters is on top. Frequency in Hertz is on the bottom.Take-home message of this slideThe higher the frequency (or the shorter the wavelength) the more energy the light has. So, UV light has more energy than IR light. (It is the UV light causes sunburn.)The ‘Do Now’ I have for this has problems like “Determine the type of radiation with a wavelength of 1 m.” - Answer: RadioGo over a couple of more of these type of questions (and frequency). Good for “Do Now” and test.
I’m pushing this a little more than last year. Keep equations and math simple. Save complicated ones for bonus questions.
Use an example of throwing something into a swimming pool: With large items K.E. is large and water (like the electrons) will “splash out”. With smaller energy items like a “pin” water (like the electrons) will not splash out.Light has to be a certain frequency to eject the electron from the metal. The intensity (amplitude) doesn’t matter.
Start of Chapter 5.2Ch. 5.2VocabGround StateDe Broglie EquationHeisenberg Uncertainty PrincipleQuantum MechanicsMight use the POGIL packet in folder in conjunction with this along with the salts and H2 gas demo.
Hydrogen Spectrum.First seen in the mid-1800’s.
The n is the same as Balmer’s n. Only this time, Bohr defined it as something rather than just an integer.
The n is the same as Balmer’s n. Only this time, Bohr defined it as something rather than just an integer.
Principal quantum number n=1 is closest to nucleus; n=7 is farther awayWhen all electrons are in their lowest state it is called the “ground state”When electrons are hit with energy (thermal, electrical etc.) they may go to an “excited state” atoms have multiple excited statesBohr model only worked for hydrogen, but it did a great job of predicting electronic transitions.Each circle represents an allowable energy state.When the electron goes back to ground state, a photon is emitted.
Hydrogen Spectrum.Show actual hydrogen and other gases using tubes and power supply.
Call the sublevels “l” (the letter ell) this year. This is so I can be more comfortable in assigning chapters 6 & 7 for the AP summer assignment.After this slide, give out The Atom and Its Properties WS 5 – Electron OrbitalsAnd The Atom and Its Properties WS 6 – Ch. 5.3 Definitions
Vocab given in The Atom and Its Properties W S 8 – Ch. 6 Vocab
Lavoisier was on the wrong side of history. He associated with the aristocracy in France during the French Revolution and got guillotined.
Italian chemist StanislaoCannizzaro published a list of atomic weights used at the First International Chemical Conference in 1860.
Remember that protons, neutrons and electrons were not discovered until the late 19th and early 20th centuries.Mendeleev used Cannizzaro’s list to start his work.
The numbers above represent atomic masses.
Note: Number of valence electrons for Representative Group is either Group # (Using A & B notation) orGroup # (for Groups 1 & 2) or (Group # - 10) for the 1 – 18 Group # notation.Electron dot notation is just used for the Representative elements since the valence electrons are constant. The transition elements can vary how many electrons are actually used.
These are Problems #7 and 8 on p. 162. (I added the last one in the chart to have a D-block example.)Config Group Period Block[Ne]3s2 2 (2A) 3 S[He]2s2 2 (2A) 2 S[Kr]5s24d105p5 17(7A) 5 P[Ar]4s23d5 7 (7B) 4 DAfter this slide give out The Atom and Its Properties WS 9 – Periodic Table Basics as CW/HWAnd The Atom and Its Properties WS 10 – Ch. 6.3 Definitions
P. 175 – Problem #63 for practice.Which element in each pair has the larger 1st ionization energy?Na or Al Al why? Smaller atom in same period.Ar or Xe Ar why? Smaller atom in the same group.Ba or Mg Mg why? Smaller atom the same group.
Can show test tubes of carbon, silicon, tin, lead.Can even show ‘oxygen’, sulfur, maybe selenium?
This is Exercise #22 on p. 169 of text.Answers: F, Br, Br, F, Br (it’s a liquid instead of a gas, molecules closer together)After this slide, give out The Atom and Its Properties WS 11 – Periodic Trends