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Chem 115 POGIL Worksheet – Week 2
Basic Atomic Structure
Why?
The atomic theory of matter is the great organizing principle of chemistry. Atoms are the
fundamental building blocks of all matter. The mass relationships between elements and
compounds in chemical reactions ultimately relate back to the characteristics of the atoms of
which they are composed. To understand how atoms combine to form compounds, you need to
understand their basic composition and structure.
Learning Objective
• Understand the basis for atomic theory
• Understand the structure of atoms, isotopes, and ions
• Understand the relationship between the masses of isotopes and the atomic weight of an
element
• Become familiar with the periodic table
Success Criteria
• Be able to write the standard nuclide notation for an isotope
• Be able to determine the numbers of fundamental particles in atoms and ions
• Be able to calculate atomic mass of a mixture of isotopes and percent isotopic composition
• Be able to categorize elements by position in the periodic table
Prerequisite
• Have read Chapter 2 in the text
• Understand percentage, scientific units, and unit analysis
Information
Dalton’s Atomic Theory was based in part on the work of the French scientist Joseph Louis
Proust in 1799. Proust discovered what is now called the Law of Definite Proportions (also
called the Law of Constant Composition): A compound is always composed of the same
elements in a fixed ratio by weight.
Chemistry 115 POGIL Worksheet Assignment
Example: When 200.59 g of mercury reacts completely with 32.066 g sulfur, 232.66 g of red
mercury sulfide is produced. What is the percent composition by weight of each element in red
mercury sulfide?
For every sample of red mercury sulfide the same percent composition by weight is found. (The
mineral cinnabarite is this compound.) It follows from this that a compound of mercury and
sulfur with any other percent composition by weight must be a different substance.
Proust’s discovery suggested to Dalton that the elements from which compounds are formed must
be composed of indivisible units, which combine in specific ways. From this idea, he proposed
an atomic theory, which in modern terminology consists of the following points:
(1) All matter is composed of atoms.
(2) All atoms of an element have the same mass (atomic weight).
(3) All atoms of different elements have different masses (i.e., different atomic weights).
(4) Atoms are indestructible and indivisible.
(5) Compounds are formed when atoms of two or more elements combine.
(6) In a compound the relative numbers and kinds of atoms are constant.
Points (2), (3), and (4) are now known to be incorrect, in light of the following later discovered
facts:
(2) Many elements are composed of a mixture of isotopes, atoms of the same element with
different masses.
(3) Some atoms of two different elements may have virtually the same mass; these are called
isobars.
(4) Atoms can be split (fission) or merged (fusion) in nuclear reactions. Some of the mass of
atoms is converted to energy in nuclear reactions.
Chemistry 115 POGIL Worksheet Assignment
Key Questions
1. What is the Law of Definite Proportions?
Information
Dalton knew that some pairs of elements could make more than one kind of compound and that
the percentages of each element were different in each case. On the basis of his atomic theory he
predicted and experimentally verified the Law of Multiple Proportions: If two elements can
form more than one compound, then the ratios of the weights of one element in the compounds to
a fixed weight of the other element are small whole numbers.
Key Questions
2. Suppose elements X and Y can form two compounds. One compound has as many X atoms
as Y atoms (formula XY), and the other compound has twice as many X atoms as Y atoms
(formula X2Y). What mass ratios would you compare between these compounds to
demonstrate the Law of Multiple Proportions? What whole number ratio would be expected
between these ratios?
Information
Today we know that atoms may be composed of three fundamental particles: protons, neutrons,
electrons.
Particle Unit Charge Mass
Proton (p) 1+ 1.6726 x 10-24 g
Neutron (n) 0 1.6749 x 10-24 g
Electron (e) 1– 9.1095 x 10-28 g
Unit electrical charge is ±1.6022 x 10-19 coulomb (C). The nucleus at the center of the atom
contains one or more positively charged protons. All atoms of a given element have the same
number of protons, which defines the element’s atomic number, given the symbol Z. In
addition, the nucleus may contain one or more neutrons, which have approximately the same
mass as protons but have no charge. Together, protons and neutrons are known as nucleons.
Any atom with a certain number of nucleons is called a nuclide. The number of nucleons
defines the nuclide’s mass number, given the symbol A:
A = number of protons + number of neutrons
Note that the mass number is an integer count of the number of nucleons, and not a statement of
an atom’s mass. Isotopes of an element have the same atomic number (Z) but have different
mass numbers (A), because they have different numbers of neutrons. Isobars are nuclides of
different elements (different Z) with the same mass number (A). Isobars have nearly (but not
exactly) the same mass.
The standard notation for a nuclide has the form , where X is the element’s symbol; Z is its
atomic number, equal to the number of protons; and A is the mass number, equal to the total number of nucleons (protons and neutrons). An electrically neutral atom has the same number of
protons as electrons, negatively charged particles that reside outside the nucleus. Atoms may
acquire electrical charge by either gaining or losing one or more electrons, thus becoming
monatomic ions. Positive ions are cations; negative ions are anions.
Atom → [Cation]n+ + ne–
Atom + ne– → [Anion]n–
Key Questions
3. What is the basis for defining the atomic number (Z) of an element?
4. What is the basis for defining the mass number (A) of a nuclide?
5. Are Z and A exact or inexact numbers?
6. How does an atom become a cation or anion?
7. Does Z or A change in forming an ion? Why or why not?
8. In some nuclear reactions an atom’s number of protons can change. Is it the same element
after such a change?
Exercises
9. On the periodic table attached, each block shows the atomic number of the element at the top,
above the element’s symbol. With the aid of the periodic table, give the standard nuclide
notation for the following isotopes used in medicine: phosphorous-32, chromium-51, cobalt-
60, iodine-131.
10. With the aid of the periodic table, fill in the blanks in the following table:
Charge 0 0 0 3+ 2–
Symbol 56Fe
Protons
35
34
Neutrons
45
38
Electrons
79 28
Mass no.
197
79
Information
Because the masses of atoms are so small, it is more convenient to give nuclide masses in atomic
mass units, abbreviated amu or u (the latter is the official SI abbreviation), rather grams. The
atomic mass unit is defined as follows:
One atomic mass unit is defined as 1/12 of the mass of a atom.
In atomic mass units the fundamental particles have the following masses
proton 1.007277 u
neutron 1.008665 u
electron 0.0005486 u
We cannot use these data to calculate the mass of a given atom, because the mass of a nuclide is
not simply the sum of the masses of its fundamental particles. When atoms are formed from
protons, neutrons, and electron, some mass is converted into energy, called the binding energy.
The mass equivalent of this energy can be calculated from the difference between the measured
mass of the nuclide and the sum of the masses of its subatomic particles, using E = mc2, where m
is the mass converted into energy, and c is the speed of light in a vacuum.
Because of the existence of isotopes, the masses of individual atoms in a sample of an element
may not all be the same. Indeed, with a few exceptions, most naturally occurring samples of an
element are mixtures of two or more isotopes in unequal portions. We generally deal with
samples containing large numbers of atoms with the usual mix of isotopes for the element, so it
is more useful to use an average atomic mass, weighted according to isotopic abundance. By
long standing tradition, this average has been called the atomic weight, even though the quantity
is actually mass. In general, tabulated values of atomic weights for elements do not represent the
mass of a single nuclide, unless the element occurs naturally as only one isotope.
Key Questions
11. For all elements except fluorine, the atomic weight listed on the periodic table does not
correspond to the mass of any nuclide. What does the atomic mass of most elements
represent?
12. The atomic weight listed for fluorine on the periodic table (18.998403 u) does correspond to
the mass of a particular nuclide. What does that imply about the isotopic composition of
naturally occurring fluorine?
Exercises
13. Boron consists of 19.78% 10B with atomic mass 10.0129 u and 80.22% 11B with atomic mass
11.00931 u. Calculate the atomic weight of naturally occurring boron.
14. By definition, the mass of a atom is exactly 12 u. What is the sum of the masses of the
particles comprising a atom? Why is the sum not 12 u?
Information
In 1869 Dmitri Mendeleev (Russian) and Julius Lothar Meyer (German) independently
discovered that when elements are arranged in order of their atomic weights, characteristic
properties of certain elements are repeated in other heavier elements at regular intervals in the
sequence. From this emerged the first statement of periodic law:
The properties of elements are a periodic function of their atomic weights.
This ordering, however, seemed to place some elements out of sequence. A better arrangement,
based on atomic number, became possible in 1913, when Henry G. J. Moseley found that the
atomic numbers of elements could be determined experimentally from their characteristic x-ray
frequencies. Today the periodic law is based on atomic numbers, rather than atomic weights:
The properties of elements are a periodic function of their atomic numbers (Z).
In most modern periodic tables, each block for an element shows its atomic number in the first
line, the element’s symbol underneath that, and the atomic weight of a naturally occurring sample
of the element below its symbol. The following definitions are used in conjunction with the
periodic table:
group – a column in the periodic table, listing elements that tend to show similar chemical
behavior. In North America, groups have been numbered 1 through 8 (or 0) with appended
letter designations A or B (e.g., 1A, 3B). The newer I.U.P.A.C. system uses numbers 1 through
15. Although hydrogen, H, is sometimes shown in group 1 (and even group 17), it really belongs
to no group, because its chemistry is unique.
period – a row in the periodic table. Periods are numbered 1 through 7.
main group elements (or representative elements) – members of the A group elements (old
North American system); i.e., groups 1A (1) and 2A (2), and 3A (13) through 8A (18) (newer
I.U.P.A.C. system designations in parentheses).
transition elements – members of the B group elements (old North American system),
corresponding to the groups 3 through 12 in the I.U.P.A.C. system. The first, second, and third
transition series span these groups in periods 4, 5, and 6, respectively. [Element number 89
(actinium, Ac) in period 7 begins a fourth transition series that would be continued with elements
104 through 112, but these are all unstable, synthetic elements.]
lanthanides – elements 58 through 71 in the first row at the bottom of the periodic table (a
continuation of period 6). Lanthanum (La) is actually the first element of the third transition
series, not a lanthanide.
actinides – elements 90 through 103 in the second row at the bottom of the periodic table (a
continuation of period 7). Actinium (Ac) is actually the first element of an incomplete fourth
transition series, not an actinide.
There are three categories of elements: metals, nonmetals, and metalloids, defined as follows:
metals – elements in groups 1A (1) and 2A (2), the transition elements, the lanthanides and
actinides, and the heavier elements in groups 3A (13) through 5A (15) that lie below the stair
step shown on some periodic tables. At room temperature metals are shiny solids (except
mercury and gallium above 29.78oC, which are liquids) that are malleable, ductile, and
conductive of heat and electricity. Metals characteristically are cations in their ionic compounds.
nonmetals – the elements in groups 4A (14) through 7A (17) that lie above the stair step on some
periodic tables. Individual nonmetals may be either solids, liquids, or gases at room temperature.
They are poor conductors of heat and electricity. Nonmetals characteristically are anions, when
existing as monatomic ions in ionic compounds. When combined with other nonmetals, they
typically form molecular compounds or complex ions.
metalloids – the elements B, Si, Ge, As, Sb, Te, Po, At, which lie along the stair step shown on
some periodic tables. All are solids with semi-metallic properties. They show poorer
conductivity relative to metals and may be semiconductors (e.g., Si and Ge).
Key Questions
18. What information about an element is provided in the box for that element in the periodic
table?
19. What determines the sequence of elements from the first to the last?
20. What is the difference between a group and a period?
21. Where are the metals, nonmetals, and metalloids located?
22. Are the majority of elements metals, nonmetals, or metalloids?
23. Does hydrogen belong to group 1? Why or why not?
Exercises
24. Write the name, symbol, atomic number, and average mass for each of the following,
and indicate whether the element is metal, nonmetal, or metalloid:
a. The group 2 element in period 3
b. The group 16 element in period 2
c. The group 15 element in period 4
25. Write the name and symbol of the element that has 48 electrons.
26. Name the elements with properties similar to chlorine, Cl.
27. Give the symbols and names of elements 57 and 72 in the period 6. Why are they adjacent to
each other in the periodic table?
Problem
28. Chlorine consists of 35Cl with a mass of 34.96885 u and 37Cl with a mass of 36.96590 u. The
atomic weight of chlorine is 35.453 u. What is the percent abundance of each isotope?
Periodic Table of the Elements
1A 8A
1 18
1
H 1.008
2A 3A 4A 5A 6A 7A
2 13 14 15 16 17
2
He 4.003
3
Li 6.941
4
Be 9.012
3B 4B 5B 6B 7B +))Q 8B ))), 1B 2B
3 4 5 6 7 8 9 10 11 12
5
B 10.81
6
C 12.01
7
N 14.01
8
O 16.00
9
F 19.00
10
Ne 20.18
11
Na 22.99
12
Mg 24.31
13
Al 26.98
14
Si 28.09
15
P 30.97
16
S 32.07
17
Cl 35.45
18
Ar 39.95
19
K 39.10
20
Ca 40.08
21
Sc 44.96
22
Ti 47.88
23
V 50.94
24
Cr 52.00
25
Mn 54.94
26
Fe 55.85
27
Co 58.93
28
Ni 58.69
29
Cu 63.55
30
Zn 65.39
31
Ga 69.72
32
Ge 72.61
33
As 74.92
34
Se 78.96
35
Br 79.90
36
Kr 83.80
37
Rb 85.47
38
Sr 87.62
39
Y 88.91
40
Zr 91.22
41
Nb 92.91
42
Mo 95.94
43
Tc [98]
44
Ru 101.1
45
Rh 102.9
46
Pd 106.4
47
Ag 107.9
48
Cd 112.4
49
In 114.8
50
Sn 118.7
51
Sb 121.8
52
Te 127.6
53
I 126.9
54
Xe 131.3
55
Cs 132.9
56
Ba 137.3
57
La 138.9
72
Hf 178.5
73
Ta 180.9
74
W 183.9
75
Re 186.2
76
Os 190.2
77
Ir 192.2
78
Pt 195.1
79
Au 197.0
80
Hg 200.6
81
Tl 204.4
82
Pb 207.2
83
Bi 209.0
84
Po [209]
85
At [210]
86
Rn [222]
87
Fr [223]
88
Ra [226]
89
Ac [227]
104
Rf [261]
105
Db [262]
106
Sg [266]
107
Bh [264]
108
Hs [269]
109
Mt [268]
110
Ds [281]
111
Rg [272]
112
Cn [285]
113
[284]
114
Fl [289]
115
[288]
116
Lv [293]
117
[294]
118
[294]
58 59 60 61 62 63 64 65 66 67 68 69 70 71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 140.1 140.9 144.2 [145] 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0
90 91 92 93 94 95 96 97 98 99 100 101 102 103
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 232.0 231.0 238.0 [237] [244] [243] [247] [247] [251] [252] [257] [258] [259] [262]
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