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Unit 1
The Atom, Atomic Theory & the Nucleus
Chapters 3 & 21
Atoms: The Building Blocks of Matter
CHAPTER 3
Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory
The Particle Theory of Matter
• In 400 B.C. Democritus, a Greek
philosopher, first proposed the
idea of a basic particle of
matter that could not be
divided any further.
• He called this particle the atom,
based on the Greek word atomos
meaning indivisible.
• This early theory was not backed up by
experimental evidence and was ignored by the
scientific community for nearly 2000 years.
Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory
Foundations of Atomic Theory
• By the late 1700s, experiments with chemical
reactions led to the discovery of 3 basic laws:
1. The Law of Conservation of Mass – Mass
is neither created nor destroyed during
ordinary chemical reactions or physical
changes.
Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory
Foundations of Atomic Theory
The Law of Conservation of Mass
Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory
Foundations of Atomic Theory
2. The Law of Definite Proportions – A
chemical compound contains the same
elements in exactly the same proportions
by mass regardless of the size of the
sample or source of the compound.
Visual Concept
Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory
Foundations of Atomic Theory
3. The Law of Multiple Proportions – If
different compounds are composed of the
same two elements, then the ratio of the
masses of the elements is always a ratio of
small whole numbers.
Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory
Dalton’s Atomic Theory
• In 1808, John Dalton proposed an explanation
for the three laws. His atomic theory states:
1. All matter is composed of atoms.
2. Atoms of the same element are identical; atoms of
different elements are different.
3. Atoms cannot be subdivided,
created, or destroyed.
4. Atoms of different elements
combine in simple wholenumber ratios to form
chemical compounds.
5. In chemical reactions, atoms
are combined, separated, or rearranged.
Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory
Corrections to Dalton’s Theory
• Dalton turned Democritus’s idea into a scientific
theory that could be tested by experiment.
• But not all aspects of Dalton’s theory have
proven to be correct. We now know that:
– Atoms are divisible into even smaller particles.
– A given element can have atoms with different
masses.
Chapter 3 – Section 2: The Structure of the Atom
Dalton’s Atomic Model
• An atom is the smallest particle of an element
that has all the properties of that element.
• Atoms are too small
to see…even
through the most
powerful
microscope!!
• Dalton thought
atoms were solid
balls of matter and
were indivisible.
Chapter 3 – Section 2: The Structure of the Atom
Discovery of the Electron
• In 1897, Joseph John (JJ) Thomson showed
that cathode rays are composed of identical
negatively charged particles, which were
named electrons.
• The electron was the
first subatomic particle
to be discovered.
Chapter 3 – Section 2: The Structure of the Atom
Thomson’s Cathode Ray Tube Experiment
Visual Concept
Chapter 3 – Section 2: The Structure of the Atom
Charge and Mass of the Electron
• In 1909, Robert Millikan measured
the charge on the electron during
his oil drop experiment.
• Using the charge-to-mass ratio,
scientists were able to figure out
the mass of the electron: about
1/2000 the mass of a hydrogen atom.
Chapter 3 – Section 2: The Structure of the Atom
Millikan’s Oil Drop Experiment
Visual Concept
Chapter 3 – Section 2: The Structure of the Atom
Thomson’s Plum Pudding Model
• After the work of Thomson and
Millikan, the accepted model of
the atom was called the
plum pudding model.
• The atom was viewed as a ball of positivelycharged material with tiny
negatively-charged
electrons spread evenly
throughout.
Chapter 3 – Section 2: The Structure of the Atom
Discovery of the Atomic Nucleus
• More detail of the atom’s structure was
provided in 1911 by Ernest Rutherford and his
associates Hans Geiger and Ernest Marsden.
•The results of their gold foil
experiment led to the discovery
of a very densely packed bundle
of matter with a positive
electric charge.
•Rutherford called this positive
bundle of matter the nucleus.
Chapter 3 – Section 2: The Structure of the Atom
The Gold Foil Experiment
Visual Concept
Chapter 3 – Section 2: The Structure of the Atom
Rutherford’s Atomic Model
• After Rutherford’s gold foil experiment, the
accepted model of the atom looked like this:
• A small, positively-charged nucleus with negative
electrons surrounding it at some distance away.
Most of the atom is empty space.
Chapter 3 – Section 2: The Structure of the Atom
Subatomic Particles
• The nucleus is made up of at least
one positively charged particle called
a proton and usually one or more
neutral particles called neutrons.
• Protons, neutrons, and electrons are
often referred to as subatomic particles.
Chapter 3 – Section 3: Counting Atoms
Atomic Number
• The atomic number (Z) of an element is the
number of protons of each atom of that
element.
• Atoms of the
same element
all have the
same number
of protons.
Chapter 3 – Section 3: Counting Atoms
Mass Number
• The mass number is the total number of
protons and neutrons in the nucleus of
an atom.
• Atoms of the
same element
can have
different
mass numbers.
Chapter 3 – Section 3: Counting Atoms
Isotopes
•Isotopes are atoms of the same element that
have different masses.
•Isotopes have the same number of protons and
electrons but different numbers of neutrons.
•Most of the
elements
consist of
mixtures of
isotopes.
Chapter 3 – Section 3: Counting Atoms
Designating Isotopes
•Hyphen notation: The mass number is written
with a hyphen after the name of the element.
uranium-235
•Nuclear symbol: The superscript indicates the
mass number and the subscript indicates the
atomic number.
Mass number
Atomic number
235
92
U
Chapter 3 – Section 3: Counting Atoms
Calculating Neutrons
• The number of neutrons is found by
subtracting the atomic number from the
mass number.
mass number − atomic number = number of neutrons
• Nuclide is a general term for a specific
isotope of an element.
Chapter 3 – Section 3: Counting Atoms
Calculating Subatomic Particles
Sample Problem
How many protons, electrons, and neutrons are
there in an atom of chlorine-37?
Solution:
Number of protons = atomic number (on periodic table)
17
Number of electrons = number of protons
17
Number of neutrons = mass number - protons
20
Significant Figures
• the non-place-holding digits in a reported
measurement are called significant figures.
• All non-zero numbers are significant, but
some zeros in a written number are only there
to help you locate the decimal point.
• How do you remember which zeros are
significant and which are not?
Significant Figures
• The Atlantic – Pacific Rule:
Pacific (Present):
If decimal point is
present, start
with the first nonzero number on
the left.
Atlantic (Absent):
If decimal point is
absent, start with
the first non-zero
number on the
right.
Significant Figures
• Examples: How many significant figures?
99,000
2 sig figs
99,000.
0.0099
5 sig figs
2 sig figs
0.0990
3 sig figs
Calculating with Significant Figures
• When you use your measurements in calculations,
your answer may only be as exact as your least exact
measurement.
• For addition and subtraction, round to the fewest
decimal places.
Example: (3 decimals) (1 decimal) (unrounded) (rounded)
50.259 + 17.4 = 67.659
67.7
• For multiplication and division, round to the fewest
significant figures.
Example: (3 sigfigs) (1 sigfig) (unrounded) (rounded)
0.135 x 20 = 2.7
3
Chapter 3 – Section 3: Counting Atoms
The Atomic Mass Unit
• The standard used by scientists to compare
units of atomic mass is the carbon-12 atom.
•One atomic mass unit, or 1 amu, is exactly
1/12 the mass of a carbon-12 atom.
•The atomic mass of any atom
is determined by comparing it
with the mass of the
carbon-12 atom.
Chapter 3 – Section 3: Counting Atoms
Average Atomic Mass
•Average atomic mass is the weighted average of
the atomic masses of the naturally occurring
isotopes of an element.
•The average atomic mass of an element depends
on both the mass
and the relative
abundance of
each of the
element’s
isotopes.
Percent
• To convert a decimal to a percent, multiply by 100
(move the decimal point 2 places to the right.)
.205
.090
20.5%
9.0%
• To convert a percent to a decimal, divide by 100
(move the decimal point 2 places to the left.)
65%
.03%
0.65
.0003
Multiplying by a Percent
• First, convert the percent to a decimal,
then multiply.
Example:
How many grams of carbon are in a 50.0 gram
sample of a substance that is 17.5% carbon?
17.5%
0.175
50.0 g x 0.175 = 8.75 g
Chapter 3 – Section 3: Counting Atoms
Calculating a Weighted Average
If you have the following grades, what would your
marking period average be?
Category
Percent of Grade
Grade
Tests
45 ÷ 100 = 0.45
x 72 = 32.4
Quizzes
10 ÷ 100 = 0.10
x 78 = 7.8
Labs
25 ÷ 100 = 0.25
x 84 = 21.0
HW/CW
10 ÷ 100 = 0.10
x 98 =
9.8
Project
10 ÷ 100 = 0.10
x 94 =
+ 9.4
80.4
• First, change percents to decimals.
• Next, multiply each grade by its decimal percent.
• Finally, add up all the products.
Chapter 3 – Section 3: Counting Atoms
Calculating Average Atomic Mass
Sample Problem 1
Copper consists of 69.15% copper-63, which
has an atomic mass of 62.929 601 amu, and
30.85% copper-65, which has an atomic mass
of 64.927 794 amu. What is the Average
Atomic Mass of Copper?
Solution
• Change percents to decimals.
• Multiply the atomic mass of
each isotope by its relative
abundance.
• Add up all of the products.
Relative
Abundance
Mass
Cu-63 0.6915 x
62.93 = 43.52
Cu-65 0.3085 x
64.93 = 20.03
+
63.55
Chapter 3 – Section 3: Counting Atoms
Calculating Average Atomic Mass
Sample Problem 2
A student believed that she had discovered a new
element and named it mythium. Analysis found it
contained two isotopes. The composition of the
isotopes was 19.9% of atomic mass 10.013 and 80.1%
of atomic mass 11.009. What is the average atomic
mass, and do you think mythium was a new element?
Solution:
Round off to:
• Average Atomic Mass:
10.8
(.199 x 10.013) + (.801 x 11.009) = 10.811
• Because the atomic mass is the same as the atomic
mass of boron, mythium was not a new element.
Scientific Notation
• A shorthand system of writing very large or
very small numbers.
• The power of 10 (called the exponent), is the
number of times the decimal point is moved.
• The number in front of the times sign(called
the coefficient) must be greater than 1 and
less than 10.
Scientific Notation
• Examples: What would each of these
numbers be in scientific notation?
3000
3 x 103
32,000
0.05
3.2 x 104
5 x 10-2
0.0058
5.8 x 10-3
Multiplying in Scientific Notation
• First, multiply the coefficients.
• Then, add the exponents together.
• Rewrite your answer in correct scientific notation
(coefficient less than 10), and remember to round
to the correct number of significant figures.
Example: 5.5 x 103 x 3.7 x 104 =
20.35x 107 =
2.0 x 108
Dividing in Scientific Notation
• First, divide the coefficients.
• Then, subtract the exponents.
• Rewrite your answer in correct scientific notation
(coefficient less than 10), and remember to round
to the correct number of significant figures.
Example: 1.8 x 10-2 ÷ 3.2 x 10-6 =
0.5625 x 104 =
5.6 x 103
Chapter 3 – Section 3: Counting Atoms
The Mole
• A mole (mol) is the
amount of a substance
that contains as many
particles as there are atoms
in exactly 12 g of carbon-12.
• It is a counting unit, similar to a dozen. In a
dozen, there are 12 things. In a mole, there
are 6.02 x 1023 things.
Visual Concept
Chapter 3 – Section 3: Counting Atoms
Avogadro’s Number
• Avogadro’s number:
6.022 1415 × 1023— the
number of particles in exactly
one mole of a pure substance.
• Named for
nineteenth-century
Italian scientist
Amedeo Avogadro.
Visual Concept
Chapter 3 – Section 3: Counting Atoms
Molar Mass
• Molar Mass is the mass of one
mole of a pure substance.
• Molar mass units are g/mol.
• The molar mass of an element
is the same number as its atomic
mass, only the units are different.
Try some - What is the molar mass of:
• Magnesium?
24.3 g/mol
• Carbon?
12.0 g/mol
Chapter 3 – Section 3: Counting Atoms
Gram/Mole Conversions
• Chemists use molar mass as a conversion factor
in chemical calculations.
• Remember, molar mass means grams per mole.
Example:
• What is the mass of 2.5 moles of Helium gas?
This is what
we’re given:
Use molar mass as
a conversion factor
2.5 mol He x
4.0 g He = 10. g He
1 mol He
Chapter 3 – Section 3: Counting Atoms
Gram/Mole Conversions
Sample Problem 1
What is the mass in grams of 3.50 mol of the
element copper, Cu?
Solution:
Given
Conversion factor
3.5 mol Cu x
63.5 g Cu = 220 g Cu
1 mol Cu
Chapter 3 – Section 3: Counting Atoms
Gram/Mole Conversions
Sample Problem 2
A chemist produced 11.9 g of aluminum, Al.
How many moles of aluminum were
produced?
Solution:
Given
11.9 g Al
Conversion factor
x
1 mol Al = 0.441 mol Al
27.0 g Al
Chapter 3 – Section 3: Counting Atoms
Conversions with Avogadro’s Number
• Avogadro’s number can be used as a conversion
factor between atoms and moles.
• Avogadro’s number units are atoms per mole.
Example:
•How many moles of silver, Ag, are in 3.01  1023
atoms of silver?
Given
3.01 x 1023 atoms Ag x
Conversion factor
1 mol Ag
= 0.500 mol Ag
6.02 x 1023 atoms Ag
Chapter 3 – Section 3: Counting Atoms
Conversions with Avogadro’s Number
Sample Problem 1
What is the mass in grams of 1.20  108 atoms
of copper, Cu?
nd
2
Solution:
st
Given
1.20 x 108 atoms Cu x
1
Conversion factor
Conversion
factor
1 mol Cu
x 63.5 g Cu
6.02 x 1023atoms Cu
1 mol Cu
= 1.27 x 10-14 g Cu
Nuclear Chemistry
CHAPTER 21
Chapter 3 – Section 2: The Structure of the Atom
Forces in the Atom
• Electrons and protons attract because of opposite
electrical charges, but protons and protons repel
since they have the same charge.
• The nucleus is held
together by a
mysterious force
called the strong
nuclear force which
only exists between
nucleons (protons
and neutrons) which
are very close
together.
Chapter 21 – Section 2: Radioactive Decay
Naturally Radioactive Elements
• All of the elements beyond atomic number 83
are unstable and thus radioactive.
Chapter 21 – Section 1: The Nucleus
Nuclear Reactions
• Large, unstable nuclei spontaneously break
apart to form smaller, more stable nuclei.
• A nuclear reaction is a reaction that affects
the nucleus of an atom.
Example: 92238 U  42 He  90234 Th
• A transmutation is a change in the identity of
a nucleus as a result of a change in the
number of its protons.
Chapter 21 – Section 1: The Nucleus
Nuclear Reactions
Sample Problem
Identify the products that balance the following
nuclear reactions:
212
4
Po

a. 84
2 He  ____
b. 1122 Na  ____ 1022 Ne
Solution:
208
a. Atomic Mass: 212 = 4 + _____
82
Atomic Number: 84 = 2 + _____
0 = 22
b. Atomic Mass:
22 + ____
-1 = 10
Atomic Number: 11 + ____
208
82
Pb
0
1
e
Chapter 21 – Section 2: Radioactive Decay
Radioactive Decay
• Radioactive decay is the
spontaneous disintegration
of a nucleus into a lighter
nucleus, accompanied by
nuclear radiation.
• Nuclear radiation is particles and/or
electromagnetic radiation emitted from the
nucleus during radioactive decay.
Chapter 21 – Section 2: Radioactive Decay
Types of Radioactive Decay
Alpha Emission
• An alpha particle (α) is two protons and two
neutrons bound together and is emitted
from the nucleus during some kinds of
radioactive decay.
4
• 2 He
• The atomic number
decreases by two and
the mass number
decreases by 4.
Chapter 21 – Section 2: Radioactive Decay
Types of Radioactive Decay (continued)
Beta Emission
• A beta particle (β) is an electron emitted from
the nucleus during some kinds of radioactive
decay (a neutron can be converted into a
proton and an electron.)
• 01 β
• The atomic number
increases by one and
the mass number
stays the same.
Chapter 21 – Section 2: Radioactive Decay
Types of Radioactive Decay (continued)
Positron Emission
• A positron (β+) is a particle that has the
same mass as an electron, but has a
positive charge (to decrease the number of
protons, a proton can be converted into a
neutron by emitting a positron.)
0
• 1 β
• The atomic number
decreases by one and
the mass number
stays the same.
Chapter 21 – Section 2: Radioactive Decay
Types of Radioactive Decay (continued)
Electron Capture
• In electron capture, an inner orbital electron
is captured by the nucleus of its own atom.
(An inner orbital electron combines with a
proton to form a neutron.)
0
• 1
e
• The atomic number
decreases by one
and the mass number
stays the same.
Chapter 21 – Section 2: Radioactive Decay
Types of Radioactive Decay (continued)
Gamma Emission
• Gamma rays () are high-energy electromagnetic
waves emitted from an unstable nucleus.
• Atomic number and
mass number both
stay the same because
gamma rays have no
charge and no mass.
• They are pure energy,
and very dangerous to living things.
Chapter 21 – Section 2: Radioactive Decay
Comparing Alpha, Beta and Gamma
• Alpha particles - big and slow,
can’t penetrate skin or paper.
• Beta particles - about 100 x the
penetrating power of alpha, can be stopped by
clothing, wood, or aluminum foil.
• Gamma rays - the greatest penetrating
ability, only stopped by a thick layer of lead
or concrete. They will go right through a
person and damage living cells.
Visual Concept
Chapter 21 – Section 2: Radioactive Decay
Half-Life
• Half-life is the time
required for half the
atoms of a radioactive
nuclide to decay.
• Each radioactive
nuclide has its own
half-life.
• More-stable nuclides
decay slowly and have longer half-lives.
Chapter 21 – Section 2: Radioactive Decay
Half-Life
Sample Problem 1
The half-life of radon-222 is 4 days. After what
time will ¼ of a given amount of radon remain?
Solution:
Determine the number of half-lives that it takes
to cut a sample to ¼ of the original amount.
Then multiply that
number by the half-life.
2 x 4 days = 8 days
Chapter 21 – Section 2: Radioactive Decay
Half-Life
Sample Problem 2
Phosphorus-32 has a half-life of 14.3 days. How many
milligrams of phosphorus-32 remain after 57.2 days if
you start with 4.0 mg of the isotope?
Solution:
Determine the number of half-lives in 57.2 days.
57.2 days ÷ 14.3 days = 4 half-lives
For each half-life, multiply the original amount by ½:
4.0 mg x ½ x ½ x ½ x ½ = 0.25 mg
Chapter 21 – Section 3: Nuclear Radiation
Whole-Body Radiation Exposure
• Rem – The unit used to measure the biological
effects of absorbed radiation in humans.
Chapter 21 – Section 3: Nuclear Radiation
Radiation Detection
• Film badges - use exposure of film to
measure the approximate exposure
of people working with radiation.
• Geiger-Müller counters - instruments that
detect radiation by counting electric pulses
carried by gas ionized by radiation.
• Scintillation counters - instruments that
convert scintillating light to an electric signal
for detecting radiation.
Chapter 21 – Section 3: Nuclear Radiation
Uses of Radiation
• Radioactive dating – scientists can determine
the approximate age of an object based on the
amount of certain radioactive nuclides present.
Chapter 21 – Section 3: Nuclear Radiation
Uses of Radiation (continued)
•Radioactive tracers are radioactive atoms that
are incorporated into substances so that
movement of the substances can be followed
by radiation detectors.
–Radioactive tracers can
be used by doctors to
diagnose diseases.
–Radioactive tracers are
also used in agriculture
to determine the effectiveness of fertilizers.
Chapter 21 – Section 3: Nuclear Radiation
Uses of Radiation (continued)
•Irradiated Food – nuclear radiation is used to
prolong the shelf life of food.
Chapter 21 – Section 3: Nuclear Radiation
Uses of Radiation (continued)
• Nuclear Power Plants – use energy as heat from
nuclear reactors to produce electrical energy. They
have five main components:
1. Shielding – radiation-absorbing material used to
decrease exposure to radiation from nuclear reactors.
2. Fuel – usually Uranium-235.
3. Coolant – usually water, it absorbs excess heat energy.
4. Control rods – neutron-absorbing rods that limit the
number of free neutrons
5. Moderator – used to slow down the fast neutrons
produced by fission.
Chapter 21 – Section 3: Nuclear Radiation
Nuclear Power Plant
Chapter 21 – Section 3: Nuclear Radiation
Fission vs. Fusion
• Nuclear Fission – very heavy nuclei split into
smaller, more stable nuclei.
–Can occur spontaneously or when
nuclei are bombarded by particles.
–Controlled fission chain
reactions are used
in nuclear
power plants.
Uncontrolled
fission chain reactions
are used in nuclear bombs.
Chapter 21 – Section 3: Nuclear Radiation
Fission vs. Fusion (continued)
• Nuclear Fusion – low-mass nuclei combine to
form a heavier, more
stable nucleus.
–Fusion releases even more
energy per gram of fuel
than fission.
–Scientists are not yet able to control fusion
reactions, so we can’t use them
in power plants.
–Fusion is the primary process
that fuels our sun and the stars.
Chapter 21 – Section 3: Nuclear Radiation
Nuclear Waste
• Radioactive waste produced in nuclear reactors can
take hundreds of thousands of years to decay.
• Disposal of nuclear waste
is done with the intention
of never retrieving it.
• There are 77 disposal sites
around the country. A new one (Yucca Mountain)
is being developed for the permanent disposal of
much of our nuclear waste beginning in 2017.
Visual Concept