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How is the structure of the atom related to its behavior?
Chemistry: Unit 3
New Skills:
 Calculate weighted atomic masses
 Predict products of a nuclear reaction
 Calculate Half-life
Academic Language:
Dalton’s Atomic Theory
Atom
Cathode ray
Electron
Neutron
Nucleus
Proton
Atomic mass
Atomic number
Isotope
Mass number
Alpha particle
Alpha radiation
Beta particle
Beta radiation
Gamma ray
Nuclear equation
Nuclear Reaction
Radioactivity
Radiation
Radioactive decay
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Unit 3 Homework:
CALM: http://calm.indiana.edu/
3:1 Early Ideas about Matter:
 10 questions
3:2 Defining the Atom:
 10 questions
3:3 How Atoms Differ:
 10 questions
3:4 Unstable Nuclei and Radioactive Decay:
 10 questions
3:5 Accumulating Content and Skills:
 10 questions
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Chemistry Unit 3: Learning Goals and Objectives
3:1 Early Ideas about Matter - The ancient Greeks tried to explain matter, but the scientific
study of the atom began with John Dalton in the early 1800’s.


Compare and contrast the atomic models of Democritus, Aristotle and Dalton.
Understand how Dalton’s theory explains the conservation of mass.
3:2 Defining the Atom – An atom is made of a nucleus containing protons and neutrons;
electrons move around the nucleus.



Define atom
Distinguish between the subatomic particles in terms of relative charge and mass
Describe the structure of the atom, including the locations of the subatomic
particles.
3:3 How Atoms Differ – The number of protons and the mass number define the type of
atom.




Explain the role of atomic number in determining the identity of an atom
Define an isotope
Explain why atomic masses are not whole numbers
Calculate the number of electrons, protons, and neutrons in an atom given its mass
number and atomic number.
3:4 Unstable Nuclei and Radioactive Decay – Unstable atoms emit radiation to gain
stability.



Explain the relationship between unstable nuclei and radioactive decay.
Characterize alpha, beta, and gamma radiation in terms of mass and charge.
Solve problems involving radioactive decay rates.
3:5 – Accumulating Content and Skills:– Chemistry content is continuous and builds on
prior knowledge and skills. This section will combine this unit with previous units.

Apply knowledge and skills from previous units to content learned in this unit.
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3:1 Early ideas about Matter- The ancient Greeks tried to explain matter, but the scientific
study of the atom began with John Dalton in the early 1800’s.
Objective: Compare and contrast the atomic models of Democritus, Aristotle and Dalton.
Early Greek Philosophers - Many concluded that matter was composed of things such as
earth, water, air and fire

Democritus- (Greek 460-370BC)- the first to propose the idea that matter was not infinitely
divisible. He believed matter was made up of tiny individual particles called atomos. He
felt that these atoms could not be created nor destroyed or further divided.
1.
2.
3.
4.
Aristotle- (Greek 384-322BC) - One of the most influential philosophers of his time. He
disagreed with Democritus , one reason being he didn’t believe that empty space existed.
Aristotle’s views went unchallenged for 2000 years until science developed methods to test
the validity of his ideas.
1.
2.
**It was difficult for Democritus to defend his ideas because controlled experiments on
his theory were not possible.**
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Objective: Understand how Dalton’s theory explains the conservation of mass.
John Dalton – (English school teacher -1766-1844) – Responsible for the beginning of the
development of modern atomic theory. Dalton revised Democritus’s theories by
performing and studying many chemical reactions. Through careful observations and
measurements he was able to determine mass ratios of the elements involved in those
reactions.

John Dalton revived the idea of the atom in the early 1800s based on numerous
chemical reactions.

Dalton’s Atomic Theory
John Dalton’s theory, proposed in 1803 is known as Dalton’s atomic theory.
1. Matter is composed of extremely small particles called atoms.
2. Atoms are indivisible and indestructible
3. Atoms of a given element are identical in size, mass, and chemical properties.
4. Atoms of a specific element are different from those of another element.
5. Different atoms combine in simple whole-number ratios to form compounds.
6. In a chemical ration, atoms are separated, combined or rearranged.
***Even though some of the theory was incorrect, it provided the basis we have today
for atomic theory. As with science, his theory has since been revised and additional
information learned.***
Which parts of Dalton’s Atomic Theory were incorrect?
1.
Why?
2.
Why?
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3:2 Defining the Atoms - An atom is made of a nucleus containing protons and neutrons;
electrons move around the nucleus.
Objective: Define atom
Atoms – the smallest particle of an element that retains the properties of that element.

o
Objective: Distinguish between the subatomic particles in terms of relative charge and
mass
Objective: Describe the structure of the atom, including the locations of the subatomic
particles.
Cathode- Ray Tube – Glass tube from which most of the air has been removed and
electricity is passed through it. The ray originating from the cathode end (negative) to the
anode (positive) is called a cathode ray.


Electrons – The negative charged particles that are part of all forms of matter.
Thomson by measuring the effects of magnetic and electric fields on a cathode ray, a
charge to mass ratio was determined. It was determined that the mass of the charged
particle was much smaller than a H atom so therefore there is something smaller than
an atom. He won the Nobel Prize in 1906 for identifying first subatomic particle.
Research of the cathode ray determined the following:
1.
2.
3.
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Millikan’s Oil Drop Experiment - In the early 1910s, Robert Millikan used the oil-drop
apparatus shown below to determine the charge of an electron.


With the electron’s charge and charge-to-mass ratio known, Millikan calculated the
mass of a single electron.
Mass of an electron - 9.1 x 10-28 or 1/1840 the mass of H
Early Models Atomic Models–
•
Plum Pudding Model by Thomson – states that the atom is a uniform, positively
changed sphere containing electrons.

In 1911, Ernest Rutherford studied how positively charged alpha particles
interacted with solid matter.
Rutherford’s experiment and model – the firing of alpha particles through a thin gold
foil, found that alpha particles were deflected at many angles and most pass right
through. Therefore most of the atom was space with a dense center.
Although most of the alpha particles went through the gold foil, a few of them
bounced back, some at large angles.
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The repulsive force between the positively charged nucleus and positive
alpha particles caused the deflections.
Nucleus- The dense center region of the atom that is positively charged.


Almost all of the atom's positive charge and almost all of its mass is contained in a
dense region in the center of the atom called the nucleus.

Niels Bohr- 1913 atomic structure relating to electron arrangement.
The Nucleus


Proton – a subatomic particle carrying an equal but opposite charge of an electron.
By 1920, Rutherford had refined his concept and concluded that the positively
charge particles in the nucleus were protons.
Neutron – subatomic particle that has a mass nearly equal to a proton but carries no
electric charge. Chadwick – Rutherford’s coworker received Nobel in 1935 for this
discovery.
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The Nucleus


All atoms are made of three fundamental subatomic particles: the electron, the
proton, and the neutron.


Atoms are mostly empty space, and electrons travel around the nucleus held by an
attraction to the positively charged nucleus.

Scientists have determined that protons and neutrons are composed of subatomic
particles called quarks.
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3:3 How Atoms Differ -The number of protons and the mass number define the type of atom.
Objective: Explain the role of atomic number in determining the identity of an atom
Atomic number – the number of protons in the nucleus of an atom identifies the element.

Moseley (1887-1915) discovered …..
Objective: Define an isotope

All atoms of a particular element have the same number of protons and electrons
but the number of neutrons in the nucleus can differ.

Isotopes – Atoms with the same number of protons and electrons but different
numbers of neutrons.
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
Isotopes have the same chemical behavior.
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The mass number is the sum of the protons and neutrons in the nucleus.
Examples of notation:
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Objective: Explain why atomic masses are not whole numbers
Objective: Calculate the number of electrons, protons, and neutrons in an atom given its mass
number and atomic number.
AMU – atomic mass unit – based on carbon-12 atom – therefore amu is defined as 1/12th
the mass of carbon.

Atomic Mass – is a weighted average of the mass of the isotopes of that element and
therefore is not a whole number.
Example Problem: Silver has two isoptopes:
amu and a percent abundance of 52.00%, and
, which has a mass of 106.905
, which has a mass of 108.905
amu and a percent abundance of 48.00%. What is the atomic mass of silver?
Example Problem 2: Chromium’s four naturally occurring isotopes are provided.
Calculate chromium’s atomic mass.
Cr-50 4.35%
49.946amu
Cr-52 83.79%
51.941amu
Cr-53 9.50%
52.941amu
Cr-54 2.36%
53.939amu
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3:4 Unstable Nuclei and Radioactive Decay - Unstable atoms emit radiation to gain
stability.
Objective: Explain the relationship between unstable nuclei and radioactive decay
Radioactivity – is the spontaneous emission of radiation by a radioactive material

Radiation – is the rays and particles emitted during a nuclear reaction.

Nuclear Reaction – can change an element into a new element by changing an atom’s
nucleus.
o
Radioactive decay – is a spontaneous process by which unstable nuclei lose energy by
emitting radiation.
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Atoms that contain too many or too few neutrons are unstable and lose energy
through radioactive decay to form a stable nucleus.
Objective: Characterize alpha, beta, and gamma radiation in terms of mass and charge.
Transmutation- When a reaction occurs that alters an atom’s atomic number.
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Positron –
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Electron Capture -
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Types of Radiation:
Alpha radiation – is made up of positively charged particles called alpha particles


Found because they are deflected towards a negavtively charged plate.
o
or α
Example equation: nuclear equation showing the radioactive decay of radium-226 to
radon-222.
o
Beta Radiation – is radiation that has a negative charge and emits beta particles


Found because they are deflected toward the positive charged plate
o
β or e-
o Example equation:
Gamma radiation – a high energy radiation that possesses no mass and no charge.
o Cannot, by themselves, result in the formation of a new atom.
o Gamma rays account for most of the energy lost during radioactive decay.
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o γ
o Example equation:
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Objective: Solve problems involving radioactive decay rates.
Radiochemical Dating: Carbon undergoes beta decay to form nitrogen. Carbon-14 has a
half-life of 5730 years.

The amount of stable carbon in the dead organism remains constant while the
carbon-14 continues to decay, the ratio of unstable carbon-14 to stable carbon-12
and 13 decreases over time.
Half-life – is the time required for one-half of a radioisotope’s nuclei to decay into its
products.
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
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N is the remaining amount
No is the initial amount
n is the number of half-lives that have passed. (or this can be replaced with t/T
where t is the elapsed tie and T is the duration of the half-life with the same units in
time).
Example Problem: Krypton-85 is used in indicator lights of appliances. The half-life of
krypton-85 is 11years. How much of a 2.00 mg sample remains after 33 years?
Example Problem 2: Bandages can be sterilized by exposure to gamma radiation
from cobalt-60, which has a half-life of 5.27 years. How much of a 10.0mg sample of
cobalt-60 is left after one half life? Two half lives? Three half lives?
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3:5 – Accumulating Content and Skills:– Chemistry content is continuous and builds on
prior knowledge and skills. This section will combine this unit with previous units.
Objective: Apply knowledge and skills from previous units to content learned in this unit.

How does law of conservation of mass apply to nuclear reactions?

How did the scientific method affect the history of chemistry?
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