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Enloe Honors Chemistry
Unit 5
Name:_____________________________
Date:_______________Period:_________
Atomic Structure, Electromagnetic Spectrum, Quantum Theory, and Electron Configuration
Chm.1.1.1 Analyze the structure of atoms, isotopes, and ions.
Chm.1.1.2 Analyze an atom in terms of the location of electrons.
Chm.1.1.3 Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model.
Chm.1.3.2 Infer the physical properties (atomic radius, metallic and nonmetallic characteristics) of an element based on its position
on the Periodic Table.
Introduction
In this Unit we will discuss the location of the nucleus, protons, neutrons, electrons and the electron cloud. In one model of
the atom, the planetary model, electrons are described as orbiting the nucleus like planets orbiting the sun. As it turns out the
location of these electrons is much more complicated than this.
Knowing the electronic structure of an atom provides the key to understanding the physical and chemical properties of
elements. In this unit we will go through the development of this theory and learn about the currently accepted theory of electronic
structure.
Students should be able to:
1.
Characterize the protons, neutrons, electrons: location, relative charge, relative mass (p=1, n=1, e=1/2000).
2.
Use symbols: A= mass number, Z=atomic number
3.
Use notation for writing isotope symbols: 235 U or U-235
92
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons
Analyze diagrams related to the Bohr model of the hydrogen atom in terms of allowed, discrete energy levels in the
emission spectrum.
Describe the electron cloud of the atom in terms of a probability model.
Relate the electron configurations of atoms to the Bohr and electron cloud models.
Understand that energy exists in discrete units called quanta.
Describe the concepts of excited and ground state of electrons in the atom:
a. When an electron gains an amount of energy equivalent to the energy difference, it moves from its ground state to a
higher energy level.
b. When the electron moves to a lower energy level, it releases an amount of energy equal to the energy difference in
these levels as electromagnetic radiation (emissions spectrum).
Articulate that this electromagnetic radiation is given off as photons.
Understand the inverse relationship between wavelength and frequency, and the direct relationship between energy and
frequency.
Use the “Bohr Model for Hydrogen Atom” and “Electromagnetic Spectrum” diagrams from the Reference Tables to relate
color, frequency, and wavelength of the light emitted to the energy of the photon.
Explain that Niels Bohr produced a model of the hydrogen atom based on experimental observations. This model indicated
that:
a. an electron circles the nucleus only in fixed energy ranges called orbits;
b. an electron can neither gain or lose energy inside this orbit, but could move up or down to another orbit;
c. that the lowest energy orbit is closest to the nucleus.
Describe the wave/particle duality of electrons.
Write electron configurations, including noble gas abbreviations (no exceptions to the general rules). Included here are
extended arrangements showing electrons in orbitals.
Identify s, p, d, and f blocks on Periodic Table.
Identify an element based on its electron configuration. (Students should be able to identify elements which follow the
general rules, not necessarily those which are exceptions.)
Determine the number of valence electrons from electron configurations.
Predict the number of electrons lost or gained and the oxidation number based on the electron configuration of an atom.
1
Assignment 1
: Isotopes, Atomic Numbers and Mass Numbers
(Sec. 2.3 Brown and LeMay or Chang)
1.
Comparison of subatomic particles
Particle
Proton
Neutron
Electron
Charge
Positive (+1)
None (neutral)
Negative (-1)
Mass, amu
1.0073 (approximately 1)
1.0087 (approximately 1)
0.0005486 (negligible)
Location
Nucleus
Nucleus
Electron Cloud
2.
What makes an atom of one element different from an atom of another element?
3.
All atoms of the same element have the same number of _____________________.
4.
Since atoms are electrically neutral the number of ____________________ must equal the number of
__________________________.
5.
Atomic number (Z) =
6.
Mass Number (A) =
7.
Define isotopes.
8.
Define ion.
Fill in the gaps in the tables below
Symbol
Protons
Neutrons
Electrons
Mass no.
Dash Notation
20
Symbol
Protons
Neutrons
Electrons
Atomic no.
Mass no.
Dash Notation
31
F
73
As
56
83
69
51
78
196
Fluorine - 20
P
26
30
80
53
50
119
2
127
201
Assignment 2
: Atomic Structure Practice
On another sheet of paper or in your composition notebooks (depending on your teacher) complete numbers
2:11- 2.20 (2.14, omit d)
3
Assignment 3
: Ionic Structure
Ions – a former atom that has lost or gained electrons to become charged.
*When electrons are _______ the charge is positive
Example:
Symbol
Protons
Neutrons
Electrons
Mass No.
Atomic No.
Electric Charge
Atom
22
Na
11
10
Ion
22
Na1+
11
10
22
11
22
11
Notice the difference between the number of protons and electrons…
*When electrons are __________ the charge is negative
Example:
Symbol
Protons
Neutrons
Electrons
Mass No.
Atomic No.
Electric Charge
Atom
81
Br
35
46
Ion
81 1Br
35
46
81
35
81
35
Notice here, just like in the previous example, the number of protons in the atom and Ion does
not change!
4
Assignment 4
: Ionic Structure Practice
Fill in the blanks on the tables below.
1. Ions
Symbol
Protons
Neutrons
Electrons
Atomic no.
Mass no.
Net Charge
2. Ions
Symbol
Protons
Neutrons
Electrons
Mass no.
Net Charge
3. Ions
Symbol
Protons
Neutrons
Electrons
Atomic no.
Mass no.
Net Charge
4. Ions
Symbol
Protons
Neutrons
Electrons
Atomic no.
Mass no.
Net Charge
32
P3-
200
Hg2+
26
31
53
54
3+
37
Cl1-
88
50
118
2+
127
Sr2+
21
23
28
31
26
3+
17
O2-
52
3-
Cr3+
38
50
34
45
36
2+
40
33
76
36
16
18
74
54
1-
K1+
17
27
63
2+
5
13
25
3+
F1-
Assignment 5 : The Wave Nature of Light
Brown p. 200-201; Chang p. 244-247
1.
What part of the atom interacts in a chemical change?
2.
What is electronic structure?
3.
What does electronic structure relate to?
4.
What has lead to much our present day understanding of the structure of the atom?
5.
What is electromagnetic radiation?
6.
What is another name for electromagnetic radiation?
7.
Give several examples of electromagnetic radiation.
8.
Waves have 3 measurable properties or characteristics. Define and/or comment on each.
a. speed
b. wavelength
c. frequency
9.
What is the electromagnetic spectrum?
6
10.
Which type of electromagnetic radiation has the:
a. shortest wavelength?
b. longest wavelength?
11.
What are the frequency and wavelength of visible light?
12.
List the colors of the visible spectrum in order of increasing frequency.
13.
Which color has the:
a. longest wavelength
b. shortest wavelength?
14.
Write the mathematical equation that shows the relationship between light, wavelenght, and
frequency. Identify what each symbol represents.
15.
As wavelength decreases, frequency ______________________. As wavelength increase,
frequency ____________________________. Therefore, wavelength and frequency are
________________ proportional to each other.
Assignment 6 : Radiant Energy Problems
Work problems 6.1-6.10. Problems are to be worked on separate paper or comp. notebook.
7
Assignment 7 : Radiant Energy Problems 2
1.
Complete the table
Quantity
SI Unit
Other Units
Wavelength
Frequency
Speed
Energy
2.
Convert each of the following to meters and put them in scientific notation.
a. 410 nm = _______________m
b. 835 nm = _______________m
c. 147 nm = _______________m
3.
Determine the frequency of light with a wavelength of 4.257 x 10 2 m.
4.
The wavelength of light is 310 nm. Calculate its frequency.
5.
What is the wavelength of electromagnetic radiation if its frequency is 3.2 x 10 –2 Hz ?
6.
What distance does light travel in 2.50 minutes?
7.
Calculate the speed of a wave whose wavelength and frequency are 17.4 cm and 87.4 Hz,
respectively.
8.
Calculate the frequency, in Hz, of a wave whose speed and wavelength are 713 m/s and 1.14 m,
respectively.
8
9.
The wavelength of the green light from a traffic signal is centered at 522 nm. What is the
frequency of this radiation?
10.
What is the frequency of light having a wavelength of 456 nm?
11.
What is the wavelength of radiation whose frequency is 6.24 x 10 14 s-1.
12.
What is the frequency of radiation whose wavelength is 3.55m ?
13.
Would you be able to see either of the radiation specified in 10 and 11 ? If so, what color would
you see? (Use the diagram of the electromagnetic spectrum in your reference packet)
14.
A neon light emits radiation of 616 nm wavelength. What is the frequency of this radiation?
Using the diagram in your reference packet, predict the color associated with this wavelength?
15.
Excited barium atoms emit visible light whose frequency is 6.56 x 10 14 s-1. What is the
wavelength of this light? Use Figure 4-1 to predict its color.
16.
The average distance between Mars and Earth is about 1.3 x 10 8 miles. How long would it take
TV pictures transmitted from the Viking space vehicle on Mars’ surface to reach Earth?
17.
How many minutes would it take a radio wave to travel from the planet Venus to Earth?
(Average distance from Venus to Earth is 28 million miles)
9
Assignment 8 : Quantized Energy and Photons
Brown p. 202-205; Chang p. 247-250
Hot Objects and the Quantization of Energy
1.
The amount of radiant energy emitted by an object depends on
its____________________________________.
2.
Could the wave theory of light totally explain the emission of light from hot objects?
3.
What did Max Planck suggest concerning the emission of energy by hot objects?
4.
What is a quantum?
5.
Write the mathematical expression that shows the relationship between a quantum of energy
and the frequency of radiation? Identify what each symbol represents.
6.
What is the numerical value of Planck’s constant?
7.
According to Planck’s theory, energy is always absorbed or emitted in ____________
____________.
8.
Explain how walking up a staircase versus walking up a ramp can be compared to quantized
energy?
9.
Why is Planck’s theory of quantized energy not obvious in our daily lives?
The Photoelectric Effect and Photons
1.
What is the photoelectric effect?
2.
What is a photon?
3.
How did Einstein explain the photoelectric effect?
10
Assignment 9
: Planck’s Quantum Theory and Photoelectric Effect Problems
Show neat, well organized work for each problem. Include units in work and answer. Give answer to correct
number of significant digits.
1. Calculate the frequency of a quantum that has an energy of 1.88 x 10 -18 J.
2. Calculate the energy of a quantum that has a wavelength of 405 nm.
3. What type of electromagnetic radiation has quanta with energy of 3.83 x 10 -19 J ?
4. Barium has a threshold frequency of 6.07 x 10 14 Hz.
a. If barium is illuminated with light that has a wavelength of 385 nm will electrons be ejected from the
surface of the metal?
b. In which region of the electromagnetic spectrum does this radiation occur?
c. Will yellow light produce the photoelectric effect?
d. List all types of radiation that will produce the photoelectric effect.
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5. Sodium has a threshold frequency of 5.51 x 10 14 Hz.
a. Will microwaves eject electrons from the metal?
b. What is the lowest energy type of electromagnetic radiation needed to break the electrons from the
metal?
6. The energy required to release an electron from atoms on the surface of a certain metal is 7.96 x 10 -19 J.
a. What wavelength of light would be necessary to cause electrons to leave the surface of the metal?
b. To which region of the electromagnetic spectrum does this type of radiation belong?
c. Will blue light produce the photoelectric effect in this metal? Explain your answer.
d. List all types of electromagnetic radiation that will produce the photoelectric
effect in this metal.
e. What effect would subjecting this metal to x-rays have on the electrons ejected from this metal?
7. What happens if the intensity of the light increases, but the frequency remains the same?
8. What happens if light of a shorter wavelength is used? What if light of a longer wavelength is used?
12
Assignment 10 : The Hydrogen Atom Line Emission Spectrum
Brown p. 205-209; Chang p. 250-253
1.
Define:
a.
ground state
b.
excited state
2.
What happens when an excited electron returns to its ground state?
3.
When hydrogen’s excited electrons return to their ground state they emit only specific
frequencies of light that indicate the energy differences between the atom’s energy states were
fixed. What does this observation suggest about the distance of hydrogen’s one electron from
the nucleus?
4.
What makes up a line spectrum?
5.
Complete the chart in class during the demonstration.
Gas Tube
Color of Light in Tube
Colors in Line Spectrum
6.
State at least 4-5 points that describe Bohr’s Model of the hydrogen atom.
7.
What was the major shortcoming of Bohr’s model?
8.
Bohr’s model of the hydrogen atom along with all of the other experimental data of Planck and
Einstein led to the revision of the current atomic model. The new model was known as the
__________________ ___________________. To explain the line emission spectrum of
hydrogen, light not only was considered to have wavelike propertites but also to have
properties of ______________. This is known as the ____________________________ of light.
13
Assignment 11 : Bohr Model Practice
Use the Bohr model provided to complete the following table.
Electron movement
1.
n = 1 to n = 4
2.
n = 6 to n = 2
3.
n = 6 to n = 3
4.
n = 1 to n = 2
5.
n = 5 to n = 3
6.
n = 2 to n = 4
7.
n = 3 to n = 1
8.
n = 2 to n = 5
9.
n = 2 to n = 3
wavelength, nm
emitted or
absorbed
Type of Radiation
10. n = 4 to n = 3
Circle the correct answer.
11. Which of the following energy differences is greatest?
A.
n=4 to n=3
B.
n=3 to n=2
C.
n=1 to n=2
D.
n=5 to n=6
12. Which of the following transitions will emit light in the visible region of the spectrum?
A.
n =4 to n=1
B.
n=4 to n=3
C.
n=2 to n=4
D.
n = 5 to n=2
13. Which form of electromagnetic radiation has the greatest energy?
A.
visible light
B.
radiowaves
C.
infrared light
D.
ultraviolet light
14. Which form of electromagnetic radiation has the greatest frequency?
A.
blue light
B.
orange light
C.
ultraviolet light
D.
infrared light
14
Assignment 12 : The Quantum Model of the Atom
Brown p. 210; Chang p. 254-258
1.
Physicists questioned why the energies of electrons are quantized. In other words, why is the
electron in a Bohr atom restricted to orbiting the nucleus at certain fixed distances?
2.
In 1924, Louis de Broglie provided a solution to this puzzle.
3.
de Broglie reasoned that if light waves can behave like a stream of particles (photons) then
perhaps particles such as electrons can possess wave properties.
4.
de Broglie compared the behavior of Bohr’s quantized electron orbits to the known behavior of
waves.
He hypothesized that electrons are confined to the space around an atomic nucleus and that
electron waves exist only at specific energies.
5.
de Broglie suggested that the electron orbiting the nucleus possessed wavelike characteristics.
6.
He also suggested that the electron in its circular path about the nucleus has associated with it
a particular wavelength.
7.
He went on to suggest that the characteristic wavelength of the electron or any other particle
depends on its mass and velocity.
 = wavelength
h = Planck’s constant
m = mass
 = velocity
m = momentum
8.
If electrons produce a specific wavelength then they must also have a corresponding specific
frequency. According to the equation E = h, these frequencies correspond to specific energies
– the quantized energies of Bohr’s orbits.
9.
Within a few years after de Broglie published his theory, the wave properties of the electron
were experimentally demonstrated. It was found that electrons could be diffracted by crystals
just as x – rays are diffracted. This lead to the development of the electron microscope.
10.
Electron microscopes use a beam of electrons. The wavelike behavior of electrons is utilized in
the same way as a conventional microscope uses wave characteristics of light.
15
The Heisenburg Uncertainty Principle
Brown p. 211; Chang p. 258-260
1.
The idea of electrons having a dual wave-particle nature troubled scientist. They questioned if
electrons are both particles and waves, then where were they in the nucleus.
2.
A wave extends in space and its location is not precisely defined. Therefore, it stands to reason
that we can’t precisely determine its exact position, direction of motion and speed of motion at
any time.
3.
The German physicists, Werner Heisenberg, concluded that the dual nature of matter places a
limitation on how precisely we can know both the location and momentum of any object.
4.
Heisenburg’s idea involved the detection of electrons. Electrons are detected by their
interaction with photons. Because photons have about the same energy as electrons, any
attempt to locate a specific electron with a photon knocks the electron off its course. As a
result, there is always a basic uncertainty in trying to locate an electron.
5.
The Heinsenberg Uncertainty Principle states that is it impossible to determine simultaneously
both the position and velocity (or momentum) of an electron or any other particle.
6.
Thus it is not appropriate to imagine the electron as moving in well-defined circular orbits
about the nucleus.
7.
de Broglie’s hypothesis of the dual nature of the electron and Heisenberg’s Uncertainty
Principle set the stage for a new description of the electron.
The Schrodinger Wave Equation
Brown p. 212-213; Chang p. 260-261
1.
In 1926, Erwin Schrodinger proposed an equation, known as Schrodinger’s wave equation. This
equation incorporates both the wavelike and particlelike behavior of the electron.
2.
He opened a new way of dealing with subatomic particles known as quantum mechanics.
3.
Remember! Bohr’s model assumes that electrons travel in a circular orbit of some particular
radius around the nucleus. In the quantum mechanical model the electron’s location can’t be
described so simply.
4.
Remember! The Uncertainty Principle suggest that if we know the velocity (or momentum) of
the electron with high accuracy, our knowledge of its location is very uncertain.
5.
In the quantum mechanical model we speak of the probability that an electron will be in a
certain region of space at a given instant.
16
6.
This is represented by the square of the wave function,2, (from Schrodinger’s equation). 2 is
called the probability density.
7.
See Figure 4-11 page 101. Regions where there is a high probability of finding the electron are
said to be regions of high electron density.
8.
The quantum theory describes mathematically the wave properties of electrons and other very
small particles.
9.
Solutions to the Schrodinger wave equation are known as wave functions. Based on the
Heisenberg Uncertainty Principle, the early developers of quantum theory determined that
wave functions give only the probability of finding an electron at a given place around the
nucleus.
10.
Thus electrons do not travel around the nucleus in neat orbits, as Bohr suggested. Instead, they
exist in certain regions called orbitals.
11.
An orbital is a three-dimensional region around the nucleus that indicates the probable location
of an electron.
12.
An orbital has both a characteristic energy and characteristic shape.
13.
Orbitals are not the same as the orbits described by the Bohr model.
14.
The Bohr model introduced a single quantum number, n, to describe an orbit. The quantummechanical model uses 4 quantum numbers to describe orbitals. Three of these quantum
numbers result from solutions to the Schrodinger equation. They indicate the main energy
level, the shape, and the orientation of an orbital. The fourth number, the spin quantum
number, describes a fundamental state of the electron that occupies the orbital.
17
Assignment 13
: Electron Configurations
Brown p. 218-226; Chang p. 267-276
1.
How is the quantum model an improvement on the Bohr model?
2.
Define electron configuration.
3.
Do any atoms of different elements have the same electron configuration? Why or why not?
4.
Define ground-state electron configuration.
5.
a.
Electrons are added to orbitals one by one following three basic rules. Explain each of these rules.
Aufbau principle
b.
Pauli exclusion principle
c.
Hund’s rule
6.
What is an orbital notation?
7.
In writing an electron configuration, such as 1s2, what does the superscript indicate?
8.
What is meant by the highest occupied energy level in an atom?
9.
What are inner shell electrons?
10.
Valence shell electrons are the electrons that take part in bonding. Outer shell electrons are always
the valence shell electrons.
11.
Only electrons in _____ and _____ subshells will ever be part of the outer shell electrons in the ground
state of an atom.
12.
What are the noble gases?
13.
What is a noble gas configuration?
14.
If you inspect figure 6.28 in Brown-Lemay and Table 7.3 in Chang, you will see that the electron
configurations of certain elements appear to violate the rules we have just discussed. Suggest a reason
for the electron configuration for Cr, Cu, and Ag.
Problems
1.
2.
3.
4.
5.
In your composition notebook…
Write complete electron configurations for elements with atomic #’s 1-18, 20-36 even, 37 -55 odd.
Write the noble gas core configuration for elements with atomic #’s 1-18, 20-36 even, 37 -55 odd.
Write orbital notation for elements with atomic #’s 1-18, 20-36 even, 37 -55 odd.
Give the number of valence electrons for each element atomic #’s 1-18, 20-36 even, 37 -55 odd.
Predict the most likely charge for each element above based on their electron configurations.
18
Assignment 14 : Quantum Numbers (Omit for 15-16)
1. Principle Quantum Number ( n )
 Describes the distance of the electron from the nucleus
 Is proportional to the energy of the electron
 Identifies the energy level of the electron
 Allowed values: 1,2,3, …
 In an energy level, n, there are
o n sublevels,
o n2 orbitals, and
o a maximum of 2n2 electrons that can occupy it
Energy level
n=1
n=2
n=3
n=4
Subshells
Orbitals
Max # of electrons
2. Azimuthal ( Angular momentum) Quantum Number ( l )
 Describes the shapes of orbitals
 Identifies the sublevel or subshell in which the electron is located
 Allowed values: l = 0, 1, 2 …(n – 1)
 In a subshell, l , there are 2 l + 1 orbitals, and a maximum of 2 (2 l + 1 ) electrons
Subshell
l=0
l=1
l=2
l=3





“name”
s
p
d
f
# of orbitals
Max # of electrons
The letters s, p, d, f come from the words sharp, principal, diffuse, and fundamental, which
were used to describe certain features of spectra before quantum mechanics was
developed ( Brown, LeMay, Bursten Chemistry the Central Science 9th ed ).
A collection of orbitals with the same value of “n” is called a shell. A collection of orbitals
with the same values of “n” and “l ” is known as a subshell.
Example: n = 2 represents a shell or energy level
o n = 2 l = 0 represents the 2s subshell
o n = 2 l = 1 represents the 2p subshell
o n = 3 represents the ____________________
o n = 3 l = 0 represents the ________________
o n = 3 l = 1 represents the ________________
o n = 3 l = 2 represents the ________________
We can draw a representation of an orbital by drawing a boundary surface diagram that
encloses about 90% of the total electron density in an orbital.
See Page 264 and 265 (Chang) or 216 and 217 (Brown-Lemay): Boundary surface diagrams
of the hydrogen 1s, 2s, and 3s orbitals, boundary surface diagrams of the three 2p orbitals,
and the five 3d orbitals.
19
3. Magnetic Quantum number ( ml )
 Describes the orientation in the x, y, z planes
 Identifies the specific orbital in which an electron could be located
 Allowed values: - l … 0 … +l
Subshell
l=0
# of orbitals
1
Orbitals
s
Quantum numbers
____
m=0
____ ____ ____
m = -1 m = 0 m = +1
____ ____ ____ ____ ____
m = -2 m=-1 m=0 m=+1 m=+2
__ __ __ __ __ __ __
l=1
3
px py pz
l=2
5
d x2 – y2 , dz2 , dxy , dxz, dyz
l=3
7
__ __ __ __ __ __ __
4. Spin Quantum Number ( ms )
 Explained the splitting of lines in the emission spectra of hydrogen when an external magnetic
field is applied
 Electrons are assumed to act like tiny magnets that spin on their own axes (as Earth does )
 The electron has two opposite spins, clockwise and counterclockwise
 Allowed values : + ½ and - ½
 The value has no effect on the energy, size, shape, or orientation of an orbital, but it determines
how electrons are arranged in orbitals of equal energy.
5. Energies of orbitals

In hydrogen, all orbitals with the same value of n are equal in energy.

Orbitals of equal energy are called degenerate orbitals.

Thus the energies of the hydrogen orbitals increase as follows:
1s < 2s = 2p < 3s = 3p = 3d < 4s = 4p = 4d = 4f

The four quantum numbers n, l , m l , ms enable us to label completely an electron in any orbital
in any atom.

We can regard the set of four quantum numbers as the “address” of an electron in an atom,
somewhat in the same way that a street address, city, state, and ZIP code specify the address of
an individual.
20
Assignment 15 : Quantum Numbers Questions (Omit for 15-16)
Brown p. 213 – 218; Chang p. 261-267
1.
What is the principal quantum number?
2.
How do we symbolize the principal quantum number?
3.
What does the value of n tell us about the energy of the electron shell and its location in
respect to the nucleus?
4.
How can we calculate the number of sublevels that exist in a given shell, or main energy level?
5.
What information is given by the angular momentum quantum number? How do we symbolize
this number?
6.
What are sublevels?
7.
For each of the following values of n, indicate the numbers and types of sublevels possible for
the main energy level.(See Table 4-2)
n=1
n=2
n=3
n=4
n = 5 (give number only)
n = 6 (give number only)
n = 7 (give number only)
8.
What information is given by the magnetic quantum number? How is it symbolized?
9.
How many orbital orientations are possible for each of the following sublevels?
s
p
d
f
10.
How can we calculate the number of orbitals in each energy level?
11.
What information does the spin quantum number give us?
12.
13.
What values may the spin quantum number have?
How many electrons can a single orbital hold? Under what conditions must the electrons be?
21
14.
How can we calculate the number of electrons in each energy level?
15.
How many electrons could be contained in the following main energy levels with n equal to:
a.
b.
c.
d.
e.
f.
g.
16.
Principle
quantum
number; main
energy level (n)
1
2
1
2
3
4
5
6
7
Draw and diagrams for the different orientations in space that the s, p, and d orbitals occupy.
Sublevels in
main energy
level
(n sublevels)
Number of
orbitals per
sublevel
Number of
orbitals per
main energy
level (n2)
3
4
5
6
7
22
Number of
electrons per
sublevel
Number of
electrons per
main energy
level (2n2)
Assignment 16 : Quantum Numbers and Electron Configurations
1. An atom of silicon has a total of ________ p electrons.
2. What is the maximum number of electrons that can occupy the n=3 energy level?
3. How many orbitals are in the subshell l = 3 (d subshell)? _____ How many electrons can occupy this
subshell?
4. In a cobalt atom in the ground state, the total number of energy levels occupied by 1 or more electrons is
________.
5. In a cobalt atom in the ground state, there are ______ unpaired electrons.
6. In a cobalt atom in the ground state there are ______ subshells occupied by one or more electrons.
7. The abbreviated electron configuration of Titanium is _________________.
8. Give a possible set of quantum numbers for the last electron added to a gallium atom in its ground state.
n = _____
l = _____ ml = _____ ms = _____ (Omit)
9. A given orbital is labeled by the magnetic quantum number as ml = - 2. This orbital cannot be of which
type? s, p, d, f _______________ (Omit)
10. How many orbitals are filled in a ground state atom of iron? ______
11. The lowest atomic mass atom with a filled 3d subshell in the ground state is ___________.
12. The element with the ground state electron configuration 1s22s2 2p6 3s2 3p6 4s1 3d5 is ________.
13. The total number of energy levels occupied by one or more electrons in the ground state of a technetium
atom is _________.
14. The total number of orbitals occupied by one or more electrons in the ground state of a technetium atom
is _____.
15. Consider the quantum numbers n = 5 l = 1 ml = -1 ms = +½ . The lowest atomic mass element in the
ground state for which these quantum numbers are possible is _____. (Omit)
16. How many subshells are occupied by one or more electrons in an atom of mercury in the ground state?
_____
17. How many unpaired electrons are in an atom of Xenon in the ground state? _____
18. Give a set of quantum numbers for the last electron shown in the ground state orbital diagram of yttrium.
n = _____
l = _____ ml = _____ ms = ____ (Omit)
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19. Only electrons in _____ and _____ subshells will ever be part of the outer shell electrons in the ground
state of an atom.
20. How many outer shell electrons in the ground state of each of the following atoms?
1.
Fe
2.
Cl
3.
He
4.
Ca
5.
P
6.
K
7.
Ar
8.
Sn
9.
Al
21. How many shells are occupied by one or more electrons in the ground state of Arsenic? _____
22. The ground state electron configuration of Al has _____ completely filled shells, one or more electrons in a
total of _____ subshells and _____ orbitals; and a valence shell with ___ filled orbital(s) as well as _____
unpaired electrons.
23. The ground state electron configuration of Mn has _____ completely filled shells, one or more electrons in
a total of _____ subshells and _____ orbitals; and a valence shell with ___ filled orbital(s) as well as _____
unpaired electrons.
24. The ground state electron configuration of I has _____ completely filled shells, one or more electrons in a
total of _____ subshells and _____ orbitals; and a valence shell with ___ filled orbital(s) as well as _____
unpaired electrons.
25. The ground state electron configuration of _____ has 3 completely filled shells, one or more electrons in a
total of 11subshells and 27 orbitals; and a valence shell with 2 filled orbital(s) as well as 2 unpaired
electrons.
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