Download Atomic Structure, the Periodic Table, and Nuclear Radiation

Atomic Structure, the
Periodic Table, and Nuclear
Radiation
The Electromagnetic Spectrum
Basics of Periodic Table
• Each box on the table represents an
element.
• In each box…
– an element symbol
– the element’s atomic number
– the element’s average atomic mass
• Elements arranged in order of increasing
atomic number.
Isotope symbol (neutral)
Mass number
(#p + #n)
Atomic number
(#P)
14
6
Isotope name carbon-14
C
Isotope symbol (charged)
32
15
P
3−
Average Atomic Mass
• The masses found on the periodic table
are called average atomic masses.
• They represent the weighted average of all
the isotopes found in a sample of the
element
– Isotopes are atoms of the same element with
different numbers of neutrons
Example
The atomic masses of the two stable isotopes of boron,
boron-10 (19.78%) and boron-11 (80.22 %), are 10.0129
amu and 11.0093 amu, respectively. Calculate the
average atomic mass of boron.
Remember…
• Electron configuration…ie. Shorthand
– Should be able to do shorthand w/o diagonal rule
• Aufbau principle – electrons fill energy levels
and sublevels in order of increasing energy
• Pauli Exclusion principle – no two electrons can
have the same set of four quantum numbers
(which means no two electrons can be in the
same place at the same time)
• Hund’s rule – when adding electrons to
sublevels with more than one orbital, each
orbital gets its own electron first before pairing
s orbitals
p orbitals
d orbitals
f orbitals
• Metalloids
are on the
stair step
• Metals to the
left of the
stair step
• Nonmetals
to the right
of the stair
step (plus H)
Quantum Numbers
• Just as a point on an xy-graph needs a set
of two coordinates, each electron has a
unique set of four coordinates.
• These four coordinates represent shell
(energy level), subshell (sublevel), orbital,
and spin direction of the electron.
Principal Quantum number
• Represented by n
• Corresponds to the rows of the periodic
table
• Therefore n = 1, 2, 3, and so on
• Tells the size of the electron cloud
2nd Quantum number
• Represented by l
• Called the angular momentum quantum number
• Describes the shape of the orbital
• l can have the values from 0 to n-1
0 = s sublevel
1 = p sublevel
2 = d sublevel
3 = f sublevel
3rd Quantum Number
• Called the magnetic quantum number
• Describes the orientation in space of the orbital
– Whether the path of the electron lies on the x, y, or z
axis
• Represented by ml
•
ml can have values from –l
to +l
– if l = 2, then ml = -2, -1, 0, +1, +2
4th Quantum Number
•
•
•
•
Corresponds to the spin of an electron
Represented by ms
Clockwise represented by +1/2
Counterclockwise represented by -1/2
Therefore
• Mg
(3, 0, 0, -1/2)
• Bi
(6, 1, +1, +1/2)
• Co
(3, 2, -1, -1/2)
• Cf
(5, 3, -1, -1/2)
Diamagnetism/Paramagnetism
• Diamagnetic elements have all of their
electrons spin paired.
– Which means they have complete sublevels.
– Are not affected by a magnetic field
• Paramagnetic elements do not have all of
their electrons spin paired.
– Strongly affected by a magnetic field
Ground State vs Excited State
• In a ground state atom, all electrons are in
the lowest available sublevels.
• For an atom in the excited state, one or
more electrons have absorbed enough
energy to jump to higher energy levels.
– As soon as possible, those excited electrons
will release that energy in the form of a
photon, possibly as colored light.
Shorthand for ions
Ca2+
K+
ClS2P3-
1s22s22p63s23p6
1s22s22p63s23p6
1s22s22p63s23p6
1s22s22p63s23p6
1s22s22p63s23p6
All of these ions have the same configuration as
argon and are isoelectronic.
Periodic Trends
• You can make predictions about certain
behavior patters of an atom and its
electrons based on the position of the
atom in the periodic table.
• All the periodic trends can be understood
in terms of three basic rules.
1. Electrons are attracted to the protons in the nucleus of
an atom.
– The closer an electron is to the nucleus, the more
strongly it is attracted.
– The more protons in a nucleus, the more strongly an
electron is attracted.
2. Electrons are repelled by other electrons in an atom.
So, if other electrons are between a valence electron
and the nucleus, the valence electron will be less
attracted to the nucleus. That’s called shielding.
3. Completed energy levels (and to a lesser extent,
completed sublevels) are very stable. Atoms prefer to
add or subtract valence electrons to create complete
shells if possible.
Atomic Radius
• Atomic radius is the approximate distance
from the nucleus of an atom to its valence
electrons
– Atomic radius decreases left to right
– Atomic radius increases top to bottom
• Cations are metals that lose electrons.
– The cation is smaller than the original atom.
• Anions are nonmetals that gain electrons
– The anion is larger than the original atom.
Ionization Energy
• The energy required to remove an electron from an atom
is called ionization energy.
– The energy required to remove the next electron is
called the second ionization energy
• IE increases as you go left to right
– There will be some peaks and valleys
• IE decreases as you go down a group
• For each element, when the valence shell is empty, the
next electron must come from a shell that is full….much
more energy required.
– For example….aluminum. Al has 3 valence electrons. After
those first 3 electrons have been removed, it takes a lot more
energy to remove the 4th electron which comes from the core.
Electron Affinity
• Electron affinity is a measure of the
change in energy of an atom when an
electron is added to it.
– If the addition of an electron makes the atom
more stable, energy is given off.
– When the addition of an electron makes the
atom less stable, energy must be added.
• This happens to the alkaline earth metals and the
noble gases
Electronegativity
• Electronegativity refers to how strongly the
nucleus of an atom attracts the electrons
of other atoms in a bond (stinginess).
• Moving from left to right across a period,
electronegativity decreases (noble gases,
generally, do not have electronegativities)
• Moving down a group, electronegativity
decreases.
Notice that the noble gases are missing….just know
that Kr and Xe have an electronegativity and they
can bond with fluorine to make stable compounds.
Shielding Effect
• Shielding effect is the result of full energy
levels separating the outer electrons from
the nucleus.
– Going down a group, the shielding effect
increases due to the increased number of
energy levels “shielding” the outer electrons
from the nucleus.
– Going across a row, the shielding effect
remains the same.
Summary of Periodic Trends
Remains the same
Shielding Effect
Increases
Ionization Energy, Electronegativity,
Electron affinity
Decreases
Atomic size, Ion size
Decreases
Increases
Ionization Energy
Electron affinity
Electronegativity
Nuclear Charge
Atomic size
Shielding Effect
Ion size
Nuclear Decay
• A nucleus is held together by a
nonelectrical, nongravitational force called
the nuclear force.
• Some nuclei are more stable than others.
• When a nucleus is unstable, it can attempt
to increase its stability by altering its
number of neutrons and/or protons.
• This is the process of radioactive decay.
Alpha Emission
• In alpha decay, the nucleus emits a
particle that looks just like a helium
nucleus.
– 2 protons and 2 neutrons
• Alpha particles are not very penetrating
– They can be stopped by paper, skin
• Usually emitted by large nuclei
4
2
4
2
He= α
• When a nucleus undergoes alpha decay
– Subtract 4 from the mass number
– Subtract 2 from the atomic number
• Example of alpha decay
238
92
4
2
234
90
U → He+ Th
Beta Emission
• A beta particle is identical to an electron
• A nucleus that has too many neutrons,
changes a neutron into a proton and an
electron and emits the electron.
1
0
0
−1
1
1
n→ e + p
• Symbol for beta particle can be either
0
−1
e or
0
−1
β
• Need a sheet of metal or plastic to shield it
• When a nucleus undergoes beta decay
– The mass number remains the same
– The atomic number goes up by 1
• Example of beta decay
14
6
0
−1
14
7
C → e+ N
Positron Emission
• A positron is like an electron with a
positive charge.
• In positron emission, the nucleus changes
a proton into a neutron and a positron and
emits the positron.
1
1
0
+1
1
0
p → e+ n
• Symbol for positron can be either
0
+1
e or
0
+1
β
• A nucleus does not have enough neutrons
• When a nucleus undergoes positron
emission
– The mass number remains the same
– The atomic number drops by 1
• Example of positron emission
8
5
0
+1
8
4
B→ e+ Be
Electron capture
• In electron capture, the nucleus does not
have enough neutrons
• Its nucleus captures a low energy electron
and combines it with a proton to form a
neutron.
0
−1
1
1
1
0
e+ p→ n
• When a nucleus undergoes electron
capture
– The mass number remains the same
– The atomic number drops by one
• Example of electron capture
16
9
0
−1
16
8
F + e→ O
Gamma Rays
• Gamma rays are electromagnetic radiation
that have no mass and no charge.
• Gamma rays usually accompany other
forms of nuclear decay and are the most
penetrating of all.
– Need 2-3 inches of lead or 6-9 feet of
concrete to be shielded from gamma rays
• Think of those needing to give off a
gamma ray as having too much pent up
energy.
– Once they emit the gamma ray, they feel
much better.
• Example of gamma emission (m means
metastable)
90 m
43
0
0
90
43
Tc→ γ + Tc
• Notice that the element does not change
Nuclear Stability
• Nuclei undergo decay to achieve greater
stability. You can use the periodic table to
predict the kind of decay that an isotope will
undergo.
– If an isotope’s mass number is greater than its atomic
mass, the nucleus will try to gain protons and lose
neutrons….therefore beta decay.
– If an isotope’s mass number is less than its atomic
mass, the nucleus will try to lose protons and gain
neutrons….therefore either positron emission or
electron capture
– Alpha emission is for the large nuclei, usually with
atomic numbers of 80 or greater.
Stability Graph
Fusion
• Fusion involves combining two small
nuclei into one larger nucleus.
– This occurs on our sun and other stars.
– The product is usually not radioactive.
– Fusion occurs at very high temperatures.
– Need a magnetic field to contain the plasma
required for fusion to occur.
– Researchers have not as yet been able to
control a magnetic field on Earth. If they
could, then we could have fusion reactors
instead of fission reactors.
Fission
• Fission involves the splitting of one large
nucleus into two smaller nuclei.
– This occurs in nuclear power plants.
– The products are highly radioactive.
– Three fissionable nuclei.
• U-235, Pu-239, and Cf-252
– Fission occurs when one of the above nuclei
are bombarded with neutrons.
Samples of fission reactions
235
92
1
0
90
38
143
54
1
0
U + n→ Sr + Xe +3 n
7 x 108 yr
235
92
1
0
28.1 yr
511 msec
92
36
141
56
1
0
U + n→ Kr + Ba +3 n
1.84 sec
235
92
1
0
132
51
18.3 min
101
41
1
0
U + n→ Sb + Nb +3 n
2.79 min
7.1 sec
Mass Defect and Binding Energy
• When protons and neutrons come together
to form a nucleus, the mass of the nucleus
is less than the sum of the masses of its
constituent protons and neutrons.
– This difference is mass is called the mass
defect.
– The mass lost is this process is released in the
form of energy
• If we reverse the process, this is the same
amount of energy, called the binding
energy, required to decompose the
nucleus back into protons and neutrons.
• The relationship between mass and
energy is given by Einstein’s equation
E = mc2
where E = energy in Joules
m = mass in kg
c = speed of light, 3 x 108 m/sec
Half-life
• Half-life is the time required for ½ of a
radioactive isotope to decay
– Some half-lives are very short and are useful
in medicine.
– Some half-lives are very long and are useful
in dating artifacts.
– Each radioactive isotope has its own half-life
which never changes….it is like a finger print
Half-life problems
• Half-life problems concern themselves
with four ideas
– Total time – the overall time covered by the
problem
– Half-life – symbol is t1/2
– Beginning mass (or beginning # of particles)
– Ending mass (or ending # of particles)
• Two examples….
Example 1
Phosphorus-32 has a half-life of 28 days. How
much time is needed for 84 g of the
phosphorus-32 to decay to 5.25 g?
Example 2
Hydrogen-3 has a half-life of 12.3 years. If you
begin with 100 g of hydrogen-3, how much is
left after 36.9 years?