Download Structure of the Atom

Where are electrons found in the electron cloud?
I. The Bohr Atom
1. Bohr was the first to propose that the electrons were located in energy levels. A lower
case “n” is used to denote these principle energy levels (also called principle quantum
numbers). The principle energy levels are numbered, so that the level closest to the
nucleus is labeled n = 1. The next level is labeled n = 2 and so forth. Each principle energy
level had a certain energy value associated with the level. The closer the level was to the
nucleus, the lower the energy of the level. The further away from the nucleus, the higher
the energy is of that level. As long as the electrons were in these levels, the electrons do
not give off energy. The dark circle below represents the nucleus. The rings around the
nucleus represent the principle energy levels. Number the principle energy levels
starting with the one closest to the nucleus: n = 1, n = 2, n = 3 etc.
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2. Electron Configuration and the Periodic Table
Each principle energy level can only hold so many electrons before the level is full. A quick
and easy way to determine the maximum number of electrons (max e-) that a principle
energy level can hold is given by the following:
max e- = 2 n2. First square the
principle energy level number (n) then multiply by 2.
Energy Level (n)
Maximum number of electrons (max e- = 2 n2)
1
2
3
4
5
6
Electrons are arranged around the nucleus by filling up the first principle energy level
(n=1), then the second energy level, etc. This is the electron configuration given on your
periodic table. The number of electrons are listed for each level with a dash between
levels: for oxygen (O) which has a total of 8 electrons, the configuration is 2–6
(2 electrons are located in the first principle energy level and 6 electrons are located in
the second principle energy level. Look up the electron configuration on the periodic table
for the element given and fill in the chart. Ca is done as an example.
Element
Ca
n=1
2
n=2
8
n =3
8
n=4
2
Na
F
B
Al
C
H
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3. Completely Filled vs. Occupied Principle Energy Levels
Occupied means that there is at least one electron in the Principle Energy Levels
Li: 2 – 1
has 2 occupied Principle Energy Levels
Completely Filled means that each level has its maximum number of electrons which
can be determined by the 2n2 rule.
Li: 2 – 1
has only 1 Completely Filled Principle Energy Level
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To help you review the 2n rule complete the following chart
PEL (n)
Max e-
1
2
3
4
5
6
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For the Following:
a) Copy the electron configuration from the Periodic Table
b) Determine the number of Occupied Principle Energy Levels (PEL)
c) Determine the number of Completely Filled Principle Energy Levels
Element
Electron Configuation
# Occupied PEL
# Completely
Filled PEL
C
Na
O
Cl
He
F
Ne
Si
Zn
Au
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4. Drawing Bohr Diagrams of Atoms:
1) A circle is used for the nucleus- the # protons (# p or +) and the # of neutrons (#n)
are placed in the circle.
2) A ring is drawn around the nucleus for each energy level.
3) The electrons for each energy level are placed in pairs symmetrically around the nucleus
For F: atomic # = _____________
atomic mass = _____________
electron configuration: ______________
# p = _________ # n =____________
For Al : atomic # = _____________
atomic mass = _____________
electron configuration: ______________
# p = _________ # n =____________
Going Backwards: Determining the identity of an element from the Bohr diagram:
# p = _____________ # n =______________
atomic # = _____________
atomic mass = # p + # n = ________________
electron configuration:
# p = _____________ # n =______________
atomic # = _____________
atomic mass = # p + # n = ________________
electron configuration:
____________________________________
Isotopic Notation:
____________________________________
Isotopic Notation:
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II. Introduction to Light
Visible Light (energy we see with): part of the Electromagnetic Spectrum
1. Two theories to explain light’s behavior:
Waves
Particles of Packets of Energy
There was evidence for both models so the two theories were put together!!
Light: QUANTUM THEORY OF LIGHT
a) packets or bundles of energy called _________________ or ______________
b) travel in wave-like fashion
c) produced when electrons drop from ______________ energy levels to _______
energy levels (the greater the drop, the greater the energy the light has)
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2. Properties of Light
Wavelength () - ____________________________________________
Frequency (F) - ______________________________________________
(units: Cycles / second OR Hertz)
Energy (E) - __________________________________
Speed (velocity) – same for all electromagnetic radiation _______________
Relationships:
Frequency and Energy: Type _________________
F ______, E ________ or F ______, E ________
Frequency and Wavelength: Type _________________
F ______,  ________ or F ______,  ________
Wavelength and Energy: Type _________________
 ______, E ________ or  ______, E ________
3. Bright Line Spectra and Continuous Spectrum
A. The Rainbow: A Continuous Spectrum
Long 
Low F
Low E
R
O
Y G
B
I
V
Short 
High F
High E
LONG STEM RED ROSES: All “L’s” go together with RED
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Continuous Spectrum
 when radiation from the sunlight passes through a prism, a rainbow – a
spectrum of colors – is seen
 the colors are not separated from one another but blend together due to the
overlap of the line spectra of the 67 different elements in the sun
lightbulb
R
O
Y
G
B
I
V
C. Bright Line Spectrum
 when radiation from an excited atom (element) passes through a prism, the
radiation is separated into various wavelengths and colors
 Colors are not blended – spectrum is discontinuous – and you observe lines of color
at different locations
R
O
Y
G
B
I
V
Flame
D. Bright Line Spectra and the Bohr Atom
An electron must absorb energy before it can give off colors we see in the bright line
spectra. When energy is added, the electron moves to a higher energy level. The potential
energy of the electron increases. This is an unstable situation. In order for the electron
to return to a lower and more stable energy level, the added energy must be given off.
When the electrons return to the lower energy levels this decreases the PE because the
added energy is given off and the colors of the bright line spectra are seen. Moving
electrons to different energy levels requires different amounts of energy. These
different amounts of energy produce the different colors.
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Movement of an electron between the same 2 energy levels in DIFFERENT elements will
produce different colors. The energy between the energy levels depends on the
number of protons and the number of electrons that each element has.
BRIGHT LINE SPECTRA are produced when “electrons in the EXCITED STATE” fall back
to lower energy levels of the GROUND STATE. Unlike the continuous spectrum of
sunlight, only certain colors will be present in the BRIGHT LINE SPECTRA. The BRIGHT
LINE SPECTRUM is like a “fingerprint” of the element that produced the spectrum. Like
a fingerprint, the BRIGHT LINE SPECTRA can be used to identify the element. When
viewed with a spectroscope, the individual bands of colors in the BRIGHT LINE
SPECTRUM can be seen and the wavelength of each band determined.
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1. Below are the BRIGHT LINE SPECTRA of three elements. From the position of the
lines determine which element is the unknown. (HINT: Match up the lines present in
the unknown with the three known elements.) Unknown element = ____________
Element X
Element Y
Element Z
Unknown
2. Which of the two elements above are present in the BRIGHT LINE SPECTRUM given
below? (HINT: Match up the lines present with the three known elements. Only two
patterns should match perfectly.) ____________ and _____________
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5. Ground and Excited States
A) The lowest possible energy state that an electron can occupy is called the
__________ ____________. This is a very __________ condition. The principle
energy levels, which are occupied match those predicted by the electron configuration
on the periodic table. When electrons gain energy, the electrons move to higher
principle energy levels then they would normally occupy. This unstable situation is
called the _____________ _______________. The electrons will release the
absorbed energy, often seen as the bright line spectrum of the element, and fall back
to the ground state.
A) How to tell when energy will be absorbed or released
The Principle Energy Level (n) changes:
 If the number of the principle energy level (n) goes up, then energy is
_____________ or ______________ n = 1 to n = 3 OR n = 3 to n = 4

If the number of the principle energy level (n) goes down, then energy is
_____________ or ______________ n = 2 to n = 1

OR n = 5 to n = 3
If the energy is emitted, then ____________________ (colors) are seen.
Determine if energy is added/absorbed (+E) or released/emitted (-E) for the following
transitions ; circle the
1) n = 1 to n = 2 ______________
6) n = 1 to n = 5
______________
2) n = 4 to n = 3 ______________
7) n = 4 to n = 2
______________
3) n = 2 to n = 1 ______________
8) n = 2 to n = 3
______________
B) How do you tell the excited and ground state apart from the electron
configuration??
Ground State: Matched the predicted electron configuration found on the periodic
table. In other words, it follows the order given
Ground State for Oxygen (O) on PT= 2 – 6 (8 total electrons)
Possible Excited State for Oxygen = 1 – 7
(still 8 total electrons)
The first energy level is not filled before moving into the second energy level.
The KEY here is that the configuration does not MATCH the one on the PT.
Another possible excited state for oxygen: 2 – 5 – 1 (still 8 total electrons)
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For the following elements, fill in the chart and determine if the electron configuration
is in the GROUND STATE (GS) or EXCITED STATE (ES).
Element Electron
Symbol Configuration
C
2-4
F
1-8
Cl
2-8-7
B
2-1-1-1
Na
2-7-2
O
2-6
S
1-8-7
Zn
2-8-18-2
Br
1-7-17-8-1-1
Br
2-8-18-7
Principle
Energy
Levels
Ground or
excited
state?
III. UpDaTiNg ThE BoHR MoDeL: QuAnTuM oR WaVe MeChAnIcaL MoDeL
The Bohr Model was very good at explaining many things about atoms and electrons and
how they behaved. Just like your wardrobe needs updating, so did the model of the atom.
A revolution in physics occurred in the early 1900’s when experiments showed that matter,
just like light energy, could have a dual nature…it can act as a particle or a wave.
1. The Wave Mechanical Model
A. Still has a dense positive nucleus where protons and neutrons are located
B. Using complicated math, this math showed that electrons were not moving
in definite fixed orbits like planets but had distinct amounts of energy.
C. Four “Quantum Numbers” are used to describe the location of the electron.
2. Four “Quantum Numbers”
1) Principle Energy Levels (PEL) are also called the PRINCIPLE QUANTUM
NUMBER.) Same as the Bohr Energy levels (n = 1, 2, 3, 4, 5, 6, 7)
2) Sublevels: PEL get divided into sublevels. Like Different room types- gives
the 3-D shape where electrons are found.
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These sublevels are designated: s, p ,d and f
The number of sublevels is determined by the number of the
princple energy level (n): n = # sublevels
 Within the same principle energy level: the sublevels have different
energies:
s  p  d f
lowest energy

highest energy


3) Orbitals are “regions” or “areas” around the nucleus The sublevels are
divided into orbitals (rooms) where electrons are found.
Each orbital can hold 2 electrons.



The number of orbitals depends on the sublevel type:
The number of electrons in a sublevel depends on the number of
orbitals
Sublevel
# orbitals
x2=
# of electrons in sublevel
s
x2=
p
x2=
d
x2=
f
x2=
An orbital is defined as a region in which an electron with a particular
amount of energy is most likely to be found.
the probability of finding an electron near the nucleus is ________
the probability of finding an electron far away from the nucleus is
__________
This volume outside the nucleus are often described as an Electron
Cloud.
4) SPIN:  or  Within an orbital, the two electrons spin in opposite
directions to overcome the repulsion they feel for each other
(remember- like charges repel.)
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Principle Energy Level
Sublevel(s) Present
#
Orbitals
Present
#
electrons
in
Sublevel
Total # of e per PEL
1
2
3
4
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Remember the tricks for number of sublevels, orbitals and total number of electrons
For Principle Energy Level (PEL) n
n = _________________________ & n = ___________________________
n2 = _____________________
& 2 n2 = ________________________
2. Quantum Atom Electron Configuration from the Periodic Table
There are two ways to show the Quantum Atom Electron Configuration. One way shows
the number of electrons in the Principle Quantum Level (PEL) and Sublevel type (letter).
The other way, known as BOX DIAGRAMS, show all four quantum numbers by using boxes
for orbitals and arrows for electrons to show the opposite spin of the electrons. Using the
electron configuration from the periodic table, we will do both types.
For Hydrogen: EC form PT is 1
Quantum Atom
Box Diagram
1s1
PEL
1s
# of electrons
SUBLEVEL
PEL
ORBITAL
SUBLEVEL
ARROW SHOW SPIN
RULES TO REMEMBER WHEN DRAWING BOX DIAGRAMS;
* S has one orbital so draw 1 box
* P has 3 orbitals so draw 3 boxes,
* Each “P” box in a sublevel must have an electron before pairing up because each
orbital has the same energy as the other orbitals (Boxes)
* Electrons spin in opposite directions!
Element
Elec. Config.
From P T
Quantum Atom
E. C.
Box Diagrams
H
He
Li
Be
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Element
Elec. Config.
From P T
Quantum Atom
E. C.
Box Diagrams
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
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Element
Elec. Config.
From P T
Quantum Atom
E. C.
Box Diagrams
S
Cl
Ar
K
Ca
After argon (Ar) the energy levels begin to overlap and things get complicated. We will
not be doing elements above Calcium. However….
How many orbitals are in the d sublevel? _________ f sublevel? _____________
How many boxes would you draw for the d sublevel? _________ f sublevel? __________
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IV. The Kernel, Valence Electrons and Lewis (Electron) Dot Diagrams
A very powerful tool for showing how different elements bond together, called Lewis
(Electron) Dot Diagrams, was developed by a chemist named Lewis. Many times in
chemical bonding, only the electrons on the outside or in the highest energy level, called
the VALENCE SHELL, actually become involved in bonding to form compounds. We are
going to look at the Lewis (Electron) Dot Diagrams and elements in this next section.
Valence Shell: Outermost energy level of an element (highest number)
Oxygen: O 2-6 the 2nd principle energy level is the valence shell
Valence Electrons (VE): electrons located in the outermost energy level
- these are the electrons which are the furthest to the right in the electron
configuration found on the PT
- Oxygen: O 2-6 the 6 electrons are the valence electrons
- In a dot diagram, the VE are represented by dots: 
Kernel: This is the nucleus and all the other electrons located in the inner energy levels
In a dot diagram, the chemical symbol represents the kernel: O
The valence shell always contains only 4 orbitals. These orbitals are represented by the
four sides around the symbol. The top side represents s orbital which has less energy than
the other three p orbitals. The other three p orbitals have the same amount of energy.
First 2 electrons are always placed first in this top orbital. Going around clockwise, one
electron is placed on each remaining side until you need to pair the electrons up. (You don’t
draw the OVALs!)
s Orbital
lower energy than others

O

O



the same energy
p orbitals
Symbol
Li
Electron
Configuration
Dot Diagram
Li
Be
B
C
Be
B
C
the same energy
p oribitals
N
O
N
O
F
Ne
F
Ne
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Guided Practice: look for a pattern!
Symbol
Electron Configuration (Circle
the Valence Electrons (VE)
Number of VE
Dot Diagram
S
Se
Na
K
Si
Ge
Cl
Br
Ar
Kr
What was the pattern for electron dot diagrams?
___________________________________________________________________
___________________________________________________________________
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