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Chapter 2 Lecture
General, Organic, and Biological
Chemistry: An Integrated Approach
Laura Frost, Todd Deal and Karen Timberlake
by Richard Triplett
Chapter 2
Atoms and Radioactivity
Chapter Outline
2.1 Atoms and Their Components
2.2 Atomic Number and Mass Number
2.3 Isotopes and Atomic Mass
2.4 Counting Atoms: The Mole
2.5 Electron Arrangements
2.6 Radioactivity and Radioisotopes
2.7 Nuclear Equations and Radioactive Decay
2.8 Radiation Units and Half-Lives
2.9 Medical Applications for Radioisotopes
© 2011 Pearson Education, Inc.
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2.1 Atoms and Their Components
Subatomic Particles
• An atom is the smallest part of an element.
• An atom is made up of three subatomic
particles:
1. Electron
2. Proton
3. Neutron
•
•
•
An electron is located outside the nucleus of
the atom.
A proton is located in the nucleus of the atom.
A neutron is located in the nucleus of the
atom.
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2.1 Atoms and Their Components, Continued
Overall charge on an atom is zero because the
number of protons is equal to the number of
electrons.
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2.1 Atoms and Their Components, Continued
Structure of an Atom
• The nucleus is very compact. It contains all the
protons and neutrons.
• Analogy: If an atom were the size of an enclosed
football arena, the nucleus would be about the
size of a pea on the 50-yard line with the
electrons occupying the space within the
enclosed arena.
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2.1 Atoms and Their Components, Continued
• An electron cloud is a space occupied by the
electrons. Electrons are constantly moving about
this space.
• Electrons contribute very little to the mass of an
atom.
• The majority of the mass of an atom is located in
the nucleus and is a result of the relative mass
of protons and neutrons.
© 2011 Pearson Education, Inc.
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2.1 Atoms and Their Components, Continued
• The atomic mass unit, or amu, is a unit used by
chemists to identify the mass of an atom.
• A proton and a neutron have the same mass, so
each is defined as weighing 1 amu.
• Amu is defined as one-twelfth of a carbon atom
containing six protons and six neutrons.
Therefore, a carbon atom would be 12 amu.
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2.2 Atomic Number and Mass Number
Atomic Number
• An atomic number indicates the number of
protons in an atom.
• The atomic number is located at the top of the
element block above the elemental symbol.
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2.2 Atomic Number and Mass Number,
Continued
• All atoms of an element contain the same
number of protons.
• All atoms are neutral and therefore, the number
of protons equals the number of electrons.
Mass Number
• To determine the number of neutrons in an
atom, we must look at the mass number.
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2.2 Atomic Number and Mass Number,
Continued
• The mass number, located below the elemental
symbol in the element block of the periodic table, is
the number of protons plus the number of neutrons.
Mass number = number of protons + number of neutrons
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2.2 Atomic Number and Mass Number,
Continued
© 2011 Pearson Education, Inc.
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2.2 Atomic Number and Mass Number,
Continued
• If the mass number and atomic number for a
given atom is known, one can determine the
number of subatomic particles present.
• Symbolic notation is a method used to
represent an atom’s atomic symbol, mass
number, and atomic number.
© 2011 Pearson Education, Inc.
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2.3 Isotopes and Atomic Mass
• Most atoms of carbon found in nature have a
mass number of 12, a few have a mass number
of 13, and even fewer have a mass number of
14.
• All of these must have the same number of
protons to be considered carbon atoms.
• The different mass numbers of these carbon
atoms are due to a difference in the number of
protons.
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2.3 Isotopes and Atomic Mass, Continued
• Isotopes are atoms of the same element that
have the same number of protons, but different
number of neutrons. They have different mass
numbers.
• Isotopes can be written in symbolic notation as:
• Isotopes can also be written by stating the
element name followed by the mass number.
For example, carbon-12, carbon-13, and
carbon-14.
© 2011 Pearson Education, Inc.
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2.3 Isotopes and Atomic Mass, Continued
Atomic Mass
• Identification of the common isotope found in an
atom is determined from the periodic table.
• Carbon has three main isotopes:
1. Carbon-12
2. Carbon-13
3. Carbon-14
•
•
Review of the periodic table shows the average
atomic mass of carbon to be 12.01 amu. This
indicates that the major isotope for carbon is
carbon-12.
Atomic mass is the average atomic mass
weighted for all isotopes of a particular element.
© 2011 Pearson Education, Inc.
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2.3 Isotopes and Atomic Mass, Continued
© 2011 Pearson Education, Inc.
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2.4 Counting Atoms: The Mole
• How are atoms counted?
• Chemists use a unit called a mole, which relates
the mass of an element in grams to the number
of atoms it contains.
• Molar mass represents the number of grams in
one mole of an element and is numerically equal
to the atomic mass of the element.
• For example, 1 mole of carbon atoms has a
molar mass of 12.01 grams. This can be
expressed as 12.01 g/mole.
© 2011 Pearson Education, Inc.
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2.4 Counting Atoms: The Mole, Continued
• The number of atoms in 1 mole is defined as
602,000,000,000,000,000,000,000.
• Chemists use scientific notation to handle such
huge numbers.
• The number of atoms in 1 mole can be
represented in scientific notation as 6.02 x 1023
atoms.
© 2011 Pearson Education, Inc.
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2.4 Counting Atoms: The Mole, Continued
Math Matters: Scientific Notation
• The general form for scientific notation is C x 10n
where C is called the coefficient and is a number
between 1 and 10 and n is the exponent
indicating the number of places applying to the
coefficient.
• A positive coefficient indicates a number greater
than 1, and a negative number indicates a
number less than 1.
© 2011 Pearson Education, Inc.
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2.4 Counting Atoms: The Mole, Continued
• Only significant figures are shown in scientific
notation.
• For example, the scientific notation for 3060000
expressed with three significant figures would be
3.06 x 106, and 0.000306 would be expressed
as 3.06 x 10-4.
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2.4 Counting Atoms: The Mole, Continued
• Table 2.3 shows the relationship between
numbers and scientific notation.
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2.4 Counting Atoms: The Mole, Continued
Avogadro’s Number
• Avogadro’s number (N) is the number of atoms
present in 1 mole of atoms.
6.02 x 1023 atoms = 1 mole of atoms
OR
N = 6.02 x 1023 particles/mole
• One mole of anything will have 6.02 x 1023
particles. So 1 mole of eggs will have 6.02 x 1023
eggs just as 1 mole of carbon will have
6.02 x 1023 atoms.
© 2011 Pearson Education, Inc.
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2.5 Electron Arrangements
• Electrons of an atom move about the nucleus in
an area known as an electron cloud.
• Electrons possess energy because they are
constantly moving.
• Electrons are found in distinct energy levels
based on the amount of energy they possess.
• Electrons in the same energy level possess
similar energies.
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2.5 Electron Arrangements, Continued
• Electrons occupy the
lowest energy level,
first which is closet
to the nucleus.
• The maximum
number of electrons
in any energy level
can be calculated by
the formula 2n2,
where n is the
number of energy
level.
© 2011 Pearson Education, Inc.
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2.5 Electron Arrangements, Continued
• The table below shows the electron arrangement of the first 20 elements.
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2.5 Electron Arrangements, Continued
• Elements with the same number of electrons in
their highest energy level are in the same group.
• The highest energy level containing electrons is
known as the valence shell and the electrons
are known as the valence electrons.
• Valence electrons are furthest from the nucleus
and are the electrons responsible for the
chemical reactivity of elements.
© 2011 Pearson Education, Inc.
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2.5 Electron Arrangements, Continued
• The group number for an element represents
the number of valence electrons each element in
the group contains. For example, Group 1A
elements contain one valence electron, Group
2A elements contain two valence electrons, and
so on.
• The period number is the outermost energy
level containing valence electrons. For example,
H and He in Period 1 have their electrons in
energy level 1, Period 2 elements have valence
electrons in energy level 2, and so on.
© 2011 Pearson Education, Inc.
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2.6 Radioactivity and Radioisotopes
• When energy is given off spontaneously from
the nucleus of an atom, it is called nuclear
radiation.
• Radiation comes in many different types and
forms. Cosmic radiation is a natural radiation. It
is a major source of radiation to which humans
are exposed.
• Microwave radiation is an example of humanmade radiation.
© 2011 Pearson Education, Inc.
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2.6 Radioactivity and Radioisotopes,
Continued
• Spontaneously emitted radiation from the
nucleus of an element is called radioactivity.
• Some isotopes of elements are radioactive and
are called radioisotopes.
© 2011 Pearson Education, Inc.
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2.6 Radioactivity and Radioisotopes,
Continued
Radioisotopes are used for imaging and for
diagnosing diseases.
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2.6 Radioactivity and Radioisotopes,
Continued
• Not all naturally occurring radioisotopes are
radioactive because they have a stable nucleus.
• Isotopes that are not stable become stable by
spontaneously emitting radiation from their
nuclei. This process is called radioactive
decay.
© 2011 Pearson Education, Inc.
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2.6 Radioactivity and Radioisotopes,
Continued
Types of Radiation
•
There are three main types of radiation
particles:
1. An alpha particle (α) is a positively charged
particle with a 2+ charge. It is also known as
a helium particle.
2. A beta particle (β) is a negatively charged
particle with a 1- charge. It is the same as
an electron.
3. A gamma particle(γ) is a neutral particle.
© 2011 Pearson Education, Inc.
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2.6 Radioactivity and Radioisotopes,
Continued
Types of Radiation
Symbols for each type of radiation are shown in the
following table. Two additional types of radiation
shown are the positron and the neutron.
© 2011 Pearson Education, Inc.
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2.6 Radioactivity and Radioisotopes,
Continued
Biological Effect of Radiation
• Radiation emissions are dangerous to living
organisms because when emitted they interact
with any atoms they contact.
• Alpha particles, beta particles, and gamma rays
are know as ionizing radiation. When they
come in contact with atoms, they cause the
atoms to lose electrons, which leaves species
that are reactive and unstable.
© 2011 Pearson Education, Inc.
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2.6 Radioactivity and Radioisotopes,
Continued
Not all ionizing radiation has the same amount of
energy, so some are more dangerous than others.
This is due to the penetrating power of higher
energy radiation like gamma rays.
© 2011 Pearson Education, Inc.
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2.7 Nuclear Equations and Radioactive Decay
• The general form of a nuclear decay equation is:
Radioactive nucleus undergoing decay → new
nucleus formed + radiation emitted
• Uranium-238, a radioactive isotope, emits alpha
particles when it undergoes radioactive decay.
The nuclear equation for this type of decay is:
© 2011 Pearson Education, Inc.
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2.7 Nuclear Equations and Radioactive Decay,
Continued
• In a nuclear decay equation, the mass number
of the reactant (238) must equal the total mass
number of the products (234 + 4).
• In a nuclear decay equation, the atomic number
of the reactant (92) must equal the sum of the
atomic numbers of the products (90 + 2).
© 2011 Pearson Education, Inc.
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2.7 Nuclear Equations and Radioactive Decay,
Continued
Alpha Decay
• The alpha particle is a product of alpha decay.
• To solve an alpha particle decay equation,
remember that the mass and atomic numbers
must be equal on both sides of the equation.
Then you can look at the periodic table for the
element that has the appropriate atomic number.
© 2011 Pearson Education, Inc.
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2.7 Nuclear Equations and Radioactive Decay,
Continued
Beta Decay
• The high energy electron is emitted from an isotope
during beta decay. Remember this particle has no
mass and a negative charge.
• The product isotope will have the same mass as
the reactant, but its atomic number will increase by
1 since the beta particle is negatively charged.
© 2011 Pearson Education, Inc.
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2.7 Nuclear Equations and Radioactive Decay,
Continued
Gamma Decay
Gamma decay is only energy and will not result in
a change of the mass number or atomic number of
the product. Remember gamma rays have no
mass and no charge.
© 2011 Pearson Education, Inc.
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2.7 Nuclear Equations and Radioactive Decay,
Continued
Producing Radioactive Isotopes
• Some isotopes are produced in the laboratory by
bombarding stable isotopes with fast moving
alpha particle, protons, or neutrons.
• Some of these isotopes are important in
medicine. Technetium-99m is an example and is
produced as follows:
© 2011 Pearson Education, Inc.
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2.8 Radiation Units and Half-Lives
Radioactivity Units
• The activity of a radioactive sample is measured
in disintegrations/second.
• The unit of measuring disintegrations is called the
curie (Ci).
• The activity of a radioactive isotope defines how
quickly it emits radiation. A curie is a unit of
activity equal to 3.7 x 1010 disintegrations/second.
© 2011 Pearson Education, Inc.
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2.8 Radiation Units and Half-Lives, Continued
Half-Life
• Each radioactive isotope emits its radiation at a
different rate.
• Emission of radiation can be measure by its
half-life.
• The half-life of an isotope is the time it takes for
50% of the atoms in a radioactive sample to
decay.
© 2011 Pearson Education, Inc.
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2.8 Radiation Units and Half-Lives, Continued
Naturally occurring radioisotopes tend to have long
half-lives. Medically important radioisotopes have
shorter half-lives to allow radioactivity to be quickly
eliminated from the body.
© 2011 Pearson Education, Inc.
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2.8 Radiation Units and Half-Lives, Continued
• The amount of radioactivity left after a given
amount of time has passed can be determined if
we know the half-life of the sample.
• For example, radioactive iodine-131 has a halflife of 8 days. If a dose of 200 µC is given to a
patient, how much activity is left after 32 days?
• First, we have to know how many half-lives are
in 32 days.
© 2011 Pearson Education, Inc.
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2.8 Radiation Units and Half-Lives, Continued
• Now that we have determined 4 half-lives have
passed, we can determine how much iodine-131
radioactivity is left after 32 days.
© 2011 Pearson Education, Inc.
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2.9 Medical Applications for Radioisotopes
• Radioisotopes with short half-lives are used in
nuclear medicine in order to expose patients
with the smallest doses of radiation in the
shortest period of time.
• Isotopes are used to provide images of specific
body tissues.
• Iodine, used only by the thyroid gland, can be
used to obtain an image of the thyroid gland
because radioactive iodine will accumulate in
this gland.
© 2011 Pearson Education, Inc.
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2.9 Medical Applications for Radioisotopes,
Continued
• A radioisotope used to image specific body
tissue is called a tracer.
• Iodine-123 is used to diagnose thyroid function.
© 2011 Pearson Education, Inc.
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2.9 Medical Applications for Radioisotopes,
Continued
• Radioisotopes are also used to destroy diseased
and cancerous tissues.
• Emissions from radioactive sources can be used
without injecting the patient with radioactive
material.
• Cobalt-60 can be aimed directly at a cancerous
tumor to destroy tissue that affects the diseased
area.
© 2011 Pearson Education, Inc.
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Chapter Summary
2.1 Atoms and Their Components
•
Atoms consists of three subatomic particles:
1. Electrons (negatively charged)
2. Protons (positively charged)
3. Neutrons (neutral charge)
•
Mass of an atom is expressed as the atomic
mass unit (amu).
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.2 Atomic Number and Mass Number
• Atomic number of an atom is equal to the
number of protons in the atom.
• Mass number of an atom is equal to the sum of
the protons and neutrons in the atom.
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.3 Isotopes and Atomic Mass
• Isotopes of an atom contain the same number of
protons, but different number of neutrons so they
have a different mass number.
• The mass of an element on the periodic table is
the average mass of all its isotope.
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.4 Counting Atoms: The Mole
• Chemists use a unit called the mole to relate the
number of atoms to an element’s mass.
• Avogadro’s number indicates that 1 mole of an
element is equivalent to 6.02 x 1023 atoms.
• The atomic mass unit (amu) is equivalent to the
number of grams in one mole of atoms, which is
known as the molar mass.
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.5 Electron Arrangements
• Electrons exist in distinct energy levels.
• Energy levels are referred to as levels n = 1, 2,
3, and so on.
• Maximum number of electrons in an energy level
is equal to 2n2.
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.6 Radioactivity and Radioisotopes
•
Energy given off spontaneously from the nucleus
of an atom is called nuclear radiation. This
process is referred to as radioactive decay.
•
Three forms of radioactive decay are:
1. Alpha radiation
2. Beta radiation
3. Gamma radiation
•
Large doses of ionizing radiation cause tissue
damage.
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.7 Nuclear Equations and Radioactive Decay
• Radioactive decay can be represented in the
form of a nuclear decay equation.
• The number of protons and the mass number in
the reactant is equal to the number of protons
and mass number found in the products.
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.8 Radiation Units and Half-Lives
• Radioactive decay is measured in
disintegrations/second.’
• The curie (Ci) is the standard unit for measuring
radioactive decay.
• A curie is defined as 3.7 x 1010
disintegrations/sec.
• Half-life is the amount of time it takes for 50% of
the radiation of a sample to decay.
© 2011 Pearson Education, Inc.
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Chapter Summary, Continued
2.9 Medical Applications for Radioisotopes
• Radioisotopes are used in imaging body tissues.
• Radioisotopes are used in small doses to
diagnose disease states.
• Radioisotopes are used in high doses to destroy
tumors and cancerous tissue.
© 2011 Pearson Education, Inc.
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