General Chemistry
Unit 2
“Bunsen, I must tell you how excellent your study of
chemical spectroscopy is, as is your pioneer work in
photochemistry—but what really impresses me is that
cute little burner you’ve come up with.”
 History of the atomic theory
 Atomic structure
 Isotopes
 Ions
 Average atomic mass
 % abundance
1
At the conclusion of this unit, the student will be able to:
1.
2.
3.
4.
5.
6.
7.
Describe the development of the atomic theory.
Describe the law of conservation of mass/matter.
Describe the law of definite proportions.
Describe the law of multiple proportions.
Describe the location and amount of the subatomic particles for a specific isotope.
Define and distinguish between an ion and an atom.
Calculate the average atomic mass of an element or isotope percent/relative abundances.
We are looking for:
1a. Description of each scientists’ contribution to the development of the atomic theory.
2. Law of conservation of mass/matter: matter cannot be created nor destroyed (Lavoisier).
3. Law of definite proportions: atoms combine in specific whole number ratios (Proust).
4. Law of multiple proportions: two elements can combine in different ratios to form different
compounds; ex. CO and CO2 (Dalton).
5a. Isotopes are atoms of the same element that differ in the number of neutrons.
5b. The nucleus is the center of the atom and contains protons and neutrons.
5c. An atom contains the same number of protons and electrons to remain neutral.
5d. Drawing a Bohr model to represent the location and arrangement of the electrons.
6. An ion contains a different number of protons and electrons to give it an overall charge.
7. Given the isotope percent/relative abundances and masses of each, calculate the average atomic
mass of the element (round at the end to keep 2 decimal places).
2
Name ___________________________________
Date ________
Class Period _______ Clicker Number __________
Atomic Theorists
1. Who:
What: Law of Conservation of Mass
2. Who:
What: Law of Definite Proportions
When:
When:
Where:
3. Who:
What: Law of Multiple Proportions
When:
Where:
Where:
5. Who: Democritus
6. Who: John Dalton
What:
What:
When:
Where:
When:
Where:
4. Who: Aristotle
What:
When:
Where:
3
7. Who: J.J. Thomson
8. Who: Robert Millikan
9. Who: Ernest Rutherford
What:
What:
What:
When:
When:
When:
Where:
Where:
Where:
10. Who: Niels Bohr
11. Who: Erwin Schrodinger
12. Who: James Chadwick
What:
What:
What:
When:
When:
When:
Where:
Where:
Where:
4
Name __________________________________
Date ________
Class Period _______ Clicker Number ________
Problems with Dalton’s Theory
1. Matter is composed of indivisible particles
2. All atoms of a particular element are identical
3. Different elements have different atoms
4. Atoms combine in certain whole-number ratios
5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are
not created, destroyed, or changed into atoms of any other elements.
5
Name ___________________________________
Date ________
Class Period _______ Clicker Number ________
General Chemistry Worksheet
"Review Atomic Theorists"
Directions: Match the person or term on the left to the picture on the right.
Aristotle
Democritus
Erwin Schrodinger
Niels Bohr
JJ Thompson
Ernest Rutherford
John Dalton
Solid Sphere Model
Plum Pudding Model
Mass Centered/Stationary Planetary Model
Pulsating Planetary Model
Pulsating Orbital Model
(Quantum Mechanical Model)
6
Name ________________________________
Date ________
Class Period _______ Clicker Number ________
General Chemistry Worksheet
Atomic Theorists & Atomic Structure Review
1)______________________ Neutral particles.
2)_____________________ Scientist who discovered the electron. He also did
work with isotopes. The experiments he did used the
cathode ray tube.
3)______________________ Mass is neither created nor destroyed.
4)______________________ Value represented by # of protons + # of
neutrons.
5)______________________ He founded the “mathematical” model of the atom.
The electrons are located in the electron cloud.
6)______________________ The one element that has an isotope that does not
contain all of the subatomic particles.
7)______________________ Whenever two elements form more than one
compound, the different masses of one element
that combine with the same mass of the other
element are in the ratio of small whole numbers.
8)______________________ He developed the pulsating planetary model, where
electrons can change orbitals.
9)______________________ The Greek philosopher who believed in the atom.
10)______________________ A negatively charged particle that is located
outside of the nucleus.
11)______________________ This refers to electrons, neutrons, and protons.
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12)______________________ Atoms with the same number of protons and
electrons but a different number of neutrons.
13)______________________ This scientist used the oil drop experiment to
discover the mass of the electron.
14)______________________ A positively charged particle.
15)______________________ Chemical compounds have the same elements in
exactly the same ratios.
16)______________________ This English scientist is credited with discovering
the neutron.
17)______________________ It represents the number of protons in the
nucleus of an atom.
18)______________________ A formula for representing an element that uses
the element symbol, the mass number, and the
atomic number.
19)______________________ A scientist whose research lead to the discovery
that the nucleus is small, dense, and positively
charged. The electrons orbit around the nucleus.
The experiment he used is known as the gold foil
experiment.
20)______________________ Smallest particle of an element that retains the
chemical properties of that element.
21)______________________ A naming method that uses the name (or symbol)
of the element followed by a dash and its mass
number.
22)______________________ A Greek philosopher who did not believe in the
atom. His theory involved the “natural” elements
of fire, water, air and Earth.
23)______________________ His atomic theory model was represented by a
simple solid sphere.
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Isotopes
Atomic # = # protons (p+)
Mass Number= # protons (p+) + # neutrons (no) this is a whole number
Hyphen notation = Atomic symbol hyphen (not subtraction) followed by
mass number
= H-2 (Hydrogen w/ a mass of 2 amu)
Nuclear Symbol notation =
2
= 1
Mass #
Atomic #
(Atomic Symbol)
H (Hydrogen w/ a mass of 2 amu)
Use a periodic table and the notes from above to fill in the missing data in the table below:
Nuclear
Symbol notation
Number of
protons
Number of
neutrons
Number of
electrons in the
neutral atom
Name of element
Hyphen notation
65
29
Cu
86
Kr
78
6
117
46
36
Boron
Au-198
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Atom: -electrically neutral; the number of electrons and protons are equal.
Ion: -an electrically charged atom (positive or negative).
-the number of electrons and protons are not equal.
Cation: -a positively charged ion.
Ex) Na tends to lose one electron to form a 1+ ion, Na+
Ca tends to lose two electrons to form a 2+ ion, Ca2+
Anion: -a negatively charged ion.
Ex) Cl tends to gain one electron to form a 1- ion, ClO tends to gain two electrons to form a 2- ion, O2-
Fill in the missing information in the table below:
Hyphen notation
Al-27
Nuclear Symbol
notation
Number of
protons
Number of
neutrons
Number of
electrons
Atom, Cation, or
Anion
Overall charge
Atomic Number
He-4
238
92
Li-7
U
12
36
13
36
2
2
atom
0
2+
35
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Calculating Average Atomic Mass
Average atomic mass:
-The weighted average of all the naturally occurring isotopes for a
given element.
-This is NOT a mathematical average!
-This is the mass listed on the periodic table.
Percent (%) Abundances:
-the percentage of each isotope for a given element.
-the percent abundances of all the isotopes for a given element add up
to 100.
Relative Abundance:
-the percent abundance of an isotope divided by 100.
-the relative abundances of all the isotopes for a given element add up
to 1.
Relative Mass:
-the product of the relative abundance and the mass number for an
isotope of a given element.
**The Average Atomic Mass is the sum of all the relative masses for a
given element.** Do NOT divide your answer!!
Units on average atomic mass are amu (atomic mass units).
Round your final average atomic mass answer to 2 decimal places!
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Ex) Iron has 4 naturally occurring isotopes. Iron’s composition breaks down
as follows:
5.845% Fe-54
91.754% Fe-56
2.119% Fe-57
0.282% Fe-58
1) Convert all the % abundances to relative abundances.
0.05845
.91754
.02119
0.00282
2) Multiply the relative abundance times the mass number to obtain
relative mass.
0.05845 x 54= 3.1563
.91754 x 56= 51.38224
.02119 x 57= 1.20783
0.00282 x 58= 0.16356
3) Add the relative masses together.
3.1563 + 51.38224 + 1.20783 + 0.16356= 55.90993
4) Round the answer to 2 decimal places.
55.91
5) Put units on your answer; amu.
55.91 amu
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Determining Average Atomic Mass of an Element
1. Calculating Average Atomic Mass for Argon (round all answers to 2 decimal places):
Average Atomic Mass of Argon from the Periodic Table:
_______________________
Argon has 3 naturally occurring isotopes:
0.337% abundance of Ar-36
0.063% abundance of Ar-38
99.63% abundance of Ar-40
Calculated Average Atomic Mass of Argon:
_______________________
2. Determine the Average Atomic Mass of Sulfur:
Average Atomic Mass Number of Sulfur from the Periodic Table
_______________________
Sulfur has 4 naturally occurring isotopes:
95.00 % abundance of S-32
0.76 % abundance of S-33
4.22 % abundance of S-34
0.014 % abundance of S-36
Calculated Average Atomic Mass
_______________________
3. Determine the Average Atomic Mass of Zinc:
Average Atomic Mass of Zinc from the Periodic Table
_______________________
Zinc has 5 naturally occurring isotopes:
48.84 % abundance of Zn-64
27.62 % abundance of Zn-66
4.12 % abundance of Zn-67
18.71 % abundance of Zn-68
0.69 % abundance of Zn-70
Calculated Average Atomic Mass
_______________________
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Average Atomic Mass Calculation
Based on the number of atoms
1. Nitrogen has two isotopes with the following distributions:
14
7
N
15
7
N
= 99/100
= 1/100
The average atomic mass of Nitrogen is _______________________.
2. Antimony has two isotopes with the following distributions:
121
51
Sb
123
51
Sb
= 29/50
= 21/50
The average atomic mass of Antimony is _______________________.
3. Copper has two isotopes with the following distributions:
63
29
Cu
= 69/100
65
29
Cu
= 31/100
The average atomic mass of Copper is _______________________.
4. Silver has two isotopes with the following distributions:
107
47
Ag
= 29/50
109
47
Ag
= 21/50
The average atomic mass of Silver is _______________________.
14
Candium Lab
Purpose:
To analyze the isotopes of “candium” and to calculate its average atomic mass.
Materials:
sample of “candium”
balance
cups
Procedure: (record all your data & calculations in the data table!)
1) Obtain a sample of candium.
2) Separate the sample into its 3 isotopes (m&m’s, skittles, and Reese’s pieces).
3) Measure the total mass for each isotope; all the m&m’s, all the skittles, all the
Reese’s.
4) Count the number of “atoms” in each isotope.
5) Calculate the mass of one “atom” for each isotope using the total mass and number
of atoms.
6) Calculate the % abundance for each isotope using the number of atoms for each
isotope and the total number of atoms in the entire candium sample.
7) Calculate the average atomic mass for candium using the % abundance and mass of
one atom for each isotope. Show your work beneath the data table!
8) To check your answer, remember that the average atomic mass should be closest
to the most abundant isotope’s mass.
m&m’s
skittles
Reese’s pieces
Total mass
Number of “atoms”
Mass of one “atom”
(do NOT use the
balance!)
% abundance
Work for the average atomic mass calculation:
15
16
Calculating Isotope % Abundances
You will only be calculating percentages for elements that have 2 isotopes.
You will be given the 2 isotopes with their mass numbers.
Use the periodic table for the average atomic mass for the element.
1) Set one of the isotope’s relative abundance equal to X.
2) The relative abundances must add up to 1; therefore, the other
isotope’s relative abundance would be equal to 1-X.
3) Remember, the relative mass is equal to the relative abundance times
the mass number. Remember, the relative masses must add up to the
average atomic mass.
4) Setup an algebraic equation where the relative masses of each isotope
is set equal to the average atomic mass for the element. Use the full
average atomic mass from the periodic table, do NOT round.
(X • mass number) + [(1-x) • mass number] = average atomic mass
*When solving for X, keep 4 decimal places. This will allow the %
abundance to have 2 decimal places.
5) X is the relative abundance, so multiply this by 100 to make it %
abundance.
6) To obtain the other isotope’s % abundance, the 2 % abundances must
add up to 100.
17
Ex) Determine the % abundance for each isotope of antimony. Antimony
exists as Sb-121 and Sb-123.
(X • 121) + [(1-X) • 123] = 121.760
121X + 123 - 123X = 121.760
-2X = -1.24
X = 0.6200
The relative abundance of Sb-121 is 0.6200; therefore, the % abundance of
Sb-121 is 62.00% and the % abundance of Sb-123 is 38.00% (100-62=38).
Try this one:
Potassium exists as K-39 and K-41. Determine the % abundance for each
isotope of potassium.
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Name ___________________________________
Date ________
Class Period _______ Clicker Number ________
General Chemistry Worksheet
Calculating % Abundance of Isotopes of an Element
1. Determine the % Abundance of the two isotopes of Copper (Cu-63 and Cu-65)

Based on the average atomic mass of Copper, which isotope should have the higher percentage?
Circle:
Cu-63 and Cu-65
_______________________ % abundance of Cu-63
_______________________ % abundance of Cu-65
2. Determine the % Abundance of the two isotopes of Boron (B-10 and B-11):
_______________________ % abundance of B-10
_______________________ % abundance of B-11
3. Determine the % Abundance of the two isotopes of Chlorine (Cl-35 and Cl-37):
_______________________ % abundance of Cl-35
_______________________ % abundance of Cl-37
4. Determine the % Abundance of the two isotopes of Carbon (C-12 and C-13):
_______________________ % abundance of C-12
_______________________ % abundance of C-13
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Isotopes of Coinium Lab:
Name:_____________________
Penny data:
Part A:
Find the mass of 5 old pennies and then determine the average mass of just 1 old penny (pre-82
penny).
Mass of 1 old penny=___________________
Find the mass of 5 new pennies and then determine the average mass of one new penny (post-82
penny)
Mass of 1 new penny=__________________
Problems:
Sample #1 contains 3 old and 7 new pennies; calculate the average atomic mass of a
penny in this sample. Show your work below:
Average atomic mass of sample #1=_________
Is the average atomic mass for this sample closer to the mass of an old penny or a new
penny? Why does this occur?
Sample #2 contains 6 old and 4 new pennies; calculate the average atomic mass of a
penny in this sample. Show your work below.
Average atomic mass=____________
Is the average atomic mass for this sample closer to the mass of an old penny or a new
penny? Why does this occur?
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Part C: The Mystery Sample
1)
2)
3)
4)
5)
Get a canister of pennies. Don’t open it! Record its identifying number or letter:____________
Record the mass of the empty film canister, which is on the label of the canister.__________
Find the mass of the sealed film canister containing ten pennies. __________
Return the canister to your teacher.
Calculate the mass of just the ten pennies. ________________
6) Calculate the average mass of one penny in the canister. (your answer from # 5 divided by 10).
________________
7) Calculate the relative abundance for each penny in the sample (number of old pennies and number of
new pennies). Remember to set one of the relative abundances to X and the other to 1-X.
8)
a) Relative abundance of old pennies=______________
b) Relative abundance of new pennies=______________
9)
a) % abundance of old pennies=____________
b) % abundance of new pennies=___________
10) Your sample contained 10 total pennies. How many pennies were old?______
How many pennies were new?______
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Name____________________________
Fill in the Blanks!
Element
Name
Symbol Atomic Average Mass
# of
# of
# of
&
Number Atomic Number Protons Electrons Neutrons
Charge
Mass
Ru
101
34
34
45
76
114
54
77
Ce
82
H
1
1
2+
Fe
56
Cl
17
36
5
2
6
48
46
64
List the number of protons, electrons, and neutrons in each pair of isotopes:
6
3
Li
7
3
Li
42
20
Ca
44
20
Ca
# of Protons
# of Electrons
# of Neutrons
Write the Hyphen Notation and the Nuclear Symbol for the 1st 5 elements on the
chart at the top of this paper.
Element
1. Ruthenium
Hyphen Notation
Nuclear Symbol
2.
3.
4.
5.
23
Name ___________________________________
Date ________
Class Period _______ Clicker Number ________
General Chemistry Worksheet
"Review Atomic Theory"
1. Atomic # = 56
Mass Number = 137
a) # of Protons
b) # of Electrons
c) # of Neutrons
d) Hyphen notation
e) Nuclear Symbol
f) Name of element
2. Circle the two isotopes
A.
B.
131Xe
141Xe
54
53
C.
D.
133Xe
139Xe
54
52
3. Calculate the average atomic mass for Zinc:
Element
% Abundance
Zn-64
48.84
Zn-66
27.62
Zn-67
4.12
Zn-68
18.71
Zn-70
0.69
Compare your value with the Average atomic mass number for Zinc (Zn) on the Periodic Table.
4. Calculate the % abundance for Gallium
Element – Gallium
% Abundance
Ga-69
Ga-71
5. Calculate the average atomic mass for X:
Element
Number of Atoms/50
X-35
41 atoms
X-36
8 atoms
X-37
1 atom
6. Review Atomic Theorists from the Greeks to the Modern Era. Draw the picture of their atomic model
(Democritus does not have a picture) and know the name of that theory (we had a handout with all of
them on there – Example: JJ Thomson’s theory is represented by the plum pudding or jello salad model.)
Democritus
Aristotle
John Dalton
JJ Thomson
Ernest Rutherford
Niels Bohr
7. List three experiments of the atomic theory, who did them and what we learned from them.
8. List three laws that govern chemistry that were discussed in class, who wrote them and what did they say.
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