Download I. Structure of the Atom

Unit 3
I. Structure of the Atom
I. Structure of the Atom
1. Many philosophers of ancient times
concluded that matter was composed
of things such as earth, water, air, and
fire.
2. It was also commonly accepted that
matter could be endlessly divided into
smaller and smaller pieces.
3. Democritus, a Greek philosopher
(460-370 B.C.), was the first to
propose the idea that matter was
not infinitely divisible; he thought
matter was made up of tiny
individual particles called atomos.
I. Structure of the Atom
I. Structure of the Atom
4. Aristotle (384-322 B.C.), a Greek
philosopher, rejected the atomic
“theory” because it did not agree
with his own ideas on nature.
• The influence of Aristotle was so great
and the development of science so
primitive that his denial of the existence
of atoms went largely unchallenged for
2000 years.
I. Structure of the Atom
7. Dalton’s Atomic Theory
a. All matter composed of small particles called
atoms.
b. All atoms of a given element are identical,
having the same size, mass, and chemical
properties. Atoms of a specific element are
different from those of any other element.
c. Atoms can’t be created, divided, or destroyed.
d. Different atoms combine in simple wholenumber ratios to form compounds.
e. In a chemical reaction, atoms are separated,
combined, or rearranged.
5. The concept of the atom was revived
in the 18th century, but it took another
100 years for significant progress to
be made.
6. John Dalton (1766-1844), a teacher in
England, revised Democritus’s ideas
based on his own scientific research
and marked the beginning of the
modern atomic theory.
I. Structure of the Atom
8. Errors in Dalton’s Atomic Theory
a. Atoms ARE divisible into several subatomic
particles.
b. All atoms of a given element don’t always
have identical properties (atoms of an
element may have slightly different masses).
9. Atom: the smallest particle of an element
that retains the properties of the element
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I. Structure of the Atom
10. Thomson Model (1904): English physicist J.J.
Thomson inferred from his experiments that atoms
contained small, negatively charged particles. He
thought these “electrons” were evenly embedded
throughout a positively charged sphere, much like
chocolate chips in a ball of cookie dough.
11. Electron: a negatively charged, fast-moving
particle with an extremely small mass that is found
in all forms of matter and moves through the empty
space surrounding an atom’s nucleus.
I. Structure of the Atom
14. Bohr Model (1913): Danish physicist Niels Bohr
hypothesized that electrons traveled in fixed orbits
around the atom’s nucleus, which contained
positively charged particles. James Chadwick, a
student of Rutherford, concluded that the nucleus
contained positive protons and neutral neutrons.
15. Proton: a subatomic particle in an atom’s nucleus
that has a positive charge of 1+.
16. Neutron: a neutral subatomic particle in an atom’s
nucleus that has a mass nearly equal to that of a
proton.
I. Structure of the Atom
18. Quarks: particles of matter that make up
protons and neutrons.
19. Atomic Number: the number of protons in
an atom
20. Atomic # = # of protons = # of electrons (in
a neutral atom)
21. # of neutrons = Mass # – Atomic #
22. Isotopes: atoms with the same number of
protons but different numbers of neutrons
I. Structure of the Atom
12. Rutherford Model (1911): Another British
physicist, Ernest Rutherford, proposed that almost
all the mass of an atom – and all its positive
charges – were concentrated in a central atomic
nucleus surrounded by electrons.
13. Nucleus: the extremely small, positively charged,
dense center of an atom that contains positively
charged protons, neutral neutrons, and is
surrounded by empty space through which one or
more negatively charged electrons move.
I. Structure of the Atom
17. Electron Cloud Model (1926 –
present day): According to the
currently accepted model of atomic
structure, electrons do not follow fixed
orbits but tend to occur more
frequently in certain areas around the
nucleus at any given time.
I. Structure of the Atom
23. Mass Number: used to identify each of the
various isotopes of an element; the sum of
the number of protons and neutrons in the
nucleus; always a whole number
(example: carbon-13)
24. The masses of both protons and neutrons
are approximately 1.67 x 10-24 g.
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I. Structure of the Atom
25. Because these numbers are difficult to
work with, a different system is used.
26. Scientists assigned the carbon-12
atom a mass of exactly 12 atomic
mass units.
27. One atomic mass unit (amu) is
defined as 1/12 the mass of a carbon12 atom.
I. Structure of the Atom
29. Isotope Notation
Mass #
Atomic #
107
Ag
47
1-
Charge
II. Bohr Diagrams
1. A row of elements on the periodic table is
called a period.
2. The period # indicates the number of
energy levels in an atom of an element in
that period.
3. A column of elements on the periodic table
is called a group or family.
4. The group # indicates the # of valence
electrons (electrons in the outer energy
level).
I. Structure of the Atom
28. Atomic Mass: the weighted average mass
of the isotopes of that element
Example: Chlorine exists naturally as a
mixture of about 75% chlorine-35 and 25%
chlorine-37. Because atomic mass is a
weighted average, the chlorine-35 atoms,
which exist in greater abundance, have a
greater effect in determining the atomic
mass.
I. Structure of the Atom
30. The presence of a charge on an atom
means it has more or fewer electrons
than protons.
31. A negative charge means it has more
electrons than protons.
32. A positive charge means it has fewer
electrons than protons.
33. Ion: An atom with a charge
II. Bohr Diagrams
5. To draw a Bohr diagram, write the number
of protons (p+) and neutrons (n0) in the
nucleus. Then draw circles around the
nucleus – one for each energy level.
Determine the total number of electrons in
the atom (same as the atomic number,
unless it’s an ion). Fill in the electrons on
the circles, with a maximum of 2 electrons
on the first circle, 8 electrons on the second,
18 electrons on the third, and 32 electrons
on the fourth.
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II. Bohr Diagrams
6. Before placing more than 8 electrons
on any energy level above the second
energy level, you must first place the
correct number of valence electrons in
the outer energy level. Then go back
and fill the lower energy levels with
any remaining electrons.
III. Changes in the Nucleus
6. Atoms with mass numbers around 60 are
the most stable. In order to increase their
stability, heavy atoms sometimes fragment
into smaller atoms, and light atoms
sometimes combine into larger atoms.
7. Nuclear fission: the splitting of a nucleus
into fragments
8. Nuclear fusion: the combining of atomic
nuclei
III. Changes in the Nucleus
1.
2.
3.
4.
5.
Strong Force: force that acts between protons and
neutrons in an atomic nucleus and keeps them
together
Nuclear Reaction: a reaction that involves a
change in the nucleus of an atom
Radioactive Decay: a spontaneous process in
which unstable nuclei lose energy by emitting
radiation
Radioisotopes: isotopes of atoms with unstable
nuclei
Why do these atoms emit radiation? TO BECOME
MORE STABLE
III. Changes in the Nucleus
9. Transmutation: the conversion of an
atom of one element to an atom of
another element
10. Half-life: the time required for one-half
of a radioisotope’s nuclei to decay into
its products
Amount remaining = (Initial Amount)(1/2)n
n = # of half-lives
III. Changes in the Nucleus
11. Sample Problem:
The half-life of iron-59 is 44.5 days.
How much of a 2.000 mg sample will
remain after 133.5 days?
III. Changes in the Nucleus
12. Three types of radiation:
a. Alpha: made up of alpha particles, which
contain 2 protons and 2 neutrons and have a
2+ charge; equivalent to a helium-4 nucleus
b. Beta: made up of fast-moving electrons
called beta particles, which have a 1- charge
c. Gamma: consists of gamma rays, which are
high-energy radiation that possess no mass
or charge; usually accompany alpha and beta
radiation and account for most of the energy
lost during the radioactive process
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III. Changes in the Nucleus
13. Characteristics of Radiation
Radiation
Type
Symbol
Alpha
4
He
2
Beta
or
Charge
4
2+
1
1840
0
1-
α
0
β
-1
Gamma
Mass
(amu)
0
γ
0
III. Changes in the Nucleus
14. Nuclear Equation: a type of equation
that shows the atomic number and
mass number of the particles involved
15. Alpha radiation – The alpha decay of
radioactive radium-226 into radon-222
is shown below:
226
0
III. Changes in the Nucleus
16. Beta radiation – The beta decay of
radioactive carbon-14 into nitrogen-14
is shown below:
14
6
C
14
7
N
+
0
-1
18. Positron Emission: a radioactive decay
process that involves the emission of a
positron from a nucleus; a proton in the
nucleus is converted into a neutron and a
positron, and the positron is emitted
19. Positron: a particle with the same mass as
an electron but opposite in charge
0
1
+
4
2
He
17. Gamma radiation – For example,
gamma rays accompany the alpha
decay of uranium-238 as shown
below:
234
U
92
III. Changes in the Nucleus
Rn
86
III. Changes in the Nucleus
238
β
222
Ra
88
Th
90
+
4
2
He
+
2
0
0
γ
III. Changes in the Nucleus
20. Example of Positron Emission:
26
14
Si
26
Al
13
+
0
1
β
β
5
III. Changes in the Nucleus
21. Electron Capture: occurs when the
nucleus of an atom draws in a
surrounding electron; this electron
combines with a proton to form a
neutron
22. Example of an Electron Capture:
0
-1
e
+
81
Rb
37
81
36
Kr +
X-ray photon
III. Changes in the Nucleus
EXERCISE #1
Alpha decay of iridium-174
EXERCISE #2
Beta decay of platinum-199
EXERCISE #3
Nickel-60, a beta particle, and a gamma
ray are the products of the decay of
which isotope?
III. Changes in the Nucleus
EXERCISE #4
Positron emission from sulfur-31
EXERCISE #5
Beryllium-7 undergoes electron capture.
EXERCISE #6
The half-life of radium-224 is 3.66 days. What
was the original mass of this radioactive
isotope if 1.2 mg remains after 10.98 days?
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