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Unit 4 Notes: Chapter 5
Electrons in Atoms
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Section 5.1- Models of the atom
Goals for this section:
- Identify the inadequacies in Rutherford’s model
- Identify the new proposal in the Bohr model of the atom
- Describe the energies and positions of electrons according to the quantum mechanical
model of the atom
- Describe the shapes of the orbitals
Unit 4 Notes: Chapter 5
I.
II.
III.
Electrons in Atoms
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Identify the number of orbitals and electrons in each sublevel
Identify the sublevels in each principle energy level
Identify the last sublevel and principle energy level ofelectrons in an element using the
periodic table
Identify the number of electrons in the last sublevel and principle energy level in an element
using the periodic table
Know the following terms: Energy level, quantum, quantum mechanical model, atomic
orbital,
Rutherford’s model
A. Rutherford’s model did not explain
of elements
B. His model also did not explain why objects
when heated
Bohr’s Model
A. Bohr proposed that an electron is found only in
, around the nucleus.
B. Each possible electron orbit in Bohr’s model has a
.
C. The fixed energies an electron can have are called
.
D. A quantum of energy is the amount of energy required to move an electron from
E. The higher the energy level occupied by an electron, the
it takes to move from that energy level to the next higher energy level.
Current atomic model- Quantum Mechanical Model
A. Still has the nucleus containing
in the center.
B. Still has electrons outside the nucleus in a low density area
C. The quantum mechanical model determines the
an
electron can have and how likely it is to find the electron in various locations around the
nucleus.
D. This model is based on equations developed by Erwin Schrodinger
E. The probability of finding an electron within a certain volume of space surrounding the
nucleus can be represented as a fuzzy cloud. The cloud is
where the probability of finding the electron is high.
1) The fuzzy cloud is called an electron cloud
F. Principle Energy Levels of Electrons
1) Energy levels and # of electrons in each:
a) Level 1 Holds
b) Level 2 Holds
c) Level 3 Holds
d) Level 4 Holds
2) Principle energy levels are broken down into sublevels called:
3) An atomic orbital is often thought of as a region of space in which there is a
of finding an electron.
4) Each energy sublevel corresponds to an orbital of a
, which describes where the electron is likely to be found.
5) Each orbital contains up to
6) Orbitals include:
7) Summary of Energy Levels and Orbitals
Unit 4 Notes: Chapter 5
Sublevel
Electrons in Atoms
# of Orbitals
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Maximum # of
Electrons in
sublevel
S
P
D
F
8) Summary of Energy levels and orbitals
Principle
# of Sublevels
Identity of
Energy Level
Sublevels
1
2
3
4
Shape
# of orbitals
Maximum # of
electrons
The n represents the principle energy level
The values for “n” are the same as the period numbers.
Notice the s block has 2 columns because s can hold 2 electrons.
The d block has 10 columns because d can hold up to 10 electrons
The p block has 6 columns because p can hold up to 6 electrons
You can use this diagram to help you identify the last electrons in the series of
electrons in an element. For example: magnesium is in period 3 column 2 in the s
block. This means the last electrons in the series are in the principle energy level of
3 and the sublevel of s. Since it is in column 2, there are 2 electrons in the s
sublevel.
Sulfur is in period 3 in the 4th column of the p block. This means the last electrons
are in the principle energy level of 3 in the sublevel of p. There are 4 electrons in
the p sublevel.
F block is the part of the periodic table that is written below the main portion. It’s principle energy levels are
indicated by n-2.
Unit 4 Notes: Chapter 5
Electrons in Atoms
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A summary of terms:
- Principle Energy level- How far an electron is from the
o The value for the outermost s and p electrons is found by looking at the period
number of the element
o The value for the outermost d electrons is found by looking at the period number
and subtracting 1.
o The value for the outermost f electrons is found by looking at the period number
and subtracting 2.
- Sublevel- Indicates the
in which the electrons are found- includes s, p, d, and f
o The value for the outermost electrons if found by looking at the “block” of the
periodic table
- Number of electrons in the outermost sublevel
o Found by looking at the really small number above the column in which the element
is found.
- Put the principle energy level, sublevel, and number of electrons together like this:
Section 5.2- Electron Configurations
Goals for this section:
- Write electron configurations for various elements using the orbital diagram method, the
aufbau chart method and the periodic table method
- Explain why electrons fill orbitals the way they do
- List and explain the 3 rules for writing electron configurations
- Explain why some elements have electron configurations different from those predicted by
the aufbau principle
- Identify the number of valence electrons in elements using the electron configurations
- Know these terms: electron configuration, aufbau principle, Pauli exclusion principle, Hund’s
rule
I.
Stable Atoms
A. Change generally proceeds toward the
B. High energy systems are unstable and
to become more stable
1) (This is what you see when the crushed wintergreen mints emit light or when the
samples of gases glowed- electrons are falling from a high energy level to a lower
energy level)
C. Electrons are arranged with
possible
level
This arrangement is called the electron configuration
II.
Electron Configurations
A. The ways in which electrons are arranged in various
around the
nuclei of atoms are called electron configurations.
B. Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell
you how to find the electron configurations of atoms.
C. Aufbau Principle
1) According to the aufbau principle, electrons occupy the orbitals of
first.
D. According to the Pauli exclusion principle, an atomic orbital may hold at most
Unit 4 Notes: Chapter 5
Electrons in Atoms
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1) To occupy the same orbital, two electrons must have
2) The opposite spins are indicated by
pointing in
3) Spinning electrons produce
, which then allow the
electrons to
to each other
E. Hund’s rule states that electrons occupy orbitals of the same energy in a way that
makes the number of electrons with the same spin direction as large as possible.
1) In other words, electrons fill orbitals
and have
spins.
2) After all orbitals in a sublevel have one electron, added electrons double up in
orbitals and have opposite spins to electrons already there.
3) Example:
F. How to write electron configurations without making an orbital diagram:
a. Look at the periodic table to find the correct number of
i. (look at the atomic number to find protons and for neutral atoms, the
number of electrons matches the number of protons)
b. Use the Aufbau chart to figure out where the electrons go. Start with
and follow the arrows.
c. The numbers at the top of the chart indicate the
of electrons that can be placed in that sublevel.
d. Fill up each sublevel until you have reached the correct number of electrons,
which you found earlier on the periodic table
Unit 4 Notes: Chapter 5
Electrons in Atoms
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Aufbau Diagram
Check your work:
Use the periodic table to help you check your work.
The last electrons written in your electron configuration should be the same as the principle
energy level, sublevel, and number of electrons in the outermost energy level as indicated on
your periodic table!
G. Valence Electrons electrons that determine the
of elements
1) Since they are in the outer shell, you can find out how many are present by counting
the electrons in the
2) Oxygen example:
H. Shorthand electron configurations
1) find the element you want on the
2) look up
and find the
in that row
3) Write the noble gas chemical symbol in brackets
4) Look back at the periodic table and find the information above the element you’re
working with
5) Fill in the values for n and n-1 by looking at the row the element is in.
6) Make sure you pay attention to the presence of possible d and f orbitals if your
element is in the p block.
7) You’ll need to look at the aufbau chart for this
8) Example of Shorthand notation:
a)
I.
Remember to use your periodic table to help you find the electron
configuration for the electrons after the noble gas.
Exceptions to the Aufbau principle
Unit 4 Notes: Chapter 5
Electrons in Atoms
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1) According to the aufbau chart, copper should have this electron configuration:
2) Instead, it has this electron configuration:
3) Exceptions occur in groups 6B and 1B because:
a. Sublevels are most stable when they are
b. Sublevels are fairly stable when they are
c. Sublevels lack stability when they are
d. The d sublevel becomes more stable in groups 6B and 1B by stealing an
electron from the
Section 5.3- Physics and the Quantum Mechanical Model
Goals for this section:
- Describe the relationship between wavelength and frequency of light
- Identify the source of atomic emission spectra
- Explain how the frequencies of light are related to changes in electron energies
- Distinguish between quantum mechanics and classical physics
- Know these terms: amplitude, wavelength, frequency, hertz, electromagnetic radiation,
spectrum, ground state, photons, Heisenberg uncertainty principle
I. Physics and the Quantum Mechanical Model
A. Electrons and light
a. Electrons and light both exhibit
at the same time!
b. The “particle” of light is known as a
B. What is a wave?
a. Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible
light, ultraviolet waves, X-rays, and gamma rays.
b. All electromagnetic waves travel in a vacuum at a speed of
c. The wavelength and frequency of light are
to each other
C. Terms associated with a wave:
1) The amplitude of a wave is the wave’s
from zero to the crest.
2) The wavelength, represented by  (the Greek letter lambda), is the distance between the
.
3) The frequency, represented by  (the Greek letter nu), is the number of
to pass a given point per unit of time.
a) The SI unit of cycles per second is called a hertz (Hz). A hertz = 1/second
4) Equation and value of constant:
5) Example problem:
D. Separating Light
a. A prism separates light into the colors it contains.
b. When white light passes through a prism, it produces a rainbow of colors.
c. The rainbow is called a
Unit 4 Notes: Chapter 5
Electrons in Atoms
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E. Atomic emission spectrum : the
formed when
atoms absorb energy, forcing electrons into
energy levels. These electrons
the lose energy by
when they return to
energy levels.
a) The light is made up of only a few specific
,
depending on the element
b) Each frequency is a different
c) The light is emitted as electrons fall from one
to another, like from n=4 to n=1
d) They are like atomic fingerprints- every element is unique
1) An Explanation of Atomic Spectra
a) When the electron has its
, the
atom is in its ground state.
b) Excitation of the electron by
raises
the atom from the ground state to an excited state. (not the orbital predicted by
aufbau chart)
c) A quantum of energy in the form of light is emitted when the electron drops back to a
lower energy level.
a. This energy is directly proportional to frequency, which determines the light’s
color.
d) Light emitted by an electron moving from a higher to lower energy level has a
frequency
to the energy change of the
electron
1. Equation describing energy change of the electron
2. E = h x ν (h = 6.626 x 10-34J*s)
3. So different energy level drops result in different frequencies (and
colors) of light
F. Why does it matter that an electron behaves as both particle and wave?
1) The fact that electrons behave as waves leads to some odd observations, like:
2) Heisenberg’s uncertainty principle- it is impossible to know exactly both the
of a particle at the same time.
a. This limitation is critical in dealing with small particles such as electrons.
i. Just the act of observing an electron changes what that electron does!
b. This limitation does not matter for ordinary-sized object such as cars or
airplanes.
3) When photons hit metals, they can result in the
- This is the photoelectric effect
a. We use this property to create electron microscopes, which allow for clearer
images since electrons have
than light
Unit 4 Notes: Chapter 5
Electrons in Atoms
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Extra Into: Quantum Numbers:
• Principal Quantum Numbers
 Represented by: n
 Main E level of electron
 # E sublevels = # of main level
 N cannot be zero and must be a positive integer
• l Quantum number indicating orbital or sublevel – angular quantum number
 l is any integer between 0 and n-1
 So for n = 2, l can be 0, 1
 l = 0 means s
 l = 1 means p
 l = 2 means d
 l = 3 means f
 Magnetic quantum number m can be any integer from - l to + l
 So if
 l = 0, m must be 0 (indicates s with 1 orbital)
 l = 1, m must be -1, 0, or 1 (indicates p with 3 orbitals)
 l = 2, m must be -2, -1, 0, 1, 2 (indicates d with 5 orbitals)
 l = 3, m must be -3, -2, -1, 0, 1, 2, 3 (indicates f with 7 orbitals)
 Spin quantum number- s- indicates angluar momentum of electrons
 Is either +1/2 or -1/2
 Written as (n, l, m, s)
 How do orbitals relate to quantum numbers: http://en.wikipedia.org/wiki/Atomic_orbital
Principle
Possible
Orbitals
Magnetic
Spins
Quantum
quantum # (n) angular
present
quantum
numbers
quantum
numbers
written
numbers (l)
(n, l, m, s)
1
2
3
4
Unit 4 Notes: Chapter 5
Electrons in Atoms
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