Download The Atom.jet.2013 - Petal School District

Glencoe Chapters 4, 5, and 24
Atoms:
The smallest component of an element
having the chemical properties of that
element
The three best-known subatomic particles
are represented as follows:
Particle
Symbol
Mass
Charge
Neutron
n
1.6749 x 10-27kg
none
Proton
p+
1.6726 x 10-27kg
positive
Electron
e-
9.109 x 10-31kg
negative
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Protons and neutrons are in the nucleus
Electrons are in a “cloud” surrounding the
nucleus
Atomic Number:
Represents the number of protons in an atom
Example: Carbon (C) has 6 protons
Elements differ by their atomic number
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Mass Number (atomic mass):
Represents the number of protons and
neutrons in an atom
Ex: carbon has a mass number of 12.011
Isotopes:
Atoms of an element that contain different
numbers of neutrons
All elements have isotopes
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Atomic mass is an average of the masses of
all available isotopes of the element, based
on percent occurrence
Nomenclature to distinguish isotopes:
C-12 is “carbon twelve”
◦ ______ protons and ______ neutrons
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C-14 is “carbon fourteen”
◦ ______ protons and ______ neutrons
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Nuclear chemistry:
Deals with radioactivity, nuclear processes
and nuclear properties
Nuclear stability: depends on the neutron to
proton ratio
Band of stability:
Ratio from 1.0 to 1.5 where all isotopes are
stable
Any isotope outside the band is unstable
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Radioactivity:
The spontaneous emission of high energy
particles or rays by unstable atomic nuclei to
form stable nuclei
All transuranium elements are radioactive
Nucleic decay produces alpha particles, beta
particles, or gamma rays
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Alpha particles (α):
Represented by 42He
+2 charge
Are high speed helium nuclei
“Safest” of the basic forms of radioactivity
Fairly low penetrating power (can be stopped by
paper)
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Beta Particles (β or e-) :
Represented by 0-1e
High speed electrons
-1charge
Fairly high penetrating power (stopped by foil,
clothing)
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Gamma Rays (γ):
Represented by 00γ
High energy electromagnetic waves
No charge
Extremely high penetrating power (lead and concrete
will not completely block)
Very dangerous form of radiation
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Written as equations
mass numbers (superscripts)
atomic numbers (subscripts)
Superscripts and subscripts must equal on
both sides of the equation
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Examples:
Alpha decay:
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238
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92
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U→
234
90
4
Th +
2
0
He +
Beta decay:
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131

53
I →
131
54
Xe +
0
-1
e
2 0γ
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Positron emission
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Electron capture:
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Examples (continued):
Neutron bombardment*:
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235
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92
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1
U +
0
n
→
236
92
U
*Is different because it makes a less stable
nucleus while all the other types of decay
make a more stable nucleus. (used to initiate
chain reactions in nuclear power plants)
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A sequence of decay reactions that occur
spontaneously. Intermediate products are
unstable and the process will continue until a
stable product is formed.
Example:
Decay series:
◦ 1. Alpha decay of U-235 followed by
◦ 2. Beta decay of product from step 1 followed by
◦ 3. Alpha decay of product from step 2
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Each radioactive isotope has a half life
◦ The time it takes for half the mass of an available
sample to decay into another substance.
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Half-life can be used to:
◦ calculate how long it would take to result in a level of
the isotope that could be tolerated by living things
◦ calculate the mass of an isotope remaining after a period
of time.
mfinal = minitial x .5t/T
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t = amount of time that has passed
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T = length of the half-life
Example:
How much of a 25.0g sample of
radon-222 will remain after 30 days if
the half life of radon-222 is 10 days?
mfinal = minitial x .5t/T
t/T = 30days/10days = 3.0
mfinal = 25.0g x (.5)3 = 3.125g
More Examples
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Two sources:
Fission:
The splitting of nuclei to release energy
Neutrons are used to initiate the process
Fission reaction is self-sustaining
Currently used in nuclear power plants
Problems?
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Fusion:
Combining nuclei (hydrogen isotopes
combine to form helium)
Occurs on the sun and other stars
More powerful process than fission
Problems?
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Development of the modern atomic theory –
see handout. Study as part of lecture packet.
This will be part of the unit exam. A group
project is also from this material.
The quantum mechanical model of the atom is
based on the study of waves and light.
Definitions:
Amplitude – the distance from the crest to
the midway origin line of the wave
Wavelength ():
the distance between similar points in a set of
waves (crest to crest or trough to trough)
Measured in meters
Frequency ()
the number of waves that pass a given point
per unit time
Measured in cycles/second (sec-1)
Usually called a Hertz (Hz)
Frequency, wavelength, and energy are
related:
c = 
and
E = h
 = wavelength
 = frequency
c = the speed of light in a vacuum
(3.00x108 m/s)
h = Planck’s constant (6.626x10-34 J/Hz)
E = energy
Electromagnetic Radiation (EMR) – a form of energy that
travels in waves. Includes all visible light and other forms of
wave energy:
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All EMR is the result of electron movement
between energy levels within atoms
Electrons absorb energy to reach higher
(excited) energy state
Energy is absorbed and emitted in the form of
photons
Photons: discrete “packets” of energy
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Bright Line Spectrum:
Distinct lines of color that correspond to
wavelengths, frequencies, and energies
specific to each individual element
Can be used to identify elements
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Research by Louis deBroglie and Erwin
Schrodinger led to the branch of physics
called quantum mechanics.
Current atomic models are based on
quantum mechanics and probability
Probability is the statistical likelihood of an
occurrence
Heisenberg Uncertainty Principle: states that
it is not possible to know both the position
and the speed of an electron at the same
time
Schrodinger developed an equation to
describe electron location and behavior –
uses quantum numbers
n = Principle Quantum Number
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refers to the energy level location of the
electron
Can be from 1-7
 l = Orbital Quantum Number
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refers to the sublevel location of the
electron
Can be s, p, d, or f
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s orbitals:
◦
◦
◦
◦
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Spherical shape
Lowest energy
Closest to the nucleus
One per energy level
p orbitals:
◦
◦
◦
◦
Figure 8 shape
Higher energy
Further from the nucleus
Three per energy level starting with 2nd
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d orbitals:
◦
◦
◦
◦
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Cloverleaf shape
Even higher energy
Even further from the nucleus
Five per energy level starting with 3rd
f orbitals:
◦
◦
◦
◦
Very complex shape
Highest energy
Furthest from the nucleus
Seven per energy level starting with 4th
ml
= Magnetic Quantum Number
ms
= Spin Quantum Number
◦ designates the number of orbitals on each sublevel
◦ designates in which specific orbital the electron is
likely located
◦ describes the direction of spin for the electron
◦ is either clockwise or counter-clockwise
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Basic Rules:
Aufbau Principle – electrons enter lowest
energy levels first
Hund’s Rule – each orbital in a sublevel is
occupied by one electron before any orbital
obtains two electrons
All electrons in singly occupied orbitals have
the same spin directions
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Basic Rules (continued):
The Pauli Exclusion Principle:
◦ No more than 2 electrons may occupy any
one orbital
◦ Electrons that occupy the same orbital
must have opposite spins
◦ No two electrons in one atom will have the
same the four quantum numbers
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Use the basic rules and basic arrangement of
the periodic table:
◦ Row number
◦ Position in row
◦ s, p, d, and f blocks
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Electron Configuration Notation, Orbital
Diagrams, Noble Gas Shorthand, and Lewis
Dot Structures for each of the following. Include
in your lecture packet to study.
Use this:
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
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Orbital Diagrams:
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
P
S
Cl
Ca
Sc
V
Ni
Zn
Fe
Cu
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Electron Configuration Notation:
H
Na
He
Mg
Li
P
Be
S
B
Cl
C
Ca
N
Sc
O
V
F
Ni
Ne
Zn
Fe
Cu
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Noble Gas Shorthand:
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
P
S
Cl
Ca
Sc
V
Ni
Zn
Fe
Cu
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Lewis Dot Structures:
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
P
S
Cl
Ca
Sc
V
Ni
Zn
Fe
Cu