Download Atomic Mass Unit and Isotopes

Which scientist performed the Gold Foil
Experiment?
 What were the 2 main conclusions from
this experiment about the structure of
the atom?
 What two calculations were made as
result of the Oil Drop Experiment?

Over 100 subatomic particles have been
discovered. As a class, they are called
nucleons. The three best-known subatomic
particles are represented as follows:
proton
Symbol
Charge
Mass,g
p or +
+1
neutron electron
n
0
1.673x10-24 1.673x10-24
e-1
9.11x10-28
Elements differ from one another by the
number of protons they contain in their
nuclei (known as an element’s atomic
number), and are arranged left to right,
row by row, on the periodic table.
Since only protons and neutrons are
located inside the nucleus, the net
charge on any nucleus is positive. Ex:
Barium has an atomic number of 56, so a
barium nucleus has a charge of +56.
Since neutrons have about the same mass
as protons, both contribute to an atom’s
mass. An atom’s mass number is the sum
of protons and neutrons in its nucleus.
Ex: Barium- _____ has 56 protons and 81
neutrons, thus a mass number of
56 + 81 = 137.
Why is atomic mass listed on the periodic
table not in whole numbers? Because,
atomic mass is an average of the masses
of all available isotopes of the element,
based on percent occurrence. Ex:
Barium’s atomic mass is 137.33 amu
(atomic mass units) as read from the
periodic table, suggesting that barium
has isotopes other than barium-137 that
bring the average mass up slightly.
Electron mass is so low it is considered to be
negligible for the atom. Electrons do however
affect charge of an atom. The overall charge
of an atom is protons minus electrons. In a
neutral (zero-charged) atom, the number of
protons = the number of electrons.
An atom is neutrally charged unless electrons
have been removed or added through
chemical or electrical processes. Then the
atom is charged and is called an ion.
A positively charged atom is called a cation; a
negatively charged atom is called an anion.
Examples:
Element
Charge
state
Electron
movement # protons # electrons
calcium
neutral
None
20
20
calcium
+2 charge
2 lost
20
18
iodine
neutral
None
53
53
iodine
-1 charge
1 gained
53
54

On the last test 15% of students earned
100%, 75% of students earned 85% and
10% of students earned 70%. What was
the average percentage for the class on
the test?
Not all atoms of the same element contain
the same numbers of neutrons. Isotopes
are atoms of an element that contain
different numbers of neutrons, thus
altering mass number. Many elements
exist in isotopic form.
Nuclear notation (Isotopic nomenclature) for isotopes:
Barium-137 (Ba-137)
137
Ba
protons =
neutrons =
electrons =
+2
56
Since atomic number for an element never changes, it is
often dropped from isotopic notation and only mass
number and maybe charge is represented:
137
Ba
+2
137
or
Ba
Three important isotopes of hydrogen:
 protium (H-1) : most common isotope;
1 proton, 0 neutrons, 1 electron
 deuterium (H-2): found mostly on the sun and stars, in
heavy water;
1 proton, 1 neutron, 1 electron
 tritium (H-3): found mostly on the sun and stars, used
in many military weapons;
1 proton, 2 neutrons, 1 electron
Deuterium and tritium are the primary isotopes involved
in fusion on the sun. The sun converts 4 million tons of
hydrogen into helium every second!
AMU- Atomic Mass Unit is defined as 1/12
the mass of the carbon-12 isotope,
which is assigned 12 amus.
Atomic Mass- is the weighted average
mass of the atoms in a naturally
occurring sample of the element
-reflects both mass and abundance of
isotopes as they occur in nature
Sample Problem:
Radioactivity, the spontaneous emission of
high energy particles or rays by unstable
atomic nuclei, can produce subatomic
particles. The transuranium elements are
the most naturally radioactive. Nucleic
decay produces alpha particles, beta
particles, or gamma rays.
Alpha particles (α) – represented by
4
2He
,
are high speed helium nuclei with a +2
charge and fairly low penetrating power.
Beta particles (β) – represented by 0-1β , are
high speed electrons with a -1 charge and
high penetrating power.
Gamma rays (γ) – represented by 00γ , are
high energy electromagnetic waves with no
charge and extremely high penetrating
power.
A nuclide’s radioactive stability can be
predicted by its neutron to proton ratio. For
smaller atoms (At#<20), the stable ratio is
1:1. As the size of the atom increases, the
ratio does too, to a maximum of 1.5:1.
Atoms within the band of stability, indicated
by the graph on p.866 are stable and do
not undergo spontaneous radioactive
decay.
Atoms outside the band of stability, above or
below, are prone to decay to approach a
more stable state.
Nuclear Equations – Nuclear reactions are
represented in equation form, using nuclear
notation. The sum of the mass numbers
(superscripts) on one side of the equation
must equal the sum on the other. The same
applies to atomic numbers (subscripts).
Many nuclear reactions emit gamma rays,
which have no mathematical effect on the
equation, but a huge effect on energy
released. There are numerous types of
nuclear reactions. Examples include:
alpha decay – the emission of alpha
particles. Ex: alpha decay of uranium238
238
234
92U
→
4
90Th
+
0
2He
+ 20γ
beta decay – the emission of beta
particles. Ex: beta decay of iodine-131
131
53
I
131
→
54Xe
0
+
-1β
electron capture – results in the emission of xrays. Ex: electron capture by rubidium-81
81
0
37Rb
+
-1e
81
→
36Kr
+ x-ray photon
neutron bombardment – used to initiate fission
reactions in nuclear power plants
235
92U
+
1
0n
→
236
92U (a less stable nuclide)
Nuclear Half-Lives – Each radioactive
isotope has what is known as a half life,
the time it takes for half the mass of an
available sample to decay into another
substance. Half-lives can range from
microseconds to millions of years. (See
p.871.)

Half-life is used to calculate either how
long it would take to result in a level of
the isotope that could be tolerated by
living things or to calculate the mass of
an isotope remaining after a period of
time.
mfinal = minitial x .5t/T
t/T is the number of half lives a sample has
been through, where t = time lapsed and T =
half life (must have unit agreement)
Sample problems:
1. How much of a 25.0g sample of radon-222 will
remain after 31 days?
t/T = 31days/3.8days = 8.15
mfinal = 25.0g x (.5)8.15 = 0.0875g
2.
How much of a 350g sample of a radioacitve
isotope remains after 5 half lives has passed?
350g → 175g → 87.5g → 43.75g → 21.875g → 10.9375g 
Or:
mfinal = 350g x .55 = 10.9375g
11g
Fission vs. Fusion – Two nuclear processes
used to produce massive amounts of
energy.
Fission is the splitting of large nuclei,
resulting in the formation of smaller
nuclei. Neutrons are used to initiate the
process, which will continue on its own
(self-sustaining) with each fission process
producing more neutrons to initiate more
fissions. This process is currently used in
nuclear power plants… many issues…
reaction control, used fuel rod disposal,
etc.
Fusion is the combining of small nuclei into a larger one.
Fusion is the process that takes place on the sun and stars,
a more powerful process than fission. Ex: the fusion of
hydrogen isotopes to produce helium on the sun.
2
1H
2
+
1
1H
1H
3
+
1H
4
→ 2He + ENERGY
4
→ 2He + ENERGY
Research continues today to reproduce fusion in a
commercially viable way here on earth, a utopia for
energy supply. The major problem is reaction
containment, as fusion produces temperatures in excess
of 40 million Kelvin.
The quantum mechanical model of the atom
is based on the study of waves and light.
Important wave definitions and relationships:
amplitude – the distance from the crest to
the midway origin line of the wave.
wavelength (, the Greek letter lambda)
– the distance between similar points in a
set of waves, such as crest to crest or
trough to trough; normally expressed in a
meter-type unit.
frequency (, the Greek letter nu) – the
number of waves that pass a given point
per unit time; usually expressed in
cycles/second (sec-1 or 1/sec),
commonly known as Hertz, Hz.
Frequency, wavelength, and energy are related
to one another mathematically through:
c = 
and
E = h
where “c” is the speed of light (3.00x108 m/s)
and “h” is Planck’s constant (6.626x10-34 J/Hz).
(Watch for unitary agreement!)
Inspection of these relationships suggests that
frequency and wavelength are inversely
proportional and that energy and frequency
are directly proportional.
Electromagnetic Radiation (EMR) – a form of energy
that travels in waves. The EMR spectrum includes
all visible light and many other forms of wave
energy:
Fill in these blanks for the EMR spectrum:
The highest energy visible light is violet light.
The lowest energy waves in the EMR
spectrum are radiowaves.
Visible red light has shorter wavelength than
infrared light.
Visible green light has lower energy than
Xrays.
Visible indigo light has lower frequency than
ultraviolet light.
EMR travels at the speed of light in a vacuum,
3.00x108 m/s. Types of EMR differ from each
other by wavelength, frequency, and resulting
energy.
All EMR is the result of electron movement
between energy levels within atoms.
When an electron moves to a higher energy
level, farther from the nucleus, energy is
consumed and the atom is referred to as
“excited”.
As the electron returns to a lower energy level,
energy is emitted in the form of photons
discrete packets of radiant energy that travel
in waves (suggested by Einstein).
Because of unique arrangements of electrons within
atoms, each element has a characteristic light it
emits when exposed to a sufficient amount of heat
or electricity.
When examined through a spectrophotometer, a
device that breaks light into its component waves,
the element’s bright-line spectrum is observed.
(See p.126 for the BLS for hydrogen.) A BLS is not a
continuous spectrum like the EMR spectrum, but it is
distinct lines of color that correspond to very
specific wavelengths, frequencies, and energies.
The energies correspond to the energies required to
move electrons farther from the nucleus and that
emitted when they return to their original locations.
Elements also have unique and specific
energy requirements to completely remove
an electron from an atom. Ionization
energy is the amount of energy required to
remove an electron from a gaseous atom
of an element. The energy required is
equal to the energy emitted when an
electron is added to an atom from some
other source.
Of course, ionization energy is greater than
the various energies required or emitted as
electrons move within an atom.
Work to relate wave behavior, atomic energy levels,
and electrons by Louis deBroglie and Erwin
Schrodinger led to the branch of physics called
quantum mechanics.
Quantum mechanics and probability, the statistical
likelihood of an occurrence, are the basis for
current atomic models. Electrons are particles with
wave characteristics. Probability helps describe
the behaviors and positions of electrons.
One important assumption of today’s model is the
Heisenberg Uncertainty Principle which states that
it is not possible to describe both position and
speed of an electron at the same point in time.
An important distinction to make is that
between an orbit (a pathway, as in
Bohr’s Model of the atom) and an
orbital.
An orbital is a region of space around a
nucleus in which an electron is most likely
to be found. It is not a barrier or
pathway, only a model describing the
likely occupied area.
Schrodinger developed a mathematical
model for the wave behavior of an
electron.
The very complex equation contains four quantum
numbers that are used to describe electron
location and behavior…
n
Principle Quantum Number





refers to the energy level location of the
electron
n is 1,2,3,4,… n (up to 7)
the number of sublevels in the energy level =
n (up to 4)
the number of available orbitals in the
energy level = n2 (up to 16)
the greatest number of electrons contained
in any one energy level = 2n2 (up to 32). Ex:
where n=3, the maximum number of
electrons is 2(3)2 = 18.
l


Azimuthal or Orbital Quantum Number
refers to the sublevel location of the electron
l is either s, p, d, or f
s-sublevel orbitals are spherical in shape. They are
closest to the nucleus; thus electrons located here
have the lowest relative energies within the energy
level. There is only 1 s-orbital per energy level.
p-sublevel orbitals are dumbbell-shaped. They are
a little farther from the nucleus; thus electrons
located here have a little higher relative energies
within the energy level. There are 3 p-orbitals (px, py,
pz) per energy level, starting with energy level 2.
d-sublevel orbitals are cloverleafshaped, very complex. They are even a
little farther from the nucleus; thus
electrons located here have even higher
relative energies within the energy level.
There are 5 d-orbitals per energy level,
starting with energy level 3.
f-sublevel orbitals are extremely complex
in shape. They are farthest from the
nucleus; thus electrons located here
have the highest relative energies within
the energy level. There are 7 f-orbitals
per energy level, starting with energy
level 4.
ml Magnetic Quantum Number
 designates
the number of orbitals on
each sublevel and in which specific
orbital the electron is likely located
 describes the home orbital’s
geometric orientation about the x, y,
and z axes
ms Spin Quantum Number



describes the direction of spin for the electron
ms is either +1/2 (clockwise spin) or –1/2 (counterclockwise spin).
The Pauli Exclusion Principle is very important
here and throughout the quantum model. It
states that:
1) no more than 2 electrons may occupy any
one orbital;
2) if electrons occupy the same orbital, they
must have opposite spin directions;
3) thus, no two electrons in one atom will have
all four quantum numbers identical.
Electrons are arranged from orbitals with
lowest energy (closest to the nucleus) to
orbitals with highest energy (farthest from
the nucleus). Configurations can be
written for neutral atoms and their ions.
There are 3 basic rules that govern orbital filling:
1. Aufbau Principle – electrons enter lowest
energy levels first.
2. Hund’s Rule – when electrons fill orbitals
within a sublevel, each orbital is occupied
by one electron before any orbital obtains
two electrons. All electrons in singly
occupied orbitals have the same spin
directions.
3. Pauli Exclusion Principle – described earlier
Reading e-configs and orbital(energy) diagrams
from the Periodic Table:
Recognize the significance of the structure of
the Periodic Table:
 Row numbers correspond to Energy Levels
(Principle Quantum Numbers, n)
 The PT is arranged in blocks: s-block (groups 1
& 2), p-block (groups 13-18), d-block (groups
3-12), and f-block (the lanthanide and
actinide series)
Read the PT from left to right, row by row, adding
new electron(s) (indicated by superscript) each
step of the way:
1. s-block (maximum of 2 electrons) and p-block
(maximum of 6 electrons) electrons are in the
energy level indicated by current row number
2. when you reach d-block (maximum of 10
electrons), drop 1 energy level below the current
row number
3. when you reach f-block (maximum of 14
electrons), drop 2 energy levels below the
current row number
4. when you reach p-block, return to the original
energy level indicated by row number
Ex: for the element arsenic (As):
1s22s22p63s23p64s23d104p3
total up the superscripts and you get 33,
the atomic number of arsenic
Sample orbital diagrams and electron configurations:
Nitrogen (N) - ___ protons, ___ electrons
electron configuration:
1s22s22p3
orbital/energy diagram:
__
1s
__
2s
__ __ __
2p
Manganese (Mn) - ___ protons, ___ electrons
e-config:
1s22s22p63s23p64s23d5
orbital/energy diagram:
__
1s
__
2s
__ __ __
2p
__
3s
__ __ __
3p
__
4s
__ __ __ __ __
3d
Orbital/energy diagrams and electron configurations
can be written for ions too. It is best to begin with the
neutral diagram or configuration and make your
additions to or deletions from it!
Highest energy level electrons, referred to as valence
electrons, are the first and often the only electrons
involved in bonding. (Ions form as result of ionic
bonding.)
Note: Atoms tend to ionize in the easiest way for them to
end with a full highest energy level. “s” and/or “p”
sublevels will always be the highest energy level’s
sublevels. A full highest energy level would contain 8
electrons… the octet rule.
Samples:
Formation of an anion →
Sulfur (S) - ___ protons, ___ electrons
e-config:
1s22s22p63s23p4
Sulfur ion (S-2) – ______________ electrons
e-config:
1s22s22p63s23p6
Formation of a cation →
Iron (Fe) - ___ protons, ___ electrons
e-config:
1s22s22p63s23p64s23d6
energy diagram:
__ __
1s 2s
__ __ __ __ __ __ __ __ __ __ __ __ __
2p
3s
3p
4s
3d
Remember that valence electrons are always
the first and often the only electrons to
become involved in bonding.
It is imperative that you begin with the neutral
diagram or configuration and make
changes to it when dealing with cations!!
Remove electrons from highest energy levels
first, not necessarily the last filled.
Iron ion (Fe+2) - _________________ electrons
e-config:
1s22s22p63s23p64s23d6
energy diagram:
__ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
__ __ __ __ __
3d
Iron ion (Fe+3) - ___________________ electrons
e-config:
1s22s22p63s23p64s23d65
energy diagram:
__ __ __ __ __
__ __ __ __ __ __ __ __ __ __
1s 2s
3s
2p
3p
4s
3d
Develop the e-config for Nickel (II) ion (Ni+2):
Neutral Ni - ___ protons, ___ electrons
e-config:
1s22s22p63s23p64s23d8
Nickel ion (Ni+2) - ________________ electrons
ion e-config:
1s22s22p63s23p63d8
As a general rule, metals tend to lose
electrons, become positively charged,
when they ionize.
Nonmetals tend to gain electrons to
become negatively charged.
Metalloids, those elements on and below
the stairstep on the periodic table, can
do either depending on the
circumstances.
Based on energy diagrams and configurations, predict
charge(s) for the following elements when they ionize:
Nitrogen ion 1. Classify the element –
nonmetal, likely to become negatively charged
2. Draw energy diagram or e-config of the neutral
element:
energy diagram:
__ __ __ __ __
e-config: 1s22s22p3
1s 2s
2p
3. Locate valence electrons, and determine what must
happen to meet the octet rule.
4. Predict charge(s) - Nitrogen is likely to gain 3 electrons
to become -3 charged.
You predict charge(s) for Calcium ion and Arsenic ion:
Calcium –
a metal, will lose e- to become + charged
econfig: 1s22s22p63s23p64s2
valence electrons : level 4, 2 of them
charge prediction:
lose level 4 electrons to be +2 charged
Arsenic –
a metalloid, can act like a metal (lose e- to
become + charged) and can act like a nonmetal
(gain e- to become – charged).
econfig: 1s22s22p63s23p64s23d104p3
valence electrons : level 4, 5 of them
charge prediction:
as a nonmetal –
gain 3 electrons to fill level 4 electrons to be -3
charged
as a metal –
lose 4p-sublevel’s 3 electrons to be +3 charged,
or also lose the 4s-sublevel’s 2 electrons to be +5
charged
Once atomic number passes 18 or so, it becomes
cumbersome to write full configurations. Thus, it is convention
to use a shorthand format:
1.
Locate the element to be read.
2.
Find the noble gas that precedes it in atomic number
and place its symbol in brackets to represent all of its
electrons.
3.
Begin reading at the next row of the PT, following
guidelines, until you reach the target element.
Ex: for the element lead (Pb)
[Xe]6s24f145d106p2
total up the electrons represented by xenon and the
superscripts and you get 82, the atomic number of lead
You do neutral platinum, neutral selenium,
and Sn+4:
Normally, electron configurations are easy to
predict, but there are a few exceptions in
the areas of the transition metals. The most
notable exceptions are chromium and
copper. Their actual configurations are
slightly altered from predicted in order to
minimize energy.
Chromium, Cr
predicted: 1s22s22p63s23p64s23d4
actual: 1s22s22p63s23p64s13d5
Copper, Cu
predicted: 1s22s22p63s23p64s23d9
actual:
1s22s22p63s23p64s13d10
These exceptions apply to other elements
in their groups (under them) on the
periodic table, but the farther down the
table the less likely the exception
applies.
Lewis Structures depict an atom of an
element with its valence electrons.
 Valence electrons – electrons in the
highest energy level of an atom.
Because of filling order, valence
electrons will always be in “s” and/or “p”
sublevels.
 Valence electrons are normally the first,
and often the only, electrons to become
involved in bonding.)
Ex:sulfur (neutral)
e-config:
Lewis Structure:
To construct:
1. Write the electron configuration for the atom.
2. Write the element symbol to represent the nucleus
and all electrons other than valence electrons.
3. Begin placing valence electrons around the symbol
in the same manner that you fill orbital diagrams.
a.
b.
Place “s” electrons to the right of the symbol.
Proceed clockwise, placing “p” electrons one at a time
on the remaining three sides of the symbol. Then double
up to complete placement of remaining electrons.
You do Lewis diagrams for thallium,
molybdenum, and phosphorus: