Download Quantum Mechanics and Split Peas - EC Chemistry Lab 2015-16

Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Name:
Patterns in Electron Configuration
One of the many patterns contained in the periodic table is that of electron configuration. In
this activity, you will identify these patterns. Later, you will use these patterns to determine
the order in which electrons fill the orbitals of an atom. As you complete the activity, keep the
following in mind:
• Period = row, Group = column
• Use the table on your book cover, which shows only valence electrons.
• There are two number systems for the Groups. We will focus on the A/B system.
1.
Which Groups have an s-orbital as the last orbital?
2.
Which Groups have a p-orbital as the last orbital?
3.
Which Groups have a d-orbital as the last orbital?
4.
Which section of the table is left? This section corresponds to the f-orbitals.
5.
Look at Group 1A. What is the relationship between the Period number and the
energy level of the valence electrons?
6.
Look at Group 3A. What is the relationship between the Period number and the
energy level of the valence electrons?
7.
Look at Group 3B. What is the relationship between the Period number and the
energy level of the d-orbitals?
8.
Look at the Inner Transition Metals (bottom section). The Lanthanide series (58-71)
is part of Period 6. The Actinide series (90-103) is part of Period 7. What is the
relationship between the Period number and the energy level of the f-orbitals?
9.
Look at all of the A Groups. What is the relationship between the Group number (1A,
2A, etc.) and the total number of valence electrons for each element? (Add up the
exponents to find the total number of valence electrons.)
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Name:
Quantum Mechanics and Split Peas
I. Introduction
Bohr was oh-so-close to explaining the true nature of the electron. However, Bohr’s
theories were too simplistic for multi-electron systems. It was evident that a new type of
thinking would be needed to describe the “what’s” and “where’s” of the electron. Luckily,
some pretty famous scientists (including one named Einstein) were already on the task.
Quantum mechanics basically threw out Newtonian physics and applied completely
different rules to describe things. The important thing to remember is that “quantum”
describes things that are really small and really fast.
II. Purpose: To demonstrate ideas of quantum mechanics, orbitals, and probability models.
III. Materials:
* 40 ml of split peas
* funnel
* target
IV. Procedure
1. Separate the peas into equal volumes in the two graduated cylinders (20 ml in each).
2. Position the target so that the funnel is directly over the center of the target.
3. Have one member of the group place a finger over the end of the funnel and then add
20 ml of the peas into the funnel
4. Release the peas so that they fall onto the target. If they clump too much, re-do the
drop from a higher distance. Record the distance you dropped from.
5. Use the rules below to count how many peas are in each area of the target.
Record your data on the accompanying data sheet.
6. After completing the counting of the first run, collect the peas and repeat steps 1 – 5
from a different distance, recording this as the 2nd trial.
7. Repeat steps 1-5 with all 40 ml of peas for the 3rd trial.
8. When done, place all of the peas back in a container and clean up.
**Rules for counting:
o If a pea is completely within an area, it belongs to that area.
o If the pea is on a line, it belongs to the area that the greater portion of its volume
occupies.
o If the pea is on a line, and seems to be equally in two areas, it belongs to the area
nearest the center.
o Any part of a pea counts as a pea.
6
5
4
3
2
1
The areas are
numbered as this
picture.
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Name:
Quantum Mechanics and Split Peas Data Table:
Trial 1 – Distance ___________
Area
# of peas
% of total
#
Trial 2 – Distance ___________
Area # of peas % of total
#
Trial 3 – Distance ___________
Area # of peas % of total
#
1
1
1
2
2
2
3
3
3
4
4
4
5
5
5
6
6
6
7
7
7
Total
Total
Total
Answer the following questions to draw conclusions from this lab:
1. In the first run, which area did the most peas land in?
2. Did this change for the second run? The third?
3. If you were to do this again, only you had to predict where 90% of the peas would fall, what
would your prediction be?
4. How do these results mimic the orbitals described in the quantum mechanical model of the
atom?
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Name:
ELECTRON ORBITAL ACTIVITY
Background Info: The arrangement of electrons within the orbitals of an atom is known as the
electron configuration. The most stable arrangement is called the ground-state electron
configuration. This is the configuration where all of the electrons in an atom reside in the lowest
energy orbitals possible. Keeping in mind several “rules” that the electrons must obey, we are able
to predict the electron configurations of elements using the electron orbital diagrams and periodic
table.
Electron Orbital Diagram
The distributions of orbitals can be
laid out in the electron orbital diagram
pictured to the left. You read it from
the bottom (or closest to the nucleus)
to the top in order of increasing
energy.
7d
Increasing energy
6f
7p
6d
5f
7s
6p
The bottom energy level is level 1 - it
has the lowest energy.
5d
4f
6s
5p
Each " " represents an orbital.
4d
5s
4p
You can see that there is 1 orbital for
an s subshell. There are 3 orbitals for
a p subshell, 5 for a d, and 7 for an f
subshell. Each orbital can hold up to
2 electrons.
3d
4s
3p
3s
2p
2s
Therefore, the s subshell can hold 2
electrons; the p can hold 6; the d can
hold 10; and the f can hold 14. See chart below for a summary:
Nucleus
1s
Subshell
s
p
d
f
Number of
Orbitals
1
3
5
7
Shape of
orbital
sphere
dumbells
Max # of
electrons
1x2 = 2
3x2 = 6
5x2 = 10
7x2 = 14
Example of the Configuration of H:
Atomic number 1 means at ground state, 1 electron and 1
proton. Therefore, 1 electron in the first energy level, sub level
s, orbital 1 - 1s1
Recall, however, the “d-block” anomaly – the following diagram
may also be helpful in understanding how electrons fill energy
levels once the “(n – 1)” d-block comes into play.
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
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Procedure:
1.
For each of the following atoms or ions, determine the number of protons and electrons
contained and record them in the Data Table.
2.
Determine the correct number of electrons (chips)
according to the following three rules.
A. Place markers (electrons) in the energy levels from
low energy to higher energy (starting at level 1
which is closest to the nucleus)
B. Place one marker (electron) in each box (orbital)
before doubling up
C. Each box has a max capacity of two. When doubling
up within a box (orbital), be sure the one marker
(electron) has its arrow facing up while the other is
down.
3.
Record how your completed card looks (draw in arrows)
on the electron orbital diagrams.
4.
Also record the information for the electrons’ energy level, subshell, and number of
electrons in the data table.
5.
Convert the information from the table into the electron configuration. Electron
configurations give the address information for the electron’s location for the ground state
of an atom. See below for a key to writing electron configurations. The first one, carbon,
has been done for you as an example.
6.
Remove the chips and repeat for the next atom.
Atom
# of protons
Carbon
6
Energy level
Subshell
1, 2, 3, …
s, p, d or f
1
s
2
s
2
p
Electron Configuration:
Example:
# of electrons
6
# of electrons
2
2
2
1s2 2s2 2p2
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Data Tables:
1. B - Boron
Period (row)
# of protons
# of electrons
Energy level
1, 2, 3, …
Subshell
s, p, d or f
# of electrons
Name:
Electron Orbital Diagrams:
Electron Configuration:
2. Ne - Neon
Period (row)
# of protons
# of electrons
Energy level
1, 2, 3, …
Subshell
s, p, d or f
# of electrons
Electron Configuration:
3. Cl - Clorine
Period (row) # of protons # of electrons
Energy level
1, 2, 3, …
Subshell
s, p, d or f
# of electrons
Electron Configuration:
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
4. V - Vanadium
Period (row) # of protons
# of electrons
Energy level
Subshell
# of electrons
1, 2, 3, …
s, p, d or f
Name:
Electron Configuration:
5. Cu - Copper
Period (row) # of protons # of electrons
Energy level
Subshell
1, 2, 3, …
s, p, d or f
# of electrons
Electron Configuration:
6. Br - Bromine
Atom
# of protons # of electrons
Energy level
Subshell
1, 2, 3, …
s, p, d or f
# of electrons
Electron Configuration:
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Name:
Now let’s try some ions: Ions are formed when atoms lose or gain electrons in an effort to
create more stable electron arrangements.
Data Tables:
7.
Atom
# of protons # of electrons
+2
Be
Energy
Subshell
# of electrons
level
1, 2, 3, …
s, p, d or f
Electron Orbital Diagrams:
Electron Configuration:
8.
Atom
N-3
Energy
level
1, 2, 3, …
# of protons # of electrons
Subshell
# of electrons
s, p, d or f
Electron Configuration:
Mystery Elements: Identify the following elements based on their electron configurations.
Then re-write the configuration using Nobel Gas abbreviated notation.
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
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Concluding Questions
1. Electrons can be thought of “obeying” three basic rules when it comes to their location within
an atom. Match the following rules with the three from procedure THREE that you just used to
complete the lab.
Rule
Hund’s Rule
Single electrons with the same spin must occupy each equal-energy orbital
before additional electrons with opposite spins can occupy the same orbitals.
Pauli Exclusion Principle
A maximum of two electrons may occupy a single orbital, but only if the
electrons have opposite spins.
The Aufbau Principle:
Each electron occupies the lowest possible energy orbital
Letter?
2. What is the maximum number of electrons that can be in one orbital (or box)? ____
3. How many electrons are needed to fill the first energy level?____ How many elements are there
in the first period or row of the periodic table? ____
4. How many electrons are needed to fill the second energy level (2s + 2p)? ____. How many
elements are there in the second period of the periodic table? ____
5. How many electrons are needed to fill the third energy level (3s + 3p + 3d)? ____. How many
elements are there in the third period of the periodic table? ____ Wait a second…that seems
strange!
6. What element(s) did you write the configuration for that ended with an electron configuration
of 3d? ________________________________
What row or period is it in? ________
7. Something is unusual about the 4s and 3d orbitals. What do you notice about the order in
which they are filled?
8. Why do you think Groups 1 and 2 referred to as the s-block of the periodic table? (look at the
ending of the electron configurations you would write for elements K and Be)
9. Why are Groups 13 through 18 referred to as the p-block of the periodic table? (look at the
ending of the electron configurations you wrote for elements C, Ne, Al, and Cl)
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
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10. Why are Groups 3 through 12 referred to as the d-block of the periodic table? (look at the
ending of the electron configurations you wrote for element V and Cu)
11. The most stable elements have full outer sub shells. Next most stable are those with half filled
shells. Take a look at the electron configuration you wrote for Copper. Copper’s actual
configuration is: 1s2 2s2 2p6 3s2 3p6 4s1 3d10. Give one possible reason for this.
12. Using what you now know about electron configurations explain the notion that elements in
the same column in the periodic table have similar chemical and physical properties.
13. What element would have the following electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d5.
(For example: 1s2 means that in the first energy level there are two electrons in an s orbital.)
14. Write out the electron configuration of O2-. Notice that this is an ion!
15. Write out electron configuration of Na+. Notice that this is a ion!
16. Look at the electron configurations for O2- and Na+ which you figured out above. Compare
them to the other electron configurations you did. Do they have anything in common with one
or two of them? Which one(s) and Why?
17. Look at the electron configuration for Neon, a noble gas. Why do you think they are more
stable than the other elements?
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Name:
7d
6f
7p
6d
5f
7s
6p
5d
4f
6s
5p
4d
5s
4p
3d
4s
3p
3s
2p
2s
1s
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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure
Name:
Atomic Structure and Theory Magic Square
Directions: Put the number of the definition from the list below into the square with the
appropriate term. Check your answers by adding the numbers to see if all the sums of all rows,
both across and down add up to the same number, the Magic #.
Democritus
Dalton
Thomson
Chadwick
Total
_____
_____
_____
______
_____
Rutherford
Proton
Atom
Bohr
_____
_____
_____
_____
Wave Model
Neutron
Nucleus
Alpha particle
_____
_____
_____
_____
Electron
Model
Energy levels
Electron cloud
_____
______
_____
_____
Total _____
_____
_____
_____
_____
_____
______
Magic Number ______
1. Represented by a symbol; all are found on the Periodic Table
2. Made a mental model of the atom; Greek philosopher
3. Used by Rutherford in his experiment; made of two protons and two neutrons
4. The paths in which electrons circle the nucleus according to the Bohr model
5. The positive particle in the nucleus of an atom
6. The tiny positive core of an atom; contains protons and neutrons
7. Formed the atomic theory model of the atom; English schoolteacher
8. Discovered the nucleus using his gold foil experiment
9. Current explanation of where electrons might be found in the atom
10. Used by scientists to explain something we cannot see or understand
11. The smallest particle of an element that has the properties of that element
12. Discovered the neutron
13. Current model of the atom; proposed by Schrodinger
14. Mass of protons and neutrons
15. Developed the model of the atom in which electrons orbit the nucleus in energy levels
16. The negative particle that circles the nucleus
17. The neutral particle in the nucleus of an atom
18. Proposed the “plum-pudding” model of the atom; discovered the electron
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