STUDENT NOTES Pre-AP Chemistry
MARUSIK
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UNIT 5 NOTES: MODELS OF ATOMIC THEORY
STUDENT OBJECTIVES: Your fascinating teachers would like you amazing learners to be able to…
1. Perform calculations involving the speed of light (c), the wavelength (λ), and the frequency (ν) of electromagnetic
radiation (EMR).
2. Perform calculations involving the energy, the wavelength (λ), and the frequency (ν) of electromagnetic radiation.
3. Place forms of electromagnetic radiation in correct energy order on an electromagnetic spectrum.
4. Compare the energy, frequency, and wavelength of EMR.
5. Label an energy level diagram, indicating both excitation and relaxation.
6. Describe the four quantum numbers used to define the region of space that has a 90% probability of finding an
electron with a given energy.
7. Explain the information provided by each of the quantum numbers.
8. Distinguish between an energy level, sublevel, and an orbital.
9. List the allowable quantum numbers for each energy level.
10. Draw an orbital diagram using the Aufbau Principle, Hund’s Rule, and the Pauli Exclusion Principle.
11. Write the electron configuration for and atom or an ion.
12. Be able to identify the highest-energy electrons, the valence electrons, and the number of valence electrons.
13. Using the periodic table, describe the periodic trends in atomic radii, ionic radii, ionization energy, reactivity, and
electronegativity.
STUDENT NOTES Pre-AP Chemistry
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I. LIGHT AND ELECTRON CONFIGURATION:
Much of what we know about the atom has been learned through experiments with light; thus, we need to
know some fundamental concepts of light in order to understand the structure of the atom, especially the
placement of the electrons.
a. Characteristics of Light
i. Has "Dual" nature (or split personality) - there are times in which it behaves like a continuous
WAVE and other times when it behaves like a PACKET of light (a blip of light!!!)
b.
Light as a Wave
i. Speed of light in a vacuum is the most accurately known constant, c, in the universe.
c = 3.0 x 108 meters/sec
OR
3.0 x 1017 nm/sec (assuming a vacuum)
ii. Wavelength is represented by the Greek letter, ______________ (  ): It is the length between
corresponding parts of adjacent crests and can be expressed in ANY units of length you choose
(feet, inches, meters, kilometers, etc)
iii. Frequency is represented by the Greek letter, ______________ (  ): It is the number of wave
crests which pass a given point in 1 second. Its units are, most commonly, cycles / second ,
simply sec-1. The unit, Hertz (Hz) is also used as a unit label.
c. Electromagnetic Spectrum
i. Any visible light will have a wavelength of between ____________ nm and ____________ nm.
The longest wavelength in visible is red (about 700 nm) and the shortest visible wavelength is
violet (about 400 nm) The only thing that makes one color of light different from another is its 
and  - the velocity is always the same (the speed of light!). The colors of visible light in order
from longest wavelength to shortest spell out ROY G BIV.
ii. NOT ALL LIGHT IS VISIBLE! The entire range of light energy is considered the Electromagnetic
Spectrum.
Example 5-1. Label the following electromagnetic spectrum from lowest energy on the left hand side to highest
energy on the right hand side.
700 nm R O Y G B I V 400 nm
RED MARTIANS INVADE ROYGBIV USING X-RAY GIZMOS
Example 5-2. What mathematical relationship can you draw about wavelength and frequency?
Example 5-3. What mathematical relationship can you draw about energy and frequency?
STUDENT NOTES Pre-AP Chemistry
II.
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CALCULATIONS WITH LIGHT
FORMULAS ON THE PERIODIC TABLE:
c = 
Ephoton = h
Ephoton 
hc
λ
CONSTANTS ON THE PERIODIC TABLE:
c = 3.0 x 108 m/s
(speed of light in meters)
h= 6.63 x 10-34 J•sec
(Planck’s constant)
Note: BE CAREFUL OF CHOOSING SPEED OF LIGHT FOR YOUR FORMULA!!! You must make the units of
wavelength match up with your units for speed of light. Remember that “c” can also be 3.0 x 1017 nm/s.
Units of frequency can be reported as /second, second –1, or Hertz. They are interchangeable with each other!
Example 5-4. If the wavelength is known to be 550 nm, what is its frequency?
Example 5-5. If the frequency of light is known to be 9.45 x 1014 s-1, is the light visible? If not, is it UV or IR?
HINT: Calculate the wavelength in nm.
Example 5-6. If the wavelength of light is 7.23 x 10-5 nm, what is the frequency of light?
Example 5-7. What is the energy of a photon of light whose frequency is 7.85 x 1015 Hz?
Example 5-8. In the previous example, is the light visible? How do you know?
Example 5-9. If a light has a wavelength of 550 nm, what is the energy of one photon of this light?
Example 5-10. If one photon of light is known to have energy of 3.33 x 10–19 J, is it visible?
STUDENT NOTES Pre-AP Chemistry
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III. BOHR ATOM:
Don’t Forget! Bohr said that electrons are in energy levels – which is true! And his research
with light & electron energy levels is very important and lead to a lot of what we know about
the atom! But he also proposed the planetary model – which is not true.
a. Electron cloud made up of _______________________________________ which are like the
______________ of a ladder and are BUT they are not ____________________________________.
b. Amount of energy possessed by an electron determines its ___________________________________.
c. Levels farther from nucleus represent _________________________ energy.
d. Electrons like to hang out where they are most __________________, which is their _______________
__________________. This is the exact amount of energy they like to possess – it’s the electron’s home.
All electrons have their own ground state – on Level 1, Level 2… or Level 8!
e. When electrons gain a “quanta” (amount) of energy (through sources such as extreme heat or
electricity), the electron will ____________________ up energy levels in a process called
_________________________. This excited state is very _____________________________!!!
f.
Electrons can’t stay excited at the higher energy level forever! When they fall back to their ground
state, they will lose the extra energy in the form of a packet of light emitted – in a process called
________________________________. The larger the fall, the larger energy of light released!
When the electrons lose the energy they
gained, they release LIGHT (EMR)…
n= ___
n= ___
n= ___
n= ___
n= ___
e- ____ energy
Process___________
e- ____ EMR (energy)
Process___________
Ground State:__________________________
STUDENT NOTES Pre-AP Chemistry
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g. Each possible energy fall corresponds to a __________________________________________________
amount… which means that it corresponds to a different _____________________________________
of __________________. Some of the falls can be detected by the human eye (as they fall in the visible
range), but other falls that are not visible must be detected by instruments (as they are infrared or
ultraviolet).
h. Different elements have different energies associated with their energy levels. This means that similar
falls in __________________________________________ will produce __________________________
______________________.
i.
When we observe a color being produced from an element undergoing relaxation, we are observing all
of the VISIBLE falls simultaneously occurring. We can use things like a prism or a spectroscope to break
apart that light into its individual colors – resulting in a _______________________________________!
j.
We can use the resulting bright-line spectra to identify an element!
Example 5-11. Based on the Bright-Line Spectra shown above, identify the elements found in the mixture.
STUDENT NOTES Pre-AP Chemistry
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IV. SCHRODINGER MODEL: Quantum Mechanical Model (1926)
a. Schrodinger ___________________________ with Bohr about quantized energy levels.
b. However, Schrodinger ____________________ in the fact that he does not define an _______________
path called an orbit. Schrodinger’s model describes a region in space with a high ___________________
of finding an electron. Unfortunately, we still call this an ______________________.
ANALOGY: It is probable that you will be in the region of space defined by my classroom during this
specified period BUT we don’t know where you are in the room and how you are moving around the
room. We’re just 90% sure we will find you (maybe you are absent, in the office, etc.). My classroom is
the orbital!
c. ELECTRON LOCATION: Electrons are distributed around the nucleus in specific ways and each electron
has four coordinates to describe its most probable location or distribution: 4 quantum numbers, n, l, m,
and s. Think of these as being the “address” for an electron’s probable location! (State, city, street,
house number… each gets more and more specific.)
n = __________________________________________ (Principle Quantum Number)



Indicates how far from the nucleus the electron is located and its energy. It is designated by the Principle
Quantum Number, n.
The higher the value of n, the ____________________ away from the nucleus and the _________________
the energy level
The number of electrons in any energy level: ______________________________ (n=energy level)
Example 5-12.
Fill in the chart with how many electrons each energy level can hold.
First Energy Level
Second Energy Level
Third Energy Level
Fourth Energy Level
l = _______________________ (Second Quantum Number)
** Each energy level is divided into ___________________________ which are designated by the Second
Quantum Number, l. Actually not a number at all, but rather a letter!
** Indicates the _______________________ of the region.
1. Sublevels are labeled by the letters s, p, d, and f.
2. The first energy level will have 1 sublevel and its label is "s".
The 2nd energy level will have 2 sublevels; "s" and "p".
The 3rd energy level will have 3 sublevels; "s" , "p", and "d"
The 4th energy level will have 4 sublevels; "s", "p", "d", "f"
Here is a trick for remembering the order:
Snazzy People Dress Fantastically!
STUDENT NOTES Pre-AP Chemistry
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3. “s” sublevels have a _____________________ shape. “p” sublevels have a _____________________
shape. “d” and “f” sublevels have really weird shapes!!!
mℓ = ________________________ (Third/Magnetic Quantum Number)

Each sublevel contains __________________ which are regions of high ____________________________
for finding electrons… the mℓ indicates the _____________________________________ of those orbitals!

The number of orbitals is related to the sublevels (ℓ)

Each orbital can hold only ______________ electrons

The surface of the orbital is drawn to represent area where any particular electron can be found
______________ of the time.

Each sublevel has a set number of orbitals:
s sublevel - ________ orbital
p sublevel - ________ orbitals
d sublevel - ________ orbitals
f sublevel - ________ orbitals
STUDENT NOTES Pre-AP Chemistry
Example 5-13.
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Complete the table:
Energy Level Sublevel(s)
Orbitals (use lines to depict)
# e– per sublevel
# e– per energy
level
1
2
3
4
s
= _____________________ (Fourth Quantum Number)
** The electrons within an orbital must have
_____________________ spins to overcome their
repulsion, which we represent with opposite arrows.
** Spin is either + ½ (______) or - ½ (_____)
We are going to use these FOUR Quantum Numbers to describe the location of electrons, like your ticket describing
your seat for a sporting event! Each get more and more specific. Think of it like this:
n = stadium level (lower, platinum, nosebleed) = energy level
ℓ
= section number = sublevel
mℓ = row = orbital
s = seat number = spin
We will look very soon at how to use all of these numbers to describe an electron’s
probable location, just like level, section, row, and seat number would help us find
you at the game (assuming of course you are not out getting nachos)!!!
STUDENT NOTES Pre-AP Chemistry
V.
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ELECTRON CONFIGURATIONS AND ORBITAL DIAGRAMS
Knowing the rules for electron arrangement allows us to write what is called the electron configuration,
which gives a description of where in that atom the electrons are actually located. We can also draw an
orbital diagram which gives the same information with more detail.
____________________________________ provide a lot of information (since they show ALL FOUR
quantum numbers), but they become annoying to draw.
______________________________________ are a shorter way of showing arrangements of electrons
around the nucleus. They don’t provide all of the information that an orbital diagram would, but they do
show us quite a bit! Each element has its own electron configuration. Electron configuration gives
specifically the energy level, sublevel, and number of electrons in that sublevel.
WE HAVE SOME BACKGROUND INFO WE NEED TO GO OVER BEFORE WE CAN WRITE THESE ELECTRON
CONFIGURATIONS & ORBITAL DIAGRAMS…
First:





Remember that “n” provides information on both energy and distance.
The relationship between “n” and DISTANCE stays true! Electrons in higher “n” energy levels are more likely to
be found further from the nucleus.
In a perfect world we would have the following energy order: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f …etc.
UNFORTUNATELY, things get a little mixed up in terms of energy – meaning it doesn’t fill up in the order we
quite expect. We will learn the order for filling sublevels in just a minute.
FOR NOW: trust us! We are going to ultimately show you that the periodic table provides us with the perfect
arrangement to break the code of filling order!
The _________________ electrons are the electrons that are in the highest energy level, “n”.
The ________________________________ electrons are in the last sublevel that was filling.
ALSO: There are some rules we must follow for arranging electrons…
i. Aufbau Principle: Electrons are placed in orbitals of ______________ energy first.
ii. Pauli Exclusion Principle:
 An orbital may hold only ______________ electrons. Think of it like an “exclusive club” with
only 2 members!
 Why can’t we have an orbital with more electrons? Well, NO TWO ELECTRONS CAN HAVE
THE SAME 4 QUANTUM NUMBERS. The combination of 4 quantum numbers is for that
electron only!
 If they have the same energy (n), are in the same sublevel (ℓ), and are in the same orbital
(mℓ), then they must have opposite spins – meaning 4 different quantum numbers.
 REMEMBER: We designate those spins with arrows! ↑ for +1/2 ↓ for ─1/2
iii.
Hund’s Rule: When filling a sublevel (which contains orbitals of equal energy), one electron
enters each orbital until each _____________________, then they _______________________!
STUDENT NOTES Pre-AP Chemistry
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ABOUT IONS:
Most metals will ______________ electrons to become ____________________________ with the nearest
noble gas. When they lose electrons, the ion will be smaller than the original atom.
Non-metals will ______________ electrons when they are with metals to become _____________________
with the nearest noble gas. When they gain electrons, the ion will be larger than the original atom.
In Noble gases, electrons have maximized attraction to the nucleus and minimized repulsion with each other as much as
possible, due to their electron configuration. They are STABLE and do not react – aka INERT. When substances become
ions, they want to copy-cat the ideal electron configurations of the noble gases!!!
Example 5-14.
Group
1
2
13
15 nonmetals
16 nonmetals
17 nonmetals
For each of the following, predict the charge of the most likely ion and indicate which sublevel
gains or loses electrons.
Electron
Configuration of
Valence
Electrons?
Number of
Valence
Electrons?
Lose or Gain
Electrons to look
like Noble Gas?
Common Charge
of Ions?
Would resulting
ion be larger or
smaller than the
original atom?
STUDENT NOTES Pre-AP Chemistry
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PERIODIC TABLE and ELECTRON CONFIGURATION/ORBITAL DIAGRAMS
Periodic Table of the Elements
Group 1 (IA)
1
Group 18 (VIIIA)
1
2
1s1
1s2
Group 2 (IIA)
3
2
Group 13 (IIIA) Group 14 (IVA) Group 15 (VA) Group 16 (VIA) Group 17 (VIIA)
4
5
ATOMIC NUMBER
5
2s1
11
6
2p1
12
13
14
7
8
2p3
2p4
9
10
15
16
17
18
33
34
35
36
51
52
53
2p6
3s2
3
Group 3 (IIIB) Group 4 (IVB) Group 5 (VB) Group 6 (VIB) Group 7 (VIIB) Group 8 (VIIIB) Group 9 (VIIIB) Group 10 (VIIIB) Group 11 (IB) Group 12 (IIB)
Period
19
20
21
22
23
24
25
26
27
28
29
30
31
32
2
4
37
38
39
40
1
5
2
3d
41
42
43
3
5s
55
4p
4d
56
71
72
73
44
45
5
75
47
88
50
4d
76
77
78
79
80
81
82
10
104
105
106
107
108
109
110
111
112
4
57
58
59
84
113
114
115
61
62
63
64
65
66
6
6p
116
7p
67
68
69
4f1
89
86
4
7p
60
85
6p
1
6d
7
83
3
5d
103
54
5p
6s
87
49
5
2
6
48
8
4d
74
46
70
4f14
90
91
92
93
94
95
96
97
98
99
100
101
102
5f13
Example 5-15.
Fill in the missing outer configurations in the periodic table above.
Example 5-16.
Answer the following question about the periodic table above.
a) How many electrons does an s- sublevel hold? ______
b) How many elements are in a row in the “s” block? ______ Hmmm…
c) How many elements are in a row in the “p” block? ______
d) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “p” block?
______ Hmmm…
e) How many elements are in a row in the “d” block? ______
f)
Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “d” block?
______ Hmmm… are you noticing a recognizable pattern?
g) How many elements are in a row in the “f” block? ______
h) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “f” block?
______ Hmmm…
i)
What is the correlation between the period number and the energy level, “n”, for the “s” and “p” blocks?
j)
What is the correlation between the period number and the energy level, “n”, for the “d” block?
k) What is the correlation between the period number and the energy level, “n”, for the “f” block?
STUDENT NOTES Pre-AP Chemistry
5p
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4d
5s
4p
3d
4s
3s
3p
2p
2s
1s
Example 5-17.
(a) Write the electron configuration for Cobalt. Then, fill out the orbital diagram above for Cobalt.
(b) What are the highest-energy electrons in Cobalt?
(c) What are the outer-most (valence) electrons in Cobalt?
(d) How many valence electrons does Cobalt have?
STUDENT NOTES Pre-AP Chemistry
Example 5-18.
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Work the following regarding Sulfur.
(a) Write the electron configuration.
(b) Write the orbital diagram across from left to right (rather than upwards).
(b) What are the highest-energy electrons?
(c) What are the outer-most (valence) electrons?
(e) How many valence electrons are there?
(f) What is the most likely ion for Sulfur? Did we lose or gain electrons to form the ion?
(g) What would be the new electron configuration for the Sulfur ion?
(h) Would the resulting ion be larger or smaller than the original atom?
(i) With which noble gas is the Sulfur ion isoelectronic?
Example 5-19.
Work the following regarding Strontium.
(a) Write the electron configuration.
(b) Write the orbital diagram across from left to right (rather than upwards).
(c) What are the highest-energy electrons?
(d) What are the outer-most (valence) electrons?
(e) How many valence electrons are there?
(f) What is the most likely ion for Strontium? Did we lose or gain electrons to form the ion?
(g) What would be the new electron configuration for the Strontium ion?
(h) Would the resulting ion be larger or smaller than the original atom?
(i) With which noble gas is the Strontium ion isoelectronic?
STUDENT NOTES Pre-AP Chemistry
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Isotope: ALUMINUM – 27 (NO CHARGE)
Nuclear Symbol:
Atom or Ion? _____________________________________
# of Protons: _____________
# of Neutrons: _____________
______________________
# of Electrons: _____________
Circle the Highest-Energy Electrons.
Box the Valence Electrons.
NUCLEUS
Number of Valence Electrons:
_____
Electron Configuration: _____________________________________________________________________
Orbital Diagram:
Example 5-20.
Answer the following regarding the Aluminum isotope above.
(a) What is the most likely ion for Aluminum? Did we lose or gain electrons to form the ion?
(b) What would be the new electron configuration for the Aluminum ion?
(c) Would the resulting ion be larger or smaller than the original atom?
(d) With which noble gas is the Aluminum ion isoelectronic?
STUDENT NOTES Pre-AP Chemistry
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Isotope: BROMINE - 71 (NO CHARGE)
Nuclear Symbol:
Atom or Ion? _____________________________________
# of Protons: _____________
# of Neutrons: _____________
______________________
# of Electrons: _____________
Circle the Highest-Energy Electrons.
Box the Valence Electrons.
NUCLEUS
Number of Valence Electrons:
_____
Electron Configuration: _____________________________________________________________________
Orbital Diagram:
Example 5-21.
Answer the following regarding the Bromine isotope above.
(a) What is the most likely ion for Bromine? Did we lose or gain electrons to form the ion?
(b) What would be the new electron configuration for the Bromine ion?
(c) Would the resulting ion be larger or smaller than the original atom?
(d) With which noble gas is the Bromine ion isoelectronic?
STUDENT NOTES Pre-AP Chemistry
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Isotope: ZINC - 64 (NO CHARGE)
Nuclear Symbol:
Atom or Ion? _____________________________________
# of Protons: _____________
# of Neutrons: _____________
______________________
# of Electrons: _____________
Circle the Highest-Energy Electrons.
Box the Valence Electrons.
NUCLEUS
Number of Valence Electrons:
_____
Electron Configuration: _____________________________________________________________________
Orbital Diagram:
Example 5-22.
Answer the following regarding the Zinc isotope above.
(a) What is the most likely ion for Zinc? Did we lose or gain electrons to form the ion?
(b) What would be the new electron configuration for the Zinc ion?
(c) Would the resulting ion be larger or smaller than the original atom?
(d) Is there any noble gas that the Zinc ion is isoelectronic with?
STUDENT NOTES Pre-AP Chemistry
Example 5-23.
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Let’s see how to include the “F” Block!!! Write the electron configuration for the following.
a) Tungsten ( #74)
b) Gold (#79)
ABBREVIATED ELECTRON CONFIGURATIONS
Writing these long, repetitive configurations can be a pain in the neck and takes up a lot of space. Space is at a premium
on a periodic table, so scientists developed an abbreviated form. To write abbreviated electron configurations, look to
the noble gas in the prior period. Brackets around the symbol indicate that the element has the same electron
configuration of that noble gas PLUS whatever follows. Note: You cannot use these if a question says to write the
“complete” electron configuration!
For iron (Fe): 1s22s22p63s23p64s23d6
Example 5-24.
BECOMES: [Ar] 4s23d6
Write the abbreviated electron configurations for the following elements.
a) Antimony (#51)
b) Palladium (#46)
c) Mercury (#80)
IDENTIFYING ELEMENTS USING ELECTRON CONFIGURATIONS
Using either the total count of electrons (by totaling up the exponents on a complete electron configuration), or by
looking at the last component of the electron configuration, we can determine what element we have by using the
electron configuration.
Example 5-25.
Identify the neutral elements that have the following electron configurations.
a) 1s22s22p63s23p3
b) 1s22s22p63s23p64s23d104p65s24d7
c) [Xe] 6s24f145d106p3
d) [Rn]7s25f146d5
STUDENT NOTES Pre-AP Chemistry
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VI. PERIODIC TRENDS
GROUPS (aka Families) run VERTICALLY on the periodic table. Trends (such as reactivity and some physical properties)
tend to be consistent within a group.
PERIODS run HORIZONTALLY on the periodic table. As you work through the periods on the periodic table, trends tend
to repeat themselves from period to period (such as number of valence electrons).
However, there are some specific trends that can be summarized by looking at the periodic table as a whole. The
following contains a description of those trends. Realize with all of these trends, we are giving general descriptions of
how these trends work… as with anything in chemistry; there are exceptions to the rules!
ATOMIC SIZE (RADIUS)





The radius is ½ of the distance measured from the
center of one atom to the center of an adjacent
identical atom. The bulk of this distance is the “electron
playground”!
Atomic size decreases across a period, due to the
increased number of protons having a greater pull on
the valence electrons more as you move across.
Atomic size increases down a group, as you are adding
more energy levels as you go down.
This would mean that Fr (Francium) would theoretically
have the largest atomic size. He (Helium) would have
the smallest atomic size.
EX: Which would have a larger atomic size – Ca, Br, or Ba?
IONIZATION ENERGY






This refers to the amount of energy required to remove one outer-most (or VALENCE) electron.
Since 8 valence electrons indicate a “full” outer shell (s & p sublevels), the closer the number of valence
electrons is to 8, the more difficult it will be to pull off an electron - meaning it will take more ionization energy.
Ionization energy increases across a period, because the atom is more tightly held together as the valence
sublevels become full… meaning it will take more energy to pull an electron off.
Ionization energy decreases down a group, because if you have more energy levels, the outer-most electrons
are further away, meaning they aren’t held on as tight… meaning it takes less energy to pull them off.
This would mean that He (Helium) would theoretically have the highest ionization energy. (Yes, he only has two
valence electrons, but that’s all he can handle!) Fr (Francium) would have the lowest ionization energy.
EX: Which would have a higher ionization energy – Ca, Br, or Ba?
STUDENT NOTES Pre-AP Chemistry
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ELECTRONEGATIVITY






This refers to the attraction an atom has to gain a nearby electron. The closer an atom is to 8 valence electrons
(in the s & p sublevels), the higher the electronegatvitiy (because the nucleus has a stronger “+” charge!)
However, would an atom that already has 8 valence electrons want to gain another?
Electronegativity increases across a period, because as you move across, the atom has a greater pull to gain
another electron (to achieve the 8 valence electrons).
Electronegativity decreases down a group, because as the valence electrons get further from the nucleus, it has
less of a pull to gain more electrons.
Noble (Inert) gases have no electronegativity, because they already have 8 valence electrons (filled s & p
sublevels), so they have no pull to gain any more! This is also part of the reason they are considered “INERT” –
meaning unreactive.
This would mean that F (Fluorine) would theoretically have the highest electronegativity. Fr (Francium) would
have the lowest electronegativity.
EX: Which would have a higher electronegativity – Ca, Br, or Ba?
USE THE PERIODIC TABLE BELOW TO MAKE SUMMARIES ABOUT PERIODIC TRENDS!!!