semennnn

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

11. C-12
C-14
11
Mass # = #p+ + #n°
How Atoms Differ – Ch. 4.3

12. •Sometimes Atomic Symbols are
Displayed as:
12

13. •Displayed on Periodic Table
13

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

15. Isotope Examples
Isotope Atomic
Number
Mass
Number
Number
Protons
Number
Neutron
Number
Electron
52Cr
33 42
222 86
24 24
52 28 24
75As 33 75 33
86
86
222Rn 136
15

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

23. 23

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

25. •Set #_______
Names:________________
#p+ #n0
Mass #
(Isotope
mass)
Abun-
dance
%
Mass x
Abundance
(Weighted
share)
Bag
1
Bag
2
25
What is the element?
What is its average atomic mass?

26. Electrons in Atoms
Chapter 5

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

31. •Electromagnetic Spectrum
Note: All EM Radiation travels at 3.00 x 108 m/s
31

32. •Electromagnetic Spectrum
The speed of light (3.00  108 m/s) is the
product of it’s wavelength and frequency
c = λν.

33. The Electromagnetic Spectrum –
all light is energy
33

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

43. •Hydrogen Spectrum
44

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

49. •Energy Absorption/Emission
51

50. 52

51. •Atomic Emission Spectra
53

52. •Origin of Line Spectra
Balmer series
54

53. 55

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

58. 61

59. •Atomic Orbitals-Hydrogen
62

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

63. •Energy Levels and
Sublevels
n = 1
n = 2
n = 3
n = 4
66

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

66. •Types of
Atomic
Orbitals
70

67. Types of Atomic Orbitals
71

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

69. •p Orbitals
The three p orbitals lie 90o apart in space
73

70. •2px Orbital 3px Orbital
PLAY MOVIE
74

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

75. •Energy Level Diagram of
Hydrogen
79

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

77. Many-Electron Atoms
Orbitals and Their
Energies
81

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

82. •Electron Configurations
and the Periodic Table
86

83. •Lithium
Group 1 (1A)
Atomic number = 3
1s22s1  3 total electrons
1s
2s
3s
3p
2p
87
Interactive Periodic
Table

84. Ground State Electron Configurations
88

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

88. Electron Configurations
and the Periodic Table
92

89. Transition Metals
All 4th period elements have the
configuration [argon] nsx (n - 1)dy and
so are d-block elements.
Copper
Iron
Chromium
93

90. Transition Element Configurations
3d orbitals used
for Sc-Zn
94

91. 95

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

96. •Electron Dot Representation
100

97. Chapter 6
101

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

104. 108

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

107. 111

108. Periodic Table
Periods
112
Transition Elements
Inner
Transition
Elements
Halogens
Noble
Gases
Alkali Metals
Alkaline Earth Metals
Representative Elements
Columns,
Groups or
Families
Metals Metalloids Nonmetals

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

112. 1 2 13 14 15 16 17 18
116

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

114. 118
PLAY MOVIE
PLAY MOVIE
PLAY MOVIE

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

119. 123
Orbital energies “drop” as Z* increases

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

122. 126

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
-

132. 136
See Figure 6-14

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

137. 141

138. 142
1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 35
0
500
1000
1500
2000
2500
1st Ionization energy (kJ/mol)
Atomic Number
H Li Na K
He
Ne
Ar
Kr

139. 143
As Z* increases, orbital energies
“drop” and IE increases.

140. 144

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

144. 148
Lithium
Sodium Potassium
PLAY MOVIE
PLAY MOVIE PLAY MOVIE

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

146. 150

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

148. 152

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