Download Atomic Structure

Gilbert
Kirss
Foster
Chapter 2
Atomic Structure
Subatomic Particles?
• If you cut a piece of graphite from the tip of a pencil into
smaller and smaller pieces, how far could you go?
• You would eventually end up with atoms (translates to
“indivisible” in greek) of pure carbon.
• You can not divide a carbon atom into smaller
pieces and still have carbon
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Atomic Structure
• We have established that matter is comprised of atoms. But what
are atoms made of?
• There are three types of sub-atomic particles that make up the
atom are known as:
• electrons
• protons
• neutrons
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Setting up the Cathode Ray Exp.
• The electron was the first subatomic particle discovered.
• In the late 1800’s, J.J Thomson sought to understand the
strange “cathode ray” phenomena, which involved the
observation of a “strange, flowing energy” through gases.
• He developed the cathode ray experiment, which was
comprised of:
• Glass tube from which most of the air was removed
• Two metallic plates, an anode and a cathode, connected to a
high voltage power supply.
• The cathode is connected to the negative terminal of the
power source. The anode has a small hole drilled through
its center.
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Observation of Cathode Rays
• When the connections were made, these mysterious
cathode rays flow from the cathode to the anode, and
some of these rays escape through the hole in the anode.
• The rays are invisible, so a phosphorescent screen lines
the back of the tube, which exhibits a glowing spot when
struck by the beam.
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What Are Cathode Rays?
• Thompson soon realized that the cathode rays could be deflected
by electric and magnetic fields. The image below shows a cathode
ray beam being deflected upwards toward a positive pole.
• He also found that the mass of the cathode remained virtually
unchanged. What does this mean???
• The beam is not energy, but rather, charged, nearly massless
particles, and the particles are negative!
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Plum-Pudding Model
• Following the discovery of the electron, it became obvious that
positive charges, called protons, must also exist since matter is
electrically neutral
• However, scientist had no idea how these particles were
arranged in the atom. The first proposed model was the “plum
pudding model”, which described electrons as being spread out
in a proton “sea”
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Determining the Nuclear Model
• Following the discovery of radioactivity, the “gold foil experiment”
was designed to test the PP model.
• A thin sheet of gold foil was placed
in a phosphorescent ring. A
radioactive emitter of positive
particles was placed in front.
• If the PP model was correct, then
the positive α-particles would pass
through the foil unimpeded. But…
• The particles were actually
deflected!! How?
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Expected Results
Expected results from
“plum-pudding” model.
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Rutherford’s Experiment
• The experiment not only disproved
the PP model, but also suggested
that a very dense, very positive
“core” exists at the center of the
atom, in which all positive charges
are found.
• This came to be known as the
nucleus.
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Neutrons
• Rutherford’s model was incomplete. For example, a hydrogen
atom has one proton and one electron, but is only ¼th the mass of
a helium atom which has two electrons and two protons.
• If all of the mass of an atom comes from its sub-atomic particles,
how do we explain the unaccounted for mass?
• The answer is neutrons, particles that are equal in mass to
protons, but with no electrical charge. While scientists knew that
neutrons had to exist, they were not officially discovered until
1932.
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The Nuclear Atom
• Positively charged center of an atom, containing nearly all of
the atom’s mass
• About 1/10,000 the size of the atom
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About the Nucleus
• Atomic Mass Units (amu)
• Unit used to express the relative masses of atoms and
subatomic particles
• Equal to 1/12 of a carbon atom
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Elemental Symbols
6
Atomic #
C
Carbon
12.0107
• The number of protons in an atom is called the atomic number. An
element is defined by its atomic number. (ex. only carbon has 6
protons)
• For a given element, the number of protons DOES NOT CHANGE
• In a neutral atom, the number of protons is equal to the number of
electrons.
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Elemental Symbols
6
C
Carbon
12.0107
Mass #
• The mass number of an element is the sum of its protons and
neutrons.
• The mass #’s listed on the periodic table are averages, in units of
amu
• These averages are used because numerous variations of
elements called isotopes exist in nature.
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Isotopes
• Isotopes are variations of the same element having different
numbers of neutrons.
Atomic Mass = total number of “nucleons”
(protons, neutrons) in the nucleus
Atomic Number (Z) = the number of protons
A
Z
X
• Isotope symbols are shown below for the two isotopes of nitrogen
with their % abundances in nature. The 14N and 15N isotopes have
7 and 8 neutrons, respectively.
𝟏𝟒
𝟕𝑵
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(99.636%)
𝟏𝟓
𝟕𝑵
(0.346%)
Group Work
• Complete the missing information in the table.
23
?
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Transitional Page
Avg. atomic mass is obtained using the % abundance and the isotope mass.
𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 =
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𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑎𝑠𝑠 𝑥 (% 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒)
Group Work
• Using the given abundances and isotope masses, calculate the
average atomic mass of C. Does it match the reported value?
ISOTOPE
%A
Mass (amu)
𝟏𝟐
𝟔𝑪
𝟏𝟑
𝟔𝑪
𝟏𝟒
𝟔𝑪
98.93
12
1.07
13.003 354 8378
~0
14.003 2420
• Boron has two isotopes, 10B and 11B. Using the given isotope
masses, determine the % abundances of each isotope. Hint: total
abundance must equal 100%
ISOTOPE
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%A
Mass (amu)
𝟏𝟎
𝟓𝑩
10.013
𝟏𝟏
𝟓𝑩
11.009
Proton-Neutron Ratio
• The nuclei of most naturally occurring isotopes are very stable,
despite the massive repulsive forces that exist between the
protons in the nucleus.
• A strong force of attraction between neutrons and protons known
as the nuclear force counteracts this repulsion.
• As the number of protons increases, more neutrons are required
to stabilize the atom. Stable nuclei up to atomic number 20 have
equal numbers of protons and neutrons.
• For nuclei with atomic number above 20, the number of neutrons
exceeds the protons to create a stable nucleus.
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Proton-Neutron Ratio
• Radioactive isotopes are unstable (high in energy). This instability
is attributed to a neutron/proton ratio that is either too high or too
low.
• To become stable, they spontaneously release particles or
radiation to lower their energy.
• This release of energy is called radioactive decay.
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Radioactivity
• The three most common types of radioactive decay are alpha,
beta, and gamma
Property
α
β
γ
Reason for process
Too many protons,
too few neutrons
(n/p ratio too low)
Too few protons,
too many neutrons
(n/p ratio too high)
Too much energy in
nucleus
Charge
2+
1-
0
Mass
6.64 x 10-24 g
9.11 x 10-28 g
0
Emitted Radiation
Type
2 protons and 2
neutrons ( 42𝐻𝑒)
High energy
electron.
Pure energy
(Radiation)
Penetrating Power
Low. Stopped by
paper. Blocked by
skin.
Moderate. Stopped
by aluminum foil.
(10α)
High. Can
penetrate several
inches of lead.
(10000α)
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Radioactive Decay
• For example, the 238
92𝑈 isotope undergoes alpha decay to increase
its n/p ratio:
238
92𝑈
238 − 92 𝑛
= 1.58 𝑛/𝑝
92 𝑝
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→
234
90𝑇ℎ
+ 42𝐻𝑒
234 − 90 𝑛
= 1.60 𝑛/𝑝
90 𝑝
Radioactive Decay
• The Thorium-234 isotope undergoes beta decay which lowers the
n/p ratio:
234
90𝑇ℎ
→
234 − 90 𝑛
= 1.60 𝑛/𝑝
90 𝑝
234
91𝑃𝑎
+
0
−1𝑒
234 − 91 𝑛
= 1.57 𝑛/𝑝
91 𝑝
• In beta decay, a neutron is converted to a proton and an electron.
This causes the proton count to increase:
1
0𝑛
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→ 11𝑝 + −10𝑒
Ions
• Thus far, we’ve learned that each element has an exact number of
protons.
• For example, Hydrogen has only one proton. If you force a
second proton onto the atom, you no longer have hydrogen…
you now have Helium.
• We have also learned that atoms can have variable numbers of
neutrons (isotopes).
• Next, we will discuss ions.
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Ions
• Ions are electrically charged atoms, resulting from the gain
or loss of electrons.
• Positively charged ions are called cations. You form
cations when electrons are lost
• Negatively charged ions are called anions. You form anions
when electrons are gained
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Ion Nomenclature
• A cation is named by adding the word “ion” to the end of the
element name
• Anions are named by adding the suffix –ide to the end of an
element
𝑳𝒊+
Lithium ion
𝑪𝒍−
Chloride
𝑵𝒂+
Sodium ion
𝑺𝟐−
Sulfide
𝑴𝒈𝟐+
Magnesium ion
𝑶𝟐−
Oxide
𝑨𝒍𝟑+
Aluminum ion
𝑷𝟑−
Phosphide
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Group Work
• Fill in the missing information below
ISOTOPE
P
N
E
13
14
10
32
16𝑆
32 216𝑆
??
?? 4+
??𝑃𝑡
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