Download Unit 10 packet

Name: __________________________
Period: ___________
Unit 10 Packet – Atomic History/Structure and Quantum Mechanics
Packet Contents Sheet with Objectives (This Page)
Worksheet 1: Thermal energy, radiation, and color
Early Discoveries About the Atom Reading
Bohr’s Model Reading
Atomic Structure Notes
Atomic Structure Practice
What Ions will Main Block Elements Form?
Isotopes and Average Atomic Mass
Worksheet 2- Analyzing a Spectrograph
Worksheet 3- Isotopes and Molar Mass
Atomic History and Structure Study Guide
Atom, Ion, Isotope Activity
Quantum Mechanics Notes
Periodic Table (to fill in with the notes)
Orbital Diagramming Worksheet
Study Guide: Atomic Structure/Electrons
DO NOT, under any circumstances, throw this away!
This packet MUST be saved for the final exam.
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Unit 10 – Structure of the Atom - Objectives
Students can distinguish the line spectra of light emitted by atomic gases from the continuous
spectra emitted by hot metals related to thermal energy and understand quantum mechanics
(Three Main Rules) to be able to complete orbital diagrams and electron configurations for a
variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley,
Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem with
each of the previous models, based on experimental evidence, and describe the experiments
conducted to collect the supporting evidence.
Scale
Score
Score 4
Score 3
Score 2
Score 1
Score 0
Comment
Without any major errors, students can independently:

Students can analyze the line spectra of light emitted by atomic gases and the continuous
spectra emitted by hot metals related to thermal energy and use quantum mechanics (Three
Main Rules) to be able to complete orbital diagrams and electron configurations for a
variety of elements. For each of the scientists Thomson, Millikan, Rutherford, Moseley,
Chadwick, Bohr, students can: draw the models of the atom proposed, state the problem
with each of the previous models, based on experimental evidence, and describe the
experiments conducted to collect the supporting evidence.
Without any major errors, students can independently:

Students can distinguish the line spectra of light emitted by atomic gases from the
continuous spectra emitted by hot metals related to thermal energy and understand quantum
mechanics (Three Main Rules) to be able to complete orbital diagrams and electron
configurations for a variety of elements. For each of the scientists Thomson, Millikan,
Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom
proposed, state the problem with each of the previous models, based on experimental
evidence, and describe the experiments conducted to collect the supporting evidence.
With one or two major errors, students can independently:

Students can recognize the line spectra of light emitted by atomic gases and the continuous
spectra emitted by hot metals related to thermal energy and understand that quantum
mechanics (Three Main Rules) can be used to complete orbital diagrams and electron
configurations for a variety of elements. For each of the scientists Thomson, Millikan,
Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom
proposed, state the problem with each of the previous models, based on experimental
evidence, and describe the experiments conducted to collect the supporting evidence.
With help from the teacher, students can:

Students can distinguish the line spectra of light emitted by atomic gases from the
continuous spectra emitted by hot metals related to thermal energy and understand quantum
mechanics (Three Main Rules) to be able to complete orbital diagrams and electron
configurations for a variety of elements. For each of the scientists Thomson, Millikan,
Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom
proposed, state the problem with each of the previous models, based on experimental
evidence, and describe the experiments conducted to collect the supporting evidence.
Even with the teachers help, students show no understanding or ability to:

Students can distinguish the line spectra of light emitted by atomic gases from the
continuous spectra emitted by hot metals related to thermal energy and understand quantum
mechanics (Three Main Rules) to be able to complete orbital diagrams and electron
configurations for a variety of elements. For each of the scientists Thomson, Millikan,
Rutherford, Moseley, Chadwick, Bohr, students can: draw the models of the atom
proposed, state the problem with each of the previous models, based on experimental
evidence, and describe the experiments conducted to collect the supporting evidence.
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Name
Date
Pd
Chemistry – Unit 10 Worksheet 1
Thermal energy, radiation and color
1. Based on the IR security camera images, what is the relationship between
the temperature of an object and the intensity of radiation it emits?
2. Based on the demonstration of the light bulb, what other property of
radiation changes with temperature, in addition to intensity?
3. Describe the relationship between the color change of the filament with
temperature, and the colors in the observed spectrum of white light.
4. Order the colors of the rainbow according to the amount of energy needed to
produce them, from lowest to highest.
5. Based on the spectrum of the fluorescent lamp, what is the main difference
between the radiation emitted by a heated solid (filament) and the radiation
emitted by an atomic gas (mercury gas in the fluorescent tube)?
6. What can you say about the energy spread of the radiation emitted by an
atomic gas?
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Early Discoveries About the Atom
You have already learned about the concept of atoms and how they can combine to form
compounds. You have also seen that each element has its own characteristic set of properties which help
to distinguish it from all other elements. In this chapter we will study the structure of atoms and the laws
governing the behavior of the particles that make up atoms. This knowledge will lead to an explanation of
the properties of the elements and of their tendencies to form compounds.
John Dalton regarded the atom as a particle with no internal parts. He believed an atom to be the
smallest possible particle. However, certain experiments were being performed which gave definite
indications that Dalton's view was not correct and that there was some sort of internal structure to the
atom. It became apparent that atoms consisted of particles that had electrical charges and that these
particles interacted according to the laws of electromagnetism. Charged particles carry either a positive
(+) or negative (-) charge. We call two negative charges or two positive charges "like" charges. The laws
of electromagnetism state that two like charges {1}_______________ each other, while unlike charges,
{2}_____________ each other.
Experiments performed during the late 1800's and early 1900's by chemists and physicists made it
clear that atoms could, indeed, be broken into smaller parts, contrary to the ideas of {3}_____________.
In 1897, J.J. Thomson discovered that atoms could be "taken apart" when he studied the effects of
electrical discharge on atoms of various gases. In his experiments he concluded that atoms were coming
apart by yielding a stream of negatively charged particles with very small masses (compared to the
masses of the atoms). These small negative particles became known as electrons. Thomson is credited
with the discovery of electrons which were present in the atoms of all of the different gases that he
examined.
Another scientist, Ernest Rutherford, and his students performed experiments in England during
the first decade of the 20th century in an attempt to determine the size of atoms. In 1906, Rutherford had
his students direct a beam of positively-charged subatomic "alpha" particles at a very thin sheet of gold
metal. It was known as the alpha scattering experiment." Since they believed that matter was mostly
empty space, they expected the particles to pass through the thin sheet unhindered. To their surprise, they
found that a small fraction of the particles bounced right back! This led Rutherford to believe that in the
center of the atom was a small but very dense "nucleus" with which some of the alpha particles must have
collided. He concluded that most of the mass of the atom was contained in the {4}__________________.
He also concluded that the nucleus was {5}______________ charged since it repelled the positivelycharged alpha particles. He later said that "It was quite the most incredible event that has ever happened
to me in my life. It was almost as if you fired a 15- inch shell into a piece of tissue paper and it came back
and hit you."
Rutherford realized that electrons were located at a considerable distance from the nucleus. If this
were an accurate description of an atom and we could inflate the nucleus of a hydrogen atom to the size of
a basketball, the electron would orbit this "basketball" nucleus at a distance of more than 15 miles away!
Visualizing the atom like this enables you to realize that most of the atom is, indeed, nothing more than
empty space! So much for the athlete who thinks he is "solid muscle!"
With regard to the nucleus itself, it became obvious to scientists that the nucleus was composed
of small particles. One of the particles in the nucleus is the proton. In 1914, Rutherford was given credit
for discovering protons. A proton carries a charge equal to the charge of an electron, but opposite in
character. The electron carries a charge of -1 while the proton carries a charge of {6}_______. In the
nucleus of a neutral atom (that is, an atom with no overall electric charge), there must be an equal number
of protons to balance the charges carried by the electrons. Unlike the electron, the proton is a particle with
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a relatively large mass, in atomic terms. The proton has a mass equal to 1,836 times that of the electron!
Using the common unit of the gram to measure masses, the electron has a mass equal to 9.1 X 10-28
grams, and the proton has a mass equal to 1.673 X 10-24 grams.
A second component of the nucleus was discovered in 1932 by another
Englishman, James Chadwick. This particle became known as the neutron. A
neutron is an electrically neutral particle, meaning that it carries no electric
charge. The neutron was very difficult to discover. Because it has no charge of its
own, it is neither attracted to nor repelled by an electrical charge. A neutron has a
mass slightly larger than that of the proton, equal to 1.675 X 10-24 grams. The
presence of this particle accounted for the observed masses of atoms, which were
found to be greater than that predicted if only protons were present in the
nucleus. Figure 12.1 and the accompanying table present an overall summary of
the locations and properties of the components of an atom.
Each element differs from all others in that atoms of each element
contain a specific number of electrons, protons, and {7}_______________.
Indeed, the number of protons in the nucleus determines the actual identity of an
element. Determining the number of electrons, protons, and neutrons in any
given element is a relatively simple process. The atomic number of an element is
equal to the number of protons found in the nucleus. The element with atomic
number 19, potassium, has 19 protons. For atoms to be neutral, they must have
equal numbers of positive (protons) and negative (electrons) charges. This means that potassium must
have 19 protons and {8}________ electrons.
Particle
Electron
Proton
Neutron
Table 12.1
Location and Properties of Subatomic Particles
Charge
Comparative Mass
-1
1/1836
+1
1
0
1
Location
Outside nucleus
Inside nucleus
Inside nucleus
If an element has 10 protons in its nucleus, how many positive charges does it have?{9}________.
If an atom with 10 protons is neutral, how many electrons must it have?{10}__________
The atomic number of oxygen is 8. How many protons does it have?{11}__________
How many electrons does oxygen have?{12}___________
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Bohr’s Model
Niels Bohr developed a theory to account for the location of
electrons around the nucleus of an atom. He believed his theory would
explain the bright-line spectra emitted by excited atoms. He proposed
that electrons followed specific paths or orbits around the nucleus.
These paths or energy levels, as they are also called, are numbered
starting with the lowest one (closest to the nucleus) as 1, the next
farther from the nucleus as 2, the next as 3, and so forth. Figure 12.8
illustrates 4 of the energy levels of a n atom.
Bohr's model resembled a planetary system like our solar
system in which he suggested that the electrons revolve around the
nucleus. The first energy level, nearest the nucleus, is represented as
number 1. Each level thereafter is increased by one. A total of 7 energy
levels are needed to explain the structure of all of the elements.
The orbits around the nucleus are called energy levels because there are different and very
specific energies associated with each level. Bohr knew that energy was being added to an atom when it
was being heated or when an electrical current was passed through an element. This extra energy has to
go somewhere. The added energy is absorbed by the electrons that are in the outermost orbit (farthest
from the nucleus). Since there are specific amounts of energy associated with each energy level, the
electron that absorbs all of this extra energy can no longer stay in the orbit (energy level) in which it
normally belongs which is its "ground state." Instead, it will move to another energy level. It is now in an
"excited state" and is what we earlier referred to as an "excited" electron.
This electron is not doomed to spend the rest of its time in this higher energy level. Indeed, the
excited electron is now very unstable. Because this is an unstable state, the electron will soon return to its
ground state. This is where the atomic spectra enters into the picture. The energy that the electron emits
when it returns to its ground state is in the form of light and heat. Figure 12.9 gives a general picture of
this process. In "A”, the electron is excited and jumps to a new energy level that is further from the
nucleus. As the electron falls back to its ground state in "B", energy is given off in the form of light.
As electrons move to lower energy levels, the energy they emit is given off in a "burst" or
quantity of energy with a well-defined wavelength. The quantities of emitted energy are called quanta.
Bohr's theory of the atom gave birth to the quantum theory, a name that reflects the notion that atoms
must absorb and emit energy in specific amounts. Therefore, we say that the energy is "quantized." A
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quantum of energy can be defined as the amount of energy needed to move an electron from one energy
level to the next higher one. Similarly, it can be defined as the amount of energy emitted when an electron
moves from its present energy level to a lower one.
In his theory, Bohr proposed that electrons were only "allowed" to exist at certain distances from
the nucleus. These distances became his energy levels. He believed that electrons were not "allowed" to
exist between these levels. Although the reasons for this behavior were unclear, the idea did explain the
existence of bright-line spectra. The spectrum of hydrogen contains a specific set of bright lines. Of these,
only 3 or 4 are clearly visible to the naked eye. (See Figure 12.2) So Bohr asked, why only this set of
lines? Why does the excited hydrogen electron emit only this specific set of wavelengths?
According to Bohr, the fact that there was always the same set of bright lines in the hydrogen
spectrum was evidence that only certain energy changes were possible for the hydrogen electron. The
electron could only make certain "jumps" and, therefore, could only emit certain wavelengths of light.
The number of lines was limited because there were only a few "excited" energy levels to which the
electron was "allowed" to move. He explained that the electron was - for whatever reason - not
“permitted" to exist between these levels.
Since every element has a specific number of electrons, different wavelengths of light will be
emitted by excited electrons of atoms of different elements. Even without a spectroscope you can see that
the glow of a neon sign is quite different from the fluorescent lights in your classroom. You probably did
not realize that what you were seeing was the result of excited electrons at work.
To better understand how an electron behaves as it returns to its stable ground state, consider how
you would jump down a staircase. Just as an electron can only stop at certain energy levels, you can only
stop at certain levels (steps). The electron cannot exist between energy levels, and you cannot stop
between the steps.
To get down to the bottom of the staircase, you
could jump all the way down at once; or one step at a
time; or one step followed by a two-step jump; or
perhaps a two-step jump followed by one step, and then
by a three-step jump. A variety of different jumps is
possible, each yielding a different amount of energy.
Likewise, suppose an electron jumps out to level
4 in Figure 12.10. There are ten different sizes of jumps
possible as it returns to its ground state. So how many
different wavelengths of light could be emitted by this
electron? {31}__________ A sample containing many
excited atoms of this element would emit all possible
wavelengths.
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Atomic Structure Notes
Neutral Atoms
Ions
Isotopes
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Atomic Structure Notes
P = ________________ (+)
E = ________________ (-)
N = ________________ (0)
Reading a Periodic Table:
3
Li
Lithium
6.94
OR
7
3Li
Lithium = Element Name
Li = Element Symbol
3 = Atomic Number: number of protons, defines the element *
7 = Mass Number: number of protons and neutrons combined *
6.94 = Atomic Mass: weighted average mass of all isotopes of the element
*Never fractional protons, neutrons, or electrons- these must be whole #s
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Element or
Ion
Symbol/Ion
Atomic #
Mass #
P
15
30
# of
Protons
# of
Neutrons
# of
Electrons
Nitrogen
207
82
40
19
56
Oxygen Ion
82
16
10
17
5
18
2
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What Ions will the Main Block Elements Form?
If an atom loses electrons it forms a ( + or - ) ion.
If an atom gains electrons it forms a ( + or - ) ion.
In the table below:
 Draw a Bohr diagram of the first element in each family of main block elements.
 What ion will the element most likely form to get a full outer shell?
 On your periodic table write the ion sign (ex: +2) above the column of the element.
Main Group Element
Bohr Diagram of Neutral Atom
Ion
Li
Be
B
C
N
O
F
He
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Atomic Structure Worksheet Ions
Ion
H+
Li+
Na+
K+
FClBrO-2
Ca+2
Sn+4
S-2
Cu+3
N-3
Mg+2
Zn+2
Al+3
P-3
Type of Ion
Cation
Protons
1
Electrons
0
Anion
16
18
Loss of e- Gain of e1
X
X
2
Draw a model for the following atoms and ions:
Al
Al+3
S
S-2
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Name
Date
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Chemistry – Unit 10 Worksheet 2
Analyzing a Spectrograph
A mass spectrometer is an instrument used to separate an element's isotopes and to
measure their relative abundances. Within this device, a sample of an element is
vaporized, then ionized and accelerated down a tube. Near the end, the beam of ions is
passed through a strong magnetic field which exerts a force on the ions. Ions of greater
mass possess more inertia, or more of a tendency to continue to move in a straight line, and
so deviate only slightly from their projected path. Ions of lesser mass are more greatly
influenced by the field and demonstrate greater deviation. Examine the three mass
spectrograph readings illustrated below and answer the questions that follow. Note that
the upper scale of each spectrograph shows atomic mass (in amu). Below each
spectrograph, the percents of the various isotopes present are given.
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l. a. What is the molar mass of the isotope of the element represented by spectrum
A?
b. What are the name and atomic symbol of element A?
2. a. What are the symbols, including superscripts and subscripts for the isotopes
in
spectrum B?
b. Based on the experimentally obtained values of atomic mass and percent
abundance, calculate the average molar mass of this element. Show your
work.
c. Which isotope deviated most from its straight-line path?
3. a. Calculate the average molar mass of the element in Spectrum C.
Show your work.
b. What are the symbols, including superscripts and subscripts, of the isotopes of
this element?
c. Which isotope deviated the least from its straight-line path?
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Name
Date
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Chemistry – Unit 10
Worksheet 3
1. Element X has two natural isotopes. The isotope with a mass number of 6 has a
relative abundance of 7.5%. The isotope with a mass number of 7 has a relative
abundance of 92.5%. Determine the average molar mass for the element from these
figures. What is the true identity and atomic number of element X?
2. The element copper is found to contain the naturally occurring isotopes 29Cu63 and
29Cu65. The relative abundances are 69.1% and 30.9% respectively. Calculate the
average molar mass of copper.
3. Uranium has three isotopes with the following percent abundances: 92U234
(0.0058%), 92U235(0.71%), 92U238 (99.23%). What do you expect the molar mass of
uranium to be in whole numbers? Why?
4. A sample of silver as it occurs in nature is 52.0% of isotope 47Ag107 and 48.0% of
isotope 47Ag108. What is the average molar mass of silver? (Compare your result
with the value given in the periodic table).
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5. Ninety-two percent of the atoms of an element have a mass of 28.0 amu, 5.0% of the
atoms have a mass of 29.0 amu, and the remaining atoms have a mass of 30.0 amu.
Calculate the average molar mass and identify the element.
6. Use the following isotope data for lead to show that its molar mass is
207 amu,
(1.37%)
82Pb204
(26.26%)
82Pb206
82Pb207
82Pb208
(20.82%)
(51.55%)
7. Boron exists in the form to two stable isotopes, boron-10 and boron-11. These occur
in the abundance of 19.6 percent and 80.4 percent respectively. Calculate the
average molar mass of boron.
8. Precise molar masses of each isotope of magnesium are given below along with the
percent abundance of each isotope:
magnesium -24
23.98504
78.70%
magnesium-25
24.98584
10.13%
magnesium-26
25.98259
11.17%
Calculate the average molar mass of magnesium.
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Name
Date
Pd
Chemistry – Unit 10 Study Guide: Atomic History and Structure
1. Define:
a. Atom
b. Ion
c. Isotope
d. Atomic Number
e. Atomic Mass
f. Nucleus (Where is it? What’s in it? What charge?)
g. Electrons (Where are they? What charge?)
2. List the three subatomic particles that make up atoms.
3. List the charge and relative atomic mass (amu) of each of the three subatomic particles
that make up atoms.
4. Atomic mass number is the number of ________________ + _________________.
5. Which of the elements below are isotopes of each other?
a. 12C
b. 1H+
c. 14C-4
d. 2H
6. Which of the elements above are ions?
7. A new element Cranium is discovered in Chandler. Samples show that naturally
occurring Cranium is 99.2% 250Ce, 0.6% 255Ce, and 0.2% 260Ce. What will the atomic
mass of Ce be on the periodic table?
8. If Ce has an atomic number of 112, how many neutrons are in 250Ce?
9. How many electrons are in a neutral atom of Cranium?
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10. Draw the following atoms using the Bohr model.
a.
31
b.
30
c.
30 -3
d.
24
P
P
P
Mg+2
11. List the five postulates of Dalton’s Atomic Theory:
a.
b.
c.
d.
e.
12. Complete the following chart:
Chemical
Symbol
N
Atomic
number
Number of
protons
Number of
electrons
Number of
neutrons
15
19
17
Mass
number
Charge
30
40
18
56
82
O-2
16
5
2
13. Complete the attached Atom, Ion, Isotope 4-square activity for Fluorine (F).
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14. For each of the scientists listed below
a. Identify their major contribution to our understanding of the atom.
b. Draw the atomic model of each and order them chronologically from earliest to
most recent.
Dalton
Rutherford
Democritus
Thomson
Bohr
Chadwick
15. The gold foil experiment was used by ____________________ to show that the atom had
a dense central region called a __________________.
16. The name of Thomson’s experiment was __________________________ and it showed
the atom contained __________________.
17. The modern day model of the atom is called the _______________________ or
____________________. Draw a picture of the current day model.
18. Draw Bohr diagrams of the following neutral atoms and identify what ions they form.
Element
Neutral Atom Diagram
Chemical
Symbol of
Ion
K
Li
Cl
F
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Atom (Neutral)
Definition:
Complete Chemical Symbol
Bohr Diagram:
Subatomic Particles
P=
E=
N=
Ion
Definition:
Complete Chemical Symbol
Bohr Diagram:
Subatomic Particles
P=
E=
N=
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Isotope
Definition:
Complete Chemical Symbol
Bohr Diagram:
Subatomic Particles
P=
E=
N=
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Quantum Mechanics Notes
Key Words
And Questions
Quantum Mechanical

Model of the Atom
Notes

-
Electron Movement
Electrons are moving in waves, not in orderly
straight orbitals

In ___________ NOT in ______________

Basketball Court Model
Atoms
Structure of Energy
Levels
↓
Principle Quantum #
(Lowest n = 1…. Highest n = 7)
↓
Sublevels:
Orbitals:
Total # of Electrons:
s
p
_____
_____
↓
↓
↓
↓
□
□□□
_____
__________
↓
↓
↓
↓
2
6
10
_____
Each individual orbital holds _______ electrons.
# of sublevels in each
quantum level
n=1
→
s
n=2
→
s,p
n=3
→
_____, _____, _____
n=4 → _____, _____, _____, _____
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Notation
1s2
Electron Configuration
n
Sublevel
1
s
# of e2
3p6
_____
_____
_____
4d10
_____
_____
_____
The distribution of ______________ among the
____________ of an atom.

Rules
How many electrons (we’ll use this later)?
H _____
Be _____
He _____
B _____
Li _____
Mn _____
1. Aufbau Principle (Snake Diagram):
2. Pauli Exclusion Principle:
3. Hund’s Rule (Roommate Rule):
Practice:
Note- Use snake.
Exponents must total
the number of
electrons in the atom
H
1 electron = 1s1
He 2 electrons = 1s2
Li
3 electrons = 1s22s1
Be ___ electrons = __________
B
___ electrons = __________
Mn ___ electrons = __________
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Name
Date
Pd
Chemistry – Unit 10 Orbital Diagramming Worksheet
For each of the following elements draw an orbital diagram (box and arrow form) and write the
correct abbreviated electron configuration.
Element Symbol
Orbital Diagram
Electron Configuration
1. H
2. He
3. Li
4. Be
5. B
6. C
7. N
8. O
9. F
10. Ne
11. Na
12. Mg
13. Al
14. Si
15. P
16. S
17. Cl
18. Ar
19. K
20. Ca
21. Br
22. Kr
23. I
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Name
Date
Pd
Chemistry – Unit 10 Study Guide: Atomic Structure/Electrons
1. How do isotopes of an element differ from one another?
2. How is an ion formed?
3. An atom contains 3 protons, 4 neutrons, and 3 electrons. What is its atomic number,
mass number, and name?
4. What are the two names of the modern atomic model?
5. True or False: Like charges attract and unlike charges repel. Give evidence to support
your answer.
6. The distance from one crest of a light wave to the next successive crest of a light wave is
called what?
7. Light has a dual nature. List two ways in which light behaves.
8. What is the relationship between frequency, wavelength, and energy?
9. True or False: Waves on the high energy, short wavelength side of the electromagnetic
spectrum are NOT dangerous to humans. Give evidence to support your answer.
10. Describe the difference between an orbit and an orbital.
11. What is the principal quantum number of the most stable, lowest energy level and where
it is located in comparison to the nucleus?
12. True or False: Electrons must have a certain minimum amount of energy, called a
quantum of energy, in order to move from one energy level to a higher energy level.
Give evidence to support your answer.
13. Which principal energy level has the lowest energy? The highest?
14. Draw the general shape of the s and p orbitals.
15. How many s orbitals come in a group? p orbitals? d orbitals? f orbitals?
16. How are s orbitals represented in orbital diagrams? p? d? f?
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17. The following electron configurations belong to which elements?
1s22s22p63s2
1s22s22p63s23p5
18. Write the electron configurations for the following elements:
Germanium
Calcium
Titanium
Uranium
19. List and define the three rules that govern how electrons are placed in electron
configurations.
a.
b.
c.
20. How does the 3s orbital differ from the 2s orbital?
21. Which sublevels can be found in the fourth principal energy level of an atom?
22. Draw orbital diagrams for each of the following elements:
Mg
Sc
Ne
Pt
23. Memorize the family names and location on the periodic table.
(Nothing to write here, just get it done)
24. Know the regions on the periodic table for metals, non-metals, and semi-metals.
(Nothing to write here, just get it done)
Modeling Chemistry Unit 10 Packet
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