Download UNIT 1 - Grafton Public Schools

UNIT 1: ATOMIC STRUCTURE
& NUCLEAR CHEMISTRY
HONORS Chemistry
Grafton High School
UNIT OBJECTIVES:
• SWBAT:
Key Questions
• How did the concept of the atom change and develop from the
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time of Democritus to the time of John Dalton?
What is the structure of the nuclear atom?
What are the three kinds of subatomic particles?
What makes one element different from another?
How do isotopes of an element differ?
How do you calculate the atomic mass of an element?
How do nuclear reactions differ from chemical reactions?
What are the three types of nuclear radiation?
How much of a radioactive sample remains after each half life?
What are three devices used to detect radiation?
What are some practical uses of radioisotopes?
The ATOM
• An atom is the smallest particle of an
element that retains its identity in a
chemical reaction.
• Although early philosophers and
scientists could not observe individual
atoms, they were still able to propose
ideas about the structure of atoms.
HELIUM ATOM
Helium Atom
Shell
proton
+
-
N
N
+
electron
What do these particles consist of?
-
neutron
Subatomic Particles
Particle
Charge
Mass
proton
+ ve charge
1
neutron
No charge
1
electron
-ve charge
nil
Numbers on the Periodic Table
He
2
4
Atomic number
the number of protons in an atom
Mass Number
the number of protons and
neutrons in an atom
number of electrons = number of protons
Energy Levels (shells)
Electrons are arranged in Energy Levels or Shells
around the nucleus of an atom.
•
first shell
a maximum of 2 electrons
•
second shell
a maximum of 8 electrons
•
third shell
a maximum of 8 electrons
Electron Configuration
With electronic configuration elements are represented
numerically by the number of electrons in their shells and
number of shells.
For example;
Nitrogen
2 in 1st shell
5 in
2nd
shell
configuration = 2 , 5
2
+
5 = 7
N
7
14
Electron Configuration ~ Bohr Model
• Each orbital can hold a certain number of electrons.
Orbitals (shells) fill from the inside (nearest nucleus) out.
Electron Configurations
EXAMPLES of electronic configuration for the following
elements:
a)
Ca
20
b)
Na
40
2,8,8,2
d)
Cl
17
35
2,8,7
11
23
c)
2,8,1
e)
Si
14
28
2,8,4
O
8
16
2,6
f)
B
5
11
2,3
Bohr Diagrams
With Bohr diagrams, elements and compounds are
represented by Dots or Crosses to show electrons, and
circles to show the shells. For example:
X
Nitrogen
X X
N
XX
X
X
N
7
14
Examples of Bohr Diagrams
Draw the Dot & Cross diagrams for the following
elements;
X
8
X
a) O
b) Cl17
X
35
16
X
X
X
X
X
X
X X X Cl X X
X
X
X
O
X
X
X
X
X
X
X
X
X
Isotopes ~ varying # of neutrons
• Atoms of a certain element all have the same
number of protons
• Number of neutrons may vary
• The different variations in # of neutrons for a
certain element are called ISOTOPES
• Atomic Mass (decimal)= number of protons
+ average number of neutrons from all
isotopes
Isotopes of
Carbon
Looking at isotopes
• The different numbers of neutrons in isotopes gives them
different masses
• Isotopes of an element do not occur in equal amounts:
• Example:
Isotopic Abundance
• Isotopic Abundance = Percentage of an element’s
atoms that exist as each isotope
• Each isotope has its own mass that is different from other
isotopes of that element. Isotopic mass is the average
mass of the atoms of a specific isotope of an element.
• Determines the isotope number
Relative Atomic Mass
• WATCH: https://goo.gl/4y80yz
• What does the atomic mass on the periodic table
measure?
• The relative atomic mass shown as a decimal on the
periodic table is the amount of protons plus the weighted
average number of isotopes for that element:
Atomic Number
Chemical Symbol
Relative Atomic Mass
Calculating Relative Atomic Mass
• NOTE: Use the percent abundance of the isotopes as a
DECIMAL
Relative Atomic Mass (Average Weighted Mass):
Relative Atomic Mass = (% Isotope A as decimal x Isotopic Mass Isotope A)
+ (% Isotope B as decimal x Isotopic Mass Isotope B)
+ (% Isotope C as decimal x Isotopic Mass Isotope C)
EXAMPLE:
• Calculation of Relative Atomic Mass for Carbon:
Atomic Mass Unit
(AMU)
• The unit of
measurement for an
atom is an AMU
• AMU = Atomic Mass
Unit
• ONE (1) AMU = mass of
ONE (1) proton
• Why NOT grams?
Atomic Mass Unit (AMU)
• There are 6 X 1023 or
600,000,000,000,000,0
00,000,000 amus in one
gram
• (Remember that
electrons are 2000
times smaller than one
amu/one proton)
Radioactivity
• One of the pieces of evidence for the
fact that atoms are made of smaller
particles came from the work of
________ (1876-1934).
• She discovered ________, the
spontaneous disintegration of some
elements into smaller pieces.
• WATCH: https://goo.gl/IyuRj4
Nuclear Reactions vs. Normal
Chemical Changes
• Radioisotopes are isotopes with unstable nuclei and
undergo radioactive decay
• Nuclear reactions involve the nucleus
• The nucleus opens, and protons and neutrons are
rearranged
• The opening of the nucleus releases a tremendous
amount of energy that holds the nucleus together –
called binding energy
• “Normal” Chemical Reactions involve electrons, not
protons and neutrons
Mass Defect
• Nuclear reactions involve energy changes many times
the magnitude of chemical changes
• Chemical reactions – Energy is conserved
• Nuclear reactions – Small loss of matter (mass defect)
releases tremendous energy (some of the mass can be
converted into energy
• Shown by a very famous equation!
2
E=mc
Energy
Mass
Speed of light
Types of Radiation
Alpha (ά) – a positively
charged helium isotope •
we usually ignore the charge because it
involves electrons, not protons and
neutrons
e
0
1
•Beta (β) – an electron
4
2
He
0
1
e
•Gamma (γ) – pure energy; 0

0
called a ray rather than a
particle
Other Nuclear Particles
•
Neutron
• Positron – a positive
electron
•Proton – usually referred to
as hydrogen-1
•Any other elemental
isotope
1
0
n
0
1
e
1
1
H
Penetrating Ability
Balancing Nuclear Reactions
•In the reactants (starting materials –
on the left side of an equation) and
products (final products – on the right
side of an equation)
Atomic numbers must balance
and
Mass numbers must balance
•Use a particle or isotope to fill in the
missing protons and neutrons
Alpha, Beta, and Gamma Particles
• Atoms are not all stable.
• Excess energy in an unstable atom is released in a basic
particle or wave
• Greek alphabet is used to name particles (in order of
their discovery)
REMEMBER!
• The nuclear structure changes with radioactive decay,
but radioisotopes are NOT chemically different!
Alpha Particles
• Heaviest particle
• Rays, NOT waves –made of
high-energy particles that
are expelled from unstable
nuclei
• Alpha particle is a helium ion
• Not very penetrating; easily
absorbed
• SPEED: 16,000km/sec
(1/10th speed of light!)
Nuclear Reactions
• Alpha emission
Note that mass number (A) goes down by
4 and atomic number (Z) goes down by 2.
Nucleons (nuclear particles… protons
and neutrons) are rearranged but
conserved
Beta Particles
• LIGHTER energy particles
• Energetic electron given off
•
•
•
•
by unstable nuclei to restore
energy balance
Stopped by aluminum (few
mm thick) or 3 meters of air
8,000 times smaller than
alpha particle
SPEED: 270,000 km/sec
Can cause cellular damage
Nuclear Reactions
• Beta emission
• Note that mass number (A) is unchanged and
atomic number (Z) goes up by 1.
• This is because it is caused by a neutron that
breaks apart into a proton and electron
(electron is emitted).
Other Types of Nuclear Reactions
Positron (0+1b): a positive electron
207
Electron capture: the capture of an electron
207
Gamma Ray
• HIGH-ENERGY photon
(light wave)
• Same family as X-rays and
light
• MORE energetic and harmful
• Damaging to living cells
• Slows down energy by
transferring it to cell
components
Artificial Nuclear Reactions
New elements or new isotopes of known
elements are produced by bombarding an
atom with a subatomic particle such as a
proton or neutron -- or even a much heavier
particle such as 4He and 11B.
Reactions using neutrons are called
 reactions because a  ray is usually
emitted.
Radioisotopes used in medicine are often
made by  reactions.
Artificial Nuclear Reactions
Example of a
 reaction is production of
radioactive 31P for use in studies of P uptake in the
body.
31 P
15
+
1 n
0
--->
32 P
15
+ 
Transuranium Elements
Elements beyond 92 (transuranium) made starting
with an  reaction
238 U
92
+
239 U
92
239 Np
93
1 n
0
--->
239 U
92
+ 
--->
239 Np
93
+ 0-1b
--->
239 Pu
94
+
0 b
-1
Nuclear Fission
Nuclear Fission
• Fission is the splitting of atoms
• Controlled reaction
• These are usually very large, so that they are not as
stable
• Nuclear Power stations use heat released by nuclear
reaction (nuclear fission) to boil water to make steam
Fission Process - MODEL
Nuclear Fission & POWER
• Currently about 103 nuclear
power plants in the U.S. and
about 435 worldwide.
• 17% of the world’s energy
comes from nuclear.
Diagram of a nuclear power plant
Nuclear Explosion
• Explosion occurs when reaction is allowed
to run out of control
• Large amounts of ENERGY released
QUICKLY
• WATCH:
• Nuclear reactor controls rate of energy
release
• Uranium oxide held in fuel rods
• Rods lowered into reactor core
• Coolant (CO2) circulated to remove heat
• Control rods in core absorb neutrons and control
rate of chain reaction
• Raised to speed up; lowered to slow down
Nuclear Fusion
Fusion
small nuclei combine
2H
1
+
3H
1
4He
2
+ 1n +
0
Occurs in the sun and other stars
Energy
Nuclear Fusion
Fusion
• Excessive heat can not be contained
• Attempts at “cold” fusion have
FAILED.
• “Hot” fusion is difficult to contain
Half-Life
• HALF-LIFE is the time that it takes for 1/2 a sample to decompose.
• The rate of a nuclear transformation depends only on the “reactant”
concentration.
Half-Life
Decay of 20.0 mg of 15O. What remains after 3 halflives? After 5 half-lives?
Kinetics of Radioactive Decay
For each duration (half-life), one half of the
substance decomposes.
For example: Ra-234 has a half-life of 3.6 days
If you start with 50 grams of Ra-234
After 3.6 days > 25 grams
After 7.2 days > 12.5 grams
After 10.8 days > 6.25 grams
Learning Check!
The half life of I-123 is 13 hr. How much of
a 64 mg sample of I-123 is left after 39
hours?
Effects of Radiation
Geiger Counter
• Used to detect radioactive substances
Radiocarbon Dating
Radioactive C-14 is formed in the upper
atmosphere by nuclear reactions initiated by
neutrons in cosmic radiation
14N + 1 n ---> 14C + 1H
o
The C-14 is oxidized to CO2, which circulates
through the biosphere.
When a plant dies, the C-14 is not replenished.
But the C-14 continues to decay with t1/2 = 5730
years.
Activity of a sample can be used to date the
sample.
Nuclear Medicine: Imaging
Thyroid imaging using Tc-99m
Food Irradiation
•Food can be irradiated with  rays
from 60Co or 137Cs.
•Irradiated milk has a shelf life of 3 mo.
without refrigeration.
•USDA has approved irradiation of
meats and eggs.