Ferguson/Hopple
Academic Chemistry
Unit10:
Quantum Mechanics / Wave Mechanical Model of the Atom
Name:__________________________Period: _________
Teachers' Domain: Background Essay: Quantum Mechanical Atom
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Background Essay: Quantum Mechanical Atom
At first glance, the periodic table -- that chart that appears on the walls of science classrooms
everywhere -- appears to be an oddly shaped collection of chemical information about the elements. A
closer look, however, reveals the source of the table's name: The elements are arranged "periodically;"
that is, according to properties that repeat in regular, predictable patterns. This periodic arrangement
of the elements makes the table very useful, in that if you know the location of an element in the table,
you can predict its properties.
The more than 100 elements that make up the periodic table are organized in a series of 18 columns
and 7 rows. Each column is called a group, or family. Each row is called a period. Elements in the
same group have similar physical characteristics. For example, all of the elements in group 1 (at the
far left) react easily with other elements. Unlike the elements in a group, however, the elements in a
period do not share properties. Rather, the properties of the elements change as you move from left to
right across the row. But to understand why the table is organized as it is, it's helpful to understand the
structure of atoms.
An atom is the smallest particle of an element. An atom of any given element is made up of a certain
number of protons, an equal number of electrons, and approximately the same number of neutrons.
(The exception is hydrogen, which can have zero neutrons.) Protons and neutrons form the nucleus of
an atom, and electrons swarm around the nucleus. This swarming isn't completely haphazard, though.
Electrons inhabit various energy levels, or shells. The electron configuration shown in the periodic
table indicates how many electrons are found in each shell, from innermost to outermost. For
example, the electron configuration for calcium is 2, 8, 8, 2.
Electron configuration depends upon the energy state and magnetic spin of each electron, and these
qualities place electrons into particular subshells within each shell. The first shell, for example,
includes only one subshell at the lowest-level energy state and can hold no more than two electrons.
The second shell, with two subshells that contain four levels of energy states, can hold no more than
eight electrons. The subshells containing the lowest energy states fill first, and if a subshell is full,
additional electrons are found in the next higher subshell -- which is generally in the adjacent outer
shell.
Beginning in the fourth shell, there are subshells that have lower energy states than those in the
adjacent inner shell. Since the electrons fill the levels in order of energy, electrons can start filling
subshells in an outer shell before an inner shell is completely full. This explains why, for example, the
electron configuration for calcium can be 2,8,8,2 when the third shell can hold up to 18 electrons.
Elements are arranged in the periodic table according to atomic number, from left to right, top to
bottom. The atomic number of an element is equal to the number of protons found in an atom of that
element. For example, an atom of carbon has six protons in its nucleus; its atomic number is 6. The
elements are also arranged according to atomic mass. The mass of a single proton is equal to 1, while
the mass of a neutron is very close to 1. An atom's atomic mass, then, is close in number to the sum of
its protons and neutrons. An atom of carbon, with six protons and, on average, six neutrons in its
nucleus, has an atomic mass of 12.0107. With the lighter elements, the atomic mass is about double
the element's atomic number. As you move up to the heavier elements, the number of neutrons
relative to protons increases, causing the mass to be increasingly more than double the atomic
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Teachers' Domain: Background Essay: Quantum Mechanical Atom
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number.
When Dmitri Mendeleyev first devised the modern periodic table in 1869, he organized it such that
elements with similar characteristics fell into the same columns. Doing so naturally created rows
within the table. What scientists later found out was that these rows represented something very
significant. They discovered that the elements in each successive row contained an additional electron
shell. For example, the atoms of hydrogen and helium in the first row each had one electron shell;
atoms of elements listed in the second row had two electron shells, and so on to elements in the final
row, whose atoms each have seven shells.
From this, scientists learned what caused elements to have different characteristics. Each element's
physical characteristics are determined, in large part, by the number of electrons in the outermost shell
of its atoms. As with the number of protons, the number of electrons increases by one as you move
across the table from left to right, top to bottom. Atoms of elements in the left-hand column have one
electron in their outer shell, while atoms of elements in the right-hand column have eight electrons in
their outer shell. How does this determine an element's characteristics? Single electrons in an outer
shell can easily be taken away from the atom with the application of very little energy. This makes
atoms of elements in the left-hand column very reactive (and good conductors of heat and electricity).
It is very difficult, on the other hand, to add or remove electrons from an atom that has eight electrons
in its outer shell. The atoms of these elements, found in the column to the far right, are non-reactive.
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Note Taking Guide: Episode 304
Name___________________
Quantum Numbers
• ___________________________________
• Used to __________ an __________ in an __________
n
•
•
__________ __________ __________
Represents __________ energy level of __________
__________ # of __________ in an
__________ __________ = __________
Example: What is the maximum number of electrons that can be in the
_____ main energy level?
l
•
•
•
The __________ __________ __________
Describes the __________ __________ within an __________
__________
__________ of orbital __________ possible in __________
__________ = __________
Orbital Shapes
designated __________________
• level 1: __________
• level 2: __________
• level 3: __________
• level 4: __________
How many electrons can each sublevel hold?
s = 1 orbital x 2 e-/orbital = _______ep = 3 orbitals x 2 e-/orbital = _______ed = 5 orbitals x 2 e-/orbital = _______ef = 7 orbitals x 2 e-/orbital = _______e-
CHEMISTRY: A Study of Matter
© 2004, GPB
3.13
m
•
•
s
•
•
The __________ __________ __________
describes __________ of __________ in __________
The __________ __________ __________
describes __________ of __________ in __________
Ground State: __________ energy arrangement of __________
Diagonal Rule
Examples—
hydrogen _______________
lithium_________________
nitrogen ________________
Orbital Notation
Examples—
hydrogen
nitrogen
CHEMISTRY: A Study of Matter
© 2004, GPB
3.14
Hund's Rule:
__________ of __________ __________ are each __________ by one
__________ before any __________ is occupied by a __________
__________.
Pauli Exclusion Principle:
No two __________ in the __________ __________ can have the __________
__________ of __________ __________ __________.
The Chemistry Quiz
CR1._____
CR2._____
3._____
1._____
4._____
CHEMISTRY: A Study of Matter
© 2004, GPB
3.15
2._____
5._____
Worksheet: Energy Levels, Sublevels, Orbitals
Name___________________
Energy Level
n
# of e- in
Sublevel
E Sublevel
(type of orbital)
# of Orbitals in
Sublevel
1
2
3
4
CHEMISTRY: A Study of Matter
© 2004, GPB
3.16
Total # of ein E level
(2n2)
Worksheet: The Diagonal Rule
To determine the order in which electrons will fill orbitals in an atom,
use the diagonal rule below.
Start at the top
1s
and when you reach the end of one arrow,
return to the next, working your way down.
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
6f
7s
7p
7d
7f
CHEMISTRY: A Study of Matter
© 2004, GPB
3.17
Worksheet: Electron Distributions Review
Name___________________
I. Fill in the blanks:
1. The orbital shaped like a "dumb-bell" is the _____ orbital, while the orbital
shaped spherically is the ______ orbital.
2. How many sublevels are present in the third main energy level?_____
3. What is the maximum number of orbitals in the "d" sublevel?_____
4. The maximum number of electrons that can occupy an orbital is _____,
provided they have ____________ _________.
5. The maximum number of electrons that can occupy an energy level is
represented by the formula _________.
6. The highly probable location of an electron within the atom is a(n) _________.
II. Write the electron configuration for the following:
1. Mg:______________________________
2. As:______________________________
III.
In the space below, show the orbital notation for Mg:
CHEMISTRY: A Study of Matter
© 2004, GPB
3.20
Worksheet: Electron Distributions
Name___________________
1.
There are four types of orbitals:
s :
shaped like a ___________
An E level can contain only _____ s orbital, making up the “s sublevel”.
p :
shaped like ____________
An E level can contain _____ p orbitals, making up the “p sublevel”.
d:
shaped like double dumbbells
An E level can contain _____ d orbitals, making up the “d sublevel”.
f:
too complex to draw or describe
An E level can contain _____ f orbitals, making up the “f sublevel”.
2.
Each orbital can hold a maximum of _____ electrons. Since both electrons
have a __________ charge, they __________. What keeps them from
flying apart?
Each electron _______ on its axis. One spins __________
and the other spins _____________. When charged particles spin,
they act like tiny magnets. Since the two electrons spin
in ___________ directions, one acts like the north pole of a magnet
and the other acts like the south pole. This makes the electrons
_____________ .
3.
Since each orbital can hold _____ electrons:
The “s sublevel” can hold ______ electrons.
The “p sublevel” can hold ______ electrons.
The “d sublevel” can hold ______ electrons.
The “f sublevel” can hold ______ electrons.
We use this notation to describe an electron:
main _______ level
3p5
# of e- in __________
__________
How are electrons distributed within a sublevel?
According to Hund’s Rule, each __________ within a sublevel is
half-filled before any is __________.
CHEMISTRY: A Study of Matter
© 2004, GPB
3.18
We draw orbital diagrams to show the distribution of electrons in a sublevel.
Circles are used to represent the individual ________. __________ are used to
represent electrons in the orbital. The first electron in an orbital is represented
by a ! and the second by a ! .
A set of four ______________ numbers is assigned to each __________ to
describe its energy and location within the atom. The quantum numbers use the
symbols _______, _______, _______, and _______.
_______ is the principle quantum number and represents the ________ level of
the electron.
_______ represents the sublevel of the electron, which depends on the type of
_____________.
Pauli’s Exclusion Principle states that within an atom, no two electrons can have
the same set of _____________ _____________. If two electrons have the
same n, l, and m numbers, they are in the same _______ level, the same
___________, and the same _____________. They must then have
____________spins! So, the s quantum numbers must be different.
Practice: Write electron distributions and do the orbital notation for the
following:
1.
P:
2.
Ca:
Only do the electron distributions for the following:
1.
Co:
2.
Eu:
3.
Tc:
CHEMISTRY: A Study of Matter
© 2004, GPB
3.19
Worksheet 11 - Electronic Structure of Atoms
The Schroedinger equation defines wave equations which describe the
distribution of electrons around the nucleus. The wave functions that satisfy the
Schroedinger equation are called atomic orbitals. They define the allowed
energy states of the electrons.
The energy levels are described by three quantum numbers, n, l and ml.
n is the principal quantum number and has values of 1, 2, 3, ...
The value of n determines the size of the orbital and the energy of electrons in
that orbital.
l is the angular momentum quantum number as has values of 0 through n -1.
The value of l determines the shape of the orbital.
ml is the magnetic quantum number and has values of -l through +l and
determines the orientation of the orbital.
1.
Complete the first two columns of the chart shown below for n = 1
through n = 4.
n l
1 0
2 0
1
3 0
ml
0
orbital name
1s
-1, 0, 1
# e- total e- in level n
2
2
6
3d
4
The l = 0 orbitals are called s orbitals. The l = 1 orbitals are called p orbitals.
For l = 2, 3 and 4, they are called d, f and g orbitals. Each orbital can contain a
maximum of 2 electrons.
2.
Fill in the orbital names and the number of electrons per orbital
and per energy level in the chart.
3.
How many orbitals are present in each of the principal levels?
n=1
n=2
n=3
n=4
n=5
The s orbitals are spherical. They increase in size with increasing values of n.
1s
2s
3s
The p orbitals are dumbbell shaped. Each of the three p orbitals is oriented
differently in space, as shown below.
px
py
pz
Again, size of these orbitals increases with n.
The shapes of the d and f orbitals are more complex and are shown in the
textbook.
4.
Which of the following orbitals can not exist?
1s
5.
3p
7d
3f
4s
2d
8g
Name the orbitals described by the following quantum numbers
(e.g. n = 2, l = 1 would be 2p)
n = 3, l = 0
6.
1p
n = 3, l = 2
n = 3, l = 3
n = 5, l = 0
n = 3, l = 1
How many orbitals in an atom can have the designation:
5p
3s
n=4
4d
n=3
The importance of these orbitals is apparent when we look at the Periodic
Table. Period 1 (H and He) is the n = 1 energy level. This means that there is
only one orbital (1s) available. H has 1 electron and He has 2, completely filling
the 1s orbital. Electrons in the same orbital must have different spin states ( a
4th quantum number), either spin up (↑) or spin down (↓).
s block
p block
n=1
n=2
d block
n=3
n=4
n=5
n=6
n=7
So, the electron configuration for H is 1s1 and He is 1s2, since each s orbital can
hold two electrons.
These can also be drawn out as:
H
1s __
and
He 1s __
In He, the electrons are paired, one spin up and one down. Parallel spins are
not allowed in an orbital. Electron spin is the 4th quantum number, ms, with
values of +½ and -½ .
The next element is Li with 3 electrons or 1s2 2s1
The 2s orbital is higher in energy than the 1s orbital
Li
2s __
1s __
7.
The next element is ______ with ___ electrons.
What is its electron configuration?
Energy level diagram?
8.
Which element has 5 electrons?
Which block is it part of?
Write its electron configuration and draw its energy level diagram.
Remember that the number of orbitals changes with l .
9.
Carbon has 6 electrons, 1s2 2s2 2p2. When we put a second electron in
the p orbitals, Hund's rule states that the electrons should have parallel
spins (remain unpaired) if possible. Add the electrons to the energy level
diagram of C.
2p
2s
1s
10.
__ __ __
__
__
What is the electron configuration of oxygen?
Draw the energy level diagram for oxygen.
Notice that oxygen has unpaired electrons. This means that oxygen is
paramagnetic, and will interact with magnetic fields.
11.
Write out the electron configurations for Ne, Na and Al:
You may have noticed that a lot of the electron configuration is repetitive. Every
atom has 1s electrons. Comparing Ne, Na and Al shows that they are very
similar up to the configuration of Ne. There is a shorthand notation that can be
used. Na is basically [Ne] 3s1 and Al is [Ne] 3s2 3p1. The previous noble gas is
used as a summary of lower state electrons.
12.
In shorthand notation, what is the electron configuration for Ca?
13.
The next element is scandium. Which block is it in? Write the shorthand
electron configuration for Sc. (Hint: look at the table in question 1 to
determine orbitals)
Draw an energy level diagram for Sc.
14.
Write the shorthand notation for the electron configuration of arsenic, As.
15.
Arrange the following orbitals in order of increasing energy.
(Hint: use the Periodic Table for help).
1s
16.
3s
4s
3d
4f
3p
7s
5d
5p
Write the shorthand electron configuration for Cl-. (How many electrons
are present?)
Shown below are four different electronic configurations of carbon:
__
__
__
__
__
__
excited state 3
__
__
__
__
__
__
excited state 2
__
__
__
__
__
__
excited state 1
__
1s
__
2s
__
__
2p
__
__
3s
ground state
The bottom configuration is the ground state (lowest energy) configuration.
The others are excited state (higher energy) configurations.
17.
Describe how each excited state is different from the ground state.
What is wrong with the configuration shown below? Is it an excited state?
__
1s
18.
__
2s
__
__
2p
__
__
4s
Which of the following correspond to an excited state? Identify the atoms
and write the ground state configuration if needed.
a)
1s2
2s2
3p1
b)
1s2
2s2
2p6
c)
1s2
2s2
2p4
3s1
d)
[Ar]
4s2
3d5
4p1
We will only be looking at the first row of the transition metals.
19.
__
1s
20.
__
1s
Write the shorthand electron configuration and energy level
diagram for V.
__
2s
__ __ __
2p
__
3s
__ __ __
3p
__
4s
__ __ __ __ __
3d
The expected shorthand electron configuration and energy level
diagram for Cr is:
__
2s
__ __ __
2p
__
3s
__ __ __
3p
__
4s
__ __ __ __ __
3d
This is not what is observed. Instead, one of the 4s electrons occupies one of
the 3d orbitals.
21.
__
1s
Write the observed shorthand electron configuration and energy
diagram for Cr.
__
2s
__ __ __
2p
__
3s
__ __ __
3p
__
4s
__ __ __ __ __
3d
Notice that this puts 5 electrons in the 3d orbitals, leaving them ½ full. There
seems to be a special stability associated with the half-full orbitals.
This also happens when the 3d orbitals are full. Zn has 10 electrons in the 3d
orbitals. Cu should have only 9 electrons in the 3d orbitals.
22.
__
1s
Write the shorthand electron configuration and energy diagram of
Cu with 10 electrons in the 3d orbitals, which is actually observed.
__
2s
__ __ __
2p
__
3s
__ __ __
3p
__
4s
__ __ __ __ __
3d
Electron Configuration Practice Worksheet
Electron Configuration Practice Worksheet
Write the electron configurations using arrows of the following elements:
1)
Magnesium
2)
Cobalt
______________________________________________________
3)
Krypton
______________________________________________________
4)
Beryllium
5)
Scandium
______________________________________________________
______________________________________________________
______________________________________________________
Write the electron configurations of the following elements:
6)
Nickel
______________________________________________________
7)
Cadmium
______________________________________________________
8)
Selenium
______________________________________________________
9)
Strontium
______________________________________________________
10)
Lithium ______________________________________________________
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Electron Configuration Practice Worksheet
Determine what elements are denoted by the following electron configurations:
11)
1s22s22p63s23p5 ____________________
12)
1s22s22p63s23p64s23d104p65s2 _____________________
13)
[Kr] 5s24d105p4 ___________________
14)
[Xe] 6s24f145d7 ___________________
15)
[Rn] 7s25f12 ___________________
Determine whether the following electron configurations are or are not valid:
16)
1s22s22p63s23p64s24d104p6
__________________
17)
1s22s22p63s33d6
_________________
18)
[Rn] 7s25f9
__________________
19)
[Xe]
20)
[Ne] 3p5 3s2
__________________
__________________
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Name:
Block:
Electron Configurations
Give the electron configuration (orbital notation—with the arrows) for each of the following elements:
1. carbon
2. potassium
3. silicon
4. silver
For each of the following electron configurations, name the element.
5.
1s
2s
2p
6.
1s
2s
2p
3s
3p
4s
3d
Each of the following electron configurations has something wrong with it. For each one:
• State what the mistake is.
• Re-write the electron configuration correctly, keeping the total number of electrons the same.
7.
1s
2p
8.
1s
2s
9.
1s
2s
2p
3s
10.
1s
2s
2p
3s
2p
3p
3p
4s
4d
Honors Chem WS
ELECTRONS
1. Draw the electron configurations (boxes and arrows) for the following elements:
a. Nitrogen
b. Chlorine
c. Neon
2. Write the electron configurations for the following elements:
a. Aluminum
b. Silver
c. Francium
3. Write the noble gas notation for the following elements:
a. Sulfur
b. Barium
c. Iodine
4. The following questions refer the the configuration:
1s22s22p63s23p64s1
a. What is the total number of protons in this atom? Explain your reasoning:
b. The number of electrons in this element’s outer-most principle energy level is:
c. Which element is this?
5. The following questions refer to this configuration:
a.
b.
c.
d.
1s22s22p63s23p5
Which element is this?
How many protons are in this atom?
How many electrons are in this atom’s outer-most principle energy level?
Will this atom most likely lose or gain electrons? How many?
6. The following questions refer to this configuration:
1s22s22p63s23p1
a. Which element is this?
b. If this atom has 15 neutrons, what is its mass?
c. How many electrons are in its second principle energy level?
Chemistry
Name _____________________________
Period
Electron Configurations of Atoms and Ions Activity
Purpose: To make observations of metal ion solutions and relate them to electron configurations
Materials:
0.1 M NaCl
0.1 M MgSO4
0.1 M AlCl3
2 well plates
0.1 M CaCl2
0.1 M NiSO4
0.1 M NaOH
0.1 M FeCl3
0.1 M CuSO4
0.1 M ZnCl2
0.1 M AgNO3
0.1 M Na2CO3
Procedure:
1. On a separate sheet of paper, draw three grids similar to Figure A.
2. In Grid 1, add two drops of each chemical in to a well plate. Record your observations and
describe the color & clarity of each chemical.
3. In Grid 2, predict which of the metal cations in this experiment will form colored precipitates
upon the addition of NaOH.
Do not continue until the teacher has seen your predictions.
4. Using a second well plate (you will need your unreacted chemicals to compare to), carry out
an experiment and record your results in Grid 2. What color are the precipitates? How
accurate were your predictions? Wash and dry your well plate.
5. In Grid 3, predict which of the metal cations in this experiment will form colored precipitations
upon the addition of Na2CO3. Carry out an experiment and record your results in Grid 3. What
color are the precipitates? How accurate were you predictions?
6. Wash and dry both of your well plates.
Figure A
NaCl
MgSO4
AlCl3
FeCl3
CaCl2
NiSO4
CuSO4
ZnCl2
AgNO3
Analysis and Conclusions: Using your experimental data, answer, in complete sentences, the
following questions.
1. Transition-metal ions having partially-filled d orbitals usually have a color. Transition-metals usually lose
3+
2+
s orbital electrons first. Write electron configurations for (a) Fe and Fe , and for (b) Ni and Ni .
2. Write the exceptional electron configurations of (a) Cu and (b) Ag.
3. Explain why the exceptional electron configurations of Cu and Ag are more stable than the expected
configurations.
+
2+
4. The solutions of both Ag and Zn
configurations? (b) Write them.
ions have no color. (a) What does this suggest about their electron
5. Write the electron configurations for Cr
3+,
2+
2+
2+
Cd , Hg , and V .
6. Predict which of the following transition metal ions has color: Cr
3+,
2+
2+
2+
Cd , Hg , and V .
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