Chapter 5: Atomic Structure
and the Periodic Table
Section 5.1 - Atoms
 Objectives –
 Summarize Dalton’s atomic theory
 Describe the size of an atom
 Democritus - Greece 4th century BC
first proposed matter was
composed of atoms.
 John Dalton – England (1766-1844)
studied the ratios in which elements
combine in chemical reactions and
formulated Dalton’s atomic theory.
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Dalton’s Atomic Theory
1. All elements are composed of tiny individual
particles called atoms.
2. Atoms of the same element are identical. The
atoms of any one element are different from those
of any other atom.
3. Atoms of different elements can physically mix
together or can chemically combine with one
another in simple whole-number ratios to form
compounds.
4. Chemical reactions occur when atoms are
separated, joined or rearranged. Atoms of one
element, however, are never changed into atoms of
another element as a result of a chemical reaction.
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Size of an Atom
 The atom is the smallest particle of an
element that retains the properties of that
element.
 2.4 x 1022 atoms in a copper penny
 6 x 109 humans on earth
 100,000,000 (108) copper atoms lined up in a
row would only measure about 1 cm (10-2 m),
so that means the size of an atom is about
1x10-10m or about 1 Å (Angstrom) long.
 Semiconductor technology uses gate dimensions
of about 450 A long and less than 100 A thick,
so it is working at the atomic level.
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TEM to see at the atomic level
Where these two gold crystals
meet they are joined by a
complex arrangement of
atoms, forming a nanobridge
that accommodates their
different orientations. The
gold atoms are 2.3 angstroms
apart. TEAM (transmission
electron aberration-corrected
microscope) 0.5 angstrom
resolution makes it possible to
distinguish individual atoms
and, at the edges of the two
crystals, deduce their position
in three dimensions. (Credit: Image
courtesy of DOE/Lawrence Berkeley National Laboratory)
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More TEM images
 A TEM image of the
polio virus.
 The polio virus is
about 30 Angstroms
in size.
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Section 5.2 – Structure of the
Nuclear Atom
 Objectives –
 Distinguish among protons,
electrons and neutrons in terms of
relative mass and charge.
 Describe the structure of an atom,
including the location of the
protons, electrons and neutrons
with respect to the nucleus.
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Thompson Experiment
 Electrons are negatively charged subatomic
particles.
 J.J. Thompson discovered electrons in 1897.
 He passed electric current through gases at low
pressure. He sealed the gases in glass tubes
with metal electrodes at each end. When he
connected the electrodes to an electrical source,
the anode became positively charged and the
cathode negatively charged.
 A glowing beam, called a cathode ray, traveled
from the cathode to the anode. This cathode ray
was composed of electrons that were attracted to
the positive anode.
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Note that the
cathode ray
bends
upwards.
What does
that tell you
about the
nature of the
charged
particles that
make up the
cathode ray?
This diagram is like the Thompson apparatus, except it has
added positive and negative plates on the sides of the
cathode ray tube to cause the e- beam to bend. This can
be used to raster the beam across a screen, which was the
technology used for television.
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Milliken Oil Drop Experiment 1916
 Milliken discovered elementary charge, e, and mass
of an electron.
 Apparatus used force of gravity vs. electric field
force to measure mass of one (or more) charge.
 With field off, forces are D=drag force, mg =
weight. With field on, F=qE pulls oil drop upward
 He found every charge, q, was an integral multiple
of the charge on one electron, 1.6 x 10-19 coulomb.
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Protons
 Protons are the positively charged particle with
the same charge as an electron.

1.6x10-19 coulomb, except the charge is +
 Protons have 1840x the mass of an electron
 Mass electron: 9.11 x 10
 Mass proton: 1.67 x 10
-31
-27
kg or 9.11 x 10-28 g
kg or 1.67 x 10-24 g
 (your book gives masses in grams on pg. 111,
but kg is the SI standard unit)
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Neutrons
 Neutrons have the same mass as a proton,
1.67 x 10-24 g, but have no charge, they are
electrically neutral.
 The protons and neutrons together comprise
the nucleus of an atom.
 The symbols for these particles are
 Electron
e-
 Proton
p+
 Neutron
no
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Rutherford experiment
 In 1911, Rutherford used alpha particles 42𝐻𝑒
to bombard a very thin sheet of gold.
 An alpha particle is a Helium atom that has lost
its two electrons and has a double positive
charge, He2+ (2 protons, 2 neutrons, 0 electrons)
 Rutherford expected the alpha particles to pass
right through the gold because it was
hypothesized that the protons, neutrons and
electrons were evenly spread out within the
atom.
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Rutherford, continued
 Rutherford expected the alpha particles to pass
through with maybe only a slight deflection due
to the interaction of the charge between the He2+
and the protons (repulsion of like charges).
 A great majority did pass through the gold foil
without deflection, but a small percent bounced
off the gold foil at large angles as if they had
collided with a massive object, which turned out
to be the large gold nucleus.
 Rutherford said this “was about as credible as if
you had fired a 15-inch shell at a piece of tissue
paper and it came back and hit you,” but it
proved the existence of the nucleus.
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Rutherford Video
 Rutherford’s Experiment: Nuclear Atom
 http://www.youtube.com/watch?v=5pZj0u_XMbc
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Section 5.3: Distinguishing
Between Atoms
 Objectives –
 Explain how the atomic number identifies each
element
 Use the atomic number and mass number of an
element to find the number of protons, neutrons
and electrons
 Explain how isotopes differ and why the atomic
masses of elements are not whole numbers
 Calculate the average atomic mass of an element
from isotope data
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Atomic Number
 The atomic number of an element is the
number of protons in the nucleus of an atom of
that element.
 The atomic number identifies an element because
each element has a different number of protons.
 The number of electrons for a neutral atom
matches the number of protons. If the atom is
not electrically neutral (if it is an ion) then they
do not match.
 Example: Na+ is missing one electron, leaving it
positively charged. Cl- has an extra electron,
making it negatively charged.
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• How is the number of electrons for a neutral
atom related to the atomic number of that
element?
• Notice anything strange in the neutron column?
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Mass Number
 Mass number = # protons + # neutrons
 Example: Carbon has six protons and six neutrons,
so it’s mass number is 12
 Question: If Oxygen has an atomic number of 8,
and a mass number of 16, how many neutrons
does it have?
 Number of neutrons = mass number – atomic
number.

197
79𝐴𝑢
is the notation that shows gold has a mass
number of 197 and atomic number of 79. So how
many neutrons does it have?
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Practice Mass Number
 How many protons, electrons, neutrons?

108
47𝐴𝑔
 Silver
 Atomic number: 47 protons
 Mass number: 108
 108-47 = 61 so there are 61 neutrons
 Electrons: 47
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Practice Mass Number

207
82𝑃𝑏
 Lead
 Atomic number: 82 protons
 Mass number: 207
 207-82 = 125 neutrons
 Electrons: 82
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Sample Problem 5-2
 Find the number of protons, neutrons
and electrons in Be, Ne and Na.
Element
Atomic
Mass
Number
Number
Number
 Fill in
this table. Number
Find the number
of protons,
neutrons
Number
Protons
Neutrons
electrons
and electrons in Be, Ne and Na.
Be
Ne
Na
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Sample Problem 5-2
 Find the number of protons, neutrons and
electrons in Be, Ne and Na.
Element
Atomic
Number
Mass
Number
Number
Protons
Number
Neutrons
Number
electrons
Be
4
9
4
9-4 = 5
4
Ne
10
20
10
20-10 =
10
10
Na
11
23
11
23-11 =
12
11
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Isotopes
 What is an isotope?
 Two forms of the same element that differ in their
number of neutrons.
 The number of protons is the same though.
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Are these isotopes?
If the blue spheres are the neutrons and the
burgundy spheres are the protons, are these
three items isotopes of an element?
By the way,
how many
electrons
does this
atom have?
The percentage values shown are the percent
abundance of each isotope as it occurs in nature.
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Sample Problem 5-3
 Two isotopes of carbon are carbon-12 and
carbon-13. Write the symbol for each
isotope using superscripts and subscripts
to represent the mass number and atomic
number.


12
6𝐶
13
6𝐶
Carbon always has 6 protons.
Here it can have either 6 or 7
neutrons.
It has 6 electrons if neutral.
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Atomic Mass
 The actual mass of a proton or neutron is very
small: 1.67 x 10-24 g.
 However the mass of an electron is only
9.11 x 10-28g (1835 X less)
 So most of the mass lies in the protons and
neutrons.
 An atomic mass unit (amu) is defined using
Carbon-12. This isotope of carbon is defined as
having a mass of exactly 12 amu’s (6 protons and 6
neutrons).
 So that means 1 amu is equal to the mass of
one proton or neutron.
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Percent abundance example
Referring to Table 5.3 on page 119, Hydrogen has
three isotopes
Name of
Isotope
Symbol
Natural %
abundance
Mass in
amu
“Average”
atomic
mass
Hydrogen
1
1𝐻
99.985
1.0078
1.0079
Deuterium
2
1𝐻
0.015
2.0141
1.0079
Tritium
3
1𝐻
negligible
3.0160
1.0079
Average % weighted:
(.99985 x 1.0078) + (.00015 x 2.0141) = 1.0079
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Another weighted average mass sample
 Let’s say you have four atoms and you want to
find their weighted average mass:

35
17𝐶𝑙
35
17𝐶𝑙
35
17𝐶𝑙
 𝑎𝑣𝑒𝑟𝑎𝑔𝑒 =
𝑡𝑜𝑡𝑎𝑙 𝑝𝑟𝑜𝑡𝑜𝑛𝑠 𝑎𝑛𝑑 𝑛𝑒𝑢𝑡𝑟𝑜𝑛𝑠
4
 𝑎𝑣𝑒𝑟𝑎𝑔𝑒 =
17+17+17+17+18+18+18+20
4
37
17𝐶𝑙
= 35.5 amu
 Or to be more simple:
 𝑎𝑣𝑒𝑟𝑎𝑔𝑒 =
35+35+35+37
4
= 35.5 amu
 Check out the relative abundance of Chlorine-35
and -37 in the table on page 119.
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Sample Problem 5-4
 Which isotope of copper is more abundant:
copper-63 or copper 65? (The atomic mass of
copper is 63.546 amu)
 If it was 100% copper-63, the mass would be 63.0
amu. Likewise if it was copper-65, it would be
65.0 amu.
 If the abundances of copper-63 and copper-65
were equal, the average mass would be 64.0,
right?
 Since the value of 63.546 is closer to 63 than to
65, it must be the case that copper-63 is a larger
percent natural abundance.
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Review
 The atomic mass of an element is the
weighted average of the masses of its
isotopes.
 Need to know:
 The number of stable isotopes
 The mass of each isotope
 The natural % abundance for each
isotope
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Section 5-4: The Periodic Table
 Dmitri Mendeleev (1834-1907) was the first to
list the elements in a logical, systematic way.
 First he listed them in order by atomic mass,
such that the columns had chemicals with the
most similar properties.
 Henry Moseley (1887-1915) then determined the
atomic number of the atoms in the elements.
He also rearranged the table in terms of atomic
number instead of atomic mass.
 That is the way the periodic table is arranged
today.
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See the 7 periods labeled here?
The horizontal rows are
called periods. There are 7
periods.
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How many elements in period 6 ?? Did you count 32 elements?
Periodic Law
 When elements are arranged in order of
increasing atomic number, there is a
periodic repetition of their physical and
chemical properties.
 Elements that have similar chemical and
physical properties are positioned in the
same column of the periodic table.
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The vertical columns are called groups or
families. There are 18 groups,
numbered from left to right.
Group 1A
elements all
react
explosively
with water.
Group 8A
are the
noble
gases; inert
gases which
do not react
much.
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Group 1A to 8A elements are called
representative elements because they
exhibit a wide range of properties
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First column: alkali metals
Second column: alkaline earth metals
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Columns 3 through 12 are all metals,
called transition metals. The ones in
columns 13-15 that are highlighted in
blue are also metals.
Mercury is a liquid at
room temperature !!!
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These elements on the separate
rows down at the bottom are called
the inner transition metals or
rare-earth elements.
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These elements colored in yellow are
semi-metals or metalloids. Silicon and
Germanium are semiconductors. Aluminum
is not a metalloid.
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These elements colored in red are
called non-metals. They are nonlustrous and poor conductors of
electricity. Some are gases at room
temp, like N, O.
The elements in green
are called the halogens.
They are also part of the
non-metals category.
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The elements in turquoise are
called the noble gases,
sometimes called inert gases
because they are not very
reactive. They are also in the
non-metals category.
Summary
 In the periodic table, the elements are
organized into groups (vertical columns)
and periods (horizontal rows) in order of
increasing atomic number.
 Elements that have similar chemical
properties are in the same group / column.
 Elements are classified as metals,
nonmetals or metalloids.
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